Chemistry Dictionary

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1. Absolute Zero - The theoretical lowest possible temperature, defined as 0 K (kelvins) or -273.15 °C. At absolute zero, a system is thought to have minimum internal energy, with virtually all molecular motion ceasing (in classical conceptions).

   The concept of absolute zero has far-reaching implications in various fields:

   1. Thermodynamics: Foundational reference point; all kelvin-based temperature scales revolve around absolute zero.
   2. Entropy: According to the third law of thermodynamics, a perfect crystal at absolute zero would have zero entropy.
   3. Quantum effects: Real quantum systems still retain zero-point energy even at 0 K.

   In practical applications (e.g. cryogenics, superconductors), we never truly reach absolute zero, but we approach it. The interplay of activation-energy for chemical reactions at near-zero temperatures becomes extremely limited, and even arrhenius-base definitions can change in exotic states of matter. For mole-based calculations, avogadro-constant remains relevant even in cryogenic realms.

2. Absorption - The process by which one substance (the absorbate) penetrates and becomes uniformly distributed throughout another (the absorbent). In contrast to adsorption, where molecules are only held on the surface, absorption involves a full assimilation into the bulk of the material.

   • Physical absorption: Typically driven by intermolecular forces, such as van der Waals interactions (e.g. carbon dioxide dissolving in water).
   • Chemical absorption: Involves a chemical reaction, resulting in products bound or changed at the molecular level (e.g. reactive gases absorbed by alkaline solutions).

   Beer-Lambert law is a fundamental principle in absorption spectroscopy, describing the relationship between absorbance and the concentration of an absorbing species.

   A common context is light absorption in a medium:

   where

   • is the absorbance (unitless),
   • is the molar absorptivity (L·mol·cm),
   • is the path length of the sample (cm),
   • is the concentration of the absorbing species (mol·L).

   

3. Acetate - The conjugate base of acetic acid (CH₃COOH), represented as CH₃COO^-. It arises when an Arrhenius Acid (like acetic acid) donates a proton. Acetates are versatile in chemistry:

   1. Buffering capacity: Acetate solutions can act as buffers when combined with acetic acid, helping to stabilise pH.
   2. Coordination complexes: Metal acetates are common in both laboratory reagents and biological systems.
   3. Biochemical significance: In cellular metabolism, acetyl-CoA contains an acetyl group, functionally related to acetate.

   Chemical reaction:
   When acetic acid deprotonates,

   Because acetate can sometimes accept a proton under specific conditions, it exhibits limited amphoteric behaviour in special contexts.

4. Acid - An acid as per the Brønsted definition is any substance capable of giving off a proton, or Hydrogen particle. As per the Lewis definition an acid is any substance capable of accepting a pair of electrons.

   Notice that these two definitions are complementary, in that if: an acid gives off a proton, it is releasing a charged particle H+ into the environment; yet at the same time it must accept those electrons 2e- from the proton in order to release it.

   An acid is written as AH and its conjugated base as A-. The couple is written as: AH/A-.

   They are linked together by the two half reactions:

   Thus the two half reactions together form the complete reaction:

   Notice that the acid is giving up its H to the base, and in exchange the A receives the negative charge - which in fact represents the two electrons given up by the Hydrogen.

   If one were to follow the electrons they go from: H to A, and from the B to the H in order to form the bond.

5. Activation Energy - The minimum energy threshold that reacting molecules must overcome in order for a chemical reaction to proceed. The Arrhenius equation famously captures the effect of activation energy on reaction rate:

   where

   • is the rate constant (s or similar),
   • is the pre-exponential factor,
   • is the activation energy (J·mol),
   • is the gas constant (8.314 J·mol·K),
   • is the absolute temperature (K).

   Why it matters:

   • Catalysts lower the effective by providing an alternative reaction pathway.
   • Temperature dependence: As increases, grows, accelerating reactions.

   While arrhenius-acid uses “Arrhenius” in a different acid-base context, the conceptual link is that Svante Arrhenius contributed significantly to both reaction kinetics (the Arrhenius equation) and early acid-base theory.

6. Adsorption - The accumulation of molecules on the surface of a solid or liquid, creating a higher concentration at the interface than in the bulk. This phenomenon contrasts with absorption, where the material diffuses into the volume.

   Two primary types of adsorption are distinguished:

   • Physisorption: Involves weak van der Waals forces.
   • Chemisorption: Involves stronger chemical bonding at the surface.

   A simple way to model adsorption is through Langmuir isotherms, describing how the quantity of adsorbate on a surface varies with its partial pressure or concentration:

   where

   • is the fraction of the surface covered by adsorbate,
   • is the adsorption equilibrium constant,
   • is the partial pressure of the adsorbate.

   Many allotropes of carbon (like activated carbon) serve as excellent adsorbents in water purification and gas separation.

7. Alcohol - Any organic compound bearing one or more hydroxyl (-OH) groups attached to a carbon skeleton. Common examples include methanol (CH₃OH) and ethanol (C₂H₅OH).

   Alcohols exhibit several key properties:

   • Polarity: The OH group confers hydrogen-bonding capacity, affecting solubility in water.
   • Boiling Point: Tends to be higher than that of corresponding hydrocarbons due to hydrogen bonding.
   • Oxidation: Primary alcohols can oxidise to aldehydes and then further to carboxylic acids.

   Azeotropic Behaviour: some binary mixtures of alcohols (e.g. ethanol-water) form an azeotrope, complicating simple distillation.

   where:

   • is the number of carbon atoms
   • is the number of unsaturations

   Under strongly acidic (arrhenius-acid) conditions, certain alcohols dehydrate to form alkenes or ethers.

8. Aldehyde - Organic compounds featuring a formyl group (-CHO) bound to a carbon backbone. The carbonyl carbon (C=O) is at the end of the chain, distinguishing aldehydes from ketones (where the carbonyl is internal).

   Aldehydes are characterised by several key features:

   • Methanal (Formaldehyde): HCHO, the simplest aldehyde.
   • Ethanol oxidation: Oxidising an alcohol like ethanol (CH₃CH₂OH) can yield ethanal (CH₃CHO).

   Notable Reaction: Benedict's or Fehling's test can detect aldehydes via colour change, linking to redox chemistry and stoichiometric analysis. Here, avogadro-constant helps relate theoretical yield to actual moles involved.

   Under strong arrhenius-acid catalysis, aldehydes can further transform into other derivatives (e.g. acetals).

9. Alkali Metal - Any of the elements in Group 1 of the periodic table (e.g. Li, Na, K, Rb, Cs, Fr). They are characterised by:

   • Having a single electron in their outermost shell
   • Exhibiting high reactivity, especially with water
   • Forming strong bases (alkaline solutions) upon reaction with water

   Example reaction: Sodium reacting with water:

   Because of their reactivity, alkali metals are typically stored under oil. They share similarities yet are distinct from the alkaline-earth-metal group next door.

10. Alkaline Earth Metal - The elements in Group 2 of the periodic table: Be, Mg, Ca, Sr, Ba, and Ra. Key traits include:

    • Two valence electrons, contributing to their reactivity
    • Formation of oxides with the generic formula MO (where M = metal)
    • Less reactive than alkali metals, but still form hydroxides and other alkaline compounds

    Common example: Calcium reacting with water:

    Their name stems from the fact that their oxides and hydroxides were once known as “earths” and exhibit basic (alkaline) properties in solution.

11. Alkoxide - Conjugate bases derived from alcohols by deprotonation of the hydroxyl group. Formally written as R-O^-, they serve as strong bases and nucleophiles in both laboratory and industrial organic reactions.

    Alkoxides are characterised by several key features:

    • Preparation: Reacting alcohols with alkali metals or strong arrhenius-base can yield alkoxides.
    • Reactivity: Alkoxides can perform nucleophilic substitutions, facilitate esterification, and more.
    • Ionic Character: R-O^- is an anion, often stabilised by cationic counterparts like Na⁺ or K⁺.

    Alkoxides highlight the acid-base interplay, as their formation is contingent on the acidity of the parent alcohol.

12. Allotrope - Different structural forms of the same element, characterised by distinct physical and chemical properties. A classic example is carbon, which can appear as:

    • Diamond: Each carbon is sp³-hybridised in a tetrahedral network.
    • Graphite: Carbon atoms form planar hexagonal sheets; these sheets are loosely held together, allowing for lubrication and layered structures.
    • Buckminsterfullerene (C₆₀): A spherical “buckyball” arrangement.
    • Graphene: A single layer of graphite with extraordinary electrical and mechanical properties.

    Allotropes can drastically differ in hardness, conductivity, and adsorption capacity (e.g. activated carbon).

13. Amalgam - An alloy in which mercury is a significant component, often combined with metals such as silver, tin, or gold. Traditionally used in dental fillings, though usage has declined due to mercury toxicity concerns.

    Amalgams exhibit several key features:

    • Formation: Mercury dissolves or reacts with metal surfaces, creating metal-mercury bonds.
    • Physical Properties: Amalgams can be soft, brittle, or hard, depending on the composition.
    • Chemistry Context: Mercury's unique liquid state at room temperature influences the activation-energy of certain reactions; also, elemental mercury is an interesting allotrope candidate in discussions of phase transitions.

    In some amalgams, electronegative metals form stable anion-like species within the mercury matrix, highlighting complex redox interactions.

14. Amino Acid - Organic molecules containing both an amine group (-NH₂) and a carboxylic acid group (-COOH). They are considered the building blocks of proteins. A hallmark property is their amphoteric nature, meaning they can act as an acid or base.

    Amino acids are characterised by several key features:

    • Alpha Carbon: Central carbon bonded to an amino group, a carboxyl group, a hydrogen, and a distinctive side chain (R-group).
    • Isoelectric Point (pI): The pH at which the molecule has no net charge overall.

    If is the midpoint between the relevant acid dissociation constants and :

    These p values measure the tendency of the acid/base groups to donate or accept protons. To handle these calculations in bulk, scientists rely on Avogadro's Constant for conversion between microscopic (individual amino acids) and macroscopic (moles) quantities.

15. Ammonia - A colourless, pungent gas with the formula NH₃. It dissolves readily in water, forming ammonium hydroxide (NH₄OH).

    Ammonia is characterised by several key features:

    • Basic Character: Ammonia is a weak Arrhenius-base in aqueous solution, producing OH^- ions.
    • Amphoteric Behaviour: Although primarily basic, NH₃ can (under high pressures/temperatures) act as an acid, donating H⁺, thereby linking to amphoteric chemistry.
    • Industrial Relevance: Synthesised mainly via the Haber process, crucial for fertilisers, nitric acid production, and beyond.

    When ammonia dissolves in water:

    Its equilibrium lies to the left, reflecting its modest basicity.

16. Ammonium - The cation NH₄⁺, formed by protonation of ammonia. It often acts as a conjugate acid in acid-base equilibria.

    Ammonium is characterised by several key features:

    • Acidic Tendency: In the Arrhenius-acid sense, NH₄⁺ can donate a proton to water, albeit weakly.
    • Biological Significance: Excess ammonium in organisms must be detoxified (e.g. urea cycle in humans).
    • Ammonium Salts: Common in fertilisers (e.g. ammonium nitrate, NH₄NO₃).

    Small but crucial shifts in this equilibrium can have significant effects on soil pH and metabolic pathways.

17. Ampere - The SI base unit of electric current (symbol A). Current flow often underlies redox processes, electroplating, and various aspects of solution conductivity in chemistry.

    Usage of the ampere in chemistry is widespread, with several key applications:

    • Electrolysis: The magnitude of current, measured in amperes, directly ties to the rate at which a substance (e.g. ammonium conversion) is oxidised or reduced.
    • Faraday's Laws: These laws link charge passed in coulombs (C) to moles of substance. Recall .
    • Amphoteric: The concept of a current in solution can demonstrate how amphoteric species sometimes behave differently at electrodes, either losing or gaining protons/electrons.

    Given charge and time :

    Avogadro-constant often assists in bridging microscopic electron counts and macroscopic measurements of electric current.

18. Ampholyte - A substance that can ionise both as an acid and a base, similar to an amphoteric species but typically referring to molecules that can form both cations and anions depending on pH. Many amino-acids behave as ampholytes, existing in zwitterionic forms at certain pH levels.

    Ampholytes are characterised by several key features:

    • pH Dependence: Ampholytes' net charge changes with solution pH, guiding their behaviour in electrophoresis and ionic equilibria.
    • Relevance: Important in biological systems, where buffers often contain ampholytes to stabilise pH.
    • Arrhenius Perspective: As potential arrhenius-acid or -base donors, ampholytes handle H⁺ or OH^- in solution.

    Such species can be cationic, neutral, or anionic, illustrating their multifunctional acid-base roles.

19. Amphoteric - A substance capable of reacting both as an acid and as a base. Water (H₂O) is the quintessential example (autoprotolysis), sometimes donating a proton, sometimes accepting one. Beyond water:

    • Metal oxides such as ZnO and Al₂O₃ can display amphoteric behaviour.
    • Acetate can (under certain conditions) accept protons or donate them, albeit weakly.

    In an Arrhenius-acid context, an amphoteric species may furnish H⁺ ions, but in an Arrhenius-base context, the same species may yield OH^- or accept H⁺. This dual nature is central to many acid-base equilibrium processes.

20. Anhydride - A compound formed by removing a water molecule (H₂O) from another substance, often from acids. The most common are acid anhydrides, which yield the parent acid upon reaction with water.

    Anhydride types include:

    • Organic Acid Anhydrides: e.g. acetic anhydride, (CH₃CO)₂O, widely used in synthesis (e.g. to form esters or amides).
    • Inorganic Anhydrides: e.g. SO₃ (sulfur trioxide), the anhydride of H₂SO₄ (sulfuric acid).

    Hydrolysis Reaction:

    (For a generic symmetrical anhydride forming two alcohols.)

    Strongly related to arrhenius-acid chemistry, as many anhydrides react vigorously to produce or regenerate acids. They can also be synthesised from partial oxidation or dehydration of aldehydes.

21. Anion - A negatively charged ion, formed when a neutral atom or molecule gains electrons. For instance, halogen atoms become halide anions (F^-, Cl^-, etc.) upon electron gain.

    • Formation: Typically via electron transfer from metals or by the reduction of neutral species in redox reactions.
    • Size & Charge: An anion's radius is generally larger than its parent atom because of increased electron-electron repulsion.
    • Relevance: Anions often function as counter-ions to cations in salts and can affect the solution's acidity (Arrhenius-acid context), pH, and conductivity.

    Tip: Knowing an element's atomic number helps predict whether it tends to form an anion or cation by electron gain or loss.

22. Antifreeze - A substance (commonly ethylene glycol or propylene glycol) added to liquids (like water) to lower their freezing point. This property helps prevent damage to engines, pipelines, and other systems in cold environments.

    The mechanism of antifreeze action involves two key factors:

    • Colligative Properties: Dissolved antifreeze particles disrupt water's crystalline lattice formation.
    • Hydrogen Bonding: Glycols are alcohol-like in structure, forming hydrogen bonds that interfere with pure water's freezing pattern.

    Mathematically, the freezing-point depression can be estimated with:

    where:

    • is the cryoscopic constant,
    • is the molality of solute,
    • is the van't Hoff factor (for non-electrolytes, ).

    High activation-energy processes for crystal nucleation can be further elevated in mixed solutions. In automotive applications, water-antifreeze mixtures can also form non-ideal solutions or azeotrope-like behaviour under high temperatures and pressures.

23. Aromatic Compound - Organic compounds characterised by conjugated pi electron systems forming cyclic, planar rings. These structures follow Hückel's rule ( π electrons for aromatic stability). Although commonly placed under organic chemistry, many fundamental principles are relevant in general-chemistry contexts.

    The classic example is Benzene (C₆H₆):

    • Highly stable ring with delocalised π electrons.

    • Undergoes electrophilic substitution rather than addition reactions.

    • Certain forms of carbon (allotrope structures like fullerenes) have delocalised systems akin to aromatics.

