A2 Chemistry Unit 5

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1. 1 Chemistry Unit 5 Chapter 12 Thermodynamics Chapter 12.1 Enthalpy Change Hesss Law The Hesss Law states that: It basically says that if we convert reactant A into product B, the aggregate enthalpy change will be exactly the same as whichever reaction route we have used. Definitions: Enthalpy change The measure of heat change at a constant pressure. It is abbreviated as H. Standard enthalpy change It refers to the enthalpy change under standard conditions and for specified amount of chemicals in moles. Standard conditions A term used to define standard enthalpy and free energy changes. It relates to a condition where temperature is 298 kelvins, a pressure of 1 bar (100kPa) and if solution is involved, a concentration of 1 mol dm-3. Standard states The stable physical state of an element or compound under standard conditions. For example, the standard state for water is a liquid; for oxygen it would be in gaseous state. Standard molar enthalpy of formation (Hf ) It is the enthalpy change when one mole of compound is formed from its constituent elements, under standard state, and all products and reactants in their standard state. The enthalpy change accompanying a chemical change, is independent of the route by which the chemical change occurs. 2. 2 Standard enthalpy of atomisation (Ha ) It is the enthalpy change which accompanies the formation of one mole of gaseous atoms from the elements in its standard states under standard conditions. First ionisation energy (1st IE) It is the standard enthalpy change when one mole of gaseous atoms is converted into a mole of gaseous ions each with a single positive charge. Second ionisation energy (2nd IE) It relates to the standard enthalpy change to remove an electron from a mole of singly-positively charged ions. First electron affinity (EA) The standard enthalpy change when one mole of gaseous atoms is converted to a mole gaseous ions with each a single negative charge. Second electron affinity (EA) The standard enthalpy change when a mole of electrons are added to a mole of gaseous ions, each with a single negative charge to form ions each with two negative charges. Lattice formation enthalpy It is the standard enthalpy change when a mole of solid ionic compound is formed from its gaseous ions. Lattice dissociation enthalpy The opposite process of lattice formation: it has the same value as the lattice formation enthalpy, but always carry a positive value. It is the enthalpy change needed to convert 1 mole of solid crystal into its scattered gaseous ions. Enthalpy of hydration The standard enthalpy change when water molecules surround one mole of gaseous ions. Enthalpy of solution The standard enthalpy change when one mole of solute dissolves completely in sufficient solvent to form a solution in which the molecules or ions are far enough apart not to interact with each other. Mean bond enthalpy The enthalpy change when one mole of gaseous molecules each breaks a covalent bond, to form two free radicals, averaged over a range of compounds. 3. 3 Chapter 12.2 Born-Haber cycle Lattice enthalpies cannot be measured directly. A way to find the lattice enthalpy is via Born-Haber cycle, involving enthalpy changes which can be measured. Example Below shows a Born-Haber cycle for sodium chloride. Note that exothermic reaction points downwards; and vice-versa, endothermic reaction points upwards. We start with the two thick lines: NaCl (s), and the elements in their standard states, Na (s) + 0.5 Cl (g). The whole cycle is to find the formation of sodium chloride via a series of many small enthalpy changes. The +107 is the atomisation enthalpy of sodium. We have to produce gaseous atoms so that we can use the next stage in the cycle. The +496 is the first ionisation energy of sodium. Remember that first ionisation energies go from gaseous atoms to gaseous singly charged positive ions. The +122 is the atomisation enthalpy of chlorine. Again, we have to produce gaseous atoms so that we can use the next stage in the cycle. 4. 4 The -349 is the first electron affinity of chlorine. Remember that first electron affinities go from gaseous atoms to gaseous singly charged negative ions. And finally, we have the positive and negative gaseous ions that we can convert into the solid sodium chloride using the lattice formation enthalpy. And using Hesss Law: -411 = +107 + 496 + 122 - 349 + LE LE = -411 - 107 - 496 - 122 + 349 LE = -787 kJ mol-1 5. 5 Trends in lattice enthalpies There are two factors which affect the lattice enthalpy: the charges on the ions and the ionic radii (which affects the distance between the ions). Charge of ions For example, both sodium chloride (NaCl) and magnesium oxide (MgO) has exactly the same ionic arrangement in the lattice. However the latter lattice enthalpy is much greater. This is because in the case of magnesium oxide, magnesium has 2+ charge while an oxygen has a 2- charge; whereas in sodium chloride, the attraction is only between 1+ and 1- ions. As a conclusion, increasing charge difference result in a greater lattice enthalpy. 