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H1 Chemistry (Syllabus 8873)
Core Idea 1: Matter
Topic 1: Atomic Structure
The nucleus of the atom: neutrons and protons, isotopes, proton and nucleon numbers ● Protons, neutrons and electrons in terms of their relative charges and relative masses: ○ A proton has charge of +1 and a mass of 1 relative to the proton ○ A neutron has a charge of 0 and a mass of 1 relative to the proton ○ An electron has a charge of −1 and a mass of 0.0005 relative to the proton ● Deduction of the behaviour of beams of protons, neutrons and electrons in an electric field: ○ Protons, being positively charged, are deflected in the direction of the electric field ○ Neutrons, being neutral, are not affected by the presence of an electric field ○ Electrons, being negatively charged, are deflected in the opposite direction of the electric field ○ The extent of deflection of a particle is proportional to its |charge| to mass ratio ○ Hence, electrons, having much less mass than protons, are deflected to a greater extent ● Distribution of mass and charges within an atom: ○ The nucleus of protons and neutrons is positively charged and contains most of the mass of an atom ○ The electronic orbitals surrounding the nucleus are negatively charged and has negligible mass ● Deduction of the numbers of protons, neutrons and electrons present in both atoms and ions given proton and nucleon numbers (and charge): ○ The number of protons is equal to the proton (atomic) number of the atom or ion ○ The number of neutrons is equal to the difference between the nucleon number and the proton number ○ In a neutral atom, the number of electrons is equal to the number of protons such that the net charge of the atom is zero ○ In an ion, the number of electrons is different from the number of protons such that the difference is equal to the relative charge of the ion ○ Hence, the number of electrons in an ion is equal to the difference between the proton number of the element and the relative charge of the ion ● Contribution of protons and neutrons to atomic nuclei and isotopes: ○ The number of protons in the nucleus is known as the proton number
○ The total number of protons and neutrons in the nucleus is known as the nucleon number ○ Isotopes can be distinguished on the basis of different numbers of neutrons present if they have the same number of protons
Electrons: electronic energy levels, ionisation energies, atomic orbitals, extranuclear structure ● Number and relative energies of the s, p and d orbitals for the principal quantum numbers 1, 2 and 3 and also the 4s and 4p orbitals: ○ The 1s, 2s and 3s subshells have principal quantum numbers of 1, 2 and 3 respectively and contain 1 orbital each ○ The 2p and 3p subshells have principal quantum numbers of 2 and 3 respectively and contain 3 orbitals each ○ The 3d subshell has a principal quantum number of 3 and contains 5 orbitals ○ The 4s and 4p subshells contain 1 and 3 orbitals respectively ○ Orbitals with lower principal quantum numbers generally have lower relative energies ○ Among orbitals with the same principal quantum number, the s orbitals have the lowest energy, followed by the p and subsequently d orbitals ○ Orbitals within the same subshell have the same relative energy ○ As principal quantum number increases, the energy levels come closer to each other and overlapping of subshells may occur – the 4s orbitals have lower relative energies than the 3d orbitals ● Shapes of s and p orbitals: ○ The s orbital is spherical in shape and non-directional ○ The p x , p y and p z orbitals are dumbbell-shaped and lie along the x -, y - and z -axes respectively ● When stating the electronic configuration of atoms and ions given the proton number (and charge), recognise that any added electron occupies the vacant orbital with the lowest energy first and electrons are removed from the orbital with the highest energy first. ● Factors influencing the ionisation energies of elements (see also Topic 5): ○ Nuclear charge – As the number of protons increases, the nuclear charge of an atom increases and the force exerted by the nucleus on the electrons increases ○ Shielding effect – As the number of electron shells increases, the shielding effect of the inner shells increases and the electrons in the inner shells cancel more of the attractive force the nucleus exerts on the outermost electrons ○ The effective nuclear charge is the nuclear charge with the shielding effect taken into account ○ A higher effective nuclear charge means that the attraction between the nucleus and the outermost electrons increases ○ More energy is required to remove the outermost electrons, resulting in higher ionisation energies, when the effective nuclear charge is high ● Deduction of the electronic configurations of elements from successive ionisation energy data:
hydrogen chloride carbon dioxide methane ethene ○ Co-ordinate (dative covalent) bonds are formed when both the electrons in the covalent bond are provided by only one of the bonded atoms, such as in the formation of the ammonium ion and in the Al 2 Cl 6 molecule
