Thermodynamics

Energy and Heat

• Energy = the ability to do work or to produce heat– Kinetic energy: energy of motion– Potential energy: stored energy• Chemical potential energy = the energy stored in a

substance because of its composition

• Law of Conservation of Energy: in any chemical reaction or physical process, energy can be converted from one form to another, but it is neither created nor destroyed

Heat = the energy that is in the process of flowing from a warmer object to a cooler object– Measuring heat:

• Calorie (cal) = the amount of heat required to raise the temperature of one gram of pure water by 1°C

• Joule (J) = SI unit of heat and energy 1 calorie = 4.184 Joules

Heat Unit Conversions Convert the following:1 calorie = 4.184 Joules

1.230 calories = _____________ joules

2.5.87 joules = ____________ calories

3.52 kilocalories = ______________ joules

4.4.56 kilojoules = ______________ calories

Specific Heat

Specific heat (J/g x oC)= the amount of heat required to raise the temperature of one gram of that substance by one degree Celsius

q = cm∆Twhere:

q = energyc = specific heat capacity

m = mass of sample in grams ∆T = temperature change in Celsius

Calorimetry• Calorimeter = an insulated device used for

measuring the amount of heat absorbed or released during a chemical or physical process– You can use a calorimeter to determine the

specific heat of an unknown metal • Measured mass of water has an initial temperature• Piece of hot metal is added• The metal transfers heat to the water until the metal

and water attain the same temperature• Final temperature of the water is measured• Specific heat can be calculated

Thermochemistry

Thermochemistry = the study of heat changes that accompany chemical reactions and phase changes– Universe = system + surroundings• System = the specific part of the universe that contains

the reaction or process you wish to study.• Surroundings = everything in the universe other than

the system

Thermochemistry

Enthalpy (H) = the heat content of a system at constant pressure– Enthalpy (heat) of reaction (∆Hrxn) is the change in

enthalpy for a reaction

ΔHrxn = Hproducts – Hreactants

• If positive, energy has gone INTO the reaction• If negative, energy has been RELEASED by the reaction

Endothermic vs. Exothermic

• Endothermic reaction = chemical reaction that absorbs heat– A greater amount of energy is required to break

the existing bonds in the reactants than is released when the new bonds form in the product molecules

• Exothermic reaction = chemical reaction that gives off heat– More energy is released forming new bonds than

is required to break bonds in the initial reactants

Thermochemical EquationsThermochemical Equation = a balanced

chemical equation that includes the physical states of all reactants and products and the

energy change, usually expressed as change in enthalpy, ΔH.

Heat pack equation4Fe(s) + 3O2(g) 2Fe2O3(s) ΔH = -1625 kJ

Cold pack equationNH4NO3(s) NH4

+(aq) + NO3-(aq) ΔH = 27kJ

Changes of State

• Molar enthalpy (heat) of vaporization (ΔHvap)– the heat required to vaporize one mole of a liquid

H2O(l) H2O(g) ΔHvap = 40.7 kJ

• Molar enthalpy (heat) of fusion (ΔHfus) – the heat required to melt one mole of a solid substance

H2O(s) H2O(l) ΔHfus = 6.01 kJ

Combustion ReactionsA Review

Combustion reaction = a substance rapidly combines with oxygen to form either carbon

dioxide and water or metal oxides– Hydrocarbons burning carbon dioxide and

water• Example: CH4 (g) + 2O2 (g) CO2(g) + 2H2O (g)

– Metals burning metal oxides• Example: 2Mg (s) + O2 (g) 2MgO (s)

Combustion and Welding

When welding is done with an acetylene torch, acetylene combines with oxygen to form

carbon dioxide and water. This reaction is exothermic and enough energy is released to

melt metal

2C2H2(g) + 5O2(g) 4CO2(g) + 2H2O(g) + energy

Combustion and Challenger

• Hydrogen gas and oxygen gas react when hydrogen is heated forming water and releasing a large amount of energy

2H2(g) + O2 (g) 2 H2O (g)- Hydrogen tank propelledinto the oxygen tank by leak- Combustion reaction abovetook place, thus destroying the shuttle

Calculating Enthalpy Change

• Hess’s Law states that if you can add two or more thermochemical equations to produce a final equation for a reaction, then the sum of the enthalpy changes for the individual reactions is the enthalpy change for the final reaction.

