How are Energy and Chemistry Related?

Unit 13: Guided Notes

Main Idea:

Energy can change form and flow, but it is always conserved. Thermochemical equations express the amount of heat released or absorbed

by chemical reactions. The enthalpy change for a reaction can be calculated using Hess’s law. Changes in enthalpy and entropy determine whether a process is

spontaneous.

New Skills: Calculate heat absorbed and heat released as temperature changes. Calculate calorimetry problems Determine enthalpy changes with phase changes Use Hess’s Law to determine enthalpy changes Calculate heat absorbed and heat released from reactions Determine entropy changes Calculate Free Energy

Academic Language:CalorieCalorimeterChemical potential energyEnergyEnthalpyEnthalpy of fusionEnthalpy of vaporizationEntropyFree energyHeatHess’s LawJouleLaw of conservation of energySecond law of thermodynamicsSpecific heatSpontaneous reactionStandard enthalpy of formationsurroundingsSystemThermochemical equationThermochemistry

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Unit 13 Homework:

CALM: http://calm.indiana.edu/

Book: Chapter 15

13.1 Energy and Heat

CALM: 5 questions p519 #1,2 , p521 #4,5 p525 #12-14 p528 #17, 18, 21

13.2 Thermochemical Equations

CALM: 5 questions p532 #23-25 p533 #27-30

13.3 Calculating Enthalpy Change

CALM: 5 questions p537 #32-33 p541 #35-37

13.4 Reaction Spontaneity

CALM: 5 questions p548 #46-51

13.5 Accumulating Content and Skills

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How are Energy and Chemistry Related?

Chemistry Unit 13: Learning Goals and Objectives

13.1 Energy and Heat - Energy can change form and flow, but it is always conserved.

Define energy

Distinguish between potential and kinetic energy

Calculate the amount of heat absorbed or released by a substance as its temperature changes.

13.2 Thermochemical Equations - Thermochemical equations express the amount of heat released or absorbed by chemical reactions.

Explain the meaning of enthalpy and enthalpy change in chemical reactions and processes.

Write thermochemical equations

Describe how energy is lost or gained during changes of state.

Calculate the heat absorbed or released in a chemical reaction.

13.3 Calculating Enthalpy Change – The enthalpy change for a reaction can be calculated using Hess’s law.

Apply Hess’s law to calculate the enthalpy change for a reaction. Calculate the HΔ rxn using thermochemical equations. Determine the enthalpy change for a reaction using standard enthalpies of

formation.

13.4 Reaction Spontaneity - Changes in enthalpy and entropy determine whether a process is spontaneous.

Define Entropy

Differentiate between spontaneous and nonspontaneous processes.

Explain how changes in entropy and free energy determine the spontaneity of chemical reactions and other processes

13.5 Accumulating Content and Skills– Chemistry content is continuous and builds on prior knowledge and skills. This section will combine this unit with previous units.

Apply knowledge and skills from previous units to content learned in this unit.

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13.1 Energy and Heat - Energy can change form and flow, but it is always conserved.

Objective: Define energy

Energy

Energy is the ability to do work or produce heat. it exists in two basic forms:

1.

2.

Objective: Distinguish between potential and kinetic energy

Potential energy is energy due to the composition or position of an object.

Example:

Chemical Potential energy of the substance depends on its composition:

Kinetic energy is the related to the motion of an object or substance:

Example:

Kinetic energy of a substance is vibrational and random motion of its representative particles.

Law of Conservation of Energy

The law of conservation of energy states that in any chemical reaction or physical process, energy can be converted from on form to another, but it is neither created nor destroyed.

Also called the first law of thermodynamics.

As the diver jumps from the diving board, all of the potential energy changes to kinetic energy.

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Objective: Calculate the amount of heat absorbed or released by a substance as its temperature changes.

Heat

Heat is the energy that is produced or transferred between objects.

Energy transferred can be measured in different units:

calorie (cal.) = the amount of energy to raise the temperature of 1 gram of water 1 °.

joule (J) = SI unit for energy

1 joule = .2390 calories, 1 cal = 4.184 joules

Specific Heat

The specific heat of any substance is the amount of heat required to raise the temperature of one gram of that substance by one degree Celsius.

Different substances have different specific heat values.

Water = 4.184 J/g°C

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Heat Released and Heat Absorbed

Substances can absorb heat and release heat.

q=sm TΔ

q = heat (J), s = specific heat(J/gC), m= mass (g), and T= temperature (C)

Practice Problem #1: In the construction of bridges and skyscrapers, gaps must be left between adjoining steel beams to allow for the expansion and contraction of the metal due to heating and cooling. The temperature of a sample of iron with a mass of 10.0g changed from 50.4°C to 25.0°C with the release of 114 J. What is the specific heat of iron?

