chapter 15 energy and chemical change. 15.1 energy energy can change for and flow, but it is always...
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Chapter 15
Energy and Chemical Change
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15.1 Energy
• Energy can change for and flow, but it is always conserved.
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The Nature of Energy
• Energy – the ability to do work or produce heat– Potential energy
– Kinetic energy
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• Chemical systems contain both kinetic and potential energy– Kinetic energy
– Potential energy
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Law of conservation of energy
• 1st law of thermodynamics – energy cannot be created or destroyed, only transferred
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Chemical potential energy
• Energy that is stored in a substance because of its composition
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Heat
• Heat = energy• Symbol = q• Always flows from warmer object to cooler
object
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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
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• 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?
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• A big mac contains 550 Calories, express this energy in Joules.
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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
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Calculating changes in heat
• q = (m)(c)( T)
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• 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)
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• 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)
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15.2 Heat
• Calorimeter – insulated device used for measuring the amount of heat absorbed or released during a chemical or physical process.
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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
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• 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)
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• 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?
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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
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• ΔHrxn = Hproducts – Hreactants
• Exothermic reaction
• Endothermic reaction
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15.3 Thermochemical Equations
• Thermochemical equations express the amount of heat released or absorbed by chemical reactions
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Writing Thermochemical Equations
• Thermochemical equation – balanced chemical equation that includes the states of all reactants and products and the energy change
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• 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
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• How much heat is released when 54.0 g of glucose is burned? ΔHcomb = -2808 kJ
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• H2O(l) H2O(g) ΔHvap = 40.7 kJ
• H2O(s) H2O(l) ΔHfus = 6.01 kJ
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• How much heat must be absorbed to melt 150.0 g of water?
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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
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• 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
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• 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
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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
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• ΔHrxn = sum of ΔHf products – sum of ΔHf reactants
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• What is ΔHrxn for the following equation:CH4(g) +2O2(g) CO2(g) + 2H2O(l)
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• What is the ΔHrxn for the following:
4NH3(g) + 7O2(g) 4NO2(g) + 6H2O(g)
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15.5 Reaction Spontaneity
• Changes in enthalpy and entropy determine whether a process is spontaneous
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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
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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
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• Second law of thermodynamics – spontaneous processes always proceed in such a way that the entropy of the universe increases
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Predicting changes in entropy
• +ΔS = entropy increases• - ΔS = entropy decreases
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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
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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
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• 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)
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Gibbs Free Energy
• ΔG = ΔH – TΔS
- ΔG = spontaneous reaction+ ΔG = nonspontaneous reaction
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• For a process, ΔH = 145 kJ and ΔS = 322 J/K. is the process spontaneous at 382K?
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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