Hess’s Law Let’s Walk Through a Problem…

• Use thermochemical equations “a” and “b” below to determine ΔH for the decomposition of hydrogen peroxide (H2O2)

2 H2O2(l) 2H2O(l) + O2(g)

a)2H2(g) + O2(g) 2H2O(l) ΔH = -572 kJ

b)H2(g) + O2(g) H2O2(l) ΔH = -188 kJ

Heat of Formation

• Standard enthalpy (heat) of formation (∆Hf) = the change in enthalpy that accompanies the formation of one mole of the compound in its standard state from its constituent elements in their standard states

2S (s) + 3O2 (g) 2SO3 (g) ∆Hf = -396 kJ

396 kJ of heat are given off in this reaction

Heat of Reaction

• Heat of reaction (∆Hrxn)= amount of heat energy given off or absorbed in a particular chemical reaction for a given amount of reactants or products

• The standard heats of formation equations combine to produce the desired equation and its ∆Hrxn.

∆Hrxn = ∑∆Hf (products) - ∑∆Hf (reactants)

Heat of ReactionExample

• What is the heat of reaction for the following reaction?

2 Mg (s) + O2 (g) 2 MgO (s)

Given:∆Hf(Mg) = 0 kJ/mol

∆Hf(O2) = 0 kJ/mol

∆Hf(MgO) = -602 kJ/mol

Heat of ReactionExample (cont’d)

∆Hrxn = ∑∆Hf (products) - ∑∆Hf (reactants)

= 2 (∆Hf(MgO) ) – [2(∆Hf(Mg) + ∆Hf(O2)]

= 2 mol (-602 kJ/mol) – [0 + 0]

= -1204 kJ for 2 moles of MgO

1204 kJ are given off to make 2 moles of MgO

15.5 – Reaction Spontaneity

• Spontaneous process – any physical or chemical change that once begun, occurs with no outside intervention (but some outside energy may be necessary to get it started)

• Entropy (S) is a measure of the disorder of a system.

Second Law of Thermodynamics

• “Spontaneous processes always proceed in such a way that the entropy of the universe increases”

• Systems prefer to be in disorder.

Activation Energy

• Activation energy (Ea) = the minimum amount of energy that reacting particles must have to cause a chemical reaction

Activation Energy Endothermic Reactions

Activation EnergyExothermic Reactions

Factors that Influence Reaction Rates

• Nature of Reactants• Concentration• Surface Area• Temperature• Catalysts

Nature of ReactantsSome elements are more reactive than othersExample: sodium is more reactive than

calcium so the reaction of sodium and water occurs at a faster rate than calcium and water

More reactive elements faster rate of reaction

Concentration

• Higher concentration means there are more particles around in the reaction mix to collide

• Reactant A “finds” reactant B more easily if there are more A and B particles near each other

• Higher concentration of reactants faster rate of reaction

Surface Area• Think about dissolving sugar:– Sugar cube + water dissolves slowly– Granulated sugar + water dissolves faster

• Sugar cube overall has smaller surface area than sugar granules exposed to water for dissolving reaction to occur

• Larger surface area faster rate of reaction

Temperature• Think about dissolving cocoa mix – is it easier when the

water is cold or hot?• Increased temperature increases the kinetic motion of

particles particles can collide easier more collisions reaction

• Increased temperature faster rate of reaction

Catalysts

• Catalyst = a substance that increases the rate of a chemical reaction without itself being consumed in the reaction

• Catalysts reduce the activation energy, making it easier for the reaction to occur

***LeChâtelier’s Principle

• If a stress is applied to a reversible system at equilibrium, the system shifts in the direction that relieves the stress. (Think of a see-saw)

• Example:– Adding more reactants leads to production of

more products (shifts the equilibrium to the right)– Adding more of the products leads to an increase

in the reverse reaction, to produce more reactants (shifts the equilibrium to the left)