AS-Level Chemistry AS Level Chemistry Definitions
1. The relative atomic mass (Ar) of an element is defined as the average mass of one atom compared to 1/12 the mass of a 12C atom. OR
The relative atomic mass (Ar) of an element is defined as the mass of one mole of atoms compared to 1/12 the mass of one mole of 12C atoms.
2. The relative isotopic mass of an isotope (of a particular element) is defined as the mass of one isotope compared to 1/12 the mass of a 12C atom.
3. The relative molecular mass of a molecule is defined as the average mass of one molecule compared to 1/12 the mass of a 12C atom.
4. The relative formula mass of an ionic compound is defined as the average mass of one formula unit compared to 1/12 the mass of a 12C atom. OR
The relative formula mass of an ionic compound is defined as the mass of one mole of formula units compared to 1/12 the mass of one mole of 12C atoms.
5. A mole of substance is defined as the amount of substance that contains as many entities (atoms, molecules, ions, electrons or any other particles) as the number of atoms in 12g of the carbon-12.
It is equal to 6.022 X 1023, which is called the Avogadro constant or Avogadro number.
𝑛= 𝑚𝑎𝑠𝑠𝑀𝑟 ; 𝑛=𝑐𝑜𝑛𝑐𝑒𝑛𝑡𝑟𝑎𝑡𝑖𝑜𝑛 × 𝑣𝑜𝑙𝑢𝑚𝑒 ; 𝑛=𝑉 (𝑖𝑛 𝑑𝑚3)24 (for gases only)
Redox Reaction
6. Oxidation is a process where a chemical species loses electrons; (Oxidation Is Loss)
Reduction is a process where a chemical species gains electrons. (Reduction Is Gain)
A redox reaction refers to a reaction where oxidation and reduction occurs simultaneously.
7. An oxidizing agent is a species that accepts / gains electrons (is reduced) n a reaction.
A reducing agent is a species that donates / loses electrons (is oxidised) in a reaction.
8. A disproportionation reaction is a redox reaction in which one species is simultaneously oxidised and reduced.
e.g. 2Cu+ Cu + Cu2+
9. Atomic number of an element refers to the number of protons it contains.
Mass number (nucleon number) refers to the sum of the protons and neutrons it contains.
10. Isotopes refer to atoms of the same element with the same number of protons but different number of neutrons.
11. An atomic orbital is defined as a region of three-dimensional space around the nucleus, whereby there is a 95% chance of locating a particular electron. Each orbital has a characteristic energy level and shape. For ‘A’ level syllabus, you need to know the shapes of s and p orbitals.
Chemical Bonding
12. Valence-shell electron pair repulsion (VSEPR) theory is a model used to predict the shape of individual molecules based upon the extent of electron-pair electrostatic repulsion.
Lone pair – lone pair repulsion > Lone pair – bond pair repulsion > Bond pair – bond pair repulsion
13. Metallic bond is the electrostatic attraction between positively charged cations and the ‘sea’ of delocalized electrons.
14. Electrovalent (Ionic) bond is the electrostatic attraction between oppositely charged ions which have been formed by the transfer of one or more electrons to achieve the stable electronic configuration of a noble gas.
Coordination number of an ion in an ionic compound refers to the number of neighboring oppositely charged ions.
15. Covalent bond is the electrostatic force of attraction of the nuclei of the 2 atoms for the shared pair(s) of electrons between them.
16. Dative / Co-ordinate Covalent bond is a covalent bond in which a pair of electrons is shared between 2 atoms but ONLY ONE of them provides both electrons that make up the bond.
17. Electronegativity refers to the ability/tendency of an atom to attract electrons in a bond towards itself. Electronegativity increases across the period and decreases down a group.
18. Permanent dipole-permanent dipole interactions are a type of intermolecular forces between polar molecules (molecules with a net dipole moment) which have a simple covalent structure.
19. Temporary dipole-induced dipole interactions are a type of intermolecular forces between non-polar molecules (molecules with NO NET dipole moment) which have a simple covalent structure.
20. Hydrogen bonds are a special case of permanent dipole-permanent dipole interactions, whereby there is an attractive interaction of a hydrogen atom with an electronegative atom, such as nitrogen, oxygen or fluorine (typically from another molecule). Do not confuse this with a covalent bond between H and N, O or F. The Gaseous State
21. Basic Assumptions of kinetic theory of gases
- Gases consist of small particles of negligible size/volume as compared to the size of the container.
- Gas particles have negligible intermolecular forces of attraction between each other.
- Collisions between gas particles are perfectly elastic. I.e. there is no loss of kinetic energy upon collision.
22. Hess’ law states that the change in enthalpy accompanying a reaction is independent of the path taken between the initial and final states.
