The study of heat released or required by chemical reactions THERMOCHEMISTRY The study of heat released or required by chemical reactions Fuel is burnt to produce energy - combustion (e.g. when fossil fuels are burnt) CH4(g) + 2O2(g) CO2(g) + 2H2O(l) + energy

Energy due to position (stored energy) What is Energy? Energy Kinetic energy (EK) Potential energy (EP) Energy due to motion Energy due to position (stored energy)

Total Energy = Kinetic Energy + Potential Energy E = EK + EP Kinetic energy & potential energy are interchangeable Ball thrown upwards slows & loses kinetic energy but gains potential energy The reverse happens as it falls back to the ground

i.e. it is the total energy of all the atoms and molecules in a sample Law of Conservation of Energy: the total energy of the universe is constant and can neither be created nor destroyed; it can only be transformed. The internal energy, U, of a sample is the sum of all the kinetic and potential energies of all the atoms and molecules in a sample i.e. it is the total energy of all the atoms and molecules in a sample

Systems & Surroundings In thermodynamics, the world is divided into a system and its surroundings A system is the part of the world we want to study (e.g. a reaction mixture in a flask) The surroundings consist of everything else outside the system SYSTEM CLOSED OPEN ISOLATED

OPEN SYSTEM: can exchange both matter and energy with the surroundings (e.g. open reaction flask, rocket engine) CLOSED SYSTEM: can exchange only energy with the surroundings (matter remains fixed) e.g. a sealed reaction flask ISOLATED SYSTEM: can exchange neither energy nor matter with its surroundings (e.g. a thermos flask)

e.g. kettle heats on a gas flame HEAT and WORK HEAT is the energy that transfers from one object to another when the two things are at different temperatures and in some kind of contact e.g. kettle heats on a gas flame cup of tea cools down (loses energy as heat) Thermal motion (random molecular motion) is increased by heat energy i.e. heat stimulates thermal motion

i.e. a system does work when it expands against an external pressure Work is the transfer of energy that takes place when an object is moved against an opposing force i.e. a system does work when it expands against an external pressure Car engine: petrol burns & produces gases which push out pistons in the engine and transfer energy to the wheels of car Work stimulates uniform motion Heat and work can be considered as energy in transit

UNITS OF ENERGY S.I. unit of energy is the joule (J) Heat and work ( energy in transit) also measured in joules 1 kJ (kilojoule) = 103 J Calorie (cal): 1 cal is the energy needed to raise the temperature of 1g of water by 1oC 1 cal = 4.184 J

INTERNAL ENERGY (U) Internal energy changes when energy enters or leaves a system U = Ufinal - Uinitial U change in the internal energy Heat and work are 2 equivalent ways of changing the internal energy of a system

+ = U = q (heat) + w (work) Change in internal energy Energy supplied to system as heat Energy supplied to system as work U = q (heat) + w (work) q w U U like reserves of a bank: bank accepts deposits or withdrawals in two currencies (q & w) but stores them as common fund, U.

First Law of Thermodynamics: the internal energy of an isolated system is constant Signs (+/-) will tell you if energy is entering or leaving a system + indicates energy enters a system - indicates energy leaves a system

In a system that can’t expand, no work is done (w = 0) U = q + w An important form of work is EXPANSION WORK i.e. the work done when a system changes size and pushes against an external force e.g. the work done by hot gases in an engine as they push back the pistons HEAT In a system that can’t expand, no work is done (w = 0) U = q + w when w = 0, U = q (at constant volume)

A change in internal energy can be identified with the heat supplied at constant volume ENTHALPY (H) (comes from Greek for “heat inside”) the change in internal energy is not equal to the heat supplied when the system is free to change its volume some of the energy can return to the surroundings as expansion work  U < q

The heat supplied is equal to the change in another thermodynamic property called enthalpy (H) i.e. H = q this relation is only valid at constant pressure As most reactions in chemistry take place at constant pressure we can say that: A change in enthalpy = heat supplied

EXOTHERMIC & ENDOTHERMIC REACTIONS Exothermic process: a change (e.g. a chemical reaction) that releases heat. A release of heat corresponds to a decrease in enthalpy Exothermic process: H < 0 (at constant pressure) Burning fossil fuels is an exothermic reaction