ammonium ion Al 2 Cl 6 ○ The ammonium ion is formed when the lone pair on the N atom of ammonia, a Brønsted-Lowry base, forms a co-ordinate bond with a proton, the N atom providing both bonding electrons ● Covalent bonding in terms of orbital overlap (limited to s and p orbitals only) (see also Topic 9.2): ○ Covalent bonds involve the overlap of atomic orbitals to form molecular orbitals, and consist of σ and/or π bonds ○ A σ bond is formed when atomic orbitals overlap head-on and the shared pair of electrons occupy the space between the nuclei ○ A π bond is formed when atomic orbitals overlap side-on and the shared pair of electrons occupy the space above and below the nuclear axis ○ s orbitals can only form σ bonds while two p orbitals can form both σ and π bonds ○ π bonds are weaker than σ bonds and can only form when the atoms already possess a σ bond ○ Single bonds consist of 1 σ bond, double bonds consist of 1 σ bond and 1 π bond and triple bonds consist of 1 σ bond and 2 π bonds
Shapes of simple molecules and bond angles ● Explanation of the shapes of and bond angles in molecules by using the Valence Shell Electron Pair Repulsion theory: ○ The Valence Shell Electron Pair Repulsion theory states that electron pairs in the valence shell of a central atom will arrange themselves in space as far as possible from each other to minimise mutual repulsion ○ The repulsion between lone pairs is the greatest, followed by that between a lone pair and a bond pair, and lastly that between bond pairs ○ Molecules such as BF 3 have a trigonal planar shape and a bond angle of 120° as they have 3 bond pairs and 0 lone pairs surrounding the central atom
○ Molecules such as CO 2 have a linear shape and a bond angle of 180° as they have 2 bond pairs and 0 lone pairs surrounding the central atom ○ Molecules such as CH 4 have a tetrahedral shape and a bond angle of 109° as they have 4 bond pairs and 0 lone pairs surrounding the central atom ○ Molecules such as NH 3 have a trigonal pyramidal shape and a bond angle of 107° as they have 3 bond pairs and 1 lone pair surrounding the central atom ○ Molecules such as H 2 O have a bent shape and a bond angle of 105° as they have 2 bond pairs and 2 lone pairs surrounding the central atom ○ Molecules such as SF 6 have an octahedral shape and a bond angle of 90° as they have 6 bond pairs and 0 lone pairs surrounding the central atom ● By determining the number of bond pairs and lone pairs present on the central atom, the Valence Shell Electron Pair Repulsion theory can be used to predict the shapes of and bond angles in molecules analogous to BF 3 , CO 2 , CH 4 , NH 3 , H 2 O and SF 6.
Bond polarites and polarity of molecules ● Explanation and deduction of bond polarity using the concept of electronegativity: ○ The electronegativity of an atom is a measure of its ability to attract the shared pair of electrons in a covalent bond towards itself ○ When atoms of similar electronegativity form a covalent bond, the bonding electrons are shared equally between the atoms, resulting in the bond being nonpolar ○ When atoms of different electronegativities form a covalent bond, the more electronegative atom attracts the bonding electrons more strongly, forming a dipole, resulting in the bond being polar ○ The more electronegative atom in a covalent bond acquires a partial negative charge while the less electronegative atom acquires a partial positive charge ● Deduction of the polarity of a molecule using bond polarity and its molecular shape (analogous to BF 3 , CO 2 , CH 4 , NH 3 , H 2 O or SF 6 ): ○ If there are no polar bonds in the molecule, the molecule is definitely nonpolar ○ If there are polar bonds in the molecule but the dipoles cancel each other out, the molecule is nonpolar ○ If there are polar bonds in the molecule and the dipoles do not cancel each other out, the polarity of the molecule is the vector sum of the dipole moments in each covalent bond
Intermolecular forces, including hydrogen bonding ● Intermolecular forces of attraction and hydrogen bonding (electrostatic in nature): ○ Intermolecular forces are weak electrostatic forces of attraction between simple covalent molecules resulting from interactions between permanent or induced dipoles ○ Polar molecules such as liquid and gaseous CHCl 3 have net dipole moments with permanent dipole-permanent dipole interactions between oppositely charged ends of neighbouring molecules ○ In any molecule, including Br 2 and the noble gases, an asymmetrical charge distribution at a particular instant could result in temporary dipoles that induce
