thermochemistry notes2.notebook

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ThermochemistryThe study of energy changes in chemical reactions and physical changes

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Energy• The ability to do work

• Energy changes in chemical reactions and physical changes are measured in the form of heat

• There are two types of energy involved in these changes

• Potential

• Kinetic

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Potential EnergyThe energy due to composition or position of an object

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Kinetic Energy• Energy of motion

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Heat (q)• The total kinetic energy of the random motion of particles in a substance

• Will always flow from a warmer object to a cooler one

• Measured in Joules (J) or calories (cal)

• 1 cal = 4.184 J

• Heat is different from temperature

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TemperatureA measure of the average kinetic energy of random motion of particles in a sample of matter

The more particles, the greater the amount of heat must be transferred to raise the average kinetic energy of all the particles

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This flame has a high temperature, but a low heat content; it won’t keep you warm.

This bonfire has a high temperature AND a high heat content; it will keep you warm.

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Specific Heat• Relates temperature changes to heat changes

• Defined as the amount of heat energy required to increase the temperature of one gram of a substance by one degree Celsius

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Specific Heat• Specific heat can vary if pressure and temperature are not kept constant

• Specific heat is a physical property

• It varies depending on the substance

• Its symbol is Cp, and its units are J/g·°C

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Specific Heat• Must be measured! It cannot be determined by a chemical formula

• Substances with low specific heats require less energy to feel hot than those with high specific heats

• Specific heat can be used to calculate changes in heat

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SubstanceSpecific Heat

J/g·°CWater (liquid) 4.184

Water (solid) 2.03

Water (steam) 2.01

Ethanol (liquid) 2.44

Aluminum (solid) 0.897

Granite (solid) 0.803

Iron (solid) 0.449

Lead (solid) 0.129

Silver (solid) 0.235

Gold (solid) 0.129

Copper (solid) 0.385

Specific Heat Table

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Heat Calculations• Change in heat of a substance can be calculated using the following equation:

• q = mΔTCp

• q = change in heat

• m = mass of the substance

• ΔT = change in temperature of the substance

• Cp = specific heat of the substance

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Specific Heat Problems• Use the RED thermochemistry sheet which has the chart on it to find the specific heat of the

element or substance in the problem

• Solve using algebra and the equation

• q = mCpΔT

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Sample ProblemsIf the temperature of 34.4 g of ethanol increases from 25.0 °C to 78.8 °C, how much heat has been absorbed by the ethanol? The specific heat of ethanol is 2.44 J/(g⋅°C)

Jan 28­12:19 AM

Homework Answers for Section I: Calorimetry

1. ∆H = ­2900J 7. ∆H = ­3450J

2. Cp of Al = .90J/gºC 8. Cp of Ag = .234J/gºC

3. ∆H = 1300J 9. ∆H = ­3510J

4. Tf = 56ºC 10. Ti = 27ºC

5. ∆H = 18800J 11. Tf = 28.9ºC

6. Tf = 32.8ºC

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Try this Example ProblemsA 4.50 g nugget of pure gold absorbed 276 J of heat. What was the final temperature of the gold if the initial temperature was 25 °C ? The specific heat of gold is 0.129 J/(g⋅°C).

Jan 27­12:28 AM

Try this Example ProblemsA 155­g sample of an unknown substance was heated from 25.0°C to 40.0 °C. In the process, the substance absorbed

5696 J of energy. What is the specific heat of the substance?

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Calorimeter Questions• Transfer of heat is measured by measuring the difference in temperature transferred to water from an object

• Specific heat of water (4.184 J/g°C) and its mass is used to solve the problem.

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Sample ProblemA piece of metal is placed in a calorimeter, and causes the 335 g of water to increase in temperature from 21.0°C to 50.1°C. What is the amount of energy released by the piece of metal?

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Total Energy Changes• The amount of heat (q) involved in a reaction is positive (+) if the sample warms up. The sample is gaining heat.

• The amount of heat (q) involved in a reaction is negative (­) if the sample cools off. The sample is releasing, or losing, heat.

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EnthalpyDefined as the total absolute amount of energy in a system.

This cannot be measured or calculated directly, because of a poor understanding of the total energy in a system

Changes in energy, however, CAN be measured

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Enthalpy Change• Enthalpy change is represented as ΔH

• Defined as the heat energy released (­) or absorbed (+) by a system during a physical or chemical change

• System must be at a constant pressure throughout the change

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• ΔH = Hproducts – Hreactants

• Exothermic

• Heat is released

• Products have less heat than reactants, because some of the heat was lost to the surroundings during the reaction

• ΔH is NEGATIVE

• Reaction feels warm

Enthalpy Change

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Enthalpy Change• Endothermic

• Heat is absorbed during a reaction

• Products have more heat than reactants, because they absorbed heat from the surroundings during the reaction

• ΔH is POSITIVE

• Reaction feels cold

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Sample ProblemCalculate the ∆ H of the following reaction:

2SO2 (g) + O2 (g) → 2SO3 (g)

Feb 3­7:50 AM

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Activation Energy• Defined as the minimum amount of energy that must be supplied to a system to start a chemical change.

