ACJC Summary Definitions for H2 Chemistry

1 ACJC CHEMISTRY DEPARTMENT JC2 H2 Revision 2024 Summary of definitions for H2 This list of definitions may not be exhaustive. Students are to refer to the lecture notes given. Thank you, Chemistry Department Term Definition 1 Relative isotopic mass The relative isotopic mass of an element is the mass of an atom of the isotope relative to 1 12 of the mass of an atom of 12C. 2 Relative atomic mass The relative atomic mass of an element is the average mass of one atom of the element relative to 1 12 of the mass of one atom of 12C. 3 Relative molecular mass (molecules) The relative molecular mass of a molecule is the average mass of one molecule of a substance relative to 1 12 of the mass of an atom of 12C. 4 Relative formula mass (ions/ ionic compounds) The relative formula mass of an ionic compound is the mass of one formula unit of a substance relative to 1 12 of the mass of one atom of 12C. 5 Mole One mole is the a mount of substance that contains 6.02 x 10 23 particles (atoms, molecules, ions, electrons). 6 Empirical formula Empirical formula of a compound is the formula that shows the simplest whole number ratio of the atoms of the different elements in the compound. 7 Molecular formula Molecular formula of a compound is one which shows the actual number of atoms of each element present in one molecule of a compound.

2 8 Disproportionation A disproportionation reaction is a redox reaction in which one species is simultaneously oxidised and reduced. 9 Mass number/ nucleon number Defined as the total number of protons and neutrons in the nucleus of an atom 10 Atomic number / Proton number Defined as the number of protons in the nucleus of an atom 11 Isotopes Isotopes are atoms of the same element with the same number of protons but different number of neutrons. 12 First ionisation energy The first ionisation energy of an element is defined as the energy absorbed to remove one mole of electrons from one mole of gaseous atoms to form one mole of singly charged positive gaseous ions. Example: Ca (g) → Ca+ (g) + e− ∆H1 = +590 kJ mol−1 13 Second ionisation energy The second ionisation energy of an element is defined as the energy absorbed to remove one mole of electrons from one mole of singly charged positive gaseous ions to form one mole of doubly charged positive gaseous ions. Example: Ca+ (g) → Ca2+ (g) + e− ∆H2 = +1150 kJ mol−1 14 Atomic radius Atomic radius is half the shortest inter-nuclear distance found in the structure of the element. 15 Electronegativity Electronegativity is a measure of the ability of an atom in a molecule to attract the shared pair of electrons in a covalent bond. 16 Ionic radius Ionic radius of an ion is the radius of the spherical ion in an ionic compound. 17 Ionic bond Ionic bond is the strong electrostatic attraction between two oppositely charged ions, formed by the transfer of electrons between the original atoms. 18 Covalent bond Covalent bond is the strong electrostatic attraction between the positive nuclei of the two atoms and the bonding electrons shared between them.

3 19 Co-ordinate bond / dative covalent bond A co-ordinate bond (also called a dative covalent bond) is a covalent bond (a shared pair of electrons) in which both electrons come from the same atom. 20 Bond energy Bond energy is the energy absorbed when one mole of a given covalent bond in a gaseous compound is broken into gaseous atoms or molecules. 21 Metallic bond Metallic bond is the strong electrostatic forces of attraction between the positive metal ions and mobile sea of delocalised electrons. 22 Lattice Energy Lattice energy is defined as the energy released when 1 mole of ionic solid is formed from its constituent gaseous ions (at infinite distance apart). 23 Hess’ Law The enthalpy change for a reaction is the same whether the reaction takes place in one step or in a series of steps and is determined by the initial and final states of the system. 24 Standard Enthalpy change of Formation (∆Hfθ) The energy change when one mole of substance is formed from its constituent elements in their standard states at 298 K and 1 bar. E.g. ∆Hfø (Na2CO3): 2Na(s) + C(s) + 3/2O2(g) → Na2CO3(s) Note: ∆Hfø = 0 for elements in standard states. 25 Standard Enthalpy change of Combustion (∆HCθ) The energy evolved when one mole of substance is burnt completely in excess oxygen at 298 K and 1 bar. E.g. ΔHøc (CH3COOH): CH3COOH(l) + 2O2(g) → 2CO2(g) + 2H2O(l) Note: ∆Hcø is always exothermic. DHo = - heat change nsubstance burnt c

4 26 Standard Enthalpy change of neutralisation (∆Hneuø) The energy evolved when one mole of water is formed in the reaction between an acid and a base at 298 K and 1 bar. NaOH(aq) + ½ H2SO4(aq) → ½ Na2SO4(aq) + H2O(l) Note: ∆Hneuø is always exothermic. DHo = - heat change nwater formedn Standard Enthalpy Change of Atomisation (∆Hatomϴ) Standard enthalpy change of atomisation of an element is the energy absorbed when one mole of gaseous atoms is formed from the element in its standard state at 298 K and 1 bar. Na (s) → Na (g) ∆Hatomϴ (Na) = +107 kJ mol−1 Standard Enthalpy Change of Hydration (∆Hhyd) Standard Enthalpy change of hydration is the energy evolved when one mole of gaseous ions is converted to one mole of aqueous ions at infinite dilution at 298 K and 1 bar. Mg2+ (g) + aq → Mg2+ (aq) ∆Hhyd (Mg2+) = −1926 kJ mol−1 F– (g) + aq → F– (aq) ∆Hhyd (F– ) = −524 kJ mol−1 Standard Enthalpy Change of Solution (∆Hsol) Standard Enthalpy change of solution is the energy change when one mole of substance (in its standard state) is completely dissolved in infinite dilution in an aqueous solution at 298K and 1 bar. NaCl (s) + aq → Na+ (aq) + Cl– (aq) ∆Hsol (NaCl) = +5 kJ mol−1 Electron Affinity (EA) First electron affinity is the energy change when one mole of gaseous atoms of an element gain one mole of electrons to form one mole of singly charged negative gaseous ions.

5 Second electron affinity is the energy change when one mole of singly charged negative gaseous ions of an element gain one mole of electrons to form one mole of doubly charged negative gaseous ions. O (g) + e– → O– (g) 1st EA (O) = −141 kJ mol−1 O– (g) + e– → O2– (g) 2nd EA (O) = +780 kJ mol−1 27 Bond Energy (BE) The energy absorbed when one mole of a given covalent bond in a gaseous compound is broken into gaseous atoms or molecules. E.g. Bond energy of O – H: O –H (g) → O (g) + H (g) Note: BE is always endothermic (bond breaking). Entropy, S Entropy is a measure of the disorder/ disorderliness of a system. Gibbs Free Energy Gibbs Free Energy is a function that relate s a reaction’s enthalpy and entropy terms to its spontaneity. ∆Go = ∆Ho – T∆So ∆Go: change in standard Gibbs free energy (kJ mol−1) wher e ∆Ho: standard enthalpy change of the system (kJ mol−1) ∆So: standard entropy change of the system (J mol−1 K−1) T: temperature of the system (in K) 28 Rate of reaction The rate of a reaction is defined as the change of amount or concentration of a reactant or product per unit time. Units: mol dm-3 s-1 or mol dm-3 min-1 or mol dm-3 h-1 [mol dm-3 time-1] 29 rate equation, order of reactions Rate equation (or rate law ) is the mathematical relationship between the reaction rate and the concentration of the individual reactants. Rate of reaction = Time taken for reaction to complete Final Concentration – Initial Concentration

6 (m and n), rate constant (k) aA + bB cC rate = k [A]m [B]n Rate constant (k) is the proportionality constant for a given reaction at a certain temperature. Order of reaction