MOJZA Chemistry Notes (9701) - VOCABULARY Flashcards

Vocabulary flashcards covering key concepts and terms from the MOJZA Chemistry Notes for AS/A-level topics.

Atom

The smallest unit of an element that can exist independently, composed of a nucleus and electrons surrounding it.

Nucleus

The centre of the atom containing protons and neutrons; overall positive charge.

Proton

Positively charged subatomic particle in the nucleus (mass ~1 amu).

Neutron

Electrically neutral subatomic particle in the nucleus (mass ~1 amu).

Electron

Negatively charged subatomic particle orbiting the nucleus (tiny mass).

Atomic Number (Z)

Number of protons in the nucleus; equals the number of electrons in a neutral atom.

Mass Number (A)

Total number of protons and neutrons in the nucleus.

Isotopes

Atoms with the same atomic number but different mass numbers (different neutrons); chemistries largely similar.

Relative Atomic Mass (Ar)

Weighted average mass of an element’s atoms relative to 1/12 of the mass of carbon-12.

Relative Molecular Mass (Mr)

Average mass of a molecule relative to 1/12 of the mass of carbon-12.

Relative Formula Mass (RFM)

Average mass of a formula unit of an ionic compound relative to 1/12 of carbon-12.

Avogadro’s Constant (NA)

6.02 × 10^23 particles per mole; number of particles in one mole.

Mole

Amount of substance containing NA particles; mass equals Ar or Mr for the substance.

Moles (formula)

Moles = Mass/Ar or Mr; for gases at r.t.p., 1 mole occupies 24 dm³.

Empirical Formula

The simplest whole-number ratio of atoms in a molecule.

Molecular Formula

The actual number of atoms of all elements in a molecule.

Combustion

Oxidation of a substance in oxygen to form CO₂ and H₂O; general equation: CxHy + (x + y/2)O₂ → xCO₂ + y/2 H₂O.

Empirical to Molecular Formula Steps

Divide each element’s mass by its atomic mass, divide by smallest value, convert to whole-number ratio.

Electronic Configuration

Arrangement of electrons in atoms; fills subshells in order of increasing energy.

Principal Quantum Number (n)

Energy level/shell number; higher n means farther from nucleus.

Subshells (s, p, d)

Divisions of shells with fixed capacities: s(2), p(6), d(10); energy increases s < p < d.

Orbitals

Regions within a subshell where electrons reside; max 2 electrons per orbital.

Ground State

Most stable electron arrangement with lowest energy; lower-energy subshells fill first.

Degenerate Orbitals

Orbitals in the same subshell with the same energy (e.g., px, py, pz).

Atomic Radius

Half the distance between nuclei in covalently bonded atoms of the same element; generally decreases across a period, increases down a group.

Ionic Radius

Size of a e⁻ cloud in ions; cations smaller than their atoms, anions larger.

Ionisation Energy (IE)

Energy required to remove one mole of electrons from one mole of gaseous atoms; endothermic.

First Ionisation Energy

Energy to remove the first mole of electrons to form 1+ ions.

Second Ionisation Energy

Energy required to remove a second electron after the first has been removed.

Isotopes (chemistry)

Atoms of the same element with different numbers of neutrons; same chemical behaviour, different mass.

Electron Shells

Concentric regions around the nucleus where electrons reside; defined by principal quantum number n.

Subshells

Divisions of electron shells labelled s, p, d; determine electron capacity and shape.

Ground State Configurations of Transition Elements

4s fills before 3d; exceptions (Cu, Cr) stabilize half/full d subshells.

s-block / p-block / d-block / f-block

Groups in the periodic table defined by the type of outer-shell electron occupation (s, p, d, f).

Electronegativity

Ability of an atom to attract electrons in a bond; Fluorine is the highest; trend increases across a period and decreases down a group.

Bond Energy

Energy required to break one mole of covalent bonds (in gas phase); endothermic.

Bond Length

Distance between nuclei in a covalent bond; shorter bonds are typically stronger.

VSEPR Theory

Shape of molecules is determined by repulsion between electron pairs around a central atom.

Dipole Moment

Measure of bond polarity; arrow points from positive to negative end.

Hydrogen Bond

Strong dipole-dipole interaction involving H attached to F, O, or N with another electronegative atom.

Polarity

Separation of charge in a molecule; determined by bond polarity and molecular geometry.

Dipole-Dipole Forces

Intermolecular forces between polar molecules due to permanent dipoles.

Van der Waals Forces

Intermolecular forces incl. induced-dipole and permanent-dipole attractions.

Sigma Bond

End-to-end overlap of atomic orbitals; electron density between nuclei.

Pi Bond

Sideways overlap of adjacent p orbitals; part of double/triple bonds.

Ionic Lattice

Giant array of alternating cations and anions held by ionic attraction.

