Foundation Physical and Inorganic Chemistry

1. Foundation Physical and Inorganic Chemistry The energy of reactions

2. Contents • Energetics • Kinetics • Equilibria • Redox Reactions • Group VII, The Halogens

3. Energetics • Enthalpy Change (ΔH ) • Calorimetry • Hess’s Law • Calculation of Bond Enthalpies • Using Bond Enthalpies

4. Enthalpy change (ΔH ) • All chemical reaction involve a transfer of energy. Chemists call this energy change as the enthalpy change of the reaction. • Exothermic reactions have a negative enthalpy change, that is they transfer energy to their surroundings. • Endothermic reactions have a positive enthalpy change, that is they take in energy from their surroundings. • The standard enthalpy change of combustion is the enthalpy change when 1 mole of a compound is burnt completely in oxygen under standard conditions (298K and 100kPa), all reactants and products being in their standard state. • The standard enthalpy change of formation is the enthalpy change when 1 mole of a compound is formed from its elements under standard conditions (298K and 100kPa), all reactants and products being in their standard state.

5. Calorimetry • The enthalpy change can be calculated from the temperature change in a reaction using the equation: • q = mcΔT • q is the enthalpy change (J), m is the mass (g) c is the specific heat capacity J g-1K-1, ΔT is the temperature change in K. • Example: excess of magnesium was added to 100cm3 of 2M CuSO4 solution. The temperature increased by 45K. The specific heat capacity of water is 4.2 J g-1K-1 • Find q using the above equation • q = 100 x 4.2 x 45 =18810J • Find the amount in moles that reacted • 2X100/1000 =0.2 • Scale the enthalpy change for one mole of CuSO4 • 18810/0.2 = 94050J =94 kJ mol-1

6. Hess’s Law • Hess’s law states that, if a reaction can take place by more than one route and the initial and final conditions are the same for each route. • N2(g) + 3H2(g) 2NH3 (g) 2N(g) + 6H(g) • ΔHf = ΔH1 + ΔH2 route 1 ΔHf route 2 ΔH2 ΔH1

7. Calculation of Bond Enthalpies • Bond enthalpies are a measure of the energy required to break a bond. They are always negative. • The values quoted are usually average values, as the strength of a bond between two atoms is dependant on the molecule. • They are difficult to measure directly they are usually calculated from data enthalpy changes of combustion of several compounds.

8. route 1 ΔHf route 2 ΔH2 ΔH1 Using Bond Enthalpies • N2(g) + 3H2(g) 2NH3 (g) 2N(g) + 6H(g) • Bond enthalpies can be used to calculate the enthalpy change for a reaction. • ΔHf = ΔH1 + ΔH2 • ΔH1 = the enthalpy for breaking bonds • ΔH2 = the enthalpy for making bonds • For this reaction • ΔHf = 3xE(H–H) + E(NºN) – 6xE(N–H) • Where E(X–X) represents the bond enthalpy for XX

9. Kinetics • Collision Theory • Maxwell-Boltzmann Distribution

10. Collision Theory • Reactions can only occur when reacting particles collide with enough energy. This energy is called the activation energy. Most collisions do not result in reaction. • Increasing the temperature of the reaction increases the speed at which particles move. There are more collision and the collisions have a higher energy. This increases the number of effective collisions. Therefore the rate of reaction is increased • Increasing the concentration of reactants increase the number of particles available to react. More collisions occur. Therefore the rate of reaction is increased. Decreasing the size of particles increases the surface area. If there is a larger surface available for particles to collide with there will be more collisions. Therefore the rate of reaction is increased. • Catalyst provide a different route for the reaction to happen. This route has a lower activation energy. This means that more of the collisions will be effective. The rate of reaction is increased

11. Maxwell-Boltzmann Distribution • Below is a Maxwell-Boltzman distribution curve. It shows the distribution of molecular energies in a gas. • As the temperature increases the peak of the line moves lower and to the right. (From the blue to the red on the picture.) • The area under the graph represents the total number of particles and stays the same whatever the temperature. • Ea is the activation energy, only collisions occurring above this energy result in reaction. • Increasing the temperature increases the number of collisions occurring above the activation energy. • The change from blue to red is only an increase in temperature of 10°C, however the number of particles that have enough energy to react has increased significantly These particles do not have enough energy to react Number of particles with Energy E Ea These particle have enough Energy to react Energy E

