MGU Science Series — High School Chemistry

Part V: States of Matter, Energy, and Reaction Behavior

Energy · Heat · Calorimetry · Enthalpy · Energy Diagrams

Chapter 20 of 28

MGU Science Series · High School Chemistry · Library Textbook Edition

Chapter 20:
Thermochemistry

Energy in chemical systems, heat and temperature, exothermic and endothermic processes, specific heat, calorimetry, enthalpy, energy diagrams, and thermochemistry in fuels, food, and industry

■ Part V: States of Matter, Energy, and Reaction Behavior ■

■ 9 Sections • Chapter 20 of 28 ■

Follows Chapter 19: Liquids, Solids, and Phase Changes · Prepares students for Chapter 21: Reaction Rates

Chapter 20 — Front Matter

Purpose and Overview

Every chemical reaction involves energy. A candle releases heat and light as it burns. A cold pack absorbs heat when its chemicals dissolve. A car engine converts chemical energy in fuel into mechanical motion. A battery stores chemical energy and releases it as electricity. Thermochemistry is the branch of chemistry that studies these energy changes — measuring them, explaining them, and predicting them.

In this chapter, you will learn how chemists think about energy, heat, and temperature. You will explore what it means for a reaction to be exothermic or endothermic, how scientists measure energy changes using calorimeters, and how enthalpy is used to describe the heat content of chemical systems. You will also practice calculations involving specific heat and calorimetry, and connect thermochemical ideas to fuels, food, industry, and environmental chemistry.

Opening Question

When you hold a piece of ice, your hand gets cold and the ice begins to melt. When you light a match, heat rushes outward. Both events involve energy moving between chemicals and their surroundings — but in opposite directions. By the end of this chapter, you will know how to measure that energy flow, calculate how much heat is involved, and explain why some reactions release energy while others absorb it.

9Sections

7Objectives

25+Key Terms

4Worked Examples

Learning Objectives

1Explain the relationship between energy, heat, and temperature and distinguish among kinetic energy, potential energy, and thermal energy.
2Define system and surroundings and apply the sign convention for energy flow.
3Identify and explain exothermic and endothermic processes using examples, particle models, and energy flow direction.
4Use the specific heat equation (q = mcΔT) to calculate heat absorbed or released, temperature change, mass, or specific heat.
5Describe how a coffee cup calorimeter works and use calorimetry data to calculate heat of reaction.
6Define enthalpy change (ΔH), interpret thermochemical equations, and connect ΔH values to exothermic and endothermic reactions.
7Read, draw, and interpret energy diagrams showing reactants, products, activation energy, transition state, and ΔH.

Chapter 20 — Contents

Sections at a Glance

20.1Energy in Chemical SystemsWhat energy is; kinetic energy, potential energy, thermal energy; law of conservation of energy; why energy matters in chemistry
20.2Heat and TemperatureTemperature vs. heat; intensive and extensive properties; units of heat (joules, calories); temperature scales and Kelvin conversion
20.3Systems and SurroundingsDefining the system; open, closed, and isolated systems; sign convention for energy transfer
20.4Exothermic and Endothermic ProcessesExothermic definition and examples; endothermic definition and examples; comparison table; common misconceptions
20.5Specific HeatWhy materials heat at different rates; specific heat capacity definition and units; the equation q = mcΔT; worked examples; specific heat table
20.6CalorimetryWhat a calorimeter does; coffee cup calorimeter; bomb calorimeter; calorimetry calculations; worked example
20.7Enthalpy ChangeWhat enthalpy means; ΔH and its sign; thermochemical equations; enthalpy as a state function; standard enthalpies of formation
20.8Energy DiagramsWhat energy diagrams show; features of a diagram; exothermic diagrams; endothermic diagrams; how to read and draw diagrams; catalysts
20.9Thermochemistry in Fuels, Food, and IndustryFuel energy density; combustion of hydrocarbons; food as chemical fuel; caloric content; Haber process; refrigeration; energy storage; climate

Section 20.1

Energy in Chemical Systems

What Energy Is

Energy is the capacity to do work or transfer heat. It is not a substance you can hold or see — it is a property of matter in motion or matter in certain arrangements. A stretched rubber band has energy. A moving car has energy. A chemical bond between two atoms has energy. A warm cup of water has more energy than a cold cup of the same water.

Chemists care about energy because chemical reactions always involve changes in energy. Bonds break and form during reactions, and these events release or absorb energy. Understanding these energy changes allows chemists to predict how reactions behave, design efficient industrial processes, develop better fuels, understand metabolism, and assess environmental consequences.

Forms of Energy Relevant to Chemistry

Form of Energy	Description	Chemistry Example
Kinetic energy	Energy of motion; at the particle level, atoms and molecules vibrate, rotate, and translate through space	Faster-moving particles = higher temperature
Chemical potential energy	Energy stored in chemical bonds and in the arrangement of atoms relative to one another	Energy stored in a molecule of glucose or methane
Thermal energy	Total kinetic energy of all particles in a sample; depends on both temperature and amount of matter	A bathtub of warm water holds more thermal energy than a cup of boiling water

Law of Conservation of Energy

Energy cannot be created or destroyed. It can only be converted from one form to another or transferred from one place to another. This is sometimes called the first law of thermodynamics. When a reaction releases energy, that energy does not disappear — it flows into the surroundings. Whatever energy the chemicals lose, the surroundings gain, and vice versa. The total energy of the universe remains constant.

Section 20.2

Heat and Temperature

Two of the most commonly confused concepts in chemistry are heat and temperature. They are related, but they are not the same thing.

Critical Distinction

Temperature is a measure of the average kinetic energy of particles — how fast they move on average. It is an intensive property (does not depend on amount). Heat is the transfer of thermal energy between objects at different temperatures. It is an extensive property (depends on amount and temperature change). A small flame and a bonfire can be at the same temperature, yet the bonfire contains enormously more total energy.

Property	Temperature	Heat
What it measures	Average kinetic energy of particles	Transfer of thermal energy between objects
Type of property	Intensive (independent of amount)	Extensive (depends on mass and ΔT)
SI unit	Kelvin (K)	Joule (J)
Direction	Not directional	Always flows from warmer to cooler

Units of Heat Energy

In chemistry, heat is measured in joules (J) or kilojoules (kJ). One kilojoule equals 1,000 joules. An older unit, the calorie (cal), is the amount of heat needed to raise the temperature of one gram of water by one degree Celsius. The relationship is:

1 cal = 4.184 J 1 kcal = 1,000 cal = 4,184 J

The kilocalorie (kcal), or "Calorie" (capital C) on food labels, equals 1,000 small calories. A food label listing 200 Calories means 200 kilocalories — approximately 836,800 joules of chemical energy. In chemistry class, you will work primarily in joules and kilojoules.