6. 6 The radius of ions As we go down the Group 7 of the periodic table, we expect that the lattice enthalpies fall as negative ions are larger. This is because of the increased distance between the centres of the oppositely charged ions. 7. 7 Chapter 12.3 More enthalpy changes Enthalpy of solution: Ions are strongly attracted to water, due to the fact that water is a polar molecule. Ions can be dissolved when the lattice is broken up, which requires some energy lattice enthalpies. The constituent ions are then surrounded by cluster of solvent ions. Positive ions are surrounded by negative ends of the dipole water molecule, and negative ions surrounded by the positive ends of the dipole water molecule. This process is known as hydration. Size of enthalpy of hydration: The size of enthalpy of hydration is determined by the amount of attraction between the ions and the water molecules. Size of ions When the ions are smaller, there will be a stronger attraction. For example, smallest ions in each group, i.e. lithium in Group 1 and fluoride ion in group 7 has the highest hydration enthalpy. As ions get bigger, due to increased distance between the centre of ion and the water molecule, the enthalpy of hydration falls. Charge of ions The attraction is greater if the charge of the ion is greater. The hydration enthalpies of group 2 ions and much greater than those of group 1 ions (Mg2+ > Na+). Dissolving as a energy cycle: We can assume dissolving an ionic compound as a three-step process. Crystal lattice is broken up into separate gaseous ions lattice dissociation enthalpy. Hydrating the positive ions enthalpy of hydration. Hydrating the negative ions enthalpy of hydration. 8. 8 Lattice enthalpies and bonding: Purely ionic model of a lattice assumes that all ions are perfectly spherical, and have their charge evenly distributed around them. Often the theoretical value (which we worked out) agrees with the experimental value (which we found out via experiment). Sometimes there is a large discrepancy between the two values between the experimental and theoretical lattice formation enthalpy. An explanation for this would be due to the fact that the bond contains some covalent character, but the ionic bond nature still co-exist. This is called ion distortion. The positively charged ion, cation, slowly approaches the anion (negatively charged ion), and deform the electron cloud of the anion. It attracts the electron cloud towards the cation. How easily could a cation distort an anion depends on its size and charge. If an anion is huge in size, then the distance between the electrons and the nucleus increases. If the anion is rich in charge, too, then there are more electrons involved in the distortion. For example, S 2- ion is more easily distorted than the Cl- ion. If the cation is small is size, then the distance between the positively-charged nuclei and the negative ion would be reduced, thus increase the strength of distortion. With high positive charge, it increases the power in which the electron is attracted towards it. This type of positive ion has high concentration of positive charge. In any metallic groups of the periodic table, the member that forms the smallest cation has the biggest tendency to pull electrons from the anion, which form a high degree of covalent character. For example, lithium in group 1. In this scenario, the anion is said to be polarised. All ionic and covalent bonds can be transitional there is a continuous spectrum from ionic to purely covalent, depending on the strength of the ions. For example, Cs+F-, with large, singly charged positively; and a small singly negatively charged ion, has no degree of polarisation and is almost completely ionic. On the contrary, a bond formed of two identical atoms must be covalent. 9. 9 Chapter 12.4 Calculating enthalpy changes using mean bond enthalpies In any chemical reactions, energy is absorbed to break bonds, and released when bonds are formed. The overall difference between energy absorbed and released is the overall enthalpy change of reaction. Enthalpy change of reaction = Total energy absorbed Total energy released In order to find out the overall enthalpy change of reaction, we work out the total enthalpy required to break the bonds, and the total enthalpy released as the bonds are formed. The enthalpy change of reaction would be subtracting total energy released from sum of total energy absorbed. For example: Calculate the overall enthalpy change for the following reaction, CH4 + 2O2 CO2 + 2H2O Bonds broken: 4 C-H bonds broken: 4 x 413 2 O=O bonds broken: 2 x 498 Total: 2648 kJ mol-1 Bonds formed: 2 C=O bonds formed: 2 x 805 4 O-H bonds formed: 4 x 460 Total: 3450 kJ mol-1 Overall enthalpy change of reaction: 2648 3450 = -802 kJ mol-1. 10. 10 Hesss Law route: We can also use the Hesss Law approa