    • Aromatic surfaces can exhibit unique adsorption properties, e.g. polycyclic aromatic hydrocarbons on environmental particulates.

    Hückel's rule for aromaticity:

    where is a non-negative integer.

24. Arrhenius Acid - Defined as any substance that increases the concentration of H⁺ ions in aqueous solution. It dissociates to yield protons:

    This theory, proposed by Svante Arrhenius, laid the groundwork for modern acid-base chemistry.
    Examples: HCl, H₂SO₄, and HNO₃.

    Arrhenius also contributed the Arrhenius equation in kinetics, linking acid-base chemistry to activation-energy. An acid's behaviour complements that of an Arrhenius-base.

25. Arrhenius Base - A substance that increases the concentration of OH^- ions when dissolved in water. Classic examples include NaOH, KOH, and Ca(OH)₂. According to Arrhenius:

    An Arrhenius Acid introduces H⁺, while an Arrhenius base introduces OH^-. When these ions combine, they yield H₂O:

    This acid-base model works well for many aqueous systems, though more general concepts (Brønsted-Lowry, Lewis) extend the ideas further.

26. Atomic Mass - The mass of a single atom of a chemical element, often expressed in atomic mass units (u). One atomic mass unit is defined as 1/12 of the mass of a carbon-12 (C¹²) atom. Key notes:

    • Weighted average: An element's listed atomic mass on the periodic table is typically the weighted average of its naturally occurring isotopes.
    • Mole concept: Avogadro's Constant links atomic mass to macroscopic quantities.

    For illustation, if an element X has two isotopes:

    • X₁ (mass = 10 u, abundance = 20%)
    • X₂ (mass = 11 u, abundance = 80%)

    Then average atomic mass:

    Compare atomic-number, which simply counts protons without reference to mass.

27. Atomic Number - The number of protons in the nucleus of an atom, uniquely identifying the element. Symbolised as Z, it defines the position of the element in the periodic table:

    • Hydrogen (H) has
    • Helium (He) has
    • Lithium (Li) has

    Important distinctions:

    • Atomic mass includes protons and neutrons.
    • Isotopes of the same element have the same Z but different neutron counts.
    • Electron configuration and atomic orbitals are determined by Z, as it sets the electrostatic attraction between nucleus and electrons.

28. Atomic Orbital - A mathematical function describing the probability distribution of an electron around an atom's nucleus. Designated by quantum numbers (n, l, m, s), common orbital types are s, p, d, and f:

    • s-orbitals: Spherical symmetry around the nucleus.
    • p-orbitals: Dumbbell-shaped, oriented along x, y, or z axes.
    • d-orbitals: More complex shapes (clover leaves, doughnuts).

    Their shapes underlie bonding, reactivity, and the formation of various allotropes. Knowledge of the atomic number helps determine electron configuration, influencing how these orbitals are filled.

29. Autoionization - A process where identical molecules spontaneously ionise each other. The classic example is water:

    • pH Scale: Pure water's autoionization sets arrhenius-acid and arrhenius-base concepts, yielding the ion-product .
    • Amphoteric: Substances like water or certain amino-acid side chains can display autoionization or partial self-dissociation.
    • Thermodynamic Variation: Elevated temperatures can shift the equilibrium, affecting .

    Autoionization underscores water's amphoteric nature and is pivotal in discussions of buffer chemistry, acid-base equilibria, and advanced pH calculations.

30. Autoprotolysis - A chemical reaction occuring in water between water molecules themselves H2O in which a proton is transfered between two identical molecules, one playing the role of an acid and the other the base (in the Brønsted context).

    The chemical reaction is written:

    Any solvent which contains acidic hydrogen and lone pairs of electrons with which to accept H+ can demonstrate autoprotolysis, for example ammonia:

31. Avogadro Constant - Denoted or , representing the number of particles (atoms, molecules, ions) in one mole of a substance. Its value is approximately .

    Significance:

    • Bridges atomic scale (individual particles) to macroscopic scale (moles).
    • Underpins stoichiometric calculations, relating atomic mass in grams to the number of discrete particles.

    1 mole of carbon-12 (C¹²):

    • Weighs exactly 12 g.
    • Contains atoms.

    This fundamental constant is central to chemistry's quantitative framework.

32. Azeotrope - A mixture of two or more liquids whose proportions cannot be altered by simple distillation, as the vapour has the same composition as the liquid mixture. This phenomenon reflects strong intermolecular interactions that defy typical Raoult's law predictions.

    • Minimum-Boiling Azeotrope: Commonly seen with ethanol-water (~95.6% ethanol by volume).
    • Maximum-Boiling Azeotrope: Found in certain hydrochloric acid solutions where the solution vaporises at a higher temperature than either component alone.
    • Applications: Azeotrope formation can complicate purification processes; advanced techniques (e.g. pressure-swing distillation or absorption columns) are employed to break azeotropes.

    For a binary mixture, an ideal solution might follow:

    Yet real mixtures deviate significantly (excess enthalpy, unusual activation-energy considerations), leading to these azeotropic points.

33. Antiaromatic - If a molecule has 4n π electrons and has the all other characteristics for aromaticity then it is said to be antiaromatic.

    1. The molecule must be cyclic.
    2. The molecule must be planar, allows for p-orbitals to be parallel and interact.
    3. That the molecule must be formed of a continous ring of conjugated π bonds (p-orbital interaction); if a heteronome participates in conjugation then its doublet must be able to be conjugated with the other π bonds (can be composed of one or more rings).

    Example: cyclobutadiene; the number of π delocalised electrons is 4, its ionic form cyclobutadiene(2-) is aromatic (6 e-).

34. Aromaticity (aryl) - This is the property of being cyclic (a closed ring), and planar with a ring of resonance bonds. Aromatic chemical groups are called aryl. Requirements for aromaticity, known as Huckel's rule, are as follows:

    1. The molecule must be cyclic; can be composed of one or more rings.
    2. The molecule must be planar, allows for p-orbitals to be parallel and interact.
    3. That the molecule must be formed of a continous ring of conjugated π bonds (p-orbital interaction); if a heteronome participates in conjugation then its doublet must be able to be conjugated with the other π bonds.
    4. The molecule must have 4n + 2 electrons in a conjugated system (n ∈ ℕ); normally on sp2-hybridised atoms, and sometimes sp-hybridised.

    To apply the rule, count the number of π eletrons in the molecule, then set the equation equal to this number. Solve for n. If n ∈ ℕ, and the other rules are true then the molecule is indeed aromatic.

    Note

    Example: benzene, 3 double bonds, 6 π electrons.

    In order to know which are the π electrons to be counted, you must look at those which reside in the p-orbitals; the hybrid atoms have one p-orbital each. Thus if each part of the cyclic compound is hybridised sp2 then this means that the molecule is completely conjugated (each atom has one p-orbital) and the electrons in these p orbitals are in fact π electrons.

    Aromatic molcules are very stable and do not react easily with other compounds. Most common aromatic compounds are derivatives of benzene (found in petroleum).

    Aromaticity does not necessarily indicate an odour. This misnomer comes from the association initially given by August Wilhelm Hofmann in 1855 who studied a class of benzene molecules; which did indeed emit odours.

    

    Chemically aromaticity refers to a conjugated system made of alternating single and double bonds in a ring. This arrangement allows for electrons in the molecule's π system to be delocalised; increasing the molecules stability. Such molecules are represented by a resonance hybrid of different strucutres.

    Bond lengths in such a strucutre are the intermediate of the two composing representations. X-ray diffraction has shown that all six carbon-carbon bonds in benzene are of the same length: 1.4 Å. C-C double bonds are typically at 1.35 Å, and single bonds at 1.47 Å.

    

    The ring contains π bonds which are formed from the overlap of atomic p-orbitals; which can interact with each other freely as they are outside the plane of the molecule. Each electron is shared by all six atoms in the ring; there are not enough electrons to form double bonds between all carbon atoms, but the "extra" electrons strengthen the exisiting bonds and give rise to the resonance structure.

B

35. Balanced Equation - A chemical equation in which the number of atoms of each element is the same on both sides of the arrow, preserving mass and charge (when relevant). For example:

    • Ensures stoichiometric integrity: Using avogadro_constant, we can scale microscopic particles to macroscopic amounts.
    • Reactant and product coefficients must reflect actual mole ratios, which influences reaction enthalpy calculations.
    • collision_theory explains how correct proportions of reactants can affect reaction rates.
    • Nuclear equations (e.g. beta_decay) also require balancing nucleons and charge, though the rules differ slightly from typical chemical equations.

    Note

    Balancing is an essential step in quantitative chemistry, ensuring correct reactant-product relationships for further calculations.

36. Balmer Series - A set of visible spectral lines corresponding to electronic transitions in hydrogen that end at the n=2 energy level. Each line in the series is given by the Rydberg formula:

    • Explains distinct colours in hydrogen's emission spectrum, historically key to the bohr_model.
    • Photon frequency or wavelength can be converted to energy (in joules) and scaled by avogadro_constant to relate to mole-based energies.
    • Ties conceptually to binding_energy for the electron in hydrogen, though the electron is never fully removed in these transitions.

    Note

    The Balmer Series underpins atomic theory, revealing discrete energy levels in hydrogen and forming a stepping stone to modern quantum mechanics.

37. Barometer - An instrument measuring atmospheric pressure. Traditionally, a mercury-filled tube in which external air pressure supports a column of mercury.

    • The height of mercury correlates directly to atmospheric pressure.
    • Helps interpret how boiling_point changes with altitude or weather variations.
    • In partial-pressure contexts (daltons_law), knowledge of the barometric pressure is crucial for calculating total or gas_law relationships.
    • Distinct from gauge_pressure, which measures relative to ambient pressure.

    Note

    Mercury barometers set a historical standard for measuring absolute pressure. Modern digital barometers rely on sensor-based technology but functionally do the same job.

38. Base - As per the Brønsted definition it is any substance capable of taking a proton, or Hydrogen particle. As per the Lewis definition a base is any substance capable of releasing a pair of electrons.

    Notice that these two definitions are complementary, in that if a base takes on a proton, it is accepting a charged particle H+ from the environment; yet at the same time it must release its electrons 2e- and share them with the proton in order to form a bond.

    A base is written as B- and its conjugated acid as BH. The couple is written as: B-/BH.

    They are linked together by the two half reactions:

    Thus the two half reactions together form the complete reaction:

    Notice that the base is accepting an H from the acid, and in turn the A receives the negative charge - which in fact represents the two electrons given up by the Hydrogen.

    If one were to follow the electrons they go from hydrogen to the A, then from B to the H in order to form the bond.

39. Battery - A device that converts chemical energy into electrical energy via one or more electrochemical cells. Each cell typically has a positive and negative electrode separated by an electrolyte.

    • Operating Principle: Spontaneous redox reactions drive current flow, measured in ampere.
    • Types: Primary (non-rechargeable), secondary (rechargeable) cells.
    • Thermodynamics: Reaction enthalpy and free energy changes link to cell potentials. For instance, electrode_potential values help predict voltage.
    • Applications: Portable electronics, automotive batteries, large-scale energy storage.

    Important

    Batteries exemplify electrochemistry in action, showcasing how chemical processes produce electrical work.

40. Beta Decay - A radioactive process where a nucleus emits a beta particle (electron or positron ). In decay:

    • Charge conservation: Atomic number increases by 1, mass number () unchanged.
    • Nuclear binding_energy shifts can stabilise the nucleus.
    • Must be a balanced_equation for nucleons and charge, distinct from typical chemical equations.
    • Some isotopes (isotope) rely on decay for transformation. The bohr_model historically preceded full nuclear models but paved the way for atomic theory.

41. Binary Mixture - A two-component mixture (e.g. ethanol-water). Behaviour often differs from that of pure substances:

    • Vapour-Liquid Equilibria: Simple or complicated, can form an azeotrope.
    • Boiling Changes: Adding a second component can shift the boiling_point.
    • Mixing: Intermolecular forces, partial miscibility, or complete solutions can arise.
    • Kinetically, collision_theory in reacting binary mixtures can be altered by solvation or molecular interactions.

    Note

    Binary mixtures serve as the basis for studying more complex solutions with multiple components.

42. Binding Energy - The energy required to separate a system into its constituent parts. Broadly includes:

    • Nuclear Binding Energy: Holds protons and neutrons together, important in beta_decay.
    • Atomic or Electronic Transitions: Relates to spectral lines, such as those in the balmer_series.
    • Molecular: Sometimes conflated with bond_energy, though binding can refer to entire complexes.

    If binding is strong, it can stabilise species in redox contexts or hamper decomposition. Negative binding energy indicates stability upon formation.

    Note

    In nuclear chemistry, binding energy explains why heavier nuclei can release energy via fission or decay. In atomic/molecular contexts, it underpins electron orbital energies.

43. Bohr Model - A historic atomic model proposed by Niels Bohr for the hydrogen atom, featuring quantised orbits for electrons.

    • Explains the discrete spectral lines observed in the balmer_series.
    • Despite overshadowed by modern quantum_mechanics, it remains a cornerstone in atomic theory.
    • beta_decay expansions eventually revealed deeper nuclear interactions, beyond Bohr's scope.
    • Relation to bond_energy: Early shell models hinted at how electron arrangement affects chemical bonding, though not fully accurate for multi-electron atoms.

44. Boiling Point - The temperature at which a liquid's vapour pressure equals external pressure, causing vaporisation throughout the bulk.

    • Influences: Ambient pressure (see barometer), solution composition (binary_mixture), and added solutes (colligative_properties).
    • Standard boiling point: Measured at 1 bar (about 100°C for pure water). Historically, 1 atm was used.
    • Elevation: The presence of solute can cause a boiling_point_elevation, a prime example of colligative effects.

    Note

    Boiling point data are crucial in distillation, cooking science, and chemical process design. Pressure manipulation or solute addition can modify it significantly.

45. Benzene (C₆H₆) - An aromtaic organic molecule of the formula C6H6, molar mass 78.114 g·mol-1. It is cyclic and contains only Carbon and Hydrogen thus it is called a hydrocarbon.

    

    Benzene is naturally found in petroleum and is a foundational chemical in petrochemistry: production of plastics, pharmaceuticals, dyes, detergents, pesticides, rubber, and other important molecules.

    

    It is a liquid, colourless, highly volatile, flammable, and a known human carcinogen.

C

46. Calorimetry - The measurement of heat flow into or out of a system during chemical reactions or physical changes. Data from calorimetry allow calculation of enthalpy (), heat capacities, and other thermodynamic values.

    • Types:
      • Constant Pressure (Coffee-cup) Calorimeter: Tracks heat at atmospheric pressure.
      • Constant Volume (Bomb) Calorimeter: Sealed vessel for combustion or decomposition.
    • Stoichiometry: Linking mass changes to moles uses the avogadro_constant.
    • Relation: Reaction rates (collision_theory) and enthalpy changes can both influence or be measured by calorimetry.
    • Key Parameter: heat_capacity of the calorimeter is essential for accurate readings.

    Note

    Calorimetry underpins many thermodynamic measurements, such as the energy content of fuels and foods.

47. Catalyst - A substance that increases the rate of a reaction by lowering its activation_energy, without being consumed in the process. Catalysts can be solids (e.g. metals), liquids, or enzymes in biochemistry.

    • Mechanism: Alters reaction pathway, enabling more molecules to surpass the energy barrier.
    • Selectivity: Some catalysts favour specific products, crucial in industrial syntheses.
    • Coordination-based: In coordination_chemistry, metal complexes may provide unique active sites.
    • Chemical Bonds: Temporary covalent_bond formation between catalyst and substrate can facilitate reaction steps.
    • Kinetic Impact: By collision_theory, more effective collisions occur with a catalyst.

    Note

    Catalysts are indispensable in chemical manufacturing for efficiency, cost-saving, and minimising environmental impact.