○ Bond energy is the energy required to break one mole of covalent bonds between two atoms in the gaseous state under standard conditions ○ Bond length is the internuclear distance between the two bonding atoms ● Comparison of the reactivities of covalent bonds in terms of bond energy, bond length and bond polarity: ○ Covalent bonds with larger bond energies are less reactive as they require more energy to break during chemical reactions ○ Covalent bonds with shorter bond lengths have a greater degree of orbital overlap and hence a higher bond energy, making them less reactive ○ Polar covalent bonds experience extra attraction between the partial charges, giving them higher bond energies and hence lower reactivities than nonpolar covalent bonds ○ Bond length is a more important factor than bond polarity in determining the bond energy and reactivity of the covalent bond
Lattice structure of solids ● Lattice structure of a crystalline solid: ○ An ionic solid, such as sodium chloride or magnesium oxide, has a giant ionic structure consisting of a lattice of cations and anions held together by ionic bonds ○ A simple molecular solid, such as iodine, has a simple molecular structure with intermolecular forces of attraction between molecules ○ A giant molecular solid, such as graphite or diamond, has a giant molecular structure with atoms held together by covalent bonds ○ A hydrogen-bonded solid, such as ice, has a simple molecular structure with hydrogen bonds between molecules ○ A metallic solid, such as copper, has a giant metallic structure consisting of a lattice of cations and a sea of delocalised valence electrons held together by metallic bonds
Bonding and physical properties ● Effect of different types of structure and bonding on the physical properties of substances: ○ More thermal energy is required to overcome stronger chemical bonds between particles – hence, substances with giant structures have higher melting and boiling points than substances with simple structures, and giant covalent substances have higher melting and boiling points than ionic or metallic substances ○ Mobile charge carriers, in the form of mobile ions or free electrons, are required for a substance to conduct electricity – hence, metals conduct electricity in the solid state, ionic compounds do so only when molten or in aqueous solution, and most covalent substances do not conduct electricity ○ For a substance to be soluble, the energy released during interaction between solute and solvent particles must be enough to overcome the forces between solute particles and between solvent particles
○ The effect of the structure and bonding of a substance on its physical properties can be predicted by considering the behaviour of individual particles on the microscopic level and evaluating its implications on the macroscopic behaviour of the substance ● The type of structure and bonding present in a substance can be inferred from given information by determining the type of structure and bonding which will best correspond to the given physical or chemical properties.
Topic 3: Theories of Acids and Bases
Arrhenius and Brønsted-Lowry theories of acids and bases ● The Arrhenius theory of acids and bases states that an acid produces hydrogen ions in water while a base reacts with an acid to give a salt and water only, and that a base is a metal oxide or hydroxide: HA → H+^ + A−^ in water and 2HA + MO → MA 2 + H 2 O or HA + MOH → MA + H 2 O, where M is a metal – this theory can be applied to describe whether isolated substances are acids, bases or neither. ● Brønsted-Lowry theory of acids and bases: ○ The Brønsted-Lowry theory of acids and bases states that an acid is a proton donor and that a base is a proton acceptor: HA → H+^ + A−^ and B + H+^ → BH+ ○ When an acid HA loses a proton, a base A−^ is formed, and when a base B gains a proton, an acid BH+^ is formed – hence, A−^ is known as the conjugate base of HA and BH+^ is known as the conjugate acid of B ○ The more readily an acid (base) donates (accepts) a proton, the less readily its conjugate base (acid) accepts (donates) a proton – hence, the stronger an acid (base), the weaker its conjugate base (acid) ○ This theory can be applied to describe acids and bases in reactions involving the transfer of protons between reactants
Acid dissociation constants, K a; base dissociation constants, K b; the ionic product of water, K w ● Differences in behaviour between strong and weak acids and bases in terms of the extent of dissociation: ○ Strong acids and bases completely dissociate in aqueous solution to produce H+^ and OH−^ ions respectively ○ Weak acids and bases are partially dissociated in aqueous solution and an equilibrium is achieved ○ Weak acids and bases remain mostly undissociated at equilibrium ● Explanation of the terms pH; K a; K b; K w: ○ pH is the negative logarithm to the base ten of the hydrogen ion concentration, i.e. pH = −lg [H+] ○ K a is the equilibrium constant for the dissociation of a weak acid HA in aqueous solution, i.e. K a = [H+][A−] / [HA] ○ K b is the equilibrium constant for the dissociation of a weak base B in aqueous solution, i.e. K b = [BH+][OH−] / [B]
