Thermodynamics

Heat and Temperature Thermochemistry is the study of the transfers of energy as heat that accompany chemical reactions and physical changes Thermochemistry is the study of the transfers of energy as heat that accompany chemical reactions and physical changes Calorimeters measure the energy that is absorbed or released Calorimeters measure the energy that is absorbed or released

Heat and Temperature Temperature is a measure of the average kinetic energy of the particles in a sample of matter Temperature is a measure of the average kinetic energy of the particles in a sample of matter The joule (J) is the SI unit of heat and all other forms of energy The joule (J) is the SI unit of heat and all other forms of energy

Heat and Temperature Heat is the energy transferred between substances because of their temperature differences Heat is the energy transferred between substances because of their temperature differences Heat always moves from matter at a higher temperature to matter at a lower temperature Heat always moves from matter at a higher temperature to matter at a lower temperature

Specific Heat Specific Heat is the energy required to raise the temperature of one gram of substance by one degree Celsius Specific Heat is the energy required to raise the temperature of one gram of substance by one degree Celsius The energy (heat) gained or lost with a change in temperature can be calculated using q = C x m x ΔT The energy (heat) gained or lost with a change in temperature can be calculated using q = C x m x ΔT –q is heat (energy), C is specific heat, m is mass, and ΔT is temperature change

Enthalpy The energy absorbed in a reaction is represented by ΔH. H is the symbol for enthalpy The energy absorbed in a reaction is represented by ΔH. H is the symbol for enthalpy An enthalpy change is the amount of energy absorbed by a system as heat during a process at constant pressure An enthalpy change is the amount of energy absorbed by a system as heat during a process at constant pressure

Enthalpy ΔH = H products – H reactants ΔH = H products – H reactants The enthalpy of reaction is the quantity transferred as heat during a chemical reaction The enthalpy of reaction is the quantity transferred as heat during a chemical reaction Its units are joules/mole Its units are joules/mole

Enthalpy A thermochemical equation includes the energy absorbed or released during the given reaction A thermochemical equation includes the energy absorbed or released during the given reaction –2H 2 + O 2  2H 2 O kJ If ΔH is negative (products side), the reaction is exothermic. If it is positive (reactants side), it is endothermic If ΔH is negative (products side), the reaction is exothermic. If it is positive (reactants side), it is endothermic

Enthalpy The enthalpy change for reaction relates to the stoichiometric proportions for that reaction The enthalpy change for reaction relates to the stoichiometric proportions for that reaction Therefore, stoichiometry can be used to calculate enthalpy changes for a given sample amount for a particular reaction Therefore, stoichiometry can be used to calculate enthalpy changes for a given sample amount for a particular reaction

Enthalpy Example: Example: –How much heat will be released if 1.0 g of H 2 O 2 decomposes according to the following reaction (2H 2 O 2  2H 2 O + O 2 ΔH = -190kJ) –Calculate moles of sample (use formula wt)  1.0 g H 2 O 2 x (1 mol/34.0 g H 2 O 2 ) = mol –Multiply by enthalpy change  mol H 2 O 2 x (-190 kJ/2 mol H 2 O 2 ) = -2.8 kJ [the 2 mol is the coefficient in the balanced eqn]

Hess’ Law Hess’ Law – The overall enthalpy change for a reaction is equal to the sum of enthalpy changes for the individual steps of the process Hess’ Law – The overall enthalpy change for a reaction is equal to the sum of enthalpy changes for the individual steps of the process

Hess’ Law Smog is caused by NO 2, a gas formed from N 2 and O 2 in a two-step reaction. The net reaction is as follows: Smog is caused by NO 2, a gas formed from N 2 and O 2 in a two-step reaction. The net reaction is as follows: N 2 + O 2  2 NOΔH = 181 kJ N 2 + O 2  2 NOΔH = 181 kJ + 2NO + O 2  2NO 2 ΔH = -113 kJ N 2 + 2O 2 + 2NO  2NO + 2 NO 2 ΔH net = ΔH 1 + ΔH 2

Hess’ Law The net equation is: The net equation is: N 2 + 2O 2  2NO 2 According to Hess’ Law, the enthalpy change for the net reaction will be sum of the enthalpy changes for the individual steps [181 kJ + (-113 kJ) = 68 kJ] According to Hess’ Law, the enthalpy change for the net reaction will be sum of the enthalpy changes for the individual steps [181 kJ + (-113 kJ) = 68 kJ]

Hess’ Law Reactions may also be manipulated as algebra problems are using the following rules: Reactions may also be manipulated as algebra problems are using the following rules: 1.If the coefficients of an equation are multiplied by a factor, the enthalpy change is multiplied by that factor. This is because the heat absorbed/released by a reaction depends on the quantities of reactants & products

Hess’ Law 2. If an equation is reversed, the sign of ΔH changes also. This is because if a reaction releases heat in one direction, in the reverse reaction it will absorb heat (which changes the sign of the enthalpy change) 2. If an equation is reversed, the sign of ΔH changes also. This is because if a reaction releases heat in one direction, in the reverse reaction it will absorb heat (which changes the sign of the enthalpy change)