Temperature Scales

Scale	Unit	Freezing Point of Water	Boiling Point of Water	Use in Chemistry
Celsius	°C	0°C	100°C	Most laboratory work
Kelvin	K	273.15 K	373.15 K	Gas laws, thermodynamics; SI unit
Fahrenheit	°F	32°F	212°F	Everyday use in the United States; rarely used in chemistry calculations

Kelvin Conversion

K = °C + 273.15

The Kelvin scale begins at absolute zero — the theoretical temperature at which particle motion essentially stops. Absolute zero is −273.15°C. All Kelvin temperatures are positive numbers.

Section 20.3

Systems and Surroundings

To analyze energy changes clearly, chemists draw a conceptual boundary around the part of the universe they are studying. The system is the specific chemicals or objects being studied — the reactants and products of a reaction, for example. Everything outside the system is the surroundings: the reaction vessel, the air, the lab bench, the water in a calorimeter — essentially everything else.

This distinction is not just bookkeeping. It allows chemists to track energy flow precisely: energy leaving the system enters the surroundings, and energy entering the system comes from the surroundings.

System Type	Exchanges Matter?	Exchanges Energy?	Example
Open system	Yes	Yes	Pot of boiling soup — steam escapes, heat flows to room
Closed system	No	Yes	Sealed flask being heated
Isolated system	No	No	Perfectly insulated thermos (theoretical ideal)

Sign Convention for Energy Transfer

Chemists use the system's perspective to assign signs to energy changes:

Positive (+) — energy flows into the system. The system gains energy.

Negative (−) — energy flows out of the system. The system loses energy.

If a reaction releases heat into the surroundings, the reaction system loses energy: ΔH is negative.

Section 20.4

Exothermic and Endothermic Processes

Exothermic Processes

A process is exothermic if it releases energy to the surroundings. In an exothermic process, the system loses energy, and that energy flows into the surroundings as heat. The result: the surroundings get warmer. If you touch a flask where an exothermic reaction is occurring, you will feel it heat up.

In an exothermic process, the chemical potential energy stored in the bonds of the reactants is greater than the energy stored in the bonds of the products. The excess energy is released as heat.

Exothermic Process	What Releases Energy	Where You Encounter It
Combustion of fuels	Breaking C–H and C–C bonds; forming CO₂ and H₂O bonds	Stoves, furnaces, engines, campfires
Rusting of iron	Slow oxidation of iron by oxygen	Corroding metal surfaces
Neutralization reactions	Acid + base → salt + water	Antacids, industrial pH control
Condensation of water vapor	Gas → liquid phase transition	Steam burns, dew formation
Cellular respiration	Breakdown of glucose in living cells	Body heat, muscle energy
Dissolving calcium chloride	Ion–water interactions release energy	Instant hand warmers

Endothermic Processes

A process is endothermic if it absorbs energy from the surroundings. In an endothermic process, the system gains energy, and the surroundings lose energy. The result: the surroundings get cooler. The flask will feel cold to the touch.

In an endothermic process, the chemical potential energy stored in the products is greater than in the reactants. The extra energy must come from the surroundings.

Endothermic Process	What Absorbs Energy	Where You Encounter It
Melting of ice	Breaking hydrogen bonds in ice crystal lattice	Ice packs, cooling drinks
Evaporation of water	Breaking intermolecular attractions in liquid	Sweating, evaporative cooling
Photosynthesis	Light energy absorbed to convert CO₂ + H₂O → glucose	All green plants
Dissolving ammonium nitrate	Breaking ion–ion attractions requires more energy than ion–water interactions release	Instant cold packs
Cooking an egg	Protein denaturation requires continuous energy	Cooking and baking
Decomposition reactions	Breaking stable bonds requires energy input	Electrolysis of water, thermal decomposition

Comparing Exothermic and Endothermic

Feature	Exothermic	Endothermic
Energy flow direction	System → Surroundings	Surroundings → System
Temperature of surroundings	Increases (surroundings feel warm)	Decreases (surroundings feel cool)
Sign of system's energy change (ΔH)	Negative (−)	Positive (+)
Products vs. reactant energy	Products have less energy	Products have more energy
Common examples	Combustion, condensation, rusting, neutralization	Melting, evaporation, photosynthesis, dissolving NH₄NO₃

Common Misconception

Students often think that endothermic processes are "cold" and that exothermic processes are "hot." The temperature change you observe is in the surroundings, not the system itself. What defines a process as exo- or endothermic is the direction of energy flow. A reaction can be endothermic even at very high temperatures — what matters is whether it absorbs energy from its surroundings as it proceeds.

Practice 20.4Concept Check

Classify each process as exothermic or endothermic and explain your reasoning: (a) boiling water on a stove, (b) burning paper, (c) an ice cube melting in your hand, (d) a hand warmer activating in your pocket, (e) evaporation of rubbing alcohol from your skin.
A student mixes two chemicals in a beaker and notices the beaker becomes warm. Is the reaction exothermic or endothermic? In which direction did energy flow?
Compare the energy stored in the bonds of the reactants versus the products in an exothermic reaction. Which is greater? Where does the "extra" energy go?
In cellular respiration, your body breaks down glucose and releases energy. Is this process exothermic or endothermic from the body's perspective? What happens to the released energy?

Section 20.5

Specific Heat

Not all substances warm up at the same rate when you add heat. A metal spoon left in a pot of hot soup heats up very quickly and can burn your hand. The soup itself takes much longer to reach the same temperature, even though the spoon and soup are in contact with the same heat source. This difference arises because different materials have different specific heats.

The specific heat capacity (usually called specific heat, written c) of a substance is the amount of heat required to raise the temperature of one gram of that substance by one degree Celsius (or one Kelvin). Specific heat is measured in:

c = J / (g · °C)

Why Water Is Special

Water has an unusually high specific heat: 4.184 J/g·°C. It takes more than ten times as much energy to warm water as to warm the same mass of gold. This is why oceans and large lakes moderate climate — they absorb enormous amounts of heat without changing temperature dramatically. It is also why water is used as a coolant in car engines and industrial reactors.

