list of h2 chemistry definitions

List of H2 Chemistry Definitions for ‘A’ Levels Compiled by: Seetoh Zhiwei of 14S03Q *Underlined portions are keywords that must be present, or else no marks will be awarded. (Some answer schemes may accept words to that effect/wtte) Important – predicted definitions/appeared in past-year papers Not in current H2 syllabus – appeared in past-year papers Atoms, Molecules and Stoichiometry 1. Isotopes are atoms of the same element with the same number of protons but different number of neutrons. (Candidates are not required to know isotones and isobars.) 2. Relative atomic mass (Ar) of an element is the ratio of the average mass of one atom to of the mass of an atom of the isotope 12C. 3. Relative molecular mass (Mr) of a molecule is the ratio of the average mass of one molecule to

of the mass of an atom of the isotope 12C.

4. Relative formula mass of an ionic substance is the ratio of the average mass of one formula unit of that substance to

of the mass of an atom of the isotope 12C.

5. The molar volume of a gas is defined as the volume occupied by one mole of the gas at standard temperature and pressure of 273.15K and 1 atm. It is calculated to be 22.4dm3. However, at room temperature and pressure of 298K and 1 atm (i.e. standard conditions), it is calculated to be 24dm3. 6. Empirical formula is the simplest formula showing the ratio of the atoms of the different elements in a compound. 7. Molecular formula is the exact formula of a molecule, giving the type of atoms and the number of each type of element. 8. The molar mass is defined as the mass of one mole of a substance, which is numerically equal to the relative atomic mass or relative molecular mass. 9. A standard solution refers to one which is of known concentration, and is used to determine the concentration of another solution through titration. 10. The equivalence point is reached which marks the completion of an acid-base reaction according to the stoichiometric equation. 11. The completion of a titration is observed by a distinct colour change brought about by the use of a suitable indicator, and this point of colour change is called the end-point.

Atomic Structure 1. Aufbau Principle states that electrons are placed into the orbitals of lowest energy, then the orbital of the next lowest energy and so on. 2. Pauli Exclusion Principle states that each orbital may hold only a maximum of 2 electrons and they must have opposite spins. 3. Hund’s rule states that orbitals of a subshell must be occupied singly with parallel spins before pairing occurs, so as to ensure electrons are as far apart as possible to minimize inter-electronic repulsion. 4. Proton number/Atomic number refers to the number of protons in an atom of an element. Page | 1

List of H2 Chemistry Definitions for ‘A’ Levels Compiled by: Seetoh Zhiwei of 14S03Q 5. Nucleon number refers to the total number of protons and neutrons in the atomic nucleus of an element.

Chemical Bonding 1. Ionic bonds are electrostatic forces of attraction between oppositely-charged ions, which are formed by the transfer of electrons from one atom to another. (context specific) 2. Covalent bonds are attractive forces between the localized shared electrons and the positive nuclei of two bonding atoms. 3. A dative covalent bond/coordinate bond is formed when the shared pair of electrons is provided by only one of the bonding atoms. (Draw dot-and-cross diagram/structure of molecule to illustrate) 4. Metals have a giant metallic lattice which is held by metallic bonds between metal cations and the sea of delocalized electrons. Metallic bonds are strong electrostatic forces of attraction between metal cations and the sea of delocalized electrons. 5. Molecules have simple molecule structures made up of discrete molecules attracted to each other by weak intermolecular/van der Waals’ forces of attraction. 6. Covalent radius of an atom is half the inter-nuclear distance between two covalentlybonded atoms.

●

● Covalent radius = 0.5 x dcov dCOV

7. Metallic radius of a metal atom is half the inter-nuclear distance between two metal atoms in the metallic solid lattice. 8. Van der Waals’ radius of an atom is half the inter-nuclear distance between two noncovalently bonded atoms. 9. A sigma (σ) bond is formed when the valence orbitals overlap head-on (collinearly), i.e. the orbitals overlap along the inter-nuclear axis. A pi (π) bond is formed when the valence orbitals overlap sideways (collaterally), i.e. the orbitals overlap parallel to the inter-nuclear axis. Since there is a greater extent of overlap of electron cloud, a σ-bond is generally stronger than a π-bond.