• Endothermic reactions must have a source from which to draw their energy (usually their surroundings)

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Heat in a Chemical ReactionEnergy can be converted into other forms

Measures of changes in heat energy can be made in a calorimeter

Changes in heat energy can be used to calculate specific heat

These heats refer to the total flow of energy during a chemical change

Jan 27­12:28 AM

Sample Problems: Enthalpy Changes1. Calculate the amount of heat released by the combustion of 1.75 moles of

benzene (C6H6).

2C6H6 + 15O2 à 12CO2 + 6H2O ∆Ηo = ­98.0 kJ

Feb 3­8:09 AM

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Sample Problem­ using stoichiometry

1. How much heat is transferred when 100.0 g of calcium oxide reacts with carbon according to the equation below? Is this reaction endothermic or exothermic?

CaO + 3C à CaC2 + CO ∆Ηo = +464.8 kJ

Feb 4­1:11 AM

Homework Answers for Section II: Enthalpy1a. ΔHrxn = ­80.7kJ exothermicb. ΔHrxn = ­3119.656kJ exothermic for Hrxn sig figs don't matterc. ΔHrxn = ­­718,26kJ exothermicd. ΔHrxn = ­92.22kJ exothermic 1. ­79.5kJ exothermic 8. +88.0kJ endothermic2. ­155kJ exothermic 9. 548.8kJ exothermic3. 18.3 g CO 10. ­78.91kJ exothermic4. +17.4kJ endothermic 11. ­2630kJ exothermic5. ­3.70 kJ exothermic 12. ­256kJ exothermic6. 6.620 molecules H20 13. 93gC7. ­141kJ exothermic 14. 3.95kJ endothermic

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Hess’s Law• States that the total enthalpy change for a chemical or physical change is the same whether it takes one step or several steps.

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Rules for Manipulating Reactions• If the coefficients of an equation are multiplied by a factor, the enthalpy change is multiplied by the same factor

• If an equation is reversed, the sign of ΔH is reversed also

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Sample ProblemUse the thermochemical equations a and b to determine ΔH for the decomposition of hydrogen peroxide.

2H2O2(l) → 2H2O(l) + O2(g)

a. 2H2(g) + O2(g) → 2H2O (l) ΔH = ­572 kJ

b. H2(g) + O2(g) → H2O2(l) ΔH = ­188kJ

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Sample ProblemUse reactions a and b to determine ΔH for the following reaction:

2CO(g) + 2NO(g) → 2CO2 + N2(g)

a. 2CO(g) + O2(g) → 2CO2(g) ΔH = ­566.0 kJ

b. N2(g) + O2(g) → 2NO(g) ΔH = 180.6 kJ

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Sample Problem

Use reactions a and b to determine ΔH for the following reaction.

4Al(s) + 3MnO2(s) → 2Al2O3(s) + 3Mn(s)

a. 4Al(s) + 3O2(g) → 2Al2O3(s) ΔH= ­3352kJ

b. Mn(s) + O2(g) → MnO2(s) ΔH= ­521 kJ

Jan 27­12:28 AM

Sample Problem

• Use reactions a and b to determine ΔH for the following reaction. H2S(g) + 4F2(g) → 2HF(g) + SF6

a. 1/2 H2(g) + 1/2 F2(g) → HF ΔH = ­273 kJ

b. S(s) + 3F2 → SF6 ΔH = ­1220 kJ

c. H2(g) + S(s) → H2S(g) ΔH = ­21 kJ

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• A process that occurs without any outside intervention. a change that proceeds on its own.

• Spontaneous processes can be fast or slow

• Occurs primarily in one direction

Spontaneous Process

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Spontaneity• Some can occur when a small amount of energy is added to the system

• Some can be reversed if conditions change

• The more energy released, the lower enthalpy is and therefore the more likely the reaction will be spontaneous

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Entropy and Stability• Entropy (S): a measure of the degree randomness or disorder of the system

• The tendency of nature is toward more disorder

• Its effects increase with temperature

• Measured in units of J/K

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Entropy∆S = Sproducts – Sreactants

If entropy of a system increases during a reaction or process, Sproducts > Sreactants and ∆S is positive

If entropy of a system decreases, Sproducts < Sreactants and ∆S is negative

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Entropy• When a solid turns to a liquid and a liquid to a gas, entropy increases (∆S > 0);

• When a gas dissolves in a liquid, entropy decreases (∆S < 0)

Feb 10­1:52 AM

Predict the sign of ∆Ssystem for the following changes

1. ClF(g) + F2(g) → ClF3(g) negative Why?

2. NH3(g) → NH3(aq) negative Why?

3. CH3OH(l) → CH3OH(aq) positvie Why?

4. C10H8(l) → C10H8(s) negative Why?

5. CH3OH(s) → CH3OH(l) positive Why??

Jan 27­12:28 AM

Entropy and EnthalpyFree energy is a combination of entropy and enthalpy

Free energy (G): used to predict spontaneity

Free energy is the energy available to do work

It can indicate whether enthalpy or entropy has a greater effect

Jan 27­12:28 AM

2 way to calculate Free Energy• ∆G = Gproducts – Greactants

• Also, ∆G = ∆H – (T ∆S)

• ∆H = change in enthalpy

• T = temp in Kelvin

• ∆S = change in entropy

• Free energy is a measure of the competition between enthalpy and entropy

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Gibb’s Free Energy• If ∆G is negative, the reaction is spontaneous

• If ∆G is positive, the reaction is not spontaneous

• If ∆G is =0 the rxn is at equilbrium