Selected Specific Heat Values

Substance	Specific Heat (J/g·°C)
Water (liquid)	4.184
Ice	2.090
Steam	2.010
Ethanol	2.440
Aluminum	0.897
Iron	0.449
Copper	0.385
Gold	0.128
Granite	0.790
Glass	0.840

The Specific Heat Equation

q = mcΔT

q = heat transferred (joules, J)

m = mass of the substance (grams, g)

c = specific heat capacity (J/g·°C)

ΔT = change in temperature = T_final − T_initial (°C)

If ΔT is positive, the substance was heated and q is positive (gained heat). If ΔT is negative, the substance was cooled and q is negative (lost heat).

Worked Example 20.1

Calculating Heat Absorbed by Water

ProblemHow much heat is required to warm 250.0 g of water from 20.0°C to 85.0°C?
Givenm = 250.0 g | c = 4.184 J/g·°C (water) | T_initial = 20.0°C | T_final = 85.0°C
Step 1 ΔT = T_final − T_initial = 85.0°C − 20.0°C = 65.0°C
Step 2 q = mcΔT q = (250.0 g)(4.184 J/g·°C)(65.0°C) q = 68,000 J = 68.0 kJ
CheckHeating 250 g of water through 65°C is a large change — 68 kJ is reasonable. The sign is positive because the water is gaining heat. ✓

Answer: 68.0 kJ of heat must be added to warm the water.

Worked Example 20.2

Calculating Final Temperature

ProblemA 150.0 g piece of iron (c = 0.449 J/g·°C) at 25.0°C absorbs 2,500 J of heat. What is its final temperature?
Givenm = 150.0 g | c = 0.449 J/g·°C | T_initial = 25.0°C | q = 2,500 J
Step 1Solve for ΔT by rearranging q = mcΔT: ΔT = q / (mc) = 2500 J / [(150.0 g)(0.449 J/g·°C)] ΔT = 2500 / 67.35 = 37.1°C
Step 2 T_final = T_initial + ΔT = 25.0°C + 37.1°C = 62.1°C
CheckIron has a low specific heat, so 2,500 J raises it by 37°C — much more than water would be raised by the same heat. Reasonable. ✓

Answer: The iron reaches a final temperature of 62.1°C.

Worked Example 20.3

Finding Specific Heat from Experimental Data

ProblemA 200.0 g sample of an unknown metal is heated to 95.0°C and placed in 100.0 g of water at 22.0°C. The final temperature of both metal and water is 31.4°C. What is the specific heat of the metal? (Assume no heat loss to surroundings.)
ReasoningThe hot metal loses heat; the cool water gains heat. Since no heat escapes: q_metal = −q_water
Water q_water = m_w · c_w · ΔT_w = (100.0 g)(4.184 J/g·°C)(31.4 − 22.0)°C = (100.0)(4.184)(9.4) = 3,933 J
Metal q_metal = −3,933 J −3,933 = (200.0)(c_metal)(31.4 − 95.0) −3,933 = (200.0)(c_metal)(−63.6) c_metal = −3,933 / −12,720 = 0.309 J/g·°C
CheckThe value 0.309 J/g·°C falls between copper (0.385) and gold (0.128). This is close to the specific heat of several metals. More precise measurements would allow identification. ✓

Answer: The specific heat of the metal is approximately 0.309 J/g·°C.

Practice 20.5Specific Heat Calculations

How much heat is needed to warm 500.0 g of water from 25.0°C to 95.0°C?
A 75.0 g piece of copper (c = 0.385 J/g·°C) absorbs 1,200 J of heat. If the initial temperature is 20.0°C, what is the final temperature?
How many grams of iron (c = 0.449 J/g·°C) would be needed to absorb 5,000 J of heat while changing temperature from 22°C to 85°C?
A 100.0 g sample of an unknown liquid cools from 80.0°C to 30.0°C and releases 12,500 J of heat. Calculate the specific heat of the liquid. What common liquid might this be?
Which requires more heat — warming 100 g of water from 20°C to 90°C, or warming 100 g of aluminum (c = 0.897 J/g·°C) through the same temperature range? How many times more?
A metal block of mass 250 g absorbs 4,500 J of heat, raising its temperature from 15°C to 55°C. Find the specific heat and identify the likely metal using the table above.

Section 20.6

Calorimetry

A calorimeter is a device used to measure heat transfer during chemical or physical processes. In calorimetry, the heat released or absorbed by a process is measured by observing the temperature change of a surrounding substance — usually water or aqueous solution.

The central assumption in calorimetry is that all the heat released by the reaction is absorbed by the calorimeter contents, and no heat is lost to the outside environment. Real calorimeters are imperfect insulators, but simple coffee cup calorimeters used in high school chemistry give good approximate measurements.

The Coffee Cup Calorimeter

The simplest educational calorimeter is two nested polystyrene (foam) cups with a loose lid and a thermometer. Polystyrene is a good heat insulator. This design is used for reactions in aqueous solution — dissolving a salt, neutralizing an acid with a base, or running a simple double-replacement reaction. These reactions occur at constant atmospheric pressure, making this type of calorimetry compatible with measuring enthalpy changes.

Coffee Cup Calorimeter Procedure

1. Measure the initial temperature of the solution.

2. Add the second reactant or solute and stir carefully.

3. Record the maximum (or minimum) temperature reached.

4. Calculate: q_solution = mcΔT (using total mass of solution)

5. Since the solution gains what the reaction releases: q_rxn = −q_solution

Worked Example 20.4

Coffee Cup Calorimetry — Neutralization

Problem50.0 mL of 1.0 M HCl and 50.0 mL of 1.0 M NaOH are mixed in a coffee cup calorimeter. Both solutions start at 22.5°C. After mixing, the temperature rises to 29.3°C. Density = 1.00 g/mL; c = 4.18 J/g·°C. Calculate the heat released by the neutralization reaction.
GivenTotal volume = 100.0 mL | m = 100.0 g | c = 4.18 J/g·°C | T_i = 22.5°C | T_f = 29.3°C
Step 1 ΔT = 29.3°C − 22.5°C = 6.8°C
Step 2 q_solution = mcΔT = (100.0 g)(4.18 J/g·°C)(6.8°C) = 2,842 J ≈ 2.84 kJ
Step 3The solution temperature increased — the solution absorbed heat from the reaction. Therefore the reaction released heat. q_rxn = −q_solution = −2.84 kJ
InterpretThe negative sign confirms the reaction is exothermic. The neutralization of HCl and NaOH releases energy to the solution. ✓

Answer: q_rxn = −2.84 kJ. The reaction is exothermic.