23. The standard enthalpy change of reaction is the enthalpy change when molar quantities of reactants (as specified by the chemical equation) react to form products under standard conditions 25C and 1 atm.
24. The standard enthalpy change of formation of a compound, is the enthalpy change when 1 mole of a pure compound in a specified state is formed from its constituent elements in their standard states, under standard conditions 25C and 1 atm.
25. The standard enthalpy change of combustion of a compound, is the enthalpy change when 1 mole of that compound is completely burnt in oxygen under standard conditions 25C and 1 atm.
26. The standard enthalpy change of neutralization, is the enthalpy change when an acid and a base react to form 1 mole of water under standard conditions 25C, and 1 atm.
27. The standard enthalpy change of atomization of an element, is the enthalpy change when 1 mole of atoms in the gaseous state is formed from the element in its normal physical state under standard conditions 25C and 1 atm.
28. The bond dissociation energy of a bond is the energy required to break one mole of chemical bonds between two atoms in a molecule in the gaseous phase.
29. The first ionization energy of an element, is the energy required to remove 1 mole of electrons from 1 mole of gaseous atoms, to form 1 mole of gaseous singly charged cations. M(g) -> M+(g) + e−
30. The second ionization energy of an element, is the energy required to remove 1 mole of electrons from 1 mole of gaseous singly charged cations, to form 1 mol of gaseous doubly charged cations. M+(g) -> M2+(g) + e−
31. Lattice energy is the energy evolved when 1 mole of an ionic solid is formed from its constituent gaseous ions under standard conditions 25C and 1 atm.
32. The standard enthalpy change of hydration of a gaseous ion, is the enthalpy change when 1 mole of hydrated aqueous ions is formed from the gaseous ions under standard conditions 25C and 1 atm. .
33. The standard enthalpy change of solution of an ionic compound, is the enthalpy change when 1 mole of an ionic compound is dissolved in a large excess of water under standard conditions 25C and 1 atm. Δ𝐻𝑠𝑜𝑙𝑢𝑡𝑖𝑜𝑛= ΣΔ𝐻ℎ𝑦𝑑𝑟𝑎𝑡𝑖𝑜𝑛 − 𝐿𝑎𝑡𝑡𝑖𝑐𝑒 𝑒𝑛𝑒𝑟𝑔𝑦
Dynamic Equilibrium refers to a reversible reaction in which the forward and the backward reactions are taking place at the same rate and concentrations of reactants and product are constant.
34. Le Chatelier’s Principle states that if a system in equilibrium is subjected to a change which disturbs the equilibrium, the system will respond in such a manner as to reduce or counteract the effect of the change.
35. Monoprotic or monobasic acids can donate only one proton. E.g. HCl, HNO3 and CH3COOH
Diprotic or dibasic acids can donate two protons. E.g. H2SO4, H2S and H2CO3
A Bronsted acid is a proton donor.
A strong acid is one that dissociates completely in aqueous solution to give H3O+ ions.
HA (aq) + H2O (l) -> H3O+ (aq) + A- (aq)
Weak acids only dissociate partially in aqueous solution forming ionic equilibrium systems
HA (aq) + H2O (l) H3O+ (aq) + A- (aq)
36. A Bronsted base is a proton acceptor.
A strong base is one that dissociates completely in aqueous solution to give OH- ions.
B (aq) + H2O (l) -> BH+ (aq) + OH- (aq)
37. The minimum energy which colliding molecules must possess for successful collision/reaction is called the activation energy, Ea.
38. A catalyst is a substance that increases the rate of a reaction by providing an alternative reaction pathway that has lower activation energy.
39. Empirical formula is the simplest formula that shows the ratio of each kind of atom in a molecule. e.g. C2H5 is the empirical formula for C4H10
40. Molecular formula shows the actual number of each kind of atoms in a molecule. e.g. C4H10.
41. Structural formula shows how the atoms are connected to each other in a molecule. e.g.
CH3CH2CH2CH3
42. Homologous series are compounds have the same general formula and functional group and each homologue differs from its neighbor by a fixed group of atoms (e.g.–CH2). As we go down a homologous series, the chemical properties remain unchanged but there is a gradual change in physical properties. Examples of homologous series are alkanes, alkenes, alcohols…..
43. Structural isomerism refers to compounds with the same molecular formula but different structural formula. E.g. CH3COOCH3 and C2H5COOH
44. Stereoisomerism refers to compounds that have the same molecular formula but with different spatial arrangements.
Geometric isomers have same carbon skeleton with double bonds restricting free rotation. For geometric isomerism to exist, there must be two different groups of atoms bonded to each side of the C=C bond.
Optical isomers are non-superimposable mirror images of each other (enantiomers). Isomers have at least one chiral C atom, i.e. there are four different groups attached and have no plane of symmetry. An equal proportion of enantiomers forms a racemic mixture which is optically inactive.
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