Practice Problem #2: A metal rail released 250J of heat when its temperature was lowered to 25.0°C . If the specific heat of the rail is 0.189 J/g°C , what was the initial temperature of the rail?

Practice Problem #3: How much energy does it take to heat a 75,000 Liter pool from 20. °C to 28°C ?

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Calorimetry

A calorimeter is an insulated device used for measuring the amount of heat absorbed or released during a chemical or physical process.

Coffee Cup Calorimeter

Constant pressure calorimeters or coffee cup calorimeters are often used in school chemistry labs

Bomb Calorimeter

Bomb calorimeters are constant volume calorimeters and are often used for reactions involving gases.

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Calorimetry

When working a calorimetry problem, the heat transferred can be calculated by measuring the the change in temperature in the calorimeter.

-q (heat lost) = +q (heat gained)

-sm T=+sm TΔ Δ

Practice Problem #4: Initially 125g of water in a calorimeter is 23.40°C. After a 30.0 g metal sample at 115.0°C is placed into the calorimeter, the final temperature of the water is 29.30°C. What is the specific heat of the metal and how much energy was transferred?

13.2Thermochemical Equations- Thermochemical equations express the amount of heat released or absorbed by chemical reactions.

Objective: Explain the meaning of enthalpy and enthalpy change in chemical reactions and processes.

Thermochemistry

Thermochemistry is the study of heat changes that accompany chemical reactions and phase changes. In thermochemistry we classify two parts of the process.

System =

Surroundings =

Universe =

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Enthalpy

Enthalpy (H) is the heat content and the total amount energy in a system.

Therefore the change in heat of a system is represented, H.Δ

Example: 4Fe(s) + 3O2(g) à2Fe2O3(s) H= -1625 kJΔ

Standard Enthalpy

Standard enthalpy, H°, is the energy changes in a system when all reactants Δand products are in their standard state.

Standard state is the phase at 1 atm and 25°C (298K).

Objective: Write thermochemical equations

Thermochemical Equation

A thermochemical equation is a balanced chemical equation that includes the physical states of all reactants and products and the energy change, usually expressed as the change in enthalpy, H.Δ

Example: C6H12O6(s) + 6O2(g) à6CO2(g) + 6H2O(l) H= -2808kJΔ

In this example,

Objective: Describe how energy is lost or gained during changes of state.

Heat of Fusion/Vaporization

The heat required to vaporize one mole of a liquid is called its (enthalpy) heat of vaporization, ( HΔ vap).

H2O(l) àH2O(g) HΔ vap = 40.7 kJ/mol

The heat required to melt one mole of a solid substance is called its (enthalpy) heat of fusion, ( HΔ fus).

H2O(s) àH2O(l) HΔ fus = 6.01 kJ/mol

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Heat of Vaporization and Fusion

Objective: Calculate the heat absorbed or released in a chemical reaction.

Practice Problem #5: A calorimeter is used to measure the heat evolved when burning 1 mol of glucose.

C6H12O6(s) + 6O2(g) à6CO2(g) + 6H2O(l) H= -2808kJΔ

How much heat would be released if 54.0 g of glucose is burned?

Practice Problem #6: Which of the following processes is exothermic?

1. H2O(l) à H2O(g)

2. CH4(g) à CH4(l)

3. Br2(l) à Br2(s)

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13.3 Calculating Enthalpy Change-

Objective: Apply Hess’s law to calculate the enthalpy change for a reaction.

Hess’s Law

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

Hess’s Law Example:

2S(s)+2Os(g) à 2SO2(g) H = -594 kJΔ

2SO 2(g) + O2(g) à 2SO 3(g) H = -198 kJΔ

Objective: Calculate the HΔ rxn using thermochemical equations.

Hess’s Law Calculations

When combining reactions to for Hess’s law calculations, some reactions must be altered.

If the reaction is reversed, the sign of energy flow must be reversed (enthalpy sign changes).

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Practice Problem #7: Use thermochemical equations a and b below to determine H for the decomposition of hydrogen peroxide (HΔ 2O2).

2H2O2(l) à 2H2O(l) + O2(g)

a) 2H2(g) + O2(g) à 2H2O(l) H=-572 kJΔ

b) H2(g) + O2(g) à H2O2(l) H=-188 kJΔ

Objective: Determine the enthalpy change for a reaction using standard enthalpies of formation.