48. Charles' Law - States that the volume of a fixed amount of gas at constant pressure is directly proportional to its absolute temperature ():

    • Explains why balloons expand when heated at constant .
    • Incorporates idea of collision_theory in that higher temperature increases molecular speed, demanding larger volume.
    • Part of the combined gas_law equations.
    • Conversions from -related data to molar terms might use the avogadro_constant.
    • Affected also by external factors like boiling_point if the temperature is high enough to vaporise the gas or shift phases.

    Note

    Charles' Law is integral for balloonists, meteorologists, and general predictions of how gases behave with temperature changes.

49. Chemical Equilibrium - A dynamic state where forward and reverse reaction rates match, so concentrations of reactants and products remain constant. Represented by an equilibrium constant for a balanced_equation.

    • Rate Perspective: collision_theory ensures forward and reverse collisions occur at equal rates.
    • Thermodynamics: Equilibrium involves interplay of enthalpy and entropy.
    • Le Chatelier's Principle: System shifts if conditions change (concentration, pressure, temperature).

    Important

    'Constant' concentrations do not mean the reaction stops; rather, forward and reverse processes continue at the same pace.

50. Colligative Properties - Solution properties that depend on the number of solute particles, not their identity. Major colligative properties:

    • Boiling Point Elevation: Solute addition raises the boiling_point.
    • Freezing Point Depression: Solute lowers the freezing point.
    • Osmotic Pressure: Pressure needed to stop solvent flow across a semipermeable membrane.
    • Vapour Pressure Lowering: Fewer solvent molecules at surface.

    Note

    Techniques like ebulioscopy rely on colligative properties. In real solutions, strong solute-solvent interactions can modify 'ideal' behaviour.

51. Collision Theory - Proposes that molecules must collide with sufficient energy (at least activation_energy) and proper orientation to form products.

    • Rate ~ collision frequency fraction of collisions with .
    • Temperature increase (see charles_law) boosts molecular speeds, raising collision frequency and success fraction.
    • A catalyst lowers , boosting reaction rate.
    • At chemical_equilibrium, forward and reverse collisions occur at the same rate.

52. Complex Ion - An ion formed by a central metal (often a transition metal) bonded to one or more ligands via coordinate (dative) bonds. Written typically as [M(L)_n]^charge.

    • Example: , .
    • Part of coordination_chemistry, with the metal-ligand interactions sometimes approximated as a covalent_bond from the ligand side.
    • Ligands can behave like arrhenius_base donors, supplying electron pairs to the metal.
    • In-depth discussion involves dative_bond formation.

    Note

    Complex ions underlie many biochemical metalloproteins (e.g., haemoglobin) and industrial catalysts.

53. Coordination Chemistry - The study of complex ions and coordination compounds, where a metal centre bonds to electron-pair donors (ligands). It merges concepts of geometry, bonding theory, and metal-ligand reactivity.

    • Coordination Number: Number of donor sites attached to the metal.
    • Electronic Structure: quantum_mechanics and crystal field or ligand field theory describe metal-ligand orbital overlaps.
    • Oxidation States: Metals can exhibit varied oxidation_state in complexes.
    • Ligand Effects: Ability to cause delocalization of electron density, altering reactivity and spectral properties.
    • Houses key topics like chelation, bridging ligands, and isomerism.
    • For an example complex_ion, geometry can be tetrahedral, octahedral, square planar, etc.

54. Coulomb's Law - Describes the electrostatic force () between two point charges and , separated by distance :

    • .
    • In ionic solids or solutions, explains ionic_bond strength.
    • Helps interpret electron-nucleus attraction in the bohr_model, and bridging single-charge to mole-based avogadro_constant scales.
    • In electromagnetism, current measured in ampere depends on charge flow and Coulombic definitions.

55. Covalent Bond - A chemical bond where two atoms share electron pairs. This sharing can be equal (non-polar) or unequal (polar).

    • Origin: Overlapping orbitals from each atom, as explained by quantum_mechanics.
    • Bond Strength: Related to bond_energy.
    • Delocalization: Extended pi systems show electron delocalization across multiple atoms.
    • In complex_ion contexts, ligands form coordinate covalent (dative) bonds with the metal.

    Note

    From basic molecules (H₂, O₂) to giant networks (diamond), covalent bonding fosters chemical diversity.

56. Quantum Mechanics - The theoretical framework describing matter and energy at atomic and subatomic scales. Governs how electrons occupy orbitals and form chemical bonds.

    • Key Principles: Wave-particle duality, uncertainty principle, discrete energy levels.
    • Applications: Extends the bohr_model to multi-electron atoms, refining orbital concepts.
    • Explains covalent_bond formation via orbital overlap, clarifies coordination_chemistry with metal-ligand interactions.
    • In quantum terms, a wavefunction describes electrons' distribution in space.

    Important

    Quantum mechanics underpins modern chemistry, explaining phenomena from atomic spectra to reaction pathways at fundamental levels.

D

57. Dalton's Law - Also known as the law of partial pressures. It states that in a mixture of non-reactive gases, the total pressure () is the sum of the partial pressures () of each component gas:

    • Partial pressure is the pressure a gas would exert if it alone occupied the container.
    • Connects to the ideal_gas equation and gas_constant concepts when relating partial pressures to moles and temperature.
    • Converting partial pressures to mole fractions often involves the avogadro_constant.

    Note

    Understanding each gas's partial_pressure helps predict gas behaviour in mixtures, from atmospheric science to industrial gas processes.

58. Dative Bond - A coordinate covalent bond in which both electrons in the shared pair originate from one atom (the donor). This is common in complex_ion formation or coordination_chemistry.

    • Mechanism: A lone pair from a donor atom (like N, O, or a ligand) bonds to an electron-deficient acceptor (like a metal centre).
    • Distinction: Once formed, a dative bond is typically indistinguishable from a regular covalent_bond in terms of electron distribution.
    • Acid-Base Context: The donor is akin to a Lewis or arrhenius_base if it provides electron pairs.

    Note

    Dative bonds explain bonding in species like NH₄⁺ (ammonium) or metal-ligand complexes, key to advanced inorganic chemistry.

59. Decantation - A mechanical separation process where liquid is poured off from a solid or heavier liquid once it has settled at the bottom or formed a distinct layer.

    • Usage: Removal of supernatant from precipitates, or separating immiscible layers (oil-water).
    • Contrast: distillation involves phase change via heating; decantation is simpler, using gravity or minimal equipment.
    • Kinetics: In contexts of collision_theory, flocculants can promote particle aggregation to speed settling.
    • Relevance: For adjusting concentration of components before further analysis or disposal.

    Note

    Wine-making, wastewater treatment, and lab work often rely on decantation to gently remove clear liquid from settled solids.

60. Decomposition Reaction - A chemical reaction in which one compound splits into two or more simpler substances. Often requires energy input (heat, light, or electricity).

    Example:

    • Energy: Usually needs a certain activation_energy to break existing bonds.
    • Balancing: Must be a balanced_equation for mass/atom conservation.
    • Kinetics: collision_theory clarifies that effective collisions or sufficient energy triggers bond-breaking.
    • Thermodynamics: Reaction enthalpy can be positive (endothermic) or negative (less common).

61. Delocalization - The spread of electrons across multiple atoms or bonds, rather than confinement to a single site.

    • Common in conjugated pi systems (benzene ring, polyenes) or metals in coordination_chemistry.
    • Resonance structures illustrate electron resonance but physically, the electrons exist in a merged orbital framework.
    • covalent_bond models sometimes cannot capture delocalised electrons fully; quantum_mechanics wavefunctions do better.
    • Lowers overall energy, stabilising the system (aromatic stability, metallic conduction, etc.).

    Note

    Delocalisation underpins phenomena like aromaticity, metallic bonding, and certain "3-centre-2-electron" bridging in main group compounds.

62. Deposition - A phase transition where a gas transforms directly into a solid, skipping the liquid phase (opposite of sublimation).

    • Example: Frost forming from water vapour.
    • Usually exothermic; energy is released as the vapour molecules lock into a solid lattice, relevant to enthalpy measurements.
    • The drop in molecular randomness also affects entropy.
    • calorimetry can track heat flow during deposition for thermodynamic data.

    Note

    Deposition is employed in thin film technologies, depositing layers of material from vapour for semiconductors or optical coatings.

63. Diamagnetism - A weak magnetic response observed in materials where all electrons are paired, causing a slight repulsion from external magnetic fields. This effect arises from induced magnetic dipoles that oppose the applied field.

    • Contrasts with paramagnetism (unpaired electrons) or ferromagnetism (aligned domains).
    • In coordination_chemistry, the oxidation_state and electron count can lead to either dia- or paramagnetic complexes.
    • quantum_mechanics explains diamagnetism through subtle electron orbital responses.
    • Even seemingly non-magnetic substances exhibit diamagnetism to a tiny degree, including water or complex_ions.

64. Diffusion - The net movement of particles from an area of higher concentration to lower concentration, driven by random molecular motion.

    Fick's first law:

    where:

    • is flux,

    • is the diffusion coefficient,

    • is the concentration gradient.

    • collision_theory underlies the random collisions leading to net flux.

    • Affects solutions or gases obeying gas_law expansions or partial mixing.

    • Spreads out concentration gradients, often increasing entropy.

    Note

    Diffusion explains everyday processes like aroma dispersal, ink spreading in water, or gas mixing in the atmosphere.

65. Dipole Moment - A vector quantity () indicating the separation of charge in a molecule. It arises when bonded atoms have different electronegativity, creating partial charges.

    • = magnitude of partial charges

    • = distance between charges

    • Larger dipole moment typically means a more polar_molecule.

    • Intra- and intermolecular_forces often hinge on dipole interactions.

    • Single covalent_bond can have a dipole if it's polar.

66. Dissociation Constant (water) - The rate at which water dissociates from its protons. This can be derived from autoprotolysis.

    Dissociation constant:

    Since this is theoretically measured in a pure solution of water and the concentration of pure substances is not taken into account then the formula simplifies to:

    Moreover, since according to the balanced equation for autoprotolysis [H3O+] = [HO-], and thus we can simplify the formula even more:

    Thus:

    And thus we've demonstrated that water's .

67. Distillation - A separation technique where a liquid is vaporised, then condensed to isolate it from other components. Common for purifying mixtures or obtaining solvents.

    • Steps: Heat the mixture to its boiling_point; vapor passes into a condenser; purified distillate is collected.
    • Variations:
      • Simple Distillation: Large boiling differences.
      • Fractional Distillation: Closer boiling ranges, often for binary_mixtures or complex solutions.
      • Steam Distillation: Employing steam to reduce partial pressures for heat-labile substances.
    • Contrasts with decantation which separates phases without boiling.
    • colligative_properties can shift boiling behaviour, important in many distillation processes.

    Note

    Distillation is fundamental in petrochemical refinement, alcoholic beverage production, and lab-scale solvent recovery.

E

68. Ebulioscopy - A technique used to determine the molecular mass of a solute by measuring the boiling-point elevation of a solution. It relies on colligative_properties, where the temperature increase depends on how many particles are dissolved rather than their chemical nature.

    This method involves comparing the solvent's normal boiling temperature with that of the solution, then relating the difference () to the solute's molality. The solvent-specific proportionality constant is called the ebullioscopic constant , tying in with enthalpy considerations of phase transitions.

    • : Ebullioscopic constant
    • : Molality (mol solute / kg solvent)
    • : van't Hoff factor (number of particles the solute dissociates into)

69. Electrochemistry - The branch of chemistry concerned with electron transfer processes, redox reactions, and electric currents. It unites concepts of oxidation-reduction, electrode potentials, and conductance.

    It intersects with collision_theory when describing how ions and electrons collide at electrode surfaces. For instance, an arrhenius_acid can donate protons that may affect electrode equilibria. Quantifying these reactions often involves the avogadro_constant to relate coulombs of charge to moles of electrons. In processes like electrolysis, an external voltage drives non-spontaneous redox reactions.

70. Electrode Potential - The voltage associated with an electrode, measured relative to a reference (e.g. the Standard Hydrogen Electrode). It indicates how readily a species is oxidised or reduced at that electrode surface.

    Electrode potentials connect thermodynamics and electron flow: (Gibbs free energy) relates to (potential), and factors like enthalpy and entropy can shift equilibria. In electrochemistry or electrolysis cells, electrode potentials determine which half-reactions dominate. Under certain conditions, water's autoionization may also influence the local potential by altering pH.

    • : Moles of electrons
    • : Faraday constant
    • : Electrode potential

71. Electrolysis - A process where an external electrical current compels a non-spontaneous redox reaction. It often deconstructs compounds (e.g. splitting water into hydrogen and oxygen), or refines metals.

    • Consumes electrical energy to drive chemical change.
    • Calculation of required charge involves the avogadro_constant for relating electrons to moles.
    • The feasibility depends on ionisation_energy thresholds for removing or adding electrons, plus the relevant electrode_potential.
    • Central to electrochemistry applications such as metal plating or chlorine production.

72. Electron Affinity - The energy change when a neutral atom (gas-phase) gains an electron to form a negatively charged ion. It reflects how strongly the atom 'pulls in' an added electron.

    Exothermic in many non-metals, especially halogens, but sometimes endothermic in metals with electron-repulsion concerns. This property complements ionisation_energy, and arises from quantum_mechanics descriptions of orbital occupancy. Industrial processes using electrochemistry often leverage differences in electron affinities when separating elements. Data in tables can be converted to per-mole terms using the avogadro_constant.

73. Electronegativity - An atom's ability to attract electrons in a covalent_bond. Higher electronegativity differences between bonded atoms generally yield a larger dipole_moment.

    It differs from electron_affinity in scope: electronegativity applies to an atom in a bond, whereas electron affinity is an isolated atomic property. Pauling developed a numeric scale referencing bond energies and quantum insights (see bohr_model evolution).

74. Electrophile - Any electron-poor species that accepts a pair of electrons to form a bond. Akin to a Lewis acid or arrhenius_acid, but broader in scope.

    • Typical examples: .
    • Often reacts with a nucleophile by forming a covalent_bond through electron-pair donation.
    • Rate of electrophilic reactions can depend on collision_theory, since higher temperature or concentration fosters more collisions.

75. Emulsion - A mixture of two immiscible liquids, where one is dispersed in the other as tiny droplets. Surfactants help stabilise the interface.

    • Examples: Milk (fat droplets in water), mayonnaise (oil in water).
    • Physical separations like decantation or distillation can be problematic if droplets don't easily coalesce.
    • Concentrations or molecular counts use the avogadro_constant for thorough analysis.
    • The heat changes or enthalpy of emulsification can matter in industrial processes.

76. Endothermic Reaction - A reaction that absorbs heat from its surroundings. It exhibits a positive enthalpy (), meaning the system's enthalpy increases.

    • Observed through drop in external temperature or direct measurement via calorimetry.
    • Requires sufficient energy input for bond-breaking, consistent with collision_theory.
    • Entropy changes (entropy) may also drive the reaction if the system becomes more disordered, compensating for energy absorption.

77. Enthalpy of Formation - The enthalpy change () for forming one mole of a compound from its elemental states under standard conditions (25°C, 1 bar).

    • A building block for calculating reaction enthalpies via Hess' Law and a balanced_equation.
    • Often measured with calorimetry or deduced from known enthalpies plus theoretical quantum_mechanics estimations.
    • If , formation releases heat; if , the compound is higher in energy than its constituent elements.

78. Enthalpy - A thermodynamic quantity denoted by , representing the heat content of a system at constant pressure.

    Mathematically:

    • : Internal energy

    • : Pressure-volume work term

    • If , the process is endothermic_reaction; if , it's exothermic_reaction.

    • Relates to ebulioscopy when boiling point measurements involve heat exchange.

    • The standard enthalpy_of_formation of compounds is often tabulated for reference in reaction enthalpy calculations.

79. Entropy - Symbolised by , it gauges the disorder or randomness of a system. A higher entropy means more microstates are accessible.