The Bomb Calorimeter

For combustion reactions and other high-energy processes, chemists use a bomb calorimeter (constant-volume calorimeter). The reaction occurs inside a sealed, heavy-walled metal container (the "bomb") submerged in a known mass of water. Bomb calorimeters can withstand the high pressures generated by combustion. They are used to measure the caloric content of fuels and foods — the kilocalorie values on nutrition labels were originally determined by burning food samples in bomb calorimeters.

The calculation uses the heat capacity of the calorimeter (C_cal, in J/°C or kJ/°C), determined by calibration with a substance of known energy content:

q_rxn = −C_cal × ΔT

Practice 20.6Calorimetry Calculations

100.0 mL of 0.500 M HNO₃ is mixed with 100.0 mL of 0.500 M KOH in a coffee cup calorimeter. The temperature rises from 23.0°C to 26.8°C. Using m = 200.0 g and c = 4.18 J/g·°C, calculate the heat released by the reaction.
A student dissolves 10.0 g of NaOH in 100.0 mL of water (m = 100.0 g, c = 4.18 J/g·°C) in a coffee cup calorimeter. The temperature rises from 21.5°C to 40.8°C. Calculate q_rxn. Is this process exothermic or endothermic?
An instant cold pack contains ammonium nitrate and water. When activated, the temperature drops from 25.0°C to 9.0°C. If the pack contains 200.0 g of solution with c = 4.18 J/g·°C, how much heat is absorbed by the dissolution?
A bomb calorimeter has a heat capacity of 8.50 kJ/°C. When 1.00 g of glucose (C₆H₁₂O₆) is burned inside, the temperature rises from 25.00°C to 27.55°C. Calculate the heat of combustion per gram of glucose.

Section 20.7

Enthalpy Change

In chemistry, most reactions occur in open containers at constant atmospheric pressure. Chemists use a quantity called enthalpy (H) to describe the heat content of a system at constant pressure. Enthalpy itself cannot be directly measured, but the change in enthalpy (ΔH) can be measured through calorimetry:

Enthalpy Change

ΔH = H_products − H_reactants

ΔH negative (−) → Products have less enthalpy than reactants → heat was released → exothermic

ΔH positive (+) → Products have more enthalpy than reactants → heat was absorbed → endothermic

Units: kJ or kJ/mol (kilojoules per mole of reaction)

Thermochemical Equations

A thermochemical equation is a balanced chemical equation that includes the enthalpy change for the reaction. The ΔH value is written after the equation.

Reaction	Thermochemical Equation	Interpretation
Combustion of methane	`CH₄(g) + 2O₂(g) → CO₂(g) + 2H₂O(l) ΔH = −890 kJ`	Exothermic; 890 kJ released when 1 mol CH₄ burns
Decomposition of water	`2H₂O(l) → 2H₂(g) + O₂(g) ΔH = +572 kJ`	Endothermic; 572 kJ must be supplied to decompose 2 mol H₂O
Dissolving ammonium nitrate	`NH₄NO₃(s) → NH₄⁺(aq) + NO₃⁻(aq) ΔH = +25.7 kJ`	Endothermic; solution cools; used in instant cold packs
Formation of water from elements	`2H₂(g) + O₂(g) → 2H₂O(l) ΔH = −572 kJ`	Exothermic; reverse of decomposition

Enthalpy as a State Function

Enthalpy is a state function — its value depends only on the current state of the system, not on the path taken to reach that state. This means if a reaction occurs by two different routes ending at the same products, the total ΔH is the same. This property (known as Hess's Law) allows chemists to calculate ΔH for reactions that are difficult to measure directly by combining ΔH values of reactions that can be measured.

Standard Enthalpy of Formation

The standard enthalpy of formation (ΔH°_f) of a compound is the enthalpy change when one mole of that compound is formed from its elements in their standard states at 25°C and 1 atm. By convention, ΔH°_f for any element in its standard state is zero.

Compound	ΔH°_f (kJ/mol)
H₂O(l)	−285.8
H₂O(g)	−241.8
CO₂(g)	−393.5
NaCl(s)	−411.2
NH₃(g)	−46.1
CH₄(g)	−74.8

Practice 20.7Enthalpy and Thermochemical Equations

For the reaction CH₄(g) + 2O₂(g) → CO₂(g) + 2H₂O(l), ΔH = −890 kJ. (a) Is this exothermic or endothermic? (b) How much energy is released when 2 moles of methane burn? (c) What would ΔH be for the reverse reaction?
Write a thermochemical equation for the combustion of propane: C₃H₈(g) + 5O₂(g) → 3CO₂(g) + 4H₂O(l), ΔH = −2220 kJ. Interpret what the sign and value of ΔH tell you.
A reaction has ΔH = +350 kJ. (a) Is the reaction exothermic or endothermic? (b) Do products have more or less enthalpy than reactants? (c) Does the temperature of the surroundings increase or decrease as the reaction proceeds?
Explain why the standard enthalpy of formation of O₂(g) is defined as zero, while the standard enthalpy of formation of H₂O(l) is −285.8 kJ/mol.

Section 20.8

Energy Diagrams

An energy diagram (also called an enthalpy diagram or reaction coordinate diagram) is a graph that shows how the energy of a system changes as a reaction proceeds. The horizontal axis represents the reaction coordinate — the progress of the reaction from reactants to products. The vertical axis represents potential energy or enthalpy.

Features of an Energy Diagram

Feature	Location on Diagram	What It Tells You
Reactants	Left side, at a certain energy level	Starting energy of the system
Products	Right side, at a different energy level	Ending energy of the system
Transition state	Peak of the energy curve	Highest-energy arrangement of atoms during the reaction; exists momentarily
Activation energy (E_a)	Height from reactant level to peak	Minimum energy particles must have to react; determines reaction speed
ΔH	Difference between reactant and product levels	Overall energy change; negative = exothermic, positive = endothermic

Exothermic Reaction Diagram

In an exothermic reaction, the products are at a lower energy level than the reactants. The curve rises (activation energy hump) and then drops to a lower level. The difference in height between reactants and products is |ΔH|, and that energy is released to the surroundings.