The Gaseous State 1. There are five basic assumptions of the Kinetic Theory of Gases: a. Gases consist of small particles of negligible/insignificant volume. The gas particles are widely separated and can move anywhere in the container. b. The gas particles exert negligible attractive forces on one another. c. Collisions between gas particles are perfectly elastic, i.e. there is no loss of kinetic energy upon collision, and gas particles exchange energy upon collision. Page | 2

List of H2 Chemistry Definitions for ‘A’ Levels Compiled by: Seetoh Zhiwei of 14S03Q d. The gas particles are in constant, random, straight-line motion, colliding occasionally with one another and with the walls of the container. e. The average kinetic energy of particles in a gas is constant at constant temperature. The average kinetic energy is directly proportional to the absolute temperature (Kelvin) and at the same temperature, all gases possess the same average kinetic energy. 2. Boyle’s Law states that at constant temperature, the volume of a fixed mass of a gas is inversely proportional to the applied pressure i.e.

.

3. Charles’ Law states that at constant pressure, the volume of a fixed mass of a gas is proportional to its absolute temperature in Kelvin i.e. . 4. Gay-Lussac’s Law (Pressure Law) states that the pressure of a fixed mass of a gas at constant volume is directly proportional to its absolute temperature i.e. . 5. Avogadro’s Law states that at constant temperature and pressure, the volume of a gas is directly proportional to the number of moles of the gas i.e. . 6. Combined gas law states that at constant number of moles of a gas, the volume of the gas is inversely proportional to its pressure and directly proportional to its absolute temperature i.e.

.

7. An ideal gas is one which follows the ideal gas law (i.e. of pressure and temperature.

) under all conditions

Chemical Energetics 1. Standard enthalpy change of a reaction is the amount of energy absorbed or released in a chemical reaction when the molar quantities of reactants stated in the chemical equation react under standard conditions of 298K and 1atm. 2. Standard enthalpy change of formation is the energy change when 1 mole of a pure substance in a specific state is formed from its constituent elements under standard conditions of 298K and 1atm. 3. Standard enthalpy change of combustion is the energy released when 1 mole of a substance is completely burnt in excess oxygen under standard conditions of 298K and 1atm. 4. Standard enthalpy change of neutralization is the energy change when 1 mole of water is formed when an acid reacts with an alkali under standard conditions of 298K and 1 atm. 5. Standard enthalpy change of atomization is the energy absorbed to form 1 mole of gaseous atoms from an element in its standard state at 298K and 1atm. 6. Standard enthalpy change of hydration is the energy released when 1 mole of gaseous ions is hydrated under standard conditions of 298K and 1atm. 7. Standard enthalpy change of solution is the energy change when 1 mole of a substance in its standard state is completely dissolved in a solvent to form an infinitely dilute solution under standard conditions of 298K and 1atm. 8. The bond energy of an X-Y bond is the average energy absorbed when 1 mole of X-Y bonds are broken in the gaseous state to form 1 mole of X gaseous atoms and 1 mole of Y

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List of H2 Chemistry Definitions for ‘A’ Levels Compiled by: Seetoh Zhiwei of 14S03Q

9. 10.

11. 12.

13. 14.

gaseous atoms. (The standard enthalpy change of atomization of diatomic molecules such as O2 is half the bond energy) The bond dissociation energy of an X-Y bond is the energy absorbed to break 1 mole of that particular X-Y bond in a particular compound in the gaseous state. The lattice energy of an ionic compound is the energy released when 1 mole of a solid ionic compound is formed from its constituent gaseous ions (under standard conditions of 298K and 1 atm). The first ionization energy of an element is the energy required when 1 mole of gaseous atoms loses 1 mole of electrons to form 1 mole of singly-charged positive gaseous ions. The second ionization energy of an element is the energy required when 1 mole of singly-charged positive gaseous ions loses 1 mole of electrons to from 1 mole of doublycharged positive gaseous ions. Entropy is measure of the randomness/disorderliness of matter and energy of a system. Hess’ Law states that the enthalpy change of a reaction is the same regardless of the route by which the chemical change occurs, provided the initial and final states are the same.