○ A buffer solution is at maximum buffering capacity when [weak acid] = [salt] for an acidic buffer, or [weak base] = [salt] for an alkaline buffer ○ Buffer solutions are used to prevent large changes in pH from occurring where such changes affect the function of components in the system ○ An example is the role of H 2 CO 3 / HCO 3 −^ in controlling pH in blood, which is required to keep enzymes functioning efficiently ○ H 2 CO 3 is a weak acid which dissociates partially in solution to form its conjugate base, HCO 3 −: H 2 CO 3 ⇌ H+^ + HCO 3 − ○ An increase in H+^ ions is removed by the large reservoir of HCO 3 −^ by forming H 2 CO 3 : H+^ + HCO 3 −^ → H 2 CO 3 ○ An increase in OH−^ ions is removed by the large reservoir of H 2 CO 3 by forming HCO 3 −: OH−^ + H 2 CO 3 → HCO 3 −^ + H 2 O
Topic 4: The Periodic Table
Periodicity of atomic and physical properties of the elements: variation with proton number across the third period (sodium to chlorine) and down Group 17 of electronic configuration, atomic radius and ionic radius, ionisation energy, electronegativity, melting point and electrical conductivity For elements in the third period (sodium to chlorine), and in Group 17 (chlorine to iodine): ● Variation in the electronic configurations across a Period and down a Group: ○ Elements in the same Period have the same number of quantum shells (3 for Period 3) with the number of valence electrons increasing across the Period ○ Elements in the same Group have the same outermost electronic configuration ( n s^2 n p^5 for Group 17, where n is the number of quantum shells for the element) with the number of quantum shells increasing down the Group ● Trends and variations in atomic radius, ionic radius, first ionisation energy and electronegativity across a Period and down a Group: ○ Across a Period (e.g. sodium to chlorine), atomic radius decreases, ionic radius decreases for ions in the same isoelectronic series, first ionisation energy generally increases and electronegativity increases ○ As the number of protons increases across a Period or isoelectronic series, nuclear charge increases ○ As successive electrons are added to the same outermost shell, the shielding effect remains approximately the same across a Period; while members of an isoelectronic series have the same shielding effect ○ Therefore, the effective nuclear charge increases across a Period or isoelectronic series ○ The outermost shell is attracted more closely towards the nucleus and more energy is required to remove the outermost electron ○ Atoms become more able to attract the shared pair of electrons in a covalent bond towards themselves ○ Down a Group (e.g. chlorine to iodine), atomic radius increases, ionic radius increases, first ionisation energy generally decreases and electronegativity decreases
○ As the number of protons increases down a Group, nuclear charge increases ○ As the number of electronic shells increases, the distance of the outermost shell from the nucleus increases ○ The attraction between the nucleus and the outermost electron decreases and less energy is required to remove it ○ Atoms become less able to attract the shared pair of electrons in a covalent bond towards themselves ● Interpretation of the variation in melting point and in electrical conductivity across a Period in terms of structure and bonding in the elements: ○ Elements on the left side of a Period (e.g. sodium to aluminium) have high melting points and high electrical conductivity as they exist as giant metallic structures ○ A large amount of energy is required to overcome the strong metallic bonds between the metal cations and the sea of delocalised valence electrons ○ The delocalised valence electrons are able to act as mobile charge carriers to conduct electricity ○ As the number of delocalised valence electrons increases across the Period, the amount of charge carriers and the strength of the metallic bonds increases, resulting in an increase in the electrical conductivity and melting point of the metallic elements respectively ○ Elements in the middle of the Period (e.g. silicon) have high melting points as they exist as giant molecular structures ○ A large amount of energy is required to overcome the strong covalent bonds between atoms ○ Elements on the right side of the Period (e.g. phosphorus to chlorine) have low melting points and no electrical conductivity as they have simple molecular structures ○ Only a small amount of energy is required to overcome the weak instantaneous dipole-induced dipole interactions between molecules ○ Their valence electrons are localised in covalent bonds and they have no mobile charge carriers to conduct electricity ● Trend in volatility of Group 17 elements: ○ The volatility of Group 17 elements decreases down the Group as their melting and boiling points increase ○ Down the Group, the number of electrons and hence the size of the electron cloud increases, increasing the polarisability of the electron cloud ○ More energy is required to overcome the stronger instantaneous dipole-induced dipole attractions between molecules