Energy (kJ)
    |
    |        /\          ← transition state (peak)
    |       /  \
    |______/    \________  ← products (lower than reactants)
    |↑ reactants
    |
    |__________________________
          Reaction Progress →

    ΔH is NEGATIVE (energy released)
    E_a = height from reactants to peak
    Products are LOWER than reactants

Endothermic Reaction Diagram

In an endothermic reaction, the products are at a higher energy level than the reactants. The curve still shows an activation energy hump, but the final product level is higher than the starting reactant level. Energy must flow in from the surroundings for the reaction to proceed.

Energy (kJ)
    |               /\       ← transition state (peak)
    |              /  \_____ ← products (HIGHER than reactants)
    |             /
    |____________/           ← reactants
    |
    |__________________________
          Reaction Progress →

    ΔH is POSITIVE (energy absorbed)
    E_a = height from reactants to peak
    Products are HIGHER than reactants

Reading an Energy Diagram — Four Questions

1. Exo or endo? Products lower than reactants → exothermic. Products higher → endothermic.

2. How large is ΔH? Vertical distance between reactant and product energy levels.

3. How large is E_a? Vertical distance from reactant level to peak. Large E_a → slow reaction.

4. Effect of a catalyst? A catalyst lowers the peak (reduces E_a) without changing the reactant or product energy levels. ΔH is unchanged; only the height of the barrier changes.

Practice 20.8Energy Diagrams

Draw and label energy diagrams for each of the following, indicating whether each is exothermic or endothermic and showing the location of E_a and ΔH: (a) combustion of methane (ΔH = −890 kJ); (b) decomposition of water (ΔH = +572 kJ).
A reaction has an activation energy of 80 kJ and releases 50 kJ overall. (a) Is this reaction exothermic or endothermic? (b) What is the activation energy for the reverse reaction? (c) Sketch a labeled energy diagram.
Two reactions have the same ΔH = −200 kJ. Reaction A has E_a = 30 kJ; Reaction B has E_a = 150 kJ. Which reaction proceeds faster at room temperature? Which would more strongly benefit from a catalyst?
A catalyst is added to a reaction. On the same axes, sketch the energy diagram before and after the catalyst is added. What changes and what stays the same?
In an energy diagram, the reactants are at 400 kJ, the transition state is at 550 kJ, and the products are at 250 kJ. Calculate: (a) E_a for the forward reaction; (b) E_a for the reverse reaction; (c) ΔH for the forward reaction; (d) Is the forward reaction exothermic or endothermic?

Section 20.9

Thermochemistry in Fuels, Food, and Industry

Fuels and Combustion Energy

Fuels are substances that release large amounts of energy through combustion. The energy density of a fuel — the amount of energy released per gram — determines how useful it is. All complete combustion reactions of hydrocarbon fuels (containing only carbon and hydrogen) produce carbon dioxide and water:

fuel + O₂(g) → CO₂(g) + H₂O + energy

Incomplete combustion — when oxygen supply is insufficient — also produces carbon monoxide (CO) and soot (carbon particles), both of which are harmful to human health and the environment.

Fuel	Energy per Gram (kJ/g)	Notes
Hydrogen gas (H₂)	142	Highest energy density by mass; used in space rockets and fuel cells
Natural gas (methane, CH₄)	55	Home heating, cooking; burns cleanly
Gasoline (≈ C₈H₁₈)	48	Dominant transportation fuel; easy to store and transport
Ethanol (C₂H₅OH)	30	Biofuel; mixed with gasoline
Coal	24–35	Varies by type; significant CO₂ and pollutant emissions
Wood (cellulose)	17	Renewable but releases CO₂; particulate emissions

Why Hydrogen Is Attractive — and Challenging

Hydrogen has the highest energy density by mass of any fuel. This makes it attractive for space exploration and fuel cell vehicles. However, hydrogen is a gas at room temperature, making it difficult to store and transport safely in large quantities. Gasoline and natural gas are much easier to handle, which is why they dominate transportation and heating despite their environmental consequences.

Food as Chemical Fuel

Your body runs on the same thermochemical principles as an engine, just at lower temperatures and with far more sophisticated control. Food contains chemical potential energy stored in the bonds of carbohydrates, fats, and proteins. When you digest and metabolize food, your cells break these molecules down through cellular respiration — releasing energy to power muscle contractions, nerve signals, protein synthesis, and temperature regulation.

The energy content of food is measured in kilocalories (kcal), called "Calories" (capital C) on nutrition labels. These values are determined by burning food samples in bomb calorimeters.

Nutrient	Energy per Gram	Chemical Storage Form
Carbohydrates	4 kcal/g (≈17 kJ/g)	Sugars and starches; glucose is the primary cellular fuel
Proteins	4 kcal/g (≈17 kJ/g)	Amino acid chains; used primarily for building, not fuel
Fats (lipids)	9 kcal/g (≈38 kJ/g)	Long-chain fatty acids; most efficient long-term energy storage
Alcohol (ethanol)	7 kcal/g (≈29 kJ/g)	Metabolized primarily by the liver

Fats store more than twice as much energy per gram as carbohydrates or proteins. This is why fat is an efficient long-term energy storage molecule — you can store a lot of energy in relatively little mass. Your basal metabolic rate is the minimum energy your body needs to maintain basic functions at rest — roughly 1,500–2,000 kcal per day for an average adult. Physical activity requires additional energy beyond this baseline.

Industrial Thermochemistry

Industrial chemistry depends heavily on thermochemical principles. Designing a manufacturing process requires knowing not just whether a reaction occurs, but how much energy it requires or releases. Energy costs are often the largest expense in industrial chemical production.

Industrial Process	Thermochemical Details	Significance
Haber process (making NH₃)	N₂(g) + 3H₂(g) → 2NH₃(g) ΔH = −92 kJ Exothermic; requires high temp, pressure, and Fe catalyst	Basis for most nitrogen fertilizers; feeds roughly half the world's population
Steel production	Reduction of iron ore (iron oxides) is endothermic; requires coke as reducing agent and energy source	Thermochemistry determines cost and carbon footprint; driving development of "green steel"
Refrigeration and air conditioning	Refrigerant fluids evaporate endothermically (absorbing heat from interior) and condense exothermically outside	Thermochemical properties of refrigerant — heat of vaporization, boiling point — determine effectiveness

Energy Storage and Climate

As society moves toward renewable energy sources, energy storage becomes critically important — the sun doesn't always shine and the wind doesn't always blow. Thermochemistry is central to developing better batteries, hydrogen fuel systems, and thermal storage materials.