Chemical Equilibrium 1. Le Chatelier’s principle states that when a system at equilibrium is subjected to a change, the system will react in such a manner to counteract the change imposed so as to re-establish the equilibrium. 2. Dynamic equilibrium refers to a state in which the forward and backward reactions of a reversible reaction are both taking place at the same rate. As a result, there is no net change in the concentration of the reactants and products. Ionic Equilibrium 1. A strong acid/alkali is one which undergoes complete ionization in aqueous solutions e.g. HCl, HNO3, H2SO4, NaOH, Ba(OH)2 2. A weak acid/alkali is one which undergoes partial ionization in aqueous solutions e.g. CH3COOH, HCN, NH3, C6H5OH. Thus, the concentration of H+ ions in an aqueous solution of weak acid is lower than the initial concentration of the acid. 3. A buffer is a substance which can resist changes in the pH of a system when a small amount of acid or base is added to the system. (write equations to illustrate) 4. Bronsted-Lowry theory of acids and bases states that an acid is a proton donor and a base is a proton acceptor. Every Bronsted-Lowry acid (HA) has a conjugate base (A-) and every Bronsted base (B) has a conjugate acid (HB+), i.e. conjugate acid-base pair.

5. pH is the negative logarithm to base 10 of the hydrogen ion concentration (in mol dm-3) of the solution, pH = -lg[H3O+]. 6. Acid dissociation constant, Ka, is a measure of the strength of the acid. It is unaffected by concentration changes but is affected by temperature changes. Page | 4

List of H2 Chemistry Definitions for ‘A’ Levels Compiled by: Seetoh Zhiwei of 14S03Q HA(aq) + H2O(l)

A–(aq) + H3O+(aq)

Ka =

7. pKa is the negative logarithm to base 10 of the acid dissociation constant (in mol dm-3) of a weak acid in an aqueous solution, pKa = -lg(Ka). 8. pKb is the negative logarithm to base 10 of the acid dissociation constant (in mol dm-3) of a weak base in an aqueous solution, pKb = -lg(Kb). Reaction Kinetics 1. The rate of a reaction is a positive quantity that can be defined as the change in concentration of a particular reactant or product per unit time, i.e. . 2. The rate constant is the constant of proportionality in the rate equation of the reaction. It is a constant at a given temperature, and it changes with varying temperature. 3. The rate equation for a reaction is a mathematical expression that shows the exact dependence of the reaction rate on the concentrations of all the reactants and can only be obtained by experiment. 4. The order of reaction with respect to a given reactant is the power to which the concentration of the reactant is raised in the rate equation. 5. The overall order of a reaction is the sum of all the powers of the concentration terms in the rate equation. 6. The half-life of a reaction is the time taken for the concentration of a reactant to fall to half its initial concentration i.e.

.

7. An autocatalytic reaction is one in which a product of that reaction increases the rate of reaction initially, and that product is an autocatalyst. The rate of reaction will decrease eventually because of the decrease in concentrations of the reactants/reactants are used up. 8. The rate-determining step is the slowest step in the mechanism of any reaction and it determines the rate of the overall reaction. 9. A catalyst provides an alternative pathway of lower activation energy for the reaction to proceed. There is an increase in the proportion of reactant molecules with kinetic energy greater than or equal to the activation energy so there is an increase in the frequency of effective collisions, resulting in an increase in the rate of reaction. 10. Activation energy (Ea) is the minimum energy that must be input to a chemical system to cause a chemical reaction between reactants. (Draw energy profile diagram to illustrate)

Electrochemistry 1. Standard electrode potential is the potential associated with a given half-cell when all components are in their standard states (temperature of 298K, pressure of any gas at 1atm and concentration of any ions at 1 mol dm-3), measured at relative to the Standard Hydrogen Electrode (S.H.E.).

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List of H2 Chemistry Definitions for ‘A’ Levels Compiled by: Seetoh Zhiwei of 14S03Q