Giant Covalent Lattice

Network solids with covalent bonds throughout (e.g., diamond, SiO₂).

Giant Metallic Lattice

Metallic bonding with a lattice of cations in a sea of delocalised electrons.

Enthalpy (ΔH)

Total chemical energy change in a process; heat absorbed or released at constant pressure.

Activation Energy (Ea)

Minimum energy required for a reaction to occur.

Exothermic vs Endothermic

Exothermic: ΔH negative (heat released). Endothermic: ΔH positive (heat absorbed).

Standard Conditions (ΔH°)

1 atm (101 kPa), 298 K, 1.0 M solutions; ΔH° denotes standard enthalpy change.

ΔH°f

Enthalpy change when 1 mole of a compound forms from its elements under standard conditions.

ΔH°c

Enthalpy change of combustion; energy released when 1 mole of substance is burned.

ΔH°neut

Enthalpy change of neutralisation; energy released when acid and base form water.

Hess’s Law

Total enthalpy change is independent of the path; can be calculated from known steps.

Redox

Oxidation-reduction reactions involving electron transfer.

Oxidising Agent

Substance that gets reduced and oxidises another species.

Reducing Agent

Substance that gets oxidised and reduces another species.

Disproportionation

Reaction where same species is simultaneously oxidised and reduced.

Equilibrium Constant (Kc)

Ratio of product concentrations to reactant concentrations at equilibrium; temperature dependent.

Kp

Equilibrium constant expressed in terms of partial pressures.

Le Chatelier’s Principle

If a system at equilibrium is disturbed, it shifts to counteract the disturbance.

Haber Process

Industrial synthesis of NH₃ from N₂ and H₂ using iron catalyst under high pressure and moderate temperature.

pH

Negative logarithm of H⁺ concentration; measures acidity (0-14 scale).

pH Meter

Instrument for measuring pH accurately.

Titration

Technique to determine concentration by gradual addition of a titrant until equivalence point.

Rate of Reaction

Speed of a chemical reaction; change in concentration per unit time.

Collision Theory

Reaction occurs when particles collide with proper orientation and energy (Ea).

Catalyst

Substance that increases reaction rate without being consumed; lowers Ea.

Homogeneous Catalyst

Catalyst in the same phase as reactants.

Heterogeneous Catalyst

Catalyst in a different phase from the reactants.

Boltzmann Distribution

Distribution of molecular energies at a given temperature; higher temp flattens curve and shifts peak right.

Isomerism

Compounds with same formula but different arrangements; structural and stereoisomerism.

Structural Isomerism

Chain, positional, or functional-group isomerism.

Geometrical (cis-trans) Isomerism

Isomerism around C=C due to restricted rotation; cis vs trans forms.

Optical Isomerism

Chiral centers with non-superimposable mirror images (enantiomers); racemate is optically inactive.

Alkanes

Saturated hydrocarbons with general formula CnH₂n+₂.

Alkenes

Unsaturated hydrocarbons with C=C double bonds.

Halogenoalkanes

Alkanes in which one or more hydrogens are replaced by halogens.

Nucleophiles

Species that donate an electron pair to an electrophile.

Electrophiles

Species that accept an electron pair from a nucleophile.

SN1/SN2

Nucleophilic substitution mechanisms: SN1 is unimolecular, SN2 is bimolecular.

Hydration (of Alkenes)

Addition of water to alkenes to form alcohols in the presence of an acid.

Markovnikov’s Rule

In addition to alkenes, the electrophile adds to the more substituted carbon in order to form the more stable carbocation.

Hydrogenation

Addition of H₂ to alkenes to form alkanes; typically Pt/Ni catalysts.

Oxidation of Alkenes with KMnO₄

Cold KMnO₄ forms diols; hot KMnO₄ cleaves double bond to carbonyl compounds.

Halogenoalkanes (reactions)

Reactions include nucleophilic substitutions (with NaOH, KCN, NH₃) and eliminations (with base).

Hydrolysis of Halogenoalkanes

Halogenoalkanes react with water or hydroxide to give alcohols or carboxylates.

Esters

Derivatives of carboxylic acids with -COOR; formed by esterification and hydrolysed by water.

Nitriles

Organic compounds with -CN group; formed by nucleophilic substitution with KCN.

Hydroxynitriles

Compounds with both -OH and -CN groups; formed via nucleophilic addition to carbonyls.

Amines

Organic compounds containing -NH₂ group; can be primary, secondary, or tertiary.

Polymerisation

Process of combining many monomers to form a polymer; addition polymerisation involves double bonds.

IR Spectroscopy

Technique to identify functional groups by characteristic infrared absorptions.

Mass Spectrometry

Technique to identify unknown compounds by molecular ion and fragmentation patterns.