Combustion of fossil fuels releases CO₂ that has been stored in the earth for millions of years, returning it to the atmosphere. The thermochemistry of CO₂ in the atmosphere — how it absorbs and re-emits infrared radiation — is the chemical basis of the greenhouse effect and climate change. Understanding thermochemistry does not resolve political debates about energy policy, but it gives you the chemical literacy to evaluate evidence, understand tradeoffs, and think clearly about real constraints and possibilities.

Practice 20.9Real-World Applications

Hand warmers and cold packs: Explain the thermochemistry behind both types of instant products. Which involves an exothermic process? Which is endothermic? How does each affect the temperature of its surroundings?
Nutrition label calculation: A food item contains 15 g of fat, 30 g of carbohydrate, and 10 g of protein. Using the values above, estimate the total caloric content in kcal and in kJ.
Climate and water: The Pacific and Atlantic Oceans store vast amounts of heat due to water's high specific heat. Explain why coastal cities tend to have milder temperatures than inland cities at the same latitude.
Fuel comparison: A car engine burns approximately 8.0 g of gasoline per mile. Using the energy density of gasoline (48 kJ/g), how many kilojoules are released per mile? How does this compare to a hydrogen fuel cell car that uses 0.05 g of hydrogen per mile at 142 kJ/g?
Industrial chemistry: The contact process for making sulfuric acid involves the highly exothermic oxidation of SO₂ to SO₃. Why would industrial engineers need to carefully manage heat removal in this process? What are the risks of allowing too much heat to accumulate?
Body temperature: Your body maintains a core temperature of 37°C (98.6°F). If the surroundings are at 20°C, in what direction does heat flow? Is your body a net heat producer for the environment?

Chapter 20 — Back Matter

Chapter Summary

Sec 20.1Energy is the capacity to do work or transfer heat. Chemical reactions always involve energy changes because bonds break and form. Kinetic energy is energy of motion; chemical potential energy is stored in bonds; thermal energy is the total kinetic energy of all particles in a sample. The law of conservation of energy states that energy cannot be created or destroyed.
Sec 20.2Temperature measures the average kinetic energy of particles (intensive). Heat is the transfer of thermal energy between objects at different temperatures (extensive). Heat is measured in joules (J) or kilojoules (kJ). The Kelvin scale begins at absolute zero: K = °C + 273.15. 1 cal = 4.184 J; 1 kcal = 1,000 cal (the "Calorie" on nutrition labels).
Sec 20.3The system is the chemicals being studied; the surroundings are everything else. Energy flowing into the system is positive (+); energy flowing out is negative (−). Open systems exchange both matter and energy; closed systems exchange only energy; isolated systems exchange neither.
Sec 20.4Exothermic processes release energy to the surroundings (surroundings get warmer; ΔH negative). Examples include combustion, rusting, neutralization, condensation, and cellular respiration. Endothermic processes absorb energy from the surroundings (surroundings get cooler; ΔH positive). Examples include melting, evaporation, photosynthesis, and dissolving ammonium nitrate.
Sec 20.5Specific heat capacity (c) is the heat required to raise the temperature of one gram of a substance by one degree Celsius. Water's high specific heat (4.184 J/g·°C) makes it an exceptional coolant and climate regulator. The equation q = mcΔT relates heat (q), mass (m), specific heat (c), and temperature change (ΔT = T_final − T_initial).
Sec 20.6Calorimetry measures heat transfer by observing temperature changes in a surrounding substance. A coffee cup calorimeter is used for aqueous reactions at constant pressure. A bomb calorimeter measures combustion reactions at constant volume. In both cases: q_rxn = −q_calorimeter, because heat released by the reaction equals heat absorbed by the calorimeter contents.
Sec 20.7Enthalpy (H) is the heat content of a system at constant pressure. ΔH = H_products − H_reactants: negative for exothermic, positive for endothermic. Thermochemical equations include ΔH values. Enthalpy is a state function (path-independent). Standard enthalpies of formation (ΔH°_f) provide a reference for calculating ΔH for any reaction; ΔH°_f = 0 for elements in their standard states.
Sec 20.8Energy diagrams show how system energy changes during a reaction. Key features: reactants and products at different energy levels, activation energy barrier (E_a), and transition state at the peak. Exothermic diagrams show products lower than reactants; endothermic diagrams show products higher. Catalysts lower E_a without changing ΔH.
Sec 20.9Fuels release energy through combustion; energy density (kJ/g) compares their usefulness. Food is chemical fuel: carbohydrates and proteins supply 4 kcal/g; fats supply 9 kcal/g. Industrial thermochemistry affects production costs and design (Haber process, steelmaking, refrigeration). Combustion of fossil fuels releases CO₂, driving climate change; thermochemistry underlies both the problem and potential solutions.

Formula Review

Formula	Variables	When to Use
`q = mcΔT`	q = heat (J); m = mass (g); c = specific heat (J/g·°C); ΔT = T_final − T_initial (°C)	Calculating heat absorbed or released by a substance
`ΔT = T_final − T_initial`	temperatures in °C or K	Finding temperature change (always final minus initial)
`K = °C + 273.15`	K = Kelvin; °C = Celsius	Converting between temperature scales
`ΔH = H_products − H_reactants`	H = enthalpy (kJ)	Defining enthalpy change; negative = exothermic
`q_rxn = −q_solution`	q = heat (J or kJ)	Calorimetry: heat released by reaction equals heat gained by solution, opposite signs
`1 cal = 4.184 J`		Converting between calories and joules

Worked Example Review

Example A

Basic specific heat calculation

ProblemA 50.0 g piece of aluminum (c = 0.897 J/g·°C) is warmed from 15.0°C to 65.0°C. Find q.
Solveq = mcΔT = (50.0 g)(0.897 J/g·°C)(65.0 − 15.0)°C = (50.0)(0.897)(50.0) = 2,243 J = 2.24 kJ

Answer: q = 2.24 kJ (positive: aluminum gained heat)

Example B

Calorimetry calculation

ProblemA reaction in a coffee cup calorimeter raises the temperature of 200.0 g of solution from 21.3°C to 28.7°C. Given c = 4.18 J/g·°C, find q_rxn.
Solveq_solution = (200.0)(4.18)(28.7 − 21.3) = (200.0)(4.18)(7.4) = 6,187 J q_rxn = −6,187 J = −6.19 kJ

Answer: q_rxn = −6.19 kJ (exothermic: negative sign; reaction released heat to solution)

Example C

Interpreting a thermochemical equation

EquationC₃H₈(g) + 5O₂(g) → 3CO₂(g) + 4H₂O(l) ΔH = −2220 kJ
InterpretExothermic (ΔH < 0). 2220 kJ of heat is released when one mole of propane burns completely. Products have less enthalpy than reactants. In an energy diagram, products appear at a lower energy level than reactants.

