Chapter 15

Energy and Chemical Change

15.1 Energy

• Energy can change for and flow, but it is always conserved.

The Nature of Energy

• Energy – the ability to do work or produce heat– Potential energy

– Kinetic energy

• Chemical systems contain both kinetic and potential energy– Kinetic energy

– Potential energy

Law of conservation of energy

• 1st law of thermodynamics – energy cannot be created or destroyed, only transferred

Chemical potential energy

• Energy that is stored in a substance because of its composition

Heat

• Heat = energy• Symbol = q• Always flows from warmer object to cooler

object

Measuring Heat

• calorie – the amount of energy needed to raise the temperature of one gram of pure water by one degree Celsius– Calorie vs. calorie

• Joule – the SI unit of energy• 1 calorie = 4.184 J

• A chemical reaction releases 213 J of energy. How many calories has this reaction released?

• A big mac contains 550 Calories, express this energy in Joules.

Specific Heat

• Amount of heat needed to raise the temperature of 1 g of a substance by 1 oC– Symbol = c– Unit = J/g oC

• High specific heat = absorbs a lot of energy w/small changes in temperature– Water = 4.184 J/g oC– Concrete = 0.84 J/g oC

Calculating changes in heat

• q = (m)(c)( T)

• If the temperature of 34.4 g of ethanol increases from 25.0 oC to 78.8 oC, how much heat has been absorbed by the ethanol? (specific heat of ethanol = 2.44 J/g oC)

• The temperature of a 10.0 g piece of iron changed from 50.4 oC to 25.0 oC, how much energy was release by the iron? (specific heat of iron is 0.449 J/g oC)

15.2 Heat

• Calorimeter – insulated device used for measuring the amount of heat absorbed or released during a chemical or physical process.

Chemical Energy and the Universe

• Thermochemistry – the study of heat changes that accompany chemical reactions and phase changes – System – specific part of the universe that

contains the reaction or process you want to study– Surroundings – everything else in the universe

other than the system

• Energy being absorbed/released is indicated from the system’s point of view- q = energy released by system(exothermic)+ q = energy absorbed by system(endothermic)

• qsystem = -qsurroundings

• (m1)(c1)( Δ T1) = - (m2)(c2)(Δ T2)

• A 125 g sample of iron at 93.5 oC is dropped into an unknown mass of water at 25.0 oC. The final temperature of the mixture is 32.0 oC. The C of iron is 0.451 J/g oC, the C of water is 4.18 J/g oC. What is the mass of the water?

Enthalpy and enthalpy change

• Enthalpy (H) – heat content of a system at a constant pressure

• Cannot measure enthalpy directly but can measure change in enthalpy (heat absorbed or released in a chemical reaction)

• Enthalpy of reaction ( Hrxn)– change in enthalpy for a reaction

• ΔHrxn = Hproducts – Hreactants

• Exothermic reaction

• Endothermic reaction

15.3 Thermochemical Equations

• Thermochemical equations express the amount of heat released or absorbed by chemical reactions

Writing Thermochemical Equations

• Thermochemical equation – balanced chemical equation that includes the states of all reactants and products and the energy change

• Enthalpy (heat) of combustion (ΔHcomb) – the enthalpy change for the complete burning of one mole of the substance

• Enthalpy (heat) of vaporization (ΔHvap) – heat required to vaporize one mole of a liquid

• Enthalpy (heat) of fusion (ΔHfus) – heat required to melt one mole of a solid

• How much heat is released when 54.0 g of glucose is burned? ΔHcomb = -2808 kJ

• H2O(l) H2O(g) ΔHvap = 40.7 kJ

• H2O(s) H2O(l) ΔHfus = 6.01 kJ

• How much heat must be absorbed to melt 150.0 g of water?

15.4 Calculating Enthalpy Change

• Hess’s Law – 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

• What is the energy change for the following reaction:2S(s) + 3O2(g) 2SO3(g)

a. S(s) + O2(g) SO2(g) ΔH = -594 kJ

b. 2SO3(g) 2SO2(g) + O2(g) ΔH = 198 kJ

• What is the energy change for the following?H2O2(l) 2H2O(l) + O2(g)

a. 2H2(g) + O2(g) 2H2O(l) ΔH = -572kJ

b. H2(g) + O2(g) H2O2(l) ΔH = -188kJ

Standard heat of formation

• The change in enthalpy that accompanies the formation of one mole of the compound in its standard state

• ΔHf

• ΔHf of an element = 0

• ΔHrxn = sum of ΔHf products – sum of ΔHf reactants

• What is ΔHrxn for the following equation:CH4(g) +2O2(g) CO2(g) + 2H2O(l)

• What is the ΔHrxn for the following:

4NH3(g) + 7O2(g) 4NO2(g) + 6H2O(g)

15.5 Reaction Spontaneity

• Changes in enthalpy and entropy determine whether a process is spontaneous

Spontaneous processes

• Any physical or chemical change that once begun, occurs with no outside intervention– Iron rusting– Paper burning

• Often some energy from the surroundings must be supplied to get the process started

Entropy

• Entropy(S) – a measure of the number of possible was that the energy of a system can be distributed– Determined by the freedom of the systems

particles to move and the number of ways they can be arranged

– Disorder or randomness of a system

• Second law of thermodynamics – spontaneous processes always proceed in such a way that the entropy of the universe increases

Predicting changes in entropy

• +ΔS = entropy increases• - ΔS = entropy decreases

Changes resulting in +ΔS

• Changes in state that allow more molecule movement– (s)(l)– (l) (g)

• Number of particles increases in a reactionCaCO3 CaO + CO2

• Solid or liquid dissolves in solvent• Increase in temperature

Changes resulting in -ΔS

• Phase changes that decrease molecule movement– (g) (l)– (l) (s)

• Number of particles decreases in a reaction• Dissolving of gas in a solvent• Decrease in temperature

• Predict the sign of ΔS for the following:– ClF(g) + F2(g) ClF3(g)

– NH3(g) NH3(aq)

– CH3OH(l) CH3OH(aq)

– C10H8(l) C10H8(s)

Gibbs Free Energy

• ΔG = ΔH – TΔS

- ΔG = spontaneous reaction+ ΔG = nonspontaneous reaction

• For a process, ΔH = 145 kJ and ΔS = 322 J/K. is the process spontaneous at 382K?

ΔH ΔS ΔG Reaction Spontaneity

Negative Positive Negative Always spontaneous

Negative Negative Negative or positive

Spontaneous at low temperatures

Positive Positive Negative or positive

Spontaneous at high temperatures

Positive Negative Positive Never spontaneous