Periodicity of chemical properties of the elements in the third period: variation in oxidation number and bonding of the oxides (sodium to sulfur only) and of the chlorides (sodium to phosphorus only), reactions of these oxides and chlorides with water and acid/base behaviour of these oxides and the corresponding hydroxides For elements in the third period (sodium to chlorine): ● Variation in the highest oxidation number of the elements in oxides and chlorides, variation in bonding in oxides and chlorides, reactions of the oxides and chlorides
○ Mg2+^ and Al3+^ have high charge densities which enables them to polarise and break the O―H bond in H 2 O to release H+^ ions forming an acidic solution, while Na+^ has a low charge density and is unable to do so ○ The acid formed from the reaction of MgCl 2 with water is weaker than that for AlCl 3 as Mg2+^ has a lower charge density than Al3+ ○ SiCl 4 and PCl 5 undergo irreversible hydrolysis to give strongly acidic solutions (pH ≈ 1) and white fumes of hydrogen chloride due to their acidic nature: SiCl 4 (l) + 2H 2 O (l) → SiO 2 (s) + 4HCl (g) and PCl 5 (s) + 4H 2 O (l) → H 3 PO 4 (aq) + 5HCl (g) ○ Oxides and chlorides that react with water to give highly acidic solutions and have low melting and boiling points tend to have simple covalent structures with weak intermolecular forces between molecules ○ Oxides that react with acids to form a salt and water, or chlorides that undergo hydration in water, that have high melting and boiling points tend to have giant structures with ionic bonding between cations and anions ○ Oxides and chlorides that conduct electricity in aqueous solution tend to have giant structures with ionic bonding between cations and anions
Periodicity of chemical properties of the elements down the group (Group 1 and Group 17): as reducing agents (Group 1) and oxidising agents (Group 17), and thermal stability of Group 17 hydrides For elements in Group 1 (lithium to caesium) and Group 17 (chlorine to iodine): ● Relative reactivity of elements of Group 2 as reducing agents and Group 17 as oxidising agents: ○ Group 1 elements are strong reducing agents with reducing strength increasing down the Group ○ Group 1 elements react with water with increasing reactivity down the Group
- lithium reacts slowly with water while caesium reacts violently ○ Down Group 1, as the number of electronic shells increases, the distance of the outermost electron from the nucleus increases, causing the attraction between the nucleus and the outermost electron to decrease ○ Less energy is required to remove the outermost electron and the ease of loss of electrons increases ○ Group 17 elements are oxidising agents with oxidising strength decreasing down the Group ○ Down Group 17, as the number of electronic shells increases, the distance between the outermost shell increases, causing the attraction between any added electron and the nucleus to decrease and the ease of gain of electrons to decrease ● Trend in thermal stability Group 17 hydrides: ○ Down the Group, the thermal stability of Group 17 hydrides decreases and they decompose less easily into their elemental form and hydrogen gas ○ Down the Group, the bond length of the H―X bond increases, where X represents the Group 17 element, causing the overlap of atomic orbitals to become less effective
○ The bond energy of the H―X bond decreases and less energy is required to break them In addition: ● The characteristic properties of an element in a given Group can be predicted by using knowledge of chemical periodicity of these properties across Periods and down Groups. ● The nature, possible position in the Periodic Table and identity of unknown elements can be deduced from given information on physical and chemical properties by determining how the properties of these elements fit into known trends of these properties across Periods and down Groups.
Core Idea 3: Transformation
Topic 5: The Mole Concept and Stoichiometry
Relative masses of atoms and molecules ● Definition of the terms relative atomic, isotopic, molecular and formula mass based on the 12 C scale: ○ The relative atomic mass of an element is the weighted average mass of an atom of the element compared to 1/12 the mass of a 12 C atom ○ The relative isotopic mass of an isotope is the mass of an atom of the isotope compared to 1/12 the mass of a 12 C atom ○ The relative molecular mass of an element or compound is the weighted average mass of a molecule of the element or compound compared to 1/ the mass of a 12 C atom ○ The relative formula mass of a compound is the weighted average mass of a formula unit of the compound compared to 1/12 the mass of a 12 C atom
The mole, the Avogadro constant ● One mole is defined as the amount of substance that contains as many particles as there are carbon atoms in 12 grams of 12 C – this number is known as the Avogadro constant which has the value 6.02 × 10^23 mol−1. ● Given the relative abundances of the isotopes of an element, its relative atomic mass can be determined by taking the weighted average of the nucleon numbers of the isotopes.