The combustion of propane is strongly exothermic — this is why propane is used as a fuel for grills and heaters.

Key Terms Review

activation energy

The minimum energy that colliding particles must have for a reaction to occur; the height of the energy barrier in a reaction coordinate diagram.

calorimeter

A device used to measure the heat transferred during a chemical or physical process.

calorimetry

The experimental technique of measuring heat flow by observing temperature changes in a calorimeter.

chemical potential energy

Energy stored in chemical bonds and in the arrangement of atoms within molecules; changes when bonds break and form.

endothermic process

A process that absorbs heat from the surroundings; the surroundings get cooler; ΔH is positive.

energy

The capacity to do work or transfer heat; measured in joules (J) or kilojoules (kJ). Cannot be created or destroyed.

energy density

The amount of energy stored per unit mass of a fuel (kJ/g); used to compare fuel efficiency.

energy diagram

A graph showing how the potential energy or enthalpy of a system changes as a reaction proceeds from reactants to products; also called a reaction coordinate diagram.

enthalpy (H)

A thermodynamic quantity representing the heat content of a system at constant pressure.

enthalpy change (ΔH)

ΔH = H_products − H_reactants. Negative for exothermic reactions; positive for endothermic reactions. Units: kJ or kJ/mol.

exothermic process

A process that releases heat to the surroundings; the surroundings get warmer; ΔH is negative.

heat (q)

The transfer of thermal energy between objects at different temperatures; always flows from warmer to cooler objects.

joule (J)

The SI unit of energy. 1 kJ = 1,000 J. 1 cal = 4.184 J.

kilocalorie (kcal)

1,000 calories; the unit used on nutrition labels (written "Calorie" with capital C). 1 kcal = 4,184 J.

kinetic energy

Energy of motion; at the particle level, related to temperature (faster particles = higher temperature).

law of conservation of energy

Energy cannot be created or destroyed; it can only be converted between forms or transferred between objects. Also called the first law of thermodynamics.

potential energy

Stored energy associated with position or arrangement; in chemistry, primarily stored in chemical bonds.

specific heat capacity (c)

The amount of heat required to raise the temperature of one gram of a substance by one degree Celsius; measured in J/g·°C. Water's value is 4.184 J/g·°C.

standard enthalpy of formation

ΔH°_f: enthalpy change when one mole of a compound forms from its elements in their standard states at 25°C and 1 atm. Zero for elements in their standard states.

surroundings

Everything outside the defined system; typically the container, solution, calorimeter, or environment around the reaction.

system

The specific part of the universe being studied; typically the chemicals or reaction of interest.

temperature

A measure of the average kinetic energy of particles in a substance; an intensive property measured in °C, K, or °F.

thermal energy

The total kinetic energy of all particles in a sample; depends on both temperature and amount of matter.

thermochemical equation

A balanced chemical equation that includes the enthalpy change (ΔH) for the reaction.

thermochemistry

The branch of chemistry that studies the energy changes associated with chemical and physical processes.

transition state

The highest-energy arrangement of atoms reached during a reaction; corresponds to the peak of the activation energy barrier in an energy diagram.

Concept Check

What is the difference between heat and temperature? Give an example that shows they are not the same thing.
Define system and surroundings. In a reaction where iron rusts in a sealed jar, identify the system and the surroundings.
A reaction causes a temperature increase in the calorimeter fluid. Is the reaction exothermic or endothermic? Explain using the concept of energy flow.
Two substances have the same mass and are given the same amount of heat. Substance A heats up twice as fast as Substance B. Which substance has the higher specific heat? Explain.
Sketch a labeled energy diagram for an endothermic reaction. Label reactants, products, transition state, activation energy (E_a), and ΔH.
The enthalpy of combustion of ethanol is −1366 kJ/mol. Write the thermochemical equation for complete combustion of ethanol (C₂H₅OH). Is this exothermic or endothermic?
A coffee cup calorimeter experiment shows the solution temperature decreasing during a reaction. What does this tell you about the reaction?

Error Analysis

Each problem below contains an error. Identify it and write the correct answer or explanation.

Error 1

"When heat flows from the surroundings into a reaction system, the temperature of the surroundings increases."

What is wrong with this statement? What actually happens to the temperature of the surroundings when heat flows from them into the system?

Error 2

A student reports: "The temperature of the calorimeter solution fell from 25°C to 18°C, so ΔT = +7°C."

Identify the error in how ΔT was calculated. What is the correct value, and how does this error affect the sign of q?

Error 3

"An exothermic reaction has a positive ΔH because it produces heat."

What is wrong with this statement? What is the correct sign convention for an exothermic reaction and why?

Error 4

A student calculates specific heat as 4.19 J/g·°C after measuring a temperature change in 100 g of metal when 2,095 J were added, then identifies the metal as "water."

What is wrong with this reasoning? Can a metal have the same specific heat as water? What would be a more appropriate identification step?

Error 5

"In a coffee cup calorimeter, q_rxn and q_solution will have the same sign."

Explain why this is incorrect. What is the correct relationship between q_rxn and q_solution, and what is the sign convention?

Error 6

"Absolute zero is 0°C."

What is the correct value of absolute zero in °C and in K? What does absolute zero represent at the particle level?

Mixed Review

These problems combine multiple skills from this chapter. Show all work and include units in every step.