    Often summarised:

    • At constant temperature, for a reversible process.
    • Relates to enthalpy via . An endothermic_reaction may still be spontaneous if it increases entropy sufficiently. Conversely, exothermic_reaction can become non-spontaneous if entropy drastically decreases.
    • collision_theory in gases includes random kinetic motion, reminiscent of entropy at the molecular scale.

F

80. Fahrenheit Scale - A temperature scale developed by Daniel Gabriel Fahrenheit. It sets 32°F as the freezing point of water and 212°F as the boiling point (at sea level). Converting from Celsius () involves:

    • Primarily used in a few countries (e.g. United States).
    • In scientific contexts, the Kelvin or Celsius scales are more common (see charles_law).
    • Thermal properties like enthalpy changes can be tabulated in °C or K, so the Fahrenheit scale is often converted for consistency.
    • The difference between 32°F and 212°F parallels the 100° difference between 0°C and 100°C, used in the boiling-point references at sea level.

81. Faraday Constant - Denoted , it represents the electric charge per mole of electrons. Numerically about .

    • Ties into avogadro_constant and the elementary charge (), such that .
    • Crucial in electrochemistry for relating electron flow (coulombs) to moles of substance oxidised or reduced.
    • In electrolysis, passing coulombs through a cell corresponds to transferring one mole of electrons.

    Mathematically:

82. Fermentation - An anaerobic biochemical process in which microorganisms (often yeasts or bacteria) convert sugars to other products (e.g. ethanol, lactic acid) in the absence of oxygen.

    • Key for producing alcoholic beverages, bread, yoghurts.
    • Enzyme-driven pathways alter the enthalpy balance, generating less energy per glucose unit than full oxidation but sufficient for microbial survival.
    • Different microorganisms or conditions lead to varied end products: ethanol, lactate, or mixed acids.
    • Proton transfers and autoprotolysis-like steps can occur in the medium's water. Amino_acid composition and electronegativity differences in co-factors also affect enzymatic reactivity.

83. Fission - The splitting of a heavy nucleus (like uranium-235) into two or more lighter nuclei, accompanied by the release of large amounts of energy and neutrons.

    • Governed by nuclear binding_energy, where heavier nuclei can become more stable by splitting.
    • Often initiates a chain reaction if freed neutrons hit more fissionable nuclei.
    • Differs from beta_decay in scale and immediate energy yield.
    • Statistical aspects (distribution of daughter nuclei) can link to nuclear entropy. Early quantum explorations (see bohr_model antecedents) paved the way for fission theory.

84. Fluorescence - The emission of light by a substance that has absorbed electromagnetic radiation. Often occurs when an electron in an excited state relaxes to a lower energy state, emitting a photon.

    • Timescale: Typically very fast (nanoseconds).
    • Electronegativity and bonding environment shift the absorption/emission wavelengths.
    • Electron transitions are governed by quantum_mechanics and discrete energy_level spacings.
    • Coulombic interactions (coulombs_law) shape excited-state orbital energies.

85. Formal Charge - A theoretical charge assigned to an atom in a molecule or ion, based on an equal sharing of bonding electrons. Helps predict most stable resonance structures.

    • Calculation per atom:
    • Minimising total formal charges often indicates the most plausible structure.
    • In covalent_bond frameworks, formal-charge assignment is central. For a complex_ion, it's similarly used for each ligand and the metal centre.
    • Effects: Large formal charges can indicate strong electrostatic forces (see coulombs_law) or possible ampholyte reactivity.

86. Free Energy - Often referring to the Gibbs Free Energy () at constant pressure and temperature. It dictates the spontaneity of a reaction.

    • Formula: where is enthalpy, is temperature (K), and is entropy.
    • Negative : Reaction is spontaneous; positive : non-spontaneous.
    • Ties to endothermic_reaction or exothermic conditions, factoring in entropy changes.
    • At chemical_equilibrium, .

87. Free Radical - A species with one or more unpaired electrons, making it highly reactive. Commonly formed by homolytic bond cleavage or redox processes.

    • Tendency to chain-react, especially in polymerisation or combustion.
    • Reactions follow collision-theory, requiring sufficient energy to break stable bonds and produce radicals.
    • Unpaired electrons lead to unusual oxidation-state assignments relative to typical formal_charge.
    • Stability partly depends on an atom's ionisation_energy and the radical's ability to delocalise electron density.

88. Freeze-Drying - Also called lyophilisation. A preservation and drying technique where water is frozen, and ice is then sublimed under reduced pressure, bypassing the liquid phase.

    • Steps:
      1. Freezing the material.
      2. Lowering pressure.
      3. Deposition-like sublimation occurs, removing water.
    • Advantage: Minimal thermal damage since water does not pass through a typical boiling-point.
    • Influence on entropy and system's concentration: Removing water drastically alters solute or matrix composition.
    • Common in pharmaceuticals, food, and delicate biological samples.

89. Frequency - The number of waves (oscillations) per second, measured in hertz (Hz). In chemistry, it often appears in spectroscopy or reaction rate discussions.

    • Wave relation: where is the speed of light, is wavelength, and is frequency.
    • Spectral Lines: bohr_model suggests discrete electronic transitions with specific frequencies or energies. This underpins fluorescence and absorption.
    • Quantum mechanical context (quantum_mechanics): Photons carry energy , crucial for bond excitation or cleavage.
    • Macroscopic rates: collision-theory sometimes references collision frequency among molecules in a gas.

G

90. Galvanic Cell - An electrochemical cell that spontaneously converts chemical energy into electrical energy by a redox reaction. In a typical galvanic cell:

    • Oxidation occurs at the anode, reduction at the cathode.
    • Each electrode is immersed in a distinct half-cell solution, connected by a salt bridge.
    • The cell's potential relates to electrode_potential values, measured in volts.
    • Thermodynamics connects cell potential () and entropy or free energy ().
    • The link to electrochemistry is direct, and experimental mole calculations often involve the avogadro_constant.

    Note

    Galvanic cells are the basis for batteries (e.g. Zn-Cu Daniell cell), enabling portable power through spontaneous redox reactions.

91. Gas Constant - Symbolised by , a universal constant in the ideal gas equation: . Numerically about .

    • Relates pressure , volume , and temperature of a gas.
    • Combines the avogadro_constant and the Boltzmann constant .
    • Crucial for daltons_law of partial pressures and expansions under charles_law.
    • Appears widely in thermodynamics (calculations for entropy, enthalpy, etc.).

    Important

    Remember that different units (L·atm·K^-1·mol^-1, J·K^-1·mol^-1) can change numeric value but represent the same constant.

92. Gas Law - A collective term for relationships describing the behaviour of ideal gases. It includes Boyle's law ( at fixed ), charles_law ( at fixed ), and the combined or ideal gas law:

    • : Pressure
    • : Volume
    • : Number of moles
    • : gas_constant
    • : Absolute temperature

    Applications:

    • Explaining how boiling_point can shift under different pressures.
    • Extending to mixtures via daltons_law.

    Warning

    All these formulas assume gas behaves ideally, which can fail at very high pressure or very low temperature.

93. Gauge Pressure - The pressure measured relative to the surrounding atmospheric pressure. So if atmospheric pressure is 1 atm, a gauge reading of 2 atm means an absolute pressure of 3 atm.

    Often used in pressurised systems (e.g. gas cylinders, tyre pressures). For reference, a barometer measures absolute atmospheric pressure, which can influence the boiling_point or volume expansions under gas_law.

    

94. Geiger Counter - An instrument used to detect and measure ionising radiation, especially alpha and beta particles. It consists of a gas-filled tube and electrodes, producing a pulse when radiation ionises the gas.

    • Connected to beta_decay detection and nuclear fission studies.
    • Ion-pair creation in the tube correlates to charged particle flux, somewhat reminiscent of collision_theory.
    • Electrostatic interactions (coulombs_law) cause gas ionisation to trigger an avalanche discharge, counted as 'clicks'.

    Note

    A Geiger-Müller tube requires relatively low voltage to operate and is popular for quick radioactive surveys.

95. Gibbs Free Energy - Denoted , a thermodynamic potential that predicts spontaneity at constant pressure and temperature:

    • : enthalpy
    • : Temperature (kelvin)
    • : entropy

    A negative indicates a spontaneous process. In electrochemistry, ties to cell potentials; at chemical_equilibrium, .

    Important

    Because merges enthalpy and entropy, it offers a concise measure for whether a reaction can proceed on its own under given conditions.

96. Glass Transition - A pseudo-phase transition where an amorphous material (e.g. certain polymers, glass) shifts from a brittle 'glassy' state to a rubbery or viscous state. Unlike a sharp melting point, this transition covers a temperature range.

    • Typically analysed with calorimetry, detecting changes in heat capacity.
    • Involves subtle rearrangements of molecular segments, affecting enthalpy and entropy.
    • No distinct crystalline structure, so the transition is gradual.

    Warning

    Glass-transition temperature (T_g) is not a strict thermodynamic phase transition; it's an empirical threshold where mobility changes markedly.

97. Glucose - A monosaccharide (simple sugar) with the formula C₆H₁₂O₆. It's a primary energy source in biology and a crucial substrate for fermentation.

    • Often crystallises in ring form (pyranose).
    • In solution, it exhibits equilibrium between alpha and beta anomers.
    • Water solutions of glucose can show colligative_properties like boiling-point elevation or freezing-point depression.
    • The enthalpy of glucose combustion is used to measure caloric content in food science.

    Note

    Though vital for metabolism, excessive glucose intake can disrupt homeostasis. In biochemical contexts, partial oxidation or fermentation yields myriad products.

98. Gram-Equivalent - A measure defined as the mass of a substance that reacts with or replaces 1 mole of hydrogen ions (in acid-base) or 1 mole of electrons (in redox).

    • Historically used for normality (N) calculations: Normality = (grams of solute / litre) / (gram-equivalent mass).
    • In acid-base contexts, a gram-equivalent neutralises 1 mole of H⁺ or OH^-.
    • Ties to the avogadro_constant for bridging ionic or electron-based reactions.
    • Modern practice often prefers mole-based stoichiometry, but the concept still appears in older texts.

    Note

    One gram-equivalent of a base can supply 1 mole of OH^-, while for acid it typically provides 1 mole of H⁺.

99. Ground State - The lowest-energy electronic configuration for an atom, ion, or molecule. Electrons occupy orbitals based on minimised energy according to the Pauli exclusion principle and Hund's rule.

    • Contrasts with excited states, which have higher energy.
    • bohr_model introduced the concept for hydrogen, refined by modern quantum_mechanics.
    • Ion formation or ionisation_energy considerations start from the ground-state arrangement.
    • In molecules, partial delocalization of electrons can define a more stable ground state with lower overall energy.

    Important

    Most physical and chemical behaviour references the ground state, as excited states are often fleeting or need energy input.

H

100. Half-Life - The time required for half of a given amount of a substance (radioactive or otherwise) to decay or transform. Commonly used in nuclear beta_decay but may also apply to other processes (e.g. radical consumption in free_radical chemistry).

     The decay can often be modelled by first-order kinetics:

     • is the half-life
     • is the initial quantity

     Note

     In chemical kinetics, collision_theory still influences overall reaction rates. Spontaneous radioactive decay is also subject to quantum rules, adding an element of unpredictability. Thermodynamic factors like entropy can guide some non-nuclear half-life processes (e.g. slow polymer degradation).

101. Halogenation - A reaction in which one or more hydrogen atoms in an organic molecule are replaced by halogen atoms (F, Cl, Br, I). Mechanistically, this can proceed via free-radical, electrophilic, or nucleophilic substitution routes:

     • Free-Radical Halogenation: Initiation, propagation, termination steps (tie to free_radical).
     • Electrophilic Halogenation: Typically with aromatic rings, forming a new covalent_bond.
     • Energy Considerations: The required activation and reaction enthalpy often rely on bond strengths.
     • Reaction Rate: collision_theory says more collisions at higher temperature or concentration lead to faster halogenation.

102. Hard Water - Water containing high concentrations of calcium and magnesium ions (often Ca²⁺, Mg²⁺), leading to scale build-up and soap scum formation.

     • Causes: Limestone (CaCO₃) dissolution, dolomite (MgCO₃) presence.
     • Effects: Reduced solubility of soaps, scale in boilers, interference in detergents.
     • Chemistry: Complexation can reduce hardness. Chelating agents like edta-complex or forming complex_ion with Ca²⁺ help soften water.
     • Relevance: solubility equilibria and partial ionisation_energy considerations in mineral-laden environments.

103. Heat Capacity - The quantity of heat needed to raise an object's temperature by 1 degree (1 K or 1°C). Symbolised , or sometimes , for constant pressure/volume contexts.

     • Molar Heat Capacity: , relates to substance-specific properties.
     • Connection: In calorimetry, quantifies energy changes. Closely linked to changes in enthalpy and how entropy might vary with temperature.
     • Practical: ebulioscopy or any temperature-based technique must account for the system's or apparatus' heat capacity.

104. Henry's Law - States that the solubility of a gas in a liquid is directly proportional to the partial pressure of that gas above the solution:

     • : Partial pressure of the gas
     • : Mole fraction of the gas in the solution
     • : Henry's law constant (depends on gas-liquid pair)

     Essential for gas solutions, carbonated beverages, and comparing partial pressures under gas_law conditions. Changes in pressure or temperature can shift the gas solubility, sometimes affecting the boiling_point or colligative_properties in mixtures.

     Note

     Henry's law applies best under low gas solubility and near-ideal conditions. Deviations occur with chemical reactions or strong gas-liquid interactions.

105. Hess' Law - The enthalpy change () of an overall reaction is the sum of the enthalpy changes of the individual steps, regardless of the reaction path.

     • Often used to deduce by combining known enthalpies from sub-reactions.
     • Ties to enthalpy being a state function, meaning path independence.
     • Common approach: Summing enthalpy_of_formation values for products minus that of reactants in the balanced_equation.
     • Saves time vs. measuring each reaction path step experimentally.

106. Hydration - The process by which water molecules surround and interact with solute particles (ions or molecules). Hydration is a specific form of solvation focusing on water as the solvent.

     • Ionic Solutes: Water dipoles orient around ions, stabilising the solution.
     • Polar Molecules: Hydrogen bonding can drive hydration.
     • Thermodynamics: enthalpy may be exothermic or endothermic, and entropy considerations track how water structuring changes.
     • Acid-Base: Even amphoteric species rely on hydration shells influencing proton availability.

107. Hydrolysis - A reaction in which a bond is broken by reaction with water. Many biologically and industrially important processes (e.g. peptide bond cleavage, ester saponification) fall under hydrolysis.

     • If a salt reacts with water, the resulting solutions can be acidic (arrhenius_acid) or basic (arrhenius_base).
     • Water's amphoteric nature fosters partial ionisation or bond cleavage.
     • Mechanistic details involve collision_theory, as molecules must collide with water in the right orientation and energy to break the bond.

108. Hydronium Ion - An oxygen cation attached to three protons, its composition is written as H3O+. This is different from R3O+ known as a oxonium ion and instead of three protons can include any other groups. A hydronium ion is an oxonium ion, but the inverse is not true; as this term refers specifically to ions which have hydrogens.

109. Hydrophilic - Describes a substance or functional group that is water-attracting or water-soluble. Hydrophilic substances typically form hydrogen bonds or strong dipole interactions with water.

     • Examples: Polar head groups in surfactants, sugar molecules, certain amino-acid side chains.
     • Contrasts with hydrophobic (water-repelling) segments in an emulsion.
     • hydration stabilises these species in aqueous solutions, influencing concentration equilibria.
     • Even amphoteric molecules can exhibit both hydrophilic (polar) and hydrophobic regions.

     Note

     Hydrophilic interactions underpin biological membrane assembly, protein folding, and colloidal stability in water.

110. Hydrophobic - Describes a substance or functional group that repels or does not mix well with water. Typically non-polar or low in polarity, lacking hydrogen bond donors/acceptors.

     • Examples: Alkyl chains, many hydrocarbons, oils.
     • Relationship with hydrophilic components leads to micelle or bilayer formation, or to separate phases like an emulsion.
     • Thermodynamically, water molecules reorganise around hydrophobic moieties, reducing entropy unless aggregation occurs (hydrophobic effect).
     • Changes to solution composition can also modify certain colligative_properties.