An input of heat corresponds to an increase in enthalpy Endothermic process: a change (e.g. a chemical reaction) that requires (or absorbs) heat. An input of heat corresponds to an increase in enthalpy Endothermic process: H > 0 (at constant pressure) Forming Na+ and Cl- ions from NaCl is an endothermic process Photosynthesis is an endothermic reaction (requires energy input from sun)

Measuring Heat reaction Exothermic reaction, heat given off & temperature of water rises reaction Endothermic reaction, heat taken in & temperature of water drops

How do we relate change in temp. to the energy transferred? Heat capacity (J/oC) = heat supplied (J) temperature (oC) Heat Capacity = heat required to raise temp. of an object by 1oC more heat is required to raise the temp. of a large sample of a substance by 1oC than is needed for a smaller sample

Specific Heat Capacity (Cs) Specific heat capacity is the quantity of energy required to change the temperature of a 1g sample of something by 1oC Specific Heat Capacity (Cs) Heat capacity Mass = J / oC / g J / oC g =

Vaporisation Melting Freezing Energy has to be supplied to a liquid to enable it to overcome forces that hold molecules together endothermic process (H positive) Melting Energy is supplied to a solid to enable it to vibrate more vigorously until molecules can move past each other and flow as a liquid endothermic process (H positive) Freezing Liquid releases energy and allows molecules to settle into a lower energy state and form a solid exothermic process (H negative) (we remove heat from water when making ice in freezer)

All chemical reactions either release or absorb heat Reaction Enthalpies All chemical reactions either release or absorb heat Exothermic reactions: Reactants products + energy as heat (H -ve) e.g. burning fossil fuels Endothermic reactions: Reactants + energy as heat products (H +ve) e.g. photosynthesis

Bond Strengths Lattice Enthalpy Bond strengths measured by bond enthalpy HB (+ve values) bond breaking requires energy (+ve H) bond making releases energy (-ve H) Lattice Enthalpy A measure of the attraction between ions (the enthalpy change when a solid is broken up into a gas of its ions) all lattice enthalpies are positive I.e. energy is required o break up solids

Enthalpy of hydration Hhyd the enthalpy change accompanying the hydration of gas-phase ions Na+ (g) + Cl- (g) Na+ (aq) + Cl- (aq) -ve H values (favourable interaction) WHY DO THINGS DISSOLVE? If dissolves and solution heats up : exothermic If dissolves and solution cools down: endothermic

+ = + = Ions associating with water Breaking solid into ions Dissolving Lattice Enthalpy + Enthalpy of Hydration = Enthalpy of Solution Substances dissolve because energy and matter tend to disperse (spread out in disorder) 2nd law of Thermodynamics

Second Law of Thermodynamics: the disorder (or entropy) of a system tends to increase ENTROPY (S) Entropy is a measure of disorder Low entropy (S) = low disorder High entropy (S) = greater disorder hot metal block tends to cool gas spreads out as much as possible

= + Total entropy change entropy change of system entropy change of surroundings = + Dissolving disorder of solution disorder of surroundings must be an overall increase in disorder for dissolving to occur

1. If we freeze water, disorder of the water molecules decreases , entropy decreases ( -ve S , -ve H) 2. If we boil water, disorder of the water molecules increases , entropy increases (vapour is highly disordered state) ( +ve S , +ve H)

A non-spontaneous change is a change that occurs only when driven A spontaneous change is a change that has a tendency to occur without been driven by an external influence e.g. the cooling of a hot metal block to the temperature of its surroundings A non-spontaneous change is a change that occurs only when driven e.g. forcing electric current through a metal block to heat it

A chemical reaction is spontaneous if it is accompanied by an increase in the total entropy of the system and the surroundings Spontaneous exothermic reactions are common (e.g. hot metal block spontaneously cooling) because they release heat that increases the entropy of the surroundings. Endothermic reactions are spontaneous only when the entropy of the system increases enough to overcome the decrease in entropy of the surroundings

System in Dynamic Equilibrium  S  0 (zero entropy change) A + B C + D Dynamic (coming and going), equilibrium (no net change) no overall change in disorder  S  0 (zero entropy change)