The calculation of empirical and molecular formulae ● Definition of the terms empirical and molecular formula: ○ The empirical formula of a compound is the simplest formula that shows the relative number of atoms of each element present in the compound ○ The molecular formula of a compound is the formula that shows the actual number of atoms of each element present in one molecule of the compound ● Calculation of empirical and molecular formulae using combustion data or composition by mass:
● Calculations involving reacting masses, volumes of gases and volumes and concentrations of solutions: ○ Calculations involving reacting masses can be performed by taking into account the stoichiometric ratio of each element in the balanced equation, as well as any excess reagents, and applying the relationship amount of substance = mass / molar mass , where molar mass can be calculated from chemical formulae, to find the required quantity ○ Calculations involving volumes of gases, such as in the burning of hydrocarbons, can be performed by recognising that one mole of any gas occupies the same volume when at the same temperature and pressure, such that amount of gas = volume of gas / molar gas volume ○ The mole concentration of a solution is given by amount of solute / volume of solution and its mass concentration is given by mass of solute / volume of solution ; hence, mole concentration = mass concentration / molar mass of solute ○ The amount of solute remains the same after dilution, hence initial concentration × initial volume = final concentration × final volume ○ In the case of titration, the stoichiometric ratio of the reactants n A / n B can be expressed as c A V A / c B V B where c and V represent mole concentration and volume respectively ● Stoichiometric relationships can be deduced from calculations involving reacting masses, volumes of gases and volumes and concentrations of solutions by determining the mole ratio of the substances in question.
Topic 6: Chemical Energetics: Thermochemistry
Enthalpy changes: Δ H , of formation; combustion; neutralisation; bond energy; lattice energy ● Most chemical reactions are accompanied by energy changes, principally in the form of heat usually associated with the breaking and forming of chemical bonds which absorb and release heat respectively; the reaction can be exothermic (Δ H negative) or endothermic (Δ H positive), where Δ H represents the change in heat content, or enthalpy, of the reaction. ● Construction and interpretation of an energy profile diagram (see also Topic 7):
○ The enthalpy change of the reaction is equal to the potential energy of the products minus the potential energy of the reactants
○ The activation energy is equal to the peak potential energy minus the potential energy of the reactants ● The terms enthalpy change of reaction and standard conditions, bond energy (see also Topic 2) and lattice energy: ○ The enthalpy change of reaction is the energy change when molar quantities of reactants as stated in the thermochemical equation react together ○ The standard enthalpy change of reaction is the enthalpy change of reaction when the reaction is performed under standard conditions ○ Standard conditions refers to a temperature of 298 K, a pressure of 10^5 Pa, a concentration of 1.0 mol dm−3^ (for solutions) and the use of the most stable allotropic form at these conditions ○ The enthalpy change of formation is the energy change when one mole of a compound is formed from its constituent elements ○ The enthalpy change of combustion is the energy evolved when one mole of a compound is completely burnt in excess oxygen ○ The enthalpy change of neutralisation is the energy evolved when one mole of water is formed during neutralisation of an acid and an alkali ○ Bond energy is the energy needed to break one mole of covalent bonds between two atoms in the gaseous state ○ Since bond breaking is an endothermic process, bond energy Δ H is always positive ○ Lattice energy is the energy evolved when one mole of an ionic compound is formed from its constituent gaseous ions ○ Since the formation if a solid lattice from gaseous ions is an exothermic process, lattice energy Δ H is always negative ● Enthalpy changes can be calculated from appropriate experimental results by applying the definition of the required quantity and the relationship heat change = mc Δ T , where m , c and Δ T refer to the mass, specific heat capacity and change in temperature of the substance respectively. ● The numerical magnitude of a lattice energy is directly proportional to the product of the ionic charges and inversely proportional to the sum of the ionic radii.