I

111. Ideal Gas - A theoretical gas whose molecules:

     • Occupy negligible volume,
     • Exhibit no intermolecular forces,
     • Move randomly with perfectly elastic collisions.

     Behaviour: Follows the gas_law equations precisely (), with the gas_constant. Real gases approximate ideal behaviour at high temperature and low pressure. Linking to the avogadro_constant helps relate microscopic molecule counts to macroscopic moles.

     Note

     Deviations occur near condensation points or high pressures, where molecular volume and intermolecular attractions become significant.

112. Indicator - A substance (often a weak acid or base) that changes colour in response to pH shifts, revealing if a solution is acidic, basic, or near neutral.

     • Examples: Litmus, phenolphthalein, methyl orange.
     • Mechanism: Indicator's molecular form and its conjugate differ in colour. Shifts in ph_potential_hydrogen cause one form to dominate.
     • Relation: In arrhenius_acid/base contexts, the indicator monitors autoionization of water and H⁺ changes in solution.

     Note

     Indicators are pivotal in titrations, visually signalling the endpoint where acid-base neutralisation occurs.

113. Inert Gas - Also called a noble gas, it rarely participates in chemical reactions due to a filled valence shell (e.g. He, Ne, Ar). From a bohr_model or modern quantum perspective, their ground_state electron configurations are extremely stable.

     • Minimal Reactivity: High ionisation energies, stable oxidation_state of zero.
     • Applications: Providing non-reactive atmospheres in welding, storing reactive chemicals, or controlling collision_theory conditions in experimental setups.

     Important

     Though called 'inert', heavier noble gases (e.g. xenon) can form compounds under extreme conditions (e.g. XeF₂).

114. Infrared Spectroscopy - An analytical technique that measures molecular vibrations or rotations, typically within the infrared region of the electromagnetic spectrum.

     • Absorption Bands: Molecules absorb at frequencies matching bond vibrational modes (stretching, bending).
     • Dipole Requirement: IR absorption is strongest when a bond change alters the molecule's dipole_moment.
     • Bond Insights: Observing stretching in polar covalent_bond systems, and shifts from electron delocalization.
     • Kinetics: In some reaction studies, IR can track bond formation or cleavage, complementing collision_theory ideas.

115. Inhibitor - A substance that slows or stops a chemical reaction, in contrast to a catalyst which speeds it up. It can bind to reactive sites or otherwise raise the effective activation_energy.

     • Modes:
       • Binding essential sites (competitive).
       • Forming side products that block reactivity.
     • Relevance: In industrial processes or enzyme-driven pathways, controlling reaction rates is vital.
     • Relation: collision_theory implies fewer effective collisions if an inhibitor reduces available sites or reactivity.

116. Interhalogen Compound - A molecule consisting of two or more different halogens (e.g. ClF, IBr₃, IF₇). These can be more reactive than single-element halogens due to bond polarity and unusual oxidation states.

     • Formation: Often via direct halogenation or reacting one halogen with another under controlled conditions.
     • Electronic Structure: Typically explained by advanced bohr_model expansions or MO theory with partial covalent_bond character.
     • Reactivity: collision_theory sees these as high-energy species due to bond imbalances, making them strong oxidisers or reactants.

117. Intermolecular Forces - Attractive or repulsive interactions between molecules, weaker than chemical bonds. Types include:

     • Van der Waals (London dispersion, dipole-dipole)
     • Hydrogen Bonding (particularly strong dipole-dipole)
     • Hydrophobic effect in water-based solutions (hydrophobic vs. hydration)

     Impacts boiling_point trends, solubility, viscosity, and overall molecular assembly. Also depends on dipole_moment and electron distributions.

118. Inversion Centre - A symmetry element in a molecule or crystal that inverts every point through a central location. If a molecule has an inversion centre, each point has an equivalent counterpart at relative to that centre.

     • Common in certain coordination_chemistry complexes or extended lattices.
     • Electronic delocalization patterns can reflect inversion symmetry (or lack thereof).
     • quantum_mechanics shows that inversion-symmetric wavefunctions can be even or odd with respect to the centre.

     Important

     Inversion symmetry influences selection rules in spectroscopy (transitions may be forbidden or allowed depending on parity).

119. Ionisation Energy - The energy required to remove an electron from a neutral atom in the gas phase:

     • Typically measured in kJ/mol.
     • Trends: Increases up a group (less shielding) and across a period (higher nuclear charge). Ties to electronegativity and electron_affinity.
     • bohr_model historically explained hydrogen's ionisation energy. Modern advanced collision_theory considers multi-electron interactions.

120. Isotope - Atoms of the same element (same atomic number) but different numbers of neutrons, leading to distinct mass numbers. For example, , , .

     • Can be stable or radioactive. Radioactive isotopes often relate to half_life or beta_decay processes.
     • Mass differences can influence binding_energy within the nucleus.
     • Conversions between single-atom measures and macroscopic scales use the avogadro_constant.
     • Chemical properties are generally similar (same electron configuration), but nuclear properties vary.

J

121. J-Coupling - A through-bond spin-spin coupling phenomenon in nuclear magnetic resonance (nmr-spectroscopy). Arises from indirect interactions transmitted by electrons in a covalent_bond.

     • Not reliant on dipole-dipole interactions, but electron-mediated.
     • Magnitude measured in hertz (Hz).
     • Typically explained by quantum_mechanics regarding orbital overlaps and partial delocalization of electron density.
     • Useful in determining coupling patterns, connectivity, and structural details in organic molecules.

122. Jablonski Diagram - A visual representation of electronic states and transitions in molecules, depicting absorption, fluorescence, phosphorescence, and non-radiative pathways.

     • Arrows show transitions between singlet and triplet energy_levels.
     • Explains how fluorescence or phosphorescence emerges, tying in with quantum_mechanics for spin states.
     • Internal Conversion: Non-radiative drop between states of the same spin multiplicity.
     • Inter-System Crossing: Crossing between singlet and triplet states.
     • In practice, collisional deactivation (collision_theory) and vibrational relaxation can quench or shift emission.

     Note

     A Jablonski diagram helps rationalise photophysical and photochemical events, including energy transfer in photosynthesis or laser dyes.

123. Jacobsen Catalyst - A chiral manganese(III) Schiff base complex developed by Eric Jacobsen. Used for asymmetric epoxidation of alkenes.

     • Key features: Salen-type ligand forming a coordination_chemistry complex around Mn³⁺.
     • The catalyst's oxidation_state and electronic delocalization in the macrocyclic ligand enable selective oxidations.
     • Examples of usage: Synthesis of enantiomerically enriched epoxides, complementing other catalyst approaches.

     Important

     Jacobsen and Katsuki catalysts revolutionised green oxidation processes, offering high enantioselectivity without heavy-metal by-products.

124. Jahn-Teller Distortion - A geometric distortion of a nonlinear molecular system that removes degeneracy in electronic states. Frequently seen in octahedral coordination_chemistry complexes with d⁹ ions (e.g. Cu²⁺).

     • Origin: quantum_mechanics shows electronic degeneracy is unstable, so slight structural deformation lowers overall energy.
     • Effects: Changes in bond lengths, angles, and oxidation_state distribution.
     • Though typical in metal complexes, certain molecular systems, or even exotic allotrope forms, can exhibit analogous distortions.

125. Jar Test - A simple lab procedure used in water treatment to determine the optimal dose of coagulants or flocculants. Involves mixing multiple water samples with varying coagulant concentration, stirring, and allowing solids to settle.

     • Observes floc formation and clarifies which dosage yields the clearest supernatant.
     • Ties into decantation to separate out settled flocs.
     • Physical-chemical context: collision_theory for particle aggregation and potential emulsion breaking if oils are present.

     Note

     Jar tests guide large-scale water treatment plants, minimising chemical usage while ensuring turbidity and contaminant reduction.

126. Job's Method - Also called the method of continuous variations, it determines stoichiometric ratios in complex formation or binding by varying mole fractions of reactants while maintaining total concentration constant.

     • Plotting a signal (e.g. absorbance in spectroscopy, temperature change in calorimetry) vs. mole fraction reveals the maximum, indicating the stoichiometry of the complex.
     • Applied in coordination_chemistry to confirm ligand:metal ratios or in host-guest systems.

127. Joule's First Law - Also known as the Joule-Lenz law, it states that the heat () generated by an electric current () flowing through a resistor () for time is:

     • Ties to electrochemistry in cells or circuits, where electrical energy dissipates as heat.
     • Also relevant to reaction enthalpy if the energy is used for collisions or bond disruption (collision_theory).
     • Named after James Prescott joule and Heinrich Lenz.

128. Joule-Thomson Effect - The temperature change observed in a real gas when it expands (or is compressed) adiabatically through a porous plug or throttle, with no external work done.

     • If the intermolecular forces are significant, the gas cools (or warms) upon expansion, deviating from ideal_gas behaviour.
     • Ties to enthalpy: The process is typically run at constant enthalpy (an isenthalpic expansion).
     • heat_capacity affects the magnitude of cooling or heating.
     • Influences refrigeration cycles, where controlled expansions liquefy gases. Also complements standard gas_law concepts.

129. Joule - The SI unit of energy, work, or heat, symbolised J. Defined as 1 newton-metre (N·m) or 1 coulomb-volt (C·V). In thermodynamics:

     • Often used for enthalpy, internal energy changes, and measuring a system's energy_level.
     • gas_constant in J·K^-1·mol^-1 is a prime example of joule usage.
     • Specific heat_capacity is typically expressed in J·K^-1·g^-1 or J·K^-1·mol^-1.

     

130. Junction Potential - A voltage difference that forms at the interface between two electrolyte solutions (or between a metal and an electrolyte) due to ion concentration or mobility differences.

     • A small, often unwanted offset in cells, complicating electrochemistry measurements.
     • Minimised by using salt bridges or carefully chosen reference electrodes.
     • Ties into electrode_potential corrections for accurate readings.
     • collision_theory on ionic transport plus solution drift can intensify or reduce this potential.

     Note

     Accurate pH or redox measurements depend on reducing junction potentials. Analytical methods must account for or calibrate against them.

K

131. Kekulé Structure - A classical representation of benzene (or other aromatic compounds) as a cyclic molecule with alternating single and double bonds. Proposed by Friedrich August Kekulé.

     • Issue: Actual benzene (C₆H₆) exhibits electron delocalization across all six carbons, not discrete double/single bonds.
     • Modern Understanding: Resonance hybrids unify these covalent_bond depictions into a single structure with partial bond orders.
     • Relation: aromaticity concepts refine Kekulé’s idea, emphasising stability due to cyclic pi-electron cloud.
     • Examples: benzene_c6h6 often drawn with two Kekulé forms in resonance.

     Note

     Though historic, Kekulé’s model sparked the notion of ring structures and resonance, bridging classical valence ideas with quantum insights.

132. Kelvin Scale - The absolute temperature scale, using absolute zero () as the lowest limit. One kelvin increment equals one degree Celsius increment, but the scales offset by 273.15.

     • Relationship: .
     • Why 'Absolute': At , theoretical molecular motion ceases (though quantum zero-point energy remains).
     • Role: Fundamental in thermodynamics and gas laws (e.g. charles_law).
     • Statistical: entropy definitions rely on absolute temperature in J·K^-1.
     • The avogadro_constant ties microstates to macroscale measurements, commonly at standard temperatures in Kelvin.

     Note

     Scientific temperature calculations typically use kelvins, ensuring correct thermodynamic relations without negative absolute T values.

133. Ketone - An organic compound featuring a carbonyl group () bonded to two other carbons. General formula: R–C(=O)–R'.

     • Comparison: Differs from an aldehyde, where one side of C=O is hydrogen.
     • Physical Properties: Typically moderate boiling_point, higher than alkanes but lower than alcohols due to the polar carbonyl_group.
     • Bonding: The uses a strong covalent_bond, often reactive in nucleophilic addition.
     • Found widely in sugars (fructose, etc.) and many industrial solvents (acetone).

     Note

     Ketones can be reduced to secondary alcohols, or further converted in oxidation or condensation reactions.

134. Kinetic Energy - The energy an object or particle has due to its motion. In chemistry, vital for understanding molecular speeds and reaction rates.

     • Gas molecules: The average KE relates directly to temperature.
     • collision_theory depends on whether kinetic energy exceeds the activation threshold.
     • Some gas_law formulations assume ideal molecules with KE distribution per Maxwell-Boltzmann.
     • Changes in KE can reflect bond or phase changes, connecting with enthalpy in macroscopic processes.

     Important

     Kinetic energy drives diffusion, effusion, and collisions essential to chemical reactivity and equilibria.

135. Kinetic Molecular Theory - A conceptual framework describing gas behaviour through molecular motion. Assumptions:

     1. Gas particles have negligible volume (point masses).
     2. No intermolecular forces, only elastic collisions.
     3. Average kinetic energy depends solely on temperature.
     4. Molecular collisions are frequent, random, and explained by collision_theory.

     • Ties to the ideal_gas model.
     • Explains diffusion, effusion, pressure, and energy distributions.
     • Temperature-entropy link: As rises, molecular speeds spread, affecting entropy.

     Note

     Though approximate, kinetic molecular theory underlies basic gas laws, elucidating how volume, temperature, and pressure interconnect.

136. Kinetic Product - The product of a reaction pathway that forms fastest (i.e. requires the lowest activation_energy), though it may not always be the most stable. Often contrasted with the thermodynamic-product.

     • Low-temperature or short-reaction conditions favour kinetic product formation.
     • collision_theory emphasises that faster reaction paths yield immediate products.
     • If equilibrium or extra time/heat is allowed, the reaction may shift towards a thermodynamically more stable outcome.
     • Energy diagrams show an early (low) barrier, but a higher final enthalpy relative to the thermodynamic product.

     Note

     Knowing whether a reaction is under 'kinetic control' or 'thermodynamic control' is crucial in designing synthetic protocols.

137. Kjeldahl Method - A classical laboratory technique to determine nitrogen content in organic substances. It involves digestion of a sample in strong acid, converting nitrogen to ammonium sulphate, then distillation with strong base to release ammonia.

     • Distillation: Freed ammonia is collected in a trapping solution, then titrated to find total nitrogen.
     • Relevance: Used in protein assays, fertiliser analysis, etc.
     • Heat control or calorimetry may be used if large-scale digestion is monitored.
     • The method is robust but labour-intensive, replaced in some applications by modern automated instruments.

138. Knudsen Diffusion - A transport mechanism in porous media when pore diameter is so small that molecular collisions with the pore walls dominate over intermolecular collisions. Gas molecules behave almost independently within the pores.

     • Occurs at low pressure or with very narrow channels.
     • Different from normal diffusion (where molecules collide predominantly with each other).
     • collision_theory modifies to account for wall impacts rather than molecule-molecule collisions.
     • Common in catalytic processes through a porous_medium or microfluidic channels.

     Note

     Knudsen diffusion is vital in designing and interpreting membrane separations, microporous catalyst behaviour, and vacuum systems.

139. Kolbe Reaction - Often refers to the Kolbe electrolysis, where carboxylate salts undergo electrochemical decarboxylation to form radicals, which then couple to form hydrocarbons.

     1. Carboxylate () is oxidised at the anode.
     2. CO₂ is released, leaving an alkyl radical which can couple to produce R–R or other by-products.

     • An instance of electrolysis using an anode that might also act as a catalyst surface.
     • Reaction steps illustrate a redox cycle with radical intermediates and potential side reactions.
     • Named after Hermann Kolbe.

     Note

     Though overshadowed by modern synthetic methods, the Kolbe reaction historically demonstrated how electrochemical decarboxylation can build carbon–carbon bonds.