Organic Chemistry 1. Primary structure of a protein refers to the number and sequence of amino acids in a single polypeptide chain, which is maintained by peptide bonds. 2. Secondary structure of a protein refers to the spatial arrangement in which segments of the polypeptide backbone orientate into a regular pattern of folding and coiling, held by hydrogen bonds between the C=O group of a peptide bond and N-H group of another peptide bond. 3. Tertiary structure of a protein refers to the unique 3D conformation due to the folding of the secondary structural elements together with the spatial disposition of the side chains to form a functional protein. This structure is maintained by side chain interactions including ionic bonds, hydrogen bonds, van der Waals’ interactions/dispersion forces and disulfide (covalent) bonds. 4. Quaternary structure of a protein refers to the spatial association of 2 or more polypeptide chains (i.e. subunits) together to form a functional protein. This structure is maintained by side chain interactions including ionic bonds, hydrogen bonds, van der Waals’ forces/dispersion forces and disulfide bonds. 5. Thermal cracking is a process in which long-chain hydrocarbons are broken into smaller alkanes and alkenes (and hydrogen gas). 6. A hydrocarbon is an organic compound containing only carbon and hydrogen atoms. 7. Denaturation is a process where the secondary, tertiary and quaternary structures of proteins are broken down, resulting in the protein unfolding to give a randomly coiled polypeptide while the amino acid sequence (primary structure) remains unaffected. Denaturation is caused by mechanical agitation, heat, heavy metal ions, salt/detergents, alcohol, acids and alkalis. (Note: denaturation hydrolysis) 8. A zwitterion is a dipolar ion that contains both a cationic group and an anionic group that carries a total net charge of zero and is thus electrically neutral. It is formed when the amino acid undergoes internal acid-base reaction with the basic amino group accepting a proton from the carboxylic acid group forming a zwitterion. 9. An electrophile is an atom or group of atoms that reacts with electron-rich centres in other molecules. It possesses either a positive charge or an empty orbital in their valence shell, or a polar bond producing an atom with a partial positive charge, δ+. 10. A nucleophile is a molecule or anion with at least one lone pair of electrons. It reacts with electron-deficient centres in other molecules. 11. A radical is an atom or group of atoms that has an unpaired electron, i.e. it has an odd number of electrons. Radicals are highly reactive and are intermediates of many of the reactions of alkanes (free-radical substitution). 12. Homolytic fission is a process whereby the two shared electrons in a covalent bond are split equally between 2 atoms during bond breaking, producing radicals. 13. Heterolytic fission is a process whereby the two shared electrons in a covalent bond are split unequally between 2 atoms during bond breaking, producing a positively charged ion and a negatively charged ion. It is highly encouraged to study the following organic compounds/reactions (though not required in H2 syllabus):  Grignard reagent  Claisen Condensation  Hydride and Alkyl shifts Page | 6

List of H2 Chemistry Definitions for ‘A’ Levels Compiled by: Seetoh Zhiwei of 14S03Q     

Dehydration of Alcohols (Elimination mechanism) Cyclization of compounds (Intramolecular Nucleophilic Substitution) Acid Anhydride Formation Williamson Synthesis (Ether formation – nucleophilic substitution) Be familiar with ‘electron pushing’!

Inorganic Chemistry 1. An amphoteric compound is one that can react with both acids and bases. 2. Transition elements are d-block elements which form at least 1 stable ion which has a partially-filled d-subshell/orbital. 3. A catalyst is a substance which increases the rate of reaction without itself undergoing any permanent change and is regenerated at the end of the reaction. It provides an alternative pathway with lower activation energy for the reaction. 4. A complex ion contains a central metal ion linked to one or more surrounding ions/molecules/ligands by dative covalent/coordinate bonds. 5. A complex compound, such as Ni(CO)4, contains a central metal atom linked to one or more surrounding ligands by dative covalent bonds. 6. A ligand is an ion/molecule which contains at least one atom bearing a lone pair of electrons which can be donated to an unoccupied, low-lying orbital of a central metal atom/ion, forming a dative covalent bond. 7. Coordination number of a central metal ion/atom in a complex/compound is the total number of coordinate bonds that it formed with the ligands. 8. Monodentate ligands form only one coordinate bond with a central metal ion/atom. 9. Bidentate ligands form two coordinate bonds with a central metal ion/atom simultaneously. 10. Disproportionation is a redox reaction in which an element of a molecule/atom/ion is simultaneously oxidized and reduced. 11. Comproportionation is a redox reaction in which two reactants, each containing the same element with different oxidation states, will react to form a product in which the element involved reaches the same oxidation number. 12. A homogeneous catalyst is one which is in the same physical state as the reactants. 13. A heterogeneous catalyst is one which is in a different physical state from the reactants.

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