Define thermal energy and explain how it differs from temperature. Give an example showing that a low-temperature object can have more thermal energy than a high-temperature object.
A 100.0 g sample of aluminum (c = 0.897 J/g·°C) and a 100.0 g sample of copper (c = 0.385 J/g·°C) each receive 500 J of heat starting from 25.0°C. Calculate the final temperature of each. Which heats up more? Explain using specific heat.
Consider the thermochemical equation: 2H₂(g) + O₂(g) → 2H₂O(l) ΔH = −572 kJ. (a) Is this exothermic or endothermic? (b) How much energy is released when 1 mole of H₂O(l) is formed? (c) What is ΔH for the reverse reaction? (d) Where is this energy stored in the reactants, and where does it go when the reaction occurs?
A student mixes 50.0 mL of 2.0 M HCl with 50.0 mL of 2.0 M NaOH in a coffee cup calorimeter. The temperature rises from 21.0°C to 34.5°C. Given m = 100.0 g and c = 4.18 J/g·°C, calculate: (a) q_solution; (b) q_rxn; (c) how many moles of water were formed; (d) the heat released per mole of water formed (in kJ/mol).
In an energy diagram for a reaction, the reactants are at 400 kJ, the transition state is at 550 kJ, and the products are at 250 kJ. Calculate: (a) E_a for the forward reaction; (b) E_a for the reverse reaction; (c) ΔH for the forward reaction; (d) whether the forward reaction is exothermic or endothermic.
A nutrition label lists: fat 12 g, carbohydrate 45 g, protein 8 g. Calculate the total energy content in (a) kcal and (b) kJ. Show your work.
A bomb calorimeter (C_cal = 9.20 kJ/°C) is used to measure the heat of combustion of ethanol (C₂H₅OH, molar mass = 46.07 g/mol). When 2.30 g of ethanol is burned, the temperature rises from 24.60°C to 31.25°C. Calculate: (a) heat released per gram of ethanol; (b) heat released per mole of ethanol; (c) write the thermochemical equation with this ΔH value.

Chapter 20 Checkpoint

Complete without assistance — assesses all key skills and concepts from Chapter 20

Section 1 — Vocabulary and Concepts

Explain the difference between heat and temperature using particle-level reasoning. Which is intensive and which is extensive?
Define system and surroundings. What does a positive ΔH mean in terms of energy flow?
What is the difference between exothermic and endothermic? Give one example of each and explain the direction of energy flow.
Define specific heat capacity. Why does water's high specific heat (4.184 J/g·°C) matter for Earth's climate?

Section 2 — Calculations

How much heat is required to raise the temperature of 350.0 g of water from 18.0°C to 75.0°C?
A 120.0 g piece of aluminum (c = 0.897 J/g·°C) at 88.0°C is placed in 200.0 g of water at 18.0°C. Assuming no heat loss, calculate the final temperature of the system.
In a coffee cup calorimeter experiment, a reaction raises the temperature of 150.0 g of solution from 22.1°C to 29.6°C. Using c = 4.18 J/g·°C, calculate q_rxn. Is the reaction exothermic or endothermic?
Convert: (a) 35.0°C to Kelvin; (b) 310 K to Celsius; (c) 2,500 cal to joules; (d) 84.0 kJ to kilocalories.

Section 3 — Interpretation

Write the thermochemical equation for the combustion of methane and label it as exothermic or endothermic: CH₄(g) + 2O₂(g) → CO₂(g) + 2H₂O(l), ΔH = −890 kJ.
A reaction has reactants at 300 kJ, a transition state at 420 kJ, and products at 180 kJ. (a) Calculate E_a and ΔH. (b) Is the reaction exothermic or endothermic? (c) Sketch a labeled energy diagram.

Section 4 — Applications

A food contains 20 g fat, 50 g carbohydrate, and 15 g protein. Calculate the total energy content in kcal.
Explain the Haber process thermochemically: why does it release heat, and what practical challenges does this create for industrial operation?

Section 5 — Error Analysis

A student writes: "Because the reaction is endothermic, the products have less energy than the reactants." Identify and correct this error.
A student calculates q = mcΔT = (200 g)(4.18 J/g·°C)(22°C) = 18,392 J for an experiment where temperature rose from 22°C to 35°C. What error did the student make? What is the correct answer?

Looking Ahead

Chapter 21: Reaction Rates

Chapter 20 established how much energy is released or absorbed in a chemical reaction. Chapter 21 asks a different question: how fast does that release happen? You already know from energy diagrams that reactions must overcome an activation energy barrier. In Chapter 21, you will learn how collision theory explains why not every particle collision leads to a reaction, what factors control how quickly reactions proceed — concentration, temperature, surface area, and catalysts — and how chemists describe and measure reaction rates. The energy diagrams you studied here will reappear as tools for understanding reaction pathways, and the concept of activation energy will become central to explaining why some reactions happen instantly while others take years. The thermochemistry of Chapter 20 and the kinetics of Chapter 21 together describe not just whether a reaction releases energy, but how quickly that energy is released — a distinction with enormous practical consequences in industry, medicine, and environmental chemistry.

MGU Library Connection

Appendix H of this textbook: Gas Law and Thermochemistry Formula Sheet — includes q = mcΔT, enthalpy equations, and unit conversions
Appendix B: Measurement and Laboratory Reference — temperature scales, unit conversion, significant figures
Chapter 19: Liquids, Solids, and Phase Changes — thermochemistry of phase transitions; heating and cooling curves
Chapter 21: Reaction Rates — activation energy, collision theory, energy diagrams revisited
Chapter 22: Chemical Equilibrium — Le Châtelier's principle and the effect of temperature on equilibrium; connects ΔH to equilibrium position
Chapter 27: Environmental and Industrial Chemistry — real-world thermochemistry in fuel combustion, climate, and industrial processes
MGU Dictionary entries: enthalpy, thermochemistry, calorimetry, specific heat, exothermic, endothermic, activation energy

Chapter 20:Thermochemistry

Purpose and Overview

Learning Objectives

Sections at a Glance

Energy in Chemical Systems

What Energy Is

Forms of Energy Relevant to Chemistry

Heat and Temperature

Units of Heat Energy

Temperature Scales

Systems and Surroundings

Exothermic and Endothermic Processes

Exothermic Processes

Endothermic Processes

Comparing Exothermic and Endothermic

Specific Heat

Selected Specific Heat Values

Calorimetry

The Coffee Cup Calorimeter

The Bomb Calorimeter

Enthalpy Change

Thermochemical Equations

Standard Enthalpy of Formation

Energy Diagrams

Features of an Energy Diagram

Exothermic Reaction Diagram

Endothermic Reaction Diagram

Thermochemistry in Fuels, Food, and Industry

Fuels and Combustion Energy

Food as Chemical Fuel

Industrial Thermochemistry

Energy Storage and Climate

Chapter Summary

Formula Review

Worked Example Review

Key Terms Review

Concept Check

Error Analysis

Mixed Review

Chapter 20 Checkpoint

Chapter 20:
Thermochemistry