140. Ksp (Solubility Product) - The equilibrium constant for a sparingly soluble ionic compound dissolving in water. For a generic salt :

     Then,

     • Low means poor solubility.
     • Ties to hard_water issues (like Ca²⁺ deposits).
     • Helps predict precipitation via the common_ion_effect or changes in solutions.
     • Reflects ionic-lattice stability, bridging ionic_bond strength and aqueous solvation.

     Note

     Knowing is crucial for controlling scale formation, selective precipitation, and purification in chemical processes.

L

141. Lattice Energy - The energy released when gaseous ions combine to form an ionic solid, or equivalently, the energy required to break the ionic crystal into its constituent ions.

     • Often estimated using coulombs_law based frameworks, though real ions have finite sizes and polarisation effects.
     • Appears in the born-haber-cycle, connecting ion formation, electron affinities, and enthalpy changes.
     • Higher lattice energy usually indicates a stronger ionic_bond, leading to higher melting points and lower solubility in some solvents.

     Note

     Lattice energy trends help rationalise why smaller ions with higher charges tend to form more stable, less soluble ionic crystals.

142. Law of Mass Action - States that the rate of a chemical reaction is proportional to the product of the concentrations of the reactants, raised to their stoichiometric coefficients in the balanced_equation.

     • Forms the foundation for deriving rate laws in simple reactions (though many require experimental determination of orders).
     • Ties into chemical_equilibrium concepts and equilibrium constants.
     • In acid-base contexts (arrhenius_acid), mass action describes how changes in [H⁺] shift reaction rates.
     • collision_theory complements it by explaining the necessity of effective collisions.

     Note

     For complex mechanisms, the overall rate often depends on specific steps; the Law of Mass Action is an idealised base for deriving simpler rate laws.

143. Lewis Acid - A species (atom, ion, or molecule) that can accept a pair of electrons to form a new bond. This definition broadens arrhenius_acid to non-proton donors.

     • Often metal cations or electron-deficient molecules (e.g. BF₃, AlCl₃).
     • Forms coordinate (dative) bonds with a lewis_base, sometimes generating a complex_ion.
     • Overlaps with dative_bond concepts: the acid is the acceptor.

     Important

     Lewis acidity explains reactions beyond H⁺ donors, encompassing transition-metal catalysis and non-protonic acid-base systems.

144. Lewis Base - Any species capable of donating an electron pair, typically to form a coordinate bond. Contrasts with arrhenius_base which donates OH^- but focuses only on aqueous solutions.

     • A lewis_base + lewis_acid often yields a complex_ion.
     • Typical examples: NH₃, H₂O, halide ions, or any electron-rich site.
     • Emphasises dative_bond formation as a unifying acid-base concept across many chemical contexts.

     Note

     The Lewis definition is broad, explaining acid-base behaviour in non-aqueous solvents or advanced inorganic and organic reactions.

145. Ligand - An ion or molecule that binds to a central metal atom/ion via electron-pair donation, forming a coordination complex. Often acts as a lewis_base with a lone pair.

     • Can be monodentate (binds via one atom) or polydentate (multiple binding sites).
     • Fundamental in coordination_chemistry and forming a complex_ion.
     • Often involves a dative_bond from ligand to metal.
     • Examples: Water, ammonia, chloride, ethylenediamine, EDTA.

     Note

     Ligand identity and coordination geometry profoundly shape a metal centre’s reactivity and spectroscopic properties.

146. Limiting Reactant - The reactant that is fully consumed first in a reaction, thus determining the maximum amount of product formed. Any other reactants are in excess.

     • Found by comparing stoichiometric ratios in the balanced_equation with actual amounts available.
     • Mass or moles can be converted using the avogadro_constant if needed.
     • Once consumed, reaction halts, capping theoretical yield.
     • The leftover excess-reactant remains unreacted.

     Note

     Correctly identifying the limiting reactant is essential for accurate product predictions and resource optimisation.

147. Line Spectrum - A discrete set of wavelengths or frequencies emitted (or absorbed) by atoms or molecules. Each line correlates to an electronic transition between energy_levels.

     • Atom-specific: E.g. hydrogen's balmer_series lines.
     • Supported historically by the bohr_model, refined by modern quantum_mechanics.
     • Distinct from a continuous spectrum; each line indicates a well-defined energy difference.

     Note

     Line spectra confirm quantised energy states, vital evidence for the wave–particle dual nature of electrons.

148. Logarithm - The inverse of the exponential function, commonly used in pH or reaction rate expressions. For instance, pH :

     • If , then .
     • Reaction rates in collision_theory can appear in forms like (Arrhenius expression).
     • Different units or bases: common log (base 10) vs. natural log (base e).
     • In pH ( in mol·L^-1), a small change in scale means large changes in .

     Note

     Logarithmic scales help handle wide dynamic ranges in concentration, reaction rates, or intensities.

149. London Dispersion Forces - Weak intermolecular_forces arising from temporary dipoles induced by electron cloud fluctuations. Present in all molecules, but crucial in nonpolar substances.

     • Depend on polarizability: large or elongated electron clouds can form stronger instantaneous dipoles.
     • Influence boiling_point trends (e.g., heavier alkane = stronger dispersion = higher b.p.).
     • Contribute to the overall entropy interplay in condensed phases, though subtle.

     Note

     London dispersion forces help explain how noble gases can liquefy at low temperatures and why large nonpolar molecules have elevated boiling points.

150. Lyophilisation - Also called freeze-drying, a preservation technique where water is frozen, then removed by sublimation under reduced pressure. This avoids passing through the liquid phase.

     • Minimises thermal damage to sensitive materials (enzymes, pharmaceuticals, foods).
     • Ties to the concept of freeze_drying, though 'lyophilisation' is a more formal term.
     • colligative_properties can shift freezing behaviour in solutions with solutes present.
     • Often raises entropy of the removed vapour but benefits the dried product’s stability.

     Note

     Lyophilisation yields porous solids that can be quickly reconstituted with minimal loss of activity, crucial in biotech and food industries.

M

151. Malleability - A physical property of metals (and certain alloys) describing their capacity to be hammered or rolled into thin sheets without breaking. It stems from the ability of atoms in a metallic_bond lattice to rearrange under stress.

     • Mechanism: Layers of metal atoms can slide over each other if the bonding is non-directional.
     • Influences: Temperature, impurity content, crystal_structure, and presence of vacancies.
     • Entropy factors in slightly, as more disordered metal lattices can deform more easily.

     Note

     Malleability contrasts with brittleness. Gold is notably malleable, while ionic crystals generally shatter when struck.

152. Mass Defect - The difference between the sum of the component nucleon masses (protons + neutrons) and the actual measured mass of the nucleus. It arises because some mass is converted to nuclear binding energy (binding_energy) per .

     • Example: Helium nucleus mass is slightly less than combined masses of 2 protons + 2 neutrons.
     • Important in fission or beta_decay calculations, and describing why different isotopes have varying stability.
     • Explains how nuclear energy is stored or released.

     Note

     Mass defect underscores Einstein’s principle that mass can convert to energy, essential to understanding nuclear power and astrophysical processes.

153. Mass Number - Symbolised as , the sum of protons and neutrons in an atomic nucleus. Distinguishes isotope species of the same element with different neutron counts.

     • Not to be confused with the atomic_number (proton count).
     • Example: Carbon-14 has ( protons, neutrons).
     • Variation in can lead to differences in mass_defect and nuclear stability; some undergo beta_decay.

     Note

     Nuclear notation often appears as , where is the mass number and is the atomic number.

154. Maxwell-Boltzmann Distribution - A statistical function describing the speeds (or energies) of gas particles at a given temperature.

     • Shows that most molecules have a moderate speed, with fewer having very low or very high speeds.
     • Underpins kinetic_molecular_theory and collision_theory.
     • Ties to gas_law predictions, as average speeds scale with temperature.
     • Typically plotted as a frequency vs. speed curve, shifting to higher speeds and broader distribution at higher temperatures.

     Note

     This distribution is crucial for understanding how only a fraction of molecules exceed the activation threshold in chemical reactions.

155. Metal-Ligand Bond - The coordinate bond formed between a metal centre and a ligand in coordination_chemistry. Often described as a dative_bond with the metal accepting electron pairs.

     • Stability and properties depend on metal’s oxidation_state and ligand type.
     • Defines geometry in a complex_ion: octahedral, tetrahedral, square planar, etc.
     • Impacts reactivity, catalytic behaviour, and spectroscopic characteristics.

     Note

     Metal-ligand interactions underlie catalytic cycles, electron transfers, and biological processes like oxygen transport in haemoglobin.

156. Metallic Bond - A bonding model for metals where valence electrons form a 'sea of electrons' surrounding positive metal ions. This non-directional bonding grants metals distinct properties:

     • Malleability and ductility (malleability).
     • High electrical conductivity.
     • Partial correlation with electron_affinity and metal’s ability to release electrons into the communal pool.
     • enthalpy of vaporisation helps gauge metallic bond strength.

     Note

     In metallic bonding, delocalised electrons move freely, explaining metals' lustre, thermal conduction, and mechanical behaviour.

157. Metathesis Reaction - Also known as a double displacement reaction where two compounds exchange ions or components to form two new compounds:

     • Common in precipitation (e.g., insoluble salts) or acid_base neutralisations.
     • Must be a balanced_equation for stoichiometric correctness.
     • Reaction extent can hinge on solubility rules, forming a precipitate or non-volatile product that drives equilibrium forward.

     Note

     Metathesis underscores how certain insoluble or weakly dissociating products shift equilibria, enabling rapid ion exchange in solution.

158. Micelle - A supramolecular assembly of amphiphilic molecules in water or other solvents. Hydrophobic 'tails' cluster inward, while hydrophilic 'heads' face the solvent.

     • Enables solubilisation of nonpolar substances in water (detergents or surfactants).
     • Contrasts with an emulsion where discrete droplets form.
     • Reduces surface_tension by orienting amphiphile heads outward, hydrophobic tails inward.

     Note

     Micelles are pivotal in cleaning, emulsifying, and delivering hydrophobic drugs or nutrients in aqueous environments.

159. Molarity - A measure of concentration denoted M, defined as moles of solute per litre of solution:

     • Relates to colligative_properties calculations (e.g. boiling-point elevation, freezing-point depression).
     • Ties with avogadro_constant to switch between mole-based and particle-based counts.
     • Variation from other concentration units like molality, mass percent, or normality.
     • In preparing solutions, molarity is standard for stoichiometric calculations.

     Note

     Molarity is temperature-dependent since the solution volume changes with temperature, affecting the final concentration.

160. Mole - The SI base unit for amount of substance, symbolised mol. One mole contains entities (atoms, molecules, ions), where is the avogadro_constant, approximately .

     • Bridges microscopic scales (particles) to macroscopic measurements (grams, litres).
     • Connects atomic_mass in grams to stoichiometric balancing and yield predictions.
     • stoichiometry fundamentally depends on mole ratios in balanced reactions.

     Important

     The mole concept unifies chemical measurements, enabling direct calculations of reactants and products from mass or volume data.

N

161. Natural Abundance - The relative proportion of different isotopes of an element as they occur in nature, usually expressed as a percentage.

     • Measured precisely using mass_spectrometry.
     • Contributes to calculating average atomic_mass of elements.
     • Reflects nuclear_stability and cosmic nucleosynthesis history.
     • Often used in environmental and geological studies as natural tracers.

     Note

     Carbon-12 has ~98.9% natural abundance, while Carbon-13 is ~1.1%, crucial for understanding isotopic dating and tracing.

162. Neutralization - A chemical reaction between an acid_base pair producing water and a salt, typically resulting in a neutral pH:

     • Results in salt_formation through ionic combination.
     • Involves consumption of hydronium_ions and hydroxide ions.
     • Often exothermic, releasing heat energy.
     • Used in analytical chemistry for titrations.

     Note

     Complete neutralization doesn't always result in pH 7, as some salts undergo hydrolysis affecting final pH.

163. Noble Gas - Elements in Group 18 (He, Ne, Ar, Kr, Xe, Rn) with complete outer electron_configurations, exhibiting extreme chemical stability.

     • Demonstrates periodicity in physical properties.
     • Highest ionization_energy in their periods.
     • Full octet of valence_electrons (except He with 2).
     • Some form compounds under extreme conditions (mainly Xe, Kr).

     Note

     Noble gas configuration stability explains why other elements tend to gain, lose, or share electrons to achieve similar arrangements.

164. Nuclear Fission - The splitting of heavy atomic nuclei into lighter elements, releasing energy and often neutrons. Common example:

     • Energy release explained by mass_defect and binding_energy differences.
     • Can initiate chain_reactions in nuclear reactors.
     • Probability depends on target nucleus nuclear_stability.
     • Primary source of nuclear power generation.

     Note

     Controlled fission in reactors versus uncontrolled in weapons illustrates importance of neutron moderation.

165. Nuclear Fusion - The combination of light atomic nuclei to form heavier elements, releasing substantial energy. Example:

     • Requires extreme temperature to overcome nuclear repulsion.
     • Energy output explained by mass_defect.
     • binding_energy per nucleon peaks around iron.
     • nuclear_stability determines feasible fusion pathways.

     Note

     Fusion powers stars and holds promise for future clean energy, though controlling plasma remains challenging.

166. Nuclear Magnetic Resonance (NMR) - A spectroscopy technique analyzing atomic nuclei behavior in magnetic fields, crucial for structure determination.

     • Measures chemical_shift of nuclei in different environments.
     • Requires NMR-active isotopes (e.g., ¹H, ¹³C, ¹⁹F, ³¹P).
     • Different from electron spin resonance.
     • Provides detailed molecular structure information.

     Note

     NMR forms the basis of MRI medical imaging, though "magnetic resonance imaging" name avoids "nuclear" term.

167. Nucleophile - A chemical species that donates an electron_pair to form a chemical bond, literally "nucleus-loving."

     • Reacts with electrophiles in many organic reactions.
     • Common in substitution_reactions.
     • Often has lone pair valence_electrons.
     • Examples: OH⁻, NH₃, CN⁻, R₃P.

     Note

     Nucleophilicity differs from basicity; soft nucleophiles (I⁻) may be weak bases while strong bases (OH⁻) can be hard nucleophiles.

168. Nucleus - The dense central region of an atom containing protons and neutrons, determining the element's atomic_number and mass_number.

     • Comprises > 99.9% of atomic mass in < 0.01% of volume.
     • nuclear_stability depends on proton/neutron ratio.
     • Held together by strong nuclear force and binding_energy.
     • Size roughly proportional to cube root of mass number.

     Note

     Nuclear radius is about 10⁻¹⁵ m (fermi), vastly smaller than atomic radius (~10⁻¹⁰ m).

O

169. Orbital Hybridisation - The mathematical combination of atomic orbitals to form new hybrid orbitals that better explain molecular structure and bonding.

     Common hybridisations and geometries:

     • sp³: tetrahedral, 109.5° (e.g., CH₄)
     • sp²: trigonal planar, 120° (e.g., C₂H₄)
     • sp: linear, 180° (e.g., C₂H₂)

     Key concepts:

     • Number of hybrid orbitals equals number of combined atomic orbitals
     • Hybrid orbital energy = average of component orbital energies
     • Explains sigma bonds in covalent_bond formation
     • Foundation of valence_bond_theory

170. Orbital - A mathematical function (ψ) describing the quantum state of an electron in an atom or molecule. The square of this wave_function (ψ²) gives the probability of finding an electron at any point.

     Key characteristics:

     • Defined by four quantum_numbers: n, l, ml, ms
     • s orbital: spherical (l=0)
     • p orbitals: dumbbell-shaped (l=1), three orientations
     • d orbitals: more complex (l=2), five shapes
     • f orbitals: most complex (l=3), seven shapes

     Determines electron_density distribution and electron_configuration.

171. Oxidation State - The apparent charge on an atom in a molecule after assigning electrons to the more electronegativity element.