Hess’ Law ● Application of Hess’ Law to construct simple energy cycles and carry out calculations involving such cycles and relevant energy terms: ○ Hess’ Law states that the enthalpy change of a chemical reaction is the same regardless of the route the reaction takes, provided the reactants, products, initial and final conditions are the same ○ Enthalpy changes that cannot be found by direct experiment, such as the enthalpy change of formation, can be determined by applying Hess’ Law to construct simple energy cycles involving measured quantities, such as the enthalpy change of combustion, and equating the total enthalpy change of the different reaction routes ○ Calculations involving average bond energies can be performed by recognising that the enthalpy change of a reaction is equal to the difference between the energy required to break bonds and that to form bonds
○ The activation energy is the minimum energy the colliding molecules must possess before a collision results in a reaction ○ In any system, the particles present have a wide range of kinetic energies which can be modelled by the Boltzmann distribution ○ Only those particles with a kinetic energy greater than the activation energy and which collide in the correct orientation are able to react during collision
Effect of concentration, temperature, and catalysts on reaction rate ● Effect of temperature change on a rate constant (and hence, on the rate) of a reaction: ○ A rise in temperature will lead to an increase in the average kinetic energy of reactant molecules ○ The peak of the Boltzmann distribution is shifted to the right and lowered, resulting in more molecules possessing an energy greater than the activation energy ○ The frequency of effective collisions increases, resulting in an increase in the rate constant and hence the rate of reaction ● Effect of a catalyst on a rate constant: ○ In the presence of a catalyst, a reaction follows a different pathway with a lower activation energy, resulting in a larger rate constant ○ In the presence of a catalyst, more molecules will possess an energy greater than the lowered activation energy in accordance with the Boltzmann distribution ○ The frequency of effective collisions increases, resulting in an increase in the rate constant and hence the rate of reaction
Heterogeneous catalysts ● The mode of action of heterogeneous catalysis: ○ During heterogeneous catalysis, reactant molecules are diffused and adsorbed on an active site of the catalyst, which is in a different phase, forming weak temporary bonds between reactant molecules and the catalyst surface ○ The process of adsorption increases the concentration of reactants at the catalyst surface, weakens the bonds in reactant molecules and allows them to be orientated in the correct positions for reaction, thus increasing the number of effective collisions ○ Adjacent reactant molecules react to form products before being desorbed and diffusing away from the catalyst surface, regenerating the active site on the catalyst ○ An example of heterogeneous catalysis is the catalytic removal of oxides of nitrogen in the exhaust gases from car engines by palladium and platinum catalysts, which converts nitrogen monoxide to harmless nitrogen and oxygen gas: 2NO → N 2 + O 2 (see also Topic 9.1)
Enzymes as biological catalysts ● Enzymes:
○ Enzymes are homogeneous biological catalysts that catalyse biochemical reactions ○ Enzymes may have highly specific activity – they only catalyse one or a group of closely related reactions, and only work well over a narrow pH and temperature range
Topic 8: Chemical Equilibria
Chemical equilibria: reversible reactions; dynamic equilibrium – factors affecting chemical equilibria, equilibrium constants and Haber process ● What is meant by a reversible reaction and dynamic equilibrium: ○ A reversible reaction is a reaction that can proceed in both the forward and reverse directions, with the reaction mixture containing quantities of all reaction species in the system ○ In a reversible system, dynamic equilibrium occurs when the rates of the forward and reverse reactions are the same, and there is no net change in the concentration of reactants and products ● Le Chatelier’s Principle and its application to deduce qualitatively (from appropriate information) the effects of changes in concentration, pressure or temperature, on a system at equilibrium: ○ Le Chatelier’s Principle states that when a stress is applied to a reversible system at equilibrium, the position of equilibrium will shift so as to minimise the stress ○ When the concentration of a reactant (product) is increased [decreased] in a system at dynamic equilibrium, it reacts to reduce [increase] the concentration of the reactant (product), causing the forward [reverse] [(forward)] (reverse) reaction to be favoured and the position of equilibrium to shift to the right [left] [(right)] (left) ○ When the pressure is increased (decreased) in a gaseous system at dynamic equilibrium, it reacts to decrease (increase) the increased (decreased) pressure by favouring the reaction that forms a smaller (greater) amount of gas, causing the position of equilibrium to shift to the side with less (more) moles of gas ○ When the temperature of a system at dynamic equilibrium is increased (decreased), the system reacts to reduce (regenerate) the added (removed) heat by favouring the endothermic (exothermic) reaction, causing the position of equilibrium to shift to the side with more (less) total enthalpy ● Deduction of whether changes in concentration, pressure or temperature or the presence of a catalyst affect the value of the equilibrium constant for a reaction: ○ When the concentration or pressure of a reactant or product in a reaction mixture is changed, the position of equilibrium will shift such that the change in the numerator of the equilibrium constant is equal to the change in the denominator ○ When the temperature of a reaction mixture is changed, the rates of both the forward and reverse reactions are changed, but since one reaction is