Enthalpy of Formation

Standard enthalpy (heat) of formation ( H°Δ f) is defined as the change in enthalpy that accompanies the formation of one mole of the compound in its standard state from its elements in their standard states.

H°Δ f = -396 kJ

Enthalpy of formation is based relative to a point of reference. The enthalpy of formation of an element in its standard state is assumed to be 0.00kJ. Any changes to that element in a chemical reaction has an energy change associated with it.

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Enthalpies of Formation in Calculations

An example problem:

H2S(g) + 4F2(g) à 2HF(g) + SF6(g) H°Δ f =?

• ½H2(g) + ½F2(g) àHF(g) H°Δ f = -273 kJ

• S(s) + 3F2(g) à SF6(g) H°Δ f = -1220 kJ

• H2(g) + S(s) à H2S(g) H°Δ f = -21 kJ

Enthalpy Summation Equation

H°Δ rxn = H°ΣΔ f(products) - H°ΣΔ f(reactants)

Example problem:

H2S(g) + 4F2(g) à 2HF(g) + SF6(g) H°Δ f =?

• ½H2(g) + ½F2(g) àHF(g) H°Δ f = -273 kJ

• S(s) + 3F2(g) à SF6(g) H°Δ f = -1220 kJ

• H2(g) + S(s) à H2S(g) H°Δ f = -21 kJ

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Practice Problem #8: Use standard enthalpies of formation to calculate H°Δ rxn

for the combustion of methane.

CH4(g) + 2O2(g) àCO2(g) + 2H2O(l)

13.4 Reaction Spontaneity- Changes in enthalpy and entropy determine whether a process is spontaneous.

Objective: Define Entropy.

Entropy

Entropy (S) is a measure of the number of possible ways that the energy of a system can be distributed.

This is related to the “positional probability” or the freedom of the system’s particles to move and the number of ways they can be arranged.

Solid –

Gas –

Second Law of Thermodynamics

The second law of thermodynamics states that spontaneous processes always proceed in such a way that the entropy of the universe increases.

Entropy is sometimes considered to be a measure of the disorder and randomness of the particles that make up a system.

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Changes in Entropy

Changes in entropy are calculated in a similar way that enthalpy is:

SΔ system = Sproducts – Sreactants

If the entropy of a system decreases during a reaction, SΔ system is negative. Sproducts < Sreactants

1. Entropy changes associated with changes in state can be predicted:

a. H2O(l) à H2O(g) SΔ system

2. The dissolving of a gas in a solvent always results in a decrease in entropy

a. O2(g) àO2(aq) SΔ system

3. Assuming no change in physical state occurs, the entropy of a system usually increases when total number of gas molecules increases.

a. 2SO3(g) à2SO2(g) + O2(g) SΔ system

4. With a few exceptions, entropy increases when a solid or a liquid dissolves in a solvent.

a. NaCl(s) à Na+(aq) + Cl-

(aq) SΔ system

5. The random motion of the particles of a substance increases as its temperature increases.

a. When temperature increase, SΔ system

Practice Problem #9: What is the sign of SΔ system for each of the following changes:

1. CH3OH(l) à CH3OH(g)

2. CaCO3(s) àCaO(s) + CO2(g)

3. CuS(s) + 2O2(g) à CuSO4(s)

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Objective: Differentiate between spontaneous and nonspontaneous processes.

Spontaneous Process

A reaction that occurs with no outside intervention is a spontaneous process.

Some spontaneous processes require some outside energy to get the process started.

example: a car rusts spontaneously.

Spontaneity is dependent on two things:

The change in enthalpy of a reaction

Rust: 4Fe(s) + 3O2(g) à 2Fe2O3(s) H= -1625 kJΔ

The change in entropy of a reaction

Objective: Explain how changes in entropy and free energy determine the spontaneity of chemical reactions and other processes

Free Energy

Willard Gibbs, a Yale physicist, combined the idea of enthalpy and entropy in order to determine if a reaction is spontaneous. He called the equation Gibbs Free Energy:

GΔ system = HΔ system – T SΔ system

H is enthalpy change, T is temperature in Kelvin, and S is entropy Δ Δchange

If GΔ system

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Practice Problem #10: Determine the reaction spontaneity for the following reaction at 382K:

HΔ system =145 kJ and SΔ system = 322 kJ

13.5 Accumulating Content and Skills– Chemistry content is continuous and builds on prior knowledge and skills. This section will combine this unit with previous units.

Apply knowledge and skills from previous units to content learned in this unit.

o Predict the typical enthalpy and entropy changes for the four main types of reactions:

Synthesis

Decomposition

Combustion

Replacement

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