     Rules for assignment:

     1. Free elements = 0
     2. Group 1/2 metals = +1/+2
     3. F always -1
     4. O usually -2 (except peroxides -1)
     5. H usually +1 (except metal hydrides -1)
     6. Sum equals total charge

     Common ranges:

     • N: -3 to +5
     • S: -2 to +6
     • Mn: +2 to +7
     • Cr: +2 to +6

172. Oxidation - Loss of electrons by a chemical species, resulting in increased oxidation_state. Always paired with reduction.

     Identification:

     • Increase in oxidation number
     • Loss of electrons
     • Loss of hydrogen
     • Gain of oxygen

     Common oxidation half-reactions:

173. Oxidising Agent - Species that accepts electrons, causing oxidation of another species while being reduced itself. Strength measured by reduction potential (E°).

     Strong oxidisers (E° > 0.9V):

     • F₂ (2.87V)
     • O₃ (2.07V)
     • MnO₄⁻ (1.52V)
     • Cl₂ (1.36V)
     • Cr₂O₇²⁻ (1.33V)

     Reaction pattern:

     Important

     Higher E° values indicate stronger oxidising agents.

174. Oxonium Ion - An oxonium ion is any oxygen cation with three bonds, the simplest form is H3O+ but in reality can be R3O+; R meaning any other groups. An oxonium ion is not a hydronium ion, this term refers specifically to ions which have hydrogens.

175. Ozone - Triatomic oxygen (O₃) with bent structure. Strong oxidising_agent with resonance stabilisation.

     Structure:

     • O-O bond length: 1.278 Å
     • bond_angle: 116.8°
     • Dipole moment: 0.53 D

     Key reactions:

     Resonance structures contribute to unusual stability and reactivity.

P

176. pH - Negative logarithm of hydronium_ion concentration. Measures acidity/basicity on 0-14 scale.

     Key relationships:

     • pOH = -log[OH⁻]
     • pH + pOH = 14 (at 25°C)
     • pH = ½[pKa - log(salt/acid)]
     • buffer_solution: pH ≈ pKa when [acid]=[base]

     Common values (25°C):

     • Strong acid: pH < 2
     • Weak acid: pH 2-6
     • Neutral: pH 7
     • Weak base: pH 8-11
     • Strong base: pH > 12

177. Paramagnetism - Property where substances are attracted to magnetic fields due to unpaired_electrons. Strength determined by number of unpaired electrons.

     Key aspects:

     • magnetic_moment μ = √(n(n+2)) BM
     • Temperature dependent (Curie's Law)
     • Common in transition metal compounds
     • Explained by crystal_field_theory

     Examples (unpaired e⁻):

     • O₂: 2
     • Cu²⁺: 1
     • Fe³⁺: 5
     • Cr³⁺: 3

178. Partial Pressure - The pressure exerted by one component of a gas mixture. Follows daltons_law: total pressure equals sum of partial pressures.

     Relationships:

     • P₍ᵢ₎ = X₍ᵢ₎P_total where X = mole_fraction
     • Influences vapour_pressure equilibria
     • Key in gas_law calculations
     • Used in gas solubility (Henry's Law)

179. Pascal - SI unit of pressure (Pa), defined as one newton per square metre (N/m²).

     Common conversions:

     • 1 atm = 101,325 Pa
     • 1 bar = 100,000 Pa
     • 1 torr ≈ 133.3 Pa
     • 1 psi ≈ 6,895 Pa

     Key reference points:

     • standard_pressure: 100 kPa
     • atmospheric_pressure: 101.325 kPa
     • Typical tyre pressure: 200-300 kPa

180. pH (Potential Hydrogen) - Loosely speaking a measure of the concentration of free protons in an aqueous solution. A more precise definition is the activity of protons in the solution or a measure of [H3O+]. This allows us to determine the acidity or basicity of a solution. These measurements are all put in the context of water, and relative to water; thus water's pH is equal to 7; acids are less than 7, bases are greater than 7.

     In a base:

     In an acid:

     • Hydroxide ions = [OH-]
     • Hydronium ions = [H3O+]

     pH is measured on a logarithmic scale, it is measured in units of moles per liter of hydrogen ions often called hydronium ions sometimes less precisely called oxonium ions.

     Precisely speaking this measurement of hydrogen ions in a solution can be done experimentally. In fact an experimental measure is designated as p[H]. Moreover, the difference between pH and p[H] is so small that they are sometimes used interchangeably. The way to find this would be by use of the Nernst equation as shown below:

     The strict definition of pH does not have associated units, and is dimensionless. It is described as the log of the inverse of hydrogen activity in a solution:

     But most often used in the following form with [H3O+] being the concentration of hydronium ions in mol/L:

181. Phenolphthalein - Common indicator in acid_base titrations. Changes from colourless to pink.

     Properties:

     • pH range: 8.2-10.0
     • Colourless pH < 8.2
     • Pink pH > 10.0
     • Sharp endpoint
     • Formula: C₂₀H₁₄O₄

     Best for:

     • Strong acid/strong base titrations
     • Weak acid/strong base titrations
     • Carbonate determinations

182. Photoelectric Effect - Emission of electrons from matter by light absorption. Demonstrates photon nature of light.

     Einstein's equation:

     Where:

     • Eₖ = kinetic energy of ejected electron
     • h = Planck's constant
     • f = frequency of light
     • φ = work_function

     Key points:

     • Supports quantum_theory
     • Shows wave_particle_duality
     • Threshold frequency required
     • Instantaneous process

183. pOH (Potential Hydroxide) - The inverse of pH, this measures the potential hydroxide activity of a solution: HO-.

     Note that the sum of pOH along with pH is equal to the total of the scale.

184. Polymer - Macromolecule composed of many repeating monomer units joined by covalent bonds.

     Types:

     • Addition (chain-growth)
     • Condensation (step-growth)
     • Natural (proteins, cellulose)
     • Synthetic (PVC, nylon)

     Key properties:

     • Degree of polymerisation
     • molecular_mass distribution
     • cross_linking extent
     • Crystallinity percentage

     Common structures:

     • Linear (PE)
     • Branched (LDPE)
     • Network (phenol-formaldehyde)

185. Precipitate - Solid formed from solution during a chemical reaction. Formation governed by solubility_product Ksp.

     Formation condition:

     Key concepts:

     • ionic_equation representation
     • crystal_growth factors
     • supersaturation effects
     • Temperature dependence

     Common examples:

     • AgCl: Ksp = 1.8 × 10⁻¹⁰
     • BaSO₄: Ksp = 1.1 × 10⁻¹⁰
     • PbI₂: Ksp = 7.9 × 10⁻⁹

Q

186. Quantum Number - Four numbers describing an electron's quantum state in an electron_shell. Derived from wave_function solutions.

     Principal (n):

     • Energy level
     • n = 1, 2, 3...
     • Size of orbital

     Angular (l):

     • Subshell shape
     • l = 0 to (n-1)
     • s(0), p(1), d(2), f(3)

     Magnetic (mₗ):

     • Orbital orientation
     • mₗ = -l to +l
     • p: 3 orientations

     Spin (mₛ):

     • Electron spin
     • mₛ = ±½

187. Quantum Theory - Physical theory that energy exists in discrete units (quanta). Foundation of modern atomic theory.

     Key equations:

     Core concepts:

     • wave_particle_duality
     • photoelectric_effect
     • planck_constant h
     • Probability-based wave_functions

188. Quartz - Crystalline silica (SiO₂), a network_solid with tetrahedral structure. Shows polymorphism.

     Structure:

     • Si-O-Si angle: 144°
     • Si-O length: 1.61 Å
     • crystal_system: trigonal
     • Density: 2.65 g/cm³

     Transitions:

     • α → β at 573°C
     • Tridymite: 867°C
     • Cristobalite: 1470°C
     • Melting: 1713°C

189. Quaternary Structure - Spatial arrangement of multiple protein subunits in a functional protein complex. Builds on tertiary_structure.

     Stabilising forces:

     • hydrogen_bonds
     • Hydrophobic interactions
     • Salt bridges
     • Van der Waals forces

     Examples:

     • Haemoglobin (4 subunits)
     • Insulin (6 subunits)
     • Antibodies (4 chains)

     Influences protein_folding and function.

R

190. Radian - SI unit of angular measure, defined as arc length equal to radius. Essential in molecular_geometry and vibration analysis.

     Conversions:

     • 2π rad = 360°
     • π rad = 180°
     • 1 rad ≈ 57.3°

     Applications:

     • bond_angle measurement
     • orbital_hybridisation analysis
     • Rotational spectroscopy
     • Angular frequency (ω)

191. Radioactive Decay - Spontaneous nuclear transformation releasing particles or energy. Rate determined by decay_constant.

     First-order kinetics:

     Types:

     • α: He²⁺ emission
     • β⁻: electron emission
     • β⁺: positron emission
     • γ: photon emission
     • EC: electron capture

     half_life relation:

192. Rate Constant - Proportionality constant k in rate_law equations. Temperature dependence given by arrhenius_equation.

     Arrhenius equation:

     Where:

     • A = frequency factor
     • Eₐ = activation_energy
     • R = gas constant
     • T = temperature

     Units depend on overall reaction_rate order:

     • 0th order: mol·L⁻¹·s⁻¹
     • 1st order: s⁻¹
     • 2nd order: L·mol⁻¹·s⁻¹

193. Rate-Determining Step - Slowest elementary_step in a reaction_mechanism, controlling overall reaction rate.

     Characteristics:

     • Highest activation_energy
     • Determines rate_law
     • Cannot be bypassed
     • Focus for catalysis

     Example mechanism:

194. Reaction Quotient - Ratio Q of product to reactant concentrations at any point. Compared with equilibrium_constant K.

     For aA + bB ⇌ cC + dD:

     Relationships:

     • Q < K: reaction proceeds forward
     • Q = K: chemical_equilibrium
     • Q > K: reaction proceeds backward

     ΔG relation:

195. Redox - Coupled reduction and oxidation reactions involving electron_transfer. Changes in oxidation_state.

     Half-reactions:

     Cell potential:

     Common examples:

     • Metal corrosion
     • Battery discharge
     • Photosynthesis
     • Cellular respiration

196. Resonance - Description of delocalised electrons using multiple Lewis structures. Real structure is hybrid of resonance forms.

     Key concepts:

     • electron_delocalisation
     • Equivalent resonance forms
     • Fractional bond_orders
     • Enhanced stability

     Classic examples:

     • Benzene: Kekulé structures
     • Carbonate: three forms
     • Ozone: two forms
     • Nitrate: three forms

197. Rubber - Natural or synthetic elastomer with high elastic deformation capacity. Strengthened by vulcanisation.

     Structure:

     • cis-1,4-polyisoprene
     • polymer chains
     • cross_linking points
     • Amorphous network

     Properties:

     • Glass transition: -72°C
     • Elastic modulus: 0.1-0.5 MPa
     • Maximum strain: 500-1000%
     • Density: 0.92 g/cm³

S

198. Salt - ionic_compound formed from acid-base neutralisation. Contains cation (except H⁺) and anion (except OH⁻, O²⁻).

     Properties:

     • High melting points
     • crystal_lattice structure
     • Variable solubility
     • hydration tendency

     Common examples (formula, solubility g/100mL):

     • NaCl (36.0)
     • KNO₃ (35.7)
     • CaCO₃ (0.0013)
     • AgCl (0.000195)

199. Saturated Solution - Solution containing maximum solute at equilibrium for given temperature. Further addition causes precipitation.

     Equilibrium condition:

     Temperature effects:

     • Most solids: solubility increases with T
     • Gases: solubility decreases with T
     • Exothermic: solubility decreases with T
     • Endothermic: solubility increases with T

     Controls crystallisation behaviour.

200. Shell - Energy level occupied by electrons, designated by principal quantum_numbers n (1,2,3...).

     Capacity:

     • n=1 (K): 2e⁻
     • n=2 (L): 8e⁻
     • n=3 (M): 18e⁻
     • n=4 (N): 32e⁻

     Subshells per n:

     • n=1: 1s
     • n=2: 2s, 2p
     • n=3: 3s, 3p, 3d
     • n=4: 4s, 4p, 4d, 4f

     Follows aufbau_principle filling order.

201. Solubility Product - equilibrium_constant Ksp for dissolution of sparingly soluble salt. Controls precipitation.

     For MₓAᵧ ⇌ xM⁺ + yA⁻:

     Example values (25°C):

     • AgCl: 1.8 × 10⁻¹⁰
     • BaSO₄: 1.1 × 10⁻¹⁰
     • Ca₃(PO₄)₂: 2.1 × 10⁻³³

     Affected by common_ion_effect.

202. Spectroscopy - Study of matter through interaction with electromagnetic_spectrum. Based on absorption/emission.

     Common types (wavelength):

     • NMR: radio waves (>1m)
     • IR: 2.5-25μm
     • UV-Vis: 200-800nm
     • X-ray: < 10nm

     Energy transitions:

     • NMR: nuclear spin
     • IR: molecular vibration
     • UV-Vis: electronic
     • X-ray: core electrons

     Relates to quantum_levels.

203. Standard State - Reference state for calculating thermodynamic_properties. Basis for ΔG°, ΔH°, ΔS°.

     Conditions:

     • Pure substance
     • 1 bar standard_pressure
     • Specified temperature (usually 298K)
     • Unit activity

     States:

     • Gases: ideal at 1 bar
     • Solutions: 1M
     • Pure solids/liquids: normal state
     • Elements: most stable form

     Used in gibbs_energy calculations.

204. Stereochemistry - Study of spatial arrangement of atoms in molecules and its effect on chemical properties.

     Key concepts:

     • chirality: non-superimposable mirror images
     • optical_activity: plane polarised light rotation
     • Conformations: rotations about single bonds
     • Configuration: spatial arrangement

     Types of isomers:

     • Enantiomers (mirror images)
     • Diastereomers (non-mirror stereoisomers)
     • Conformers (rotational isomers)
     • Geometric (cis/trans)

205. Stoichiometry - Quantitative relationships between reactants and products in chemical reactions. Based on balanced_equations.

     Key calculations:

     • mole ratios
     • Mass relationships
     • limiting_reagent
     • theoretical_yield

     Example:

     • 2 mol H₂ : 1 mol O₂
     • 4g H₂ : 32g O₂
     • Yields 36g H₂O

T

206. Temperature - Measure of average molecular kinetic_energy. Determines direction of heat flow.

     Scales and conversions:

     Key points:

     • Absolute zero: 0 K
     • Triple point of water: 273.16 K
     • Standard conditions: 298.15 K
     • Controls maxwell_boltzmann_distribution

     thermal_equilibrium occurs at equal T.

207. Theoretical Yield - Maximum product amount possible from given reactants based on stoichiometry and limiting_reagent.

     Calculations:

     Factors affecting percent_yield:

     • Competing reactions
     • Reversible reactions
     • Incomplete reactions
     • Product loss in isolation

     Example yields:

     • Industrial: 80-90%
     • Laboratory: 60-80%
     • Complex synthesis: < 50%

208. Thermochemistry - Study of heat changes in chemical reactions. Branch of thermodynamics.

     Key equations:

     Methods:

     • calorimetry
     • Bomb calorimetry
     • Solution calorimetry

     Measures reaction enthalpy.

209. Thermodynamics - Study of energy transformations and their relationship to system changes.

     Laws:

     1. Energy conserved (ΔE = q + w)
     2. entropy increases (ΔS_universe > 0)
     3. S = 0 at 0 K (perfect crystal)

     Key functions:

     • enthalpy: H = E + PV
     • gibbs_energy: G = H - TS
     • Work: w = -PΔV (expansion)

     Determines reaction spontaneity.

210. Titration - Volumetric analysis determining unknown concentration via controlled reagent addition to equivalence_point.

     Types:

     • acid_base
     • Redox
     • Complexometric
     • Precipitation

     Key points:

     • Requires standardisation
     • Uses chemical indicators
     • Precise volume measurement
     • Known stoichiometry

     Common indicators (pH range):

     • Methyl orange (3.1-4.4)
     • Bromothymol blue (6.0-7.6)
     • Phenolphthalein (8.2-10.0)

211. Transition State - Highest energy configuration along reaction_coordinate. Maximum potential energy point.

     Characteristics:

     • Requires activation_energy
     • Unstable configuration
     • Partial bonds formed/broken
     • Cannot be isolated

     Different from intermediates:

     • Higher energy
     • Shorter lifetime
     • No stable structure
     • Represents maximum on energy diagram

     Key in reaction_mechanisms.