12. Equilibria • The Dynamic Nature of Equilibria • Le Chatelier’s Principle. • The Haber Process

13. The Dynamic Nature of Equilibria • Many reactions are reversible. That is they can proceed in both directions at the same time. • An equilibrium can be established if a reaction occurs in a closed system • Reactions do not stop at equilibrium, but the concentration of products and reactants remains constant because the reaction is proceeding at the same rate in both directions. • This does not mean that there are equal amounts of products and reactant though.

14. Le Chatelier’s Principle • Le Chatelier’s principle states that: • The position of the equilibrium of a system changes to minimise the effect of any imposed change in conditions. • This means: • Increasing the temperature favours the endothermic reaction • Increasing the pressure in a gaseous reaction favours the side of the reaction which has the fewest number of gaseous molecules, because this will educe the pressure. • However adding a catalyst does not effect the position of the equilibrium.

15. The Haber Process • N2(g) + 3H2(g)Â 2NH3(g) ΔH=-92kJmol-1 • Increasing the temperature will move the equilibrium in favour of the reactants, in the endothermic direction. • Increasing the pressure will favour the products because this side has the fewer number of gaseous molecules. • Lowering the temperature will favour the formation of the products, but will mean that reaction will happen slowly. A compromise temperature is used. • Too high a pressure would have high energy costs and expensive thicker walled pipes and reactor so a compromise pressure is used. • A finely divided iron catalyst is used to speed up the reaction, and the unconverted reactants are recycledd

16. Redox Reactions • Oxidation and Reduction • Oxidation States • Redox Equations.

17. Oxidation and Reduction • The term redox is used for the simultaneous processes of oxidation and reduction. • Oxidation is the loss of electrons. • Reduction ids the gain of electrons. • Any reaction which involves reduction also involves oxidation. • An oxidising agent removes electrons, a reducing agent donates electrons.

18. Oxidation States • Oxidation states show how oxidised or reduced an element is within a compound or ion. • If the oxidation state increases the substance is oxidised • If the oxidation state decreases the substance is reduced. • Rules for oxidation numbers: • For atoms in element the oxidation state is always 0 • For atoms in simple ions the oxidation state is the same as the charge • In compounds the total of all the oxidation states is 0 • In polyatomic ions the total of all the oxidation states is the same as the charge on the ion

19. Redox Equations • Sometimes it is easier to show what is being oxidised and reduced by writing half equations. • E.g Ca +Cl2 CaCl2 • This can be written as: • Ca  Ca2+ + 2e- • and Cl2 + 2e-  2Cl- • Remember only one element loses or gains electrons. • The equations must balance for atoms and charge.

20. The Halogens • Trends • Key facts

21. Trends

22. Key Facts • All halogens form diatomic molecules • A halogen will displace a halogen which is below it from its salts because reactivity decreases down the group. ie Chlorine will displace bromine. • The reaction of silver nitrate with the halide is a test for halides: AgCl –white, AgBr –cream, and AgI –yellow. • Silver halides are used in photography as they darken in the presence of light. • Chlorine is used extensively for disinfection.

23. Summary • Energetics • All chemical changes involve the transfer of energy. The overall enthalpy transfer during a chemical change is independent of the route taken. • Kinetics • The rate of reaction is effect by factors that effect the overall collision rate. We can use the Maxwell-Boltzman distribution to show the change in distribution of molecular kinetic energy with increasing temperature. • Equilibria • Many reactions are reversible and in a closed system reach an equilibrium. The position of the equilibrium of a system changes to minimise the effect of any imposed change in conditions. The nature of equilibria means that sometimes compromises need to be made in industrial processes. • Redox Reactions • These are reactions which involve simultaneous oxidation and reduction. We use the concept of oxidation number to indicate how oxidised or reduced elements are. • Group VII, The Halogens • The halogens have many similar properties. They are all reactive and this reactivity decreases down the group. Chorine is used as a disinfectant.