Calorimetry In a reaction carried out in a calorimeter, the heat of the reaction (q rxn ) is transferred to the surroundings (q surr ). Therefore, q rxn = -q surr In a reaction carried out in a calorimeter, the heat of the reaction (q rxn ) is transferred to the surroundings (q surr ). Therefore, q rxn = -q surr Remember that you can calculate q surr using q = m x C x ΔT (the temperature change is the final temp minus the initial temp) Remember that you can calculate q surr using q = m x C x ΔT (the temperature change is the final temp minus the initial temp)

Calorimetry Once q surr is determined, q rxn is determined by changing the sign of q surr Once q surr is determined, q rxn is determined by changing the sign of q surr ΔH is equal to the q rxn divided by the moles of the sample used in the reaction multiplied by the stoichiometric proportion of moles needed for the reaction ΔH is equal to the q rxn divided by the moles of the sample used in the reaction multiplied by the stoichiometric proportion of moles needed for the reaction

Calorimetry Example: Example: When a 4.25 g sample of solid NH 4 NO 3 dissolves in 60.0 g H 2 O in a calorimeter, the temperature drops from 21.0 C to 16.9 C. Calculate ΔH for the solution process. NH 4 NO 3  NH NO 3 -

Calorimetry Step 1: Calculate q surr Step 1: Calculate q surr q surr = m x C x ΔT q surr = (60.0 g)(4.184 J/g C)(16.9 C – 21.0 C) q surr = x 10 3 kJ Step 2: Determine q rxn Step 2: Determine q rxn q rxn = -q surr = -(-1.03 x 10 3 kJ) = 1.03 x 10 3 kJ

Calorimetry Step 3: Calculate ΔH Step 3: Calculate ΔH First calculate moles of your sample 4.25 g NH 4 NO 3 = mol NH 4 NO 3 ΔH = (q rxn /mol of sample) x mol from equation ΔH = (1.03 x 10 3 kJ/ mol) x (1 mol) ΔH = 19.4 kJ

Entropy Most reactions are exothermic, and tend to proceed spontaneously, making products that are more stable than the reactants Most reactions are exothermic, and tend to proceed spontaneously, making products that are more stable than the reactants Some endothermic reactions are spontaneous, which means something other than ΔH determines whether a reaction will occur. Some endothermic reactions are spontaneous, which means something other than ΔH determines whether a reaction will occur.

Entropy The 2 nd Law of Thermodynamics states that the entropy (disorder) of a system will increase over time The 2 nd Law of Thermodynamics states that the entropy (disorder) of a system will increase over time Entropy (S) is a measure of the degree of randomness of the particles in a system Entropy (S) is a measure of the degree of randomness of the particles in a system

Entropy The entropy of a pure crystalline solid is zero at absolute zero (3 rd law of thermodynamics) The entropy of a pure crystalline solid is zero at absolute zero (3 rd law of thermodynamics) Entropy change (ΔS) is the difference between the entropy of the products and the entropy of the reactants. The more random the system, the more positive this value is Entropy change (ΔS) is the difference between the entropy of the products and the entropy of the reactants. The more random the system, the more positive this value is

Gibbs Free Energy Processes in nature are driven in two directions: toward least enthalpy and greatest entropy Processes in nature are driven in two directions: toward least enthalpy and greatest entropy (Gibbs) Free Energy (G) is a relation of the enthalpy and entropy factors at a given temperature and pressure. Natural processes tend in the direction that lowers the free energy of the system (Gibbs) Free Energy (G) is a relation of the enthalpy and entropy factors at a given temperature and pressure. Natural processes tend in the direction that lowers the free energy of the system

Gibbs Free Energy Only the change in free energy can be measured using the following equation: Only the change in free energy can be measured using the following equation: ΔG = ΔH – TΔS (change in free energy = change in enthalpy minus absolute temperature times change in entropy) If ΔG is less than zero, the reaction is spontaneous

Gibbs Free Energy ΔH and ΔS can be positive or negative, which leads to the following possibilities: ΔH and ΔS can be positive or negative, which leads to the following possibilities: ΔHΔHΔHΔH ΔSΔSΔSΔS ΔGΔGΔGΔGSpontaneous? - Value (exothermic) + value (more random) Always negative Yes - Value (exothermic) - Value (less random) Negative at low temperatures At low temps + value (endothermic) + value (more random) Negative at high temps At high temps + value (endothermic) - Value (less random) Never negative No

Gibbs Free Energy Example calculation of free energy Example calculation of free energy For the reaction NH 4 Cl (s)  NH 3 (g) + HCl (g) at K, ΔH = 176 kJ/mol and ΔS = kJ/(mol ∙ K). Calculate ΔG, and tell whether the reaction is spontaneous in the forward reaction at this temperature. For the reaction NH 4 Cl (s)  NH 3 (g) + HCl (g) at K, ΔH = 176 kJ/mol and ΔS = kJ/(mol ∙ K). Calculate ΔG, and tell whether the reaction is spontaneous in the forward reaction at this temperature. Use ΔG = ΔH – TΔS Use ΔG = ΔH – TΔS ΔG = (176 kJ/mol) – ( K)[0.285 kJ/(mol ∙ K)] ΔG = (176 kJ/mol) – (84.9 kJ/mol) ΔG = 91 kJ/mol