212. Triple Point - Temperature and pressure where solid, liquid, and gas phases coexist in equilibrium. Key feature of phase_diagrams.

     Common values (T, P):

     • H₂O: 273.16 K, 611.657 Pa
     • CO₂: 216.55 K, 5.11 bar
     • N₂: 63.18 K, 0.125 bar

     Determines:

     • Phase boundaries
     • vapour_pressure
     • phase_transition conditions

213. Tungsten - transition_metal (W, Z=74). Highest melting point of metals (3422°C). refractory_metal.

     Properties:

     • Density: 19.25 g/cm³
     • Common oxidation_states: +6, +5, +4, +2
     • Strong metallic_bonding
     • Resistant to corrosion

     Key compounds:

     • WO₃: yellow oxide
     • WC: extremely hard carbide
     • W(CO)₆: organometallic
     • WS₂: lubricant

U

214. Uncertainty - Range of doubt in a measurement. Combination of random and systematic errors affecting precision and accuracy.

     Calculations:

     Rules for significant_figures:

     • Addition/subtraction: least decimal places
     • Multiplication/division: least sig figs
     • Functions: least sig figs of input

     Relates to quality of measurement.

215. Unit Cell - Smallest repeating unit of a crystal_structure. Basic building block of lattice.

     Common types (with packing_efficiency):

     • Simple cubic (52.4%)
     • Body-centred cubic (68.0%)
     • Face-centred cubic (74.0%)
     • Hexagonal close-packed (74.0%)

     Key parameters:

     • Edge lengths (a, b, c)
     • Angles (α, β, γ)
     • coordination_number
     • Atoms per cell

     Volume calculation:

216. Unpaired Electron - Single electron occupying an orbital. Causes paramagnetism and influences reactivity.

     Properties:

     • spin = ±½
     • Magnetic moment = √n(n+2) BM
     • Contributes to multiplicity
     • Affects bond formation

     Common examples (number):

     • O₂: 2
     • NO: 1
     • Fe³⁺: 5
     • Cr³⁺: 3

     Determined by electron_configuration.

217. Uranium - Heavy actinide element (U, Z=92). Primary fuel for nuclear_fission.

     Isotopes (abundance):

     • ²³⁸U (99.27%), t½ = 4.47×10⁹ years
     • ²³⁵U (0.72%), fissile
     • ²³⁴U (0.0055%)

     Properties:

     • Common oxidation_states: +6, +5, +4, +3
     • Density: 19.1 g/cm³
     • radioactive_decay: α emission
     • Forms UF₆, UO₂²⁺, U₃O₈

V

218. Valence Electron - Outermost shell electrons participating in bonding. Determines chemical behaviour and oxidation_states.

     Distribution in periodic_table:

     • Group 1 (ns¹): 1e⁻
     • Group 2 (ns²): 2e⁻
     • Group 13-18: ns²np¹⁻⁶
     • Transition metals: (n-1)d¹⁻¹⁰ns¹⁻²

     Relation to electron_configuration:

     • s-block: ns electrons
     • p-block: ns + np electrons
     • d-block: (n-1)d + ns electrons

219. Van der Waals Forces - Weak intermolecular_forces including dipole-dipole and london_dispersion forces.

     Energy relation:

     Strength factors:

     • Molecular size
     • polarisability
     • Surface contact area
     • Temperature

     Typical energies:

     • Dispersion: 0.1-1 kJ/mol
     • Dipole-dipole: 1-10 kJ/mol
     • H-bonding: 10-40 kJ/mol

220. Vapour Pressure - Pressure exerted by vapour in phase_equilibrium with its liquid/solid. Related to boiling_point.

     Clausius-Clapeyron equation:

     Factors:

     • Temperature
     • intermolecular_forces
     • Surface area
     • External pressure

     Application in clausius_clapeyron analysis.

221. Viscosity - Resistance to flow due to internal friction. Measured in Pa·s or poise (P).

     Flow equation:

     Temperature dependence:

     Common values (20°C, mPa·s):

     • Water: 1.002
     • Blood: 4.0
     • Olive oil: 84
     • Honey: ~10,000

     Influenced by:

     • molecular_mass
     • intermolecular_forces
     • temperature
     • Pressure

222. Voltaic Cell - Electrochemical cell converting chemical energy to electrical energy via spontaneous redox reactions.

     Cell potential:

     Components:

     • Anode (oxidation)
     • Cathode (reduction)
     • salt_bridge
     • External circuit

     Standard reduction potentials E° (V):

     • F₂/F⁻: +2.87
     • Cl₂/Cl⁻: +1.36
     • Cu²⁺/Cu: +0.34
     • 2H⁺/H₂: 0.00
     • Zn²⁺/Zn: -0.76

     electrode_potential determines electron_transfer.

W

223. Wavefunction - Mathematical function ψ describing quantum state of particle. Solution to schrodinger_equation.

     Properties:

     • ψ² gives probability density
     • Must be continuous
     • Must be normalisable
     • Contains all quantum_numbers

     For hydrogen atom:

     Determines electron_density in orbitals:

224. Wavelength - Distance between repeating units in wave. Key parameter in electromagnetic_spectrum.

     Key relationships:

     Spectral regions:

     • γ-rays: < 10 pm
     • X-rays: 10 pm - 10 nm
     • UV: 10-400 nm
     • Visible: 400-700 nm
     • IR: 700 nm - 1 mm

     Used in spectroscopy and energy_level analysis.

225. Weak Acid - Partial ionisation in water. Characterised by acid_dissociation_constant Ka.

     Equilibrium:

     Ka expression:

     Common examples (Ka):

     • CH₃COOH: 1.8 × 10⁻⁵
     • HF: 6.8 × 10⁻⁴
     • H₂CO₃: 4.3 × 10⁻⁷

     Important in buffer_solution and pH calculations.

226. Work - Energy transfer through force over distance. Important in thermodynamics and energy calculations.

     Types and equations:

     Sign convention:

     • Positive: work done on system
     • Negative: work done by system

     Related to pressure-volume changes in gas laws.

X

227. X-ray Crystallography - Technique determining crystal_structure using X-ray diffraction patterns.

     bragg_equation:

     Key measurements:

     • d-spacing
     • Intensity
     • Phase angles
     • Miller indices

     Determines:

     • unit_cell parameters
     • Atomic positions
     • Bond lengths/angles
     • Crystal symmetry

     Resolution limit ≈ λ/2.

228. X-ray Diffraction - Elastic scattering of X-rays by crystal lattice planes. Reveals crystal_structure.

     Conditions:

     • wavelength ≈ atomic spacing
     • Regular crystal structure
     • Coherent scattering
     • Constructive interference

     Analysis methods:

     • Powder diffraction
     • Single crystal
     • Laue patterns
     • Rotating crystal

     Uses Cu Kα (λ = 1.5418 Å) or Mo Kα (λ = 0.7107 Å).

229. X-ray Photoelectron - spectroscopy technique analysing surface composition via photoelectric_effect.

     Energy relation:

     Where:

     • EK = kinetic energy
     • h𝜈 = photon energy
     • EB = binding_energy
     • φ = work function

     Probes core_electrons:

     • 1s: 0-1000 eV
     • 2p: 50-200 eV
     • 3d: 0-50 eV

230. Xenon - noble_gas (Xe, Z=54). Forms chemical compounds despite filled electron_configuration.

     Properties:

     • Ground state: [Kr]4d¹⁰5s²5p⁶
     • Common oxidation_states: +2, +4, +6, +8
     • Density: 5.894 g/L (0°C)
     • Boiling point: 165.1 K

     Key compounds:

     • XeF₂: linear
     • XeF₄: square planar
     • XeO₃: pyramidal
     • XeF₆: distorted octahedral

     molecular_geometry follows VSEPR theory.

Y

231. Yield - Amount of product obtained from chemical reaction. Usually expressed as percentage of theoretical_yield.

     Calculation:

     Factors affecting yield:

     • Side reactions
     • Reversible reactions
     • limiting_reagent
     • Physical losses

     stoichiometry guides yield prediction.

232. Young's Modulus - Measure of material stiffness. Ratio of tensile stress to tensile strain.

     Definition:

     Typical values (GPa):

     • Diamond: 1220
     • Steel: 200
     • Aluminium: 69
     • Rubber: 0.01-0.1

     Related to:

     • crystal_structure
     • bond_strength
     • Elastic deformation
     • Material properties

233. Yttrium - transition_metal (Y, Z=39). Similar chemistry to lanthanides despite electron_configuration.

     Properties:

     • Ground state: [Kr]4d¹5s²
     • Common oxidation_state: +3
     • Atomic radius: 180 pm
     • Density: 4.472 g/cm³

     Applications:

     • YBa₂Cu₃O₇ superconductor
     • Y₃Al₅O₁₂ (YAG) lasers
     • coordination_chemistry
     • Phosphors

Z

234. Zeolite - Microporous aluminosilicate with regular crystal_structure. Acts as molecular_sieve.

     Structure:

     • SiO₄ and AlO₄ tetrahedra
     • Uniform pore size
     • Channel networks
     • Exchangeable cations

     Applications:

     • catalysis
     • ion_exchange
     • Gas separation
     • Water softening

     Common types:

     • ZSM-5
     • Zeolite Y
     • Mordenite
     • Faujasite

235. Zero Order - Reaction where reaction_rate is independent of reactant concentration. Common in surface-limited reactions.

     Integrated rate_law:

     Key features:

     • Linear concentration decay
     • half_life: t½ = [A]₀/2k
     • Rate constant units: M·s⁻¹
     • Surface-saturated reaction_mechanisms

     Examples:

     • Photochemical decomposition
     • Enzyme saturation kinetics
     • Surface catalysis

236. Zinc - transition_metal (Zn, Z=30). Full d-shell electron_configuration gives unique properties.

     Properties:

     • Ground state: [Ar]3d¹⁰4s²
     • Common oxidation_state: +2
     • E°(Zn²⁺/Zn): -0.76V
     • Amphoteric oxide ZnO

     Key compounds:

     • Zn²⁺: tetrahedral complexes
     • ZnO: semiconductor
     • Zn-finger proteins
     • Brass (Cu-Zn alloy)

     Important in coordination_chemistry.

237. Zwitterion - Molecule with equal positive and negative charges. Common form of amino_acids at isoelectric_point.

     Structure:

     Properties:

     • High dipole moment
     • pH dependent
     • Highly water soluble
     • acid_base buffer

     pH effects:

     • Low pH: cation
     • pI: zwitterion
     • High pH: anion

#

238. 1s Orbital - Spherically symmetric atomic_orbital with principal quantum_numbers n=1, l=0.

     Properties:

     • Lowest energy orbital
     • Contains up to 2 electrons
     • No angular nodes
     • Radial probability maximum at nucleus
     • Radial function: R₁₀(r) = 2(Z/a₀)³/²e⁻ᶻʳ/ᵃ⁰

     Foundation of electron_configuration.

239. 2p Orbital - Set of three degenerate orbitals with n=2, l=1. Essential in chemical bonding.

     Characteristics:

     • Three orientations (px, py, pz)
     • Contains up to 6 electrons total
     • One angular node each
     • Dumbbell-shaped
     • Important in molecular_orbital formation

     Key in electron_configuration and bonding.

240. 3d Orbital - Set of five degenerate orbitals with n=3, l=2. Critical in transition_metal chemistry.

     Shapes:

     • dz²: dumbbell + ring
     • dx²-y²: four-lobed planar
     • dxy, dxz, dyz: four-lobed diagonal

     Applications:

     • crystal_field_theory
     • Colour of complexes
     • Magnetic properties
     • electron_configuration

241. Alpha Particle - He²⁺ ion emitted in radioactive_decay. Identical to helium_nucleus.

     Properties:

     • Mass: 4.001506 u
     • Charge: +2e
     • Strong ionisation
     • Low penetration depth

     Decay equation:

     Common in heavy element nuclear_reactions.

242. Beta Particle - High-energy electron (β⁻) or positron (β⁺) emitted in radioactive_decay.

     Decay equations:

     Properties:

     • Continuous energy spectrum
     • Greater penetration than α
     • Less ionising than α
     • Involves weak nuclear_reaction

243. χ (Electronegativity) - Tendency of atom to attract bonding electrons. Usually on pauling_scale.

     Related to:

     • electron_affinity
     • ionisation_energy
     • bond_polarity

     Trends:

     • Increases left to right
     • Decreases down groups
     • F most electronegative (4.0)
     • Cs least electronegative (0.79)

244. ΔG (Gibbs Energy Change) - Change in gibbs_energy for a process. Determines spontaneity.

     Key equations:

     Interpretation:

     • ΔG < 0: spontaneous
     • ΔG = 0: equilibrium
     • ΔG > 0: non-spontaneous

     Related to enthalpy and entropy.

245. ΔH (Enthalpy Change) - Heat transferred at constant pressure. Key thermodynamics parameter.

     Types:

     • ΔH°f (formation)
     • ΔH°c (combustion)
     • ΔH°rxn (reaction)
     • ΔH°soln (solution)

     Sign convention:

     • Negative: exothermic
     • Positive: endothermic

     Follows hess_law for multi-step processes.

246. ΔS (Entropy Change) - Change in molecular disorder. Key in third_law thermodynamics.

     Calculation:

     For processes:

     • Phase changes
     • Temperature changes
     • Gas expansion
     • Mixing/dissolution

     Determines spontaneity with ΔH.

247. γ-Radiation - High-energy electromagnetic_radiation from nuclear_transitions.

     Properties:

     • No mass or charge
     • Highest penetration
     • E = hν
     • Wavelength < 10⁻¹¹ m

     Nuclear equation:

     Represents energy_level transitions in nucleus.

248. λ (Wavelength) - Distance between wave repetitions. Key in electromagnetic_spectrum analysis.

     Relationships:

     Used in:

     • spectroscopy
     • Diffraction
     • energy_level calculations
     • De Broglie relation

249. μ (Dipole Moment) - Vector measure of charge separation in molecules. Key in understanding polarity.

     Calculation:

     Units:

     • Debye (D)
     • 1 D = 3.336 × 10⁻³⁰ C·m

     Affected by:

     • electronegativity differences
     • Molecular geometry
     • Bond types
     • intermolecular_forces

250. Ω (Solid Angle) - Three-dimensional angle in steradians (sr). Important in spectroscopy and crystallography.

     Properties:

     • Full sphere = 4π sr
     • Hemisphere = 2π sr
     • Cone: Ω = 2π(1-cos θ)

     Applications:

     • Radiation patterns
     • diffraction analysis
     • Angular distribution
     • Molecular symmetry

251. π-Bond - Covalent bond formed by parallel orbital_overlap of p orbitals. Key in multiple bonds.

     Characteristics:

     • Side-to-side overlap
     • Electron density above/below axis
     • Weaker than σ bonds
     • Free rotation blocked

     Common in:

     • C=C double_bonds
     • Aromatics
     • Conjugated systems
     • molecular_orbital theory

252. ψ (Wavefunction) - Solution to schrodinger_equation describing quantum state. Foundation of quantum_mechanics.

     Properties:

     • |ψ|² gives probability_density
     • Must be continuous
     • Must be single-valued
     • Must be normalisable

     Applications:

     • orbital shapes
     • Energy levels
     • Spectroscopy
     • Bonding theory

253. σ-Bond - Covalent bond formed by head-on orbital_overlap. Strongest type of covalent_bond.

     Properties:

     • Cylindrical symmetry
     • Maximum overlap
     • Free rotation possible
     • Forms molecular framework

     Found in:

     • All single bonds
     • First bond in multiple bonds
     • molecular_orbital theory
     • Core bonding framework