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D. Energy Changes and Rates of Reaction · Energy Changes and Rates of Reaction

Energy Changes and Rates of Reaction: Unit Review

D1.1 D1.2 D2.1 D2.2 D2.3 D2.4 D2.5 D2.6 D2.7 D2.8 D3.1 D3.2 D3.3 D3.4 D3.5 D3.6 D3.7 A1.1 A1.2 A1.4 A1.5 A1.6 A1.8 A1.9 A1.10 A1.11 A1.12 A1.13 A2.1 A2.2About 40 minutes
Learning goals
  • I can calculate enthalpy changes using calorimetry, Hess's law, enthalpies of formation and bond energies.
  • I can explain reaction rates using collision theory, potential energy diagrams, catalysts and reaction mechanisms.
  • I can evaluate energy technologies and reaction conditions for efficiency and sustainability.

Words to know

enthalpy change (ΔH)
The heat absorbed or released by a reaction at constant pressure.
Hess's law
The enthalpy change of a reaction equals the sum of the enthalpy changes of the steps that add up to it.
activation energy (Ea)
The minimum energy colliding particles need for a reaction to occur.
rate-determining step
The slowest step in a reaction mechanism, which limits the overall rate.

Let's review!

This unit asked two big questions about every reaction: how much energy is released or absorbed, and how fast it happens. You compared energy changes in physical, chemical and nuclear processes; wrote thermochemical equations; measured heat with calorimeters using Q = mcΔT; and calculated enthalpy changes with Hess's law, standard enthalpies of formation and bond energies. You evaluated energy technologies and their effects on the environment. Then you measured reaction rates, explained them with collision theory and potential energy diagrams, saw how catalysts and enzymes work, ran your own rate inquiry, and studied reaction mechanisms and the rate-determining step. Finally you used all of it to make reactions more efficient and sustainable.

This review has a few questions from every lesson. Try each one, then press Show answer to check. The Chemical Equilibrium unit builds directly on this one: reactions that go both ways, and how temperature, pressure and concentration shift them.

How to use this review

  1. Work through the questions without notes first. Keep a list of the ones you miss.
  2. For each miss, find the lesson with the same title and redo its Your turn questions.
  3. Keep a one-page summary sheet of key equations: Q = mcΔT, ΔH = nΔH (molar), ΔH°rxn = ΣnΔH°f(products) − ΣnΔH°f(reactants), ΔH ≈ Σ(bonds broken) − Σ(bonds formed), rate = Δ(amount)/Δt, rate = k[A]ᵐ[B]ⁿ.
  4. Watch your signs: exothermic is negative, and reversing an equation reverses the sign of ΔH.
  5. Finish with the quiz and read every explanation, even for questions you got right.
ToolEquationUse it when
Heat transferq = mcΔT (c in J/(g·°C))a temperature change of water or another substance is measured
Calorimeter sign ruleq_reaction = −q_waterturning a temperature change into ΔH
Molar enthalpyΔH = nΔH_molarscaling ΔH to a given amount or mass
Hess's lawadd equations, add ΔHknown equations can be reversed or multiplied to match a target
Formation enthalpiesΔH° = ΣnΔH°f(products) − ΣnΔH°f(reactants)a table of ΔH°f values is available
Bond energiesΔH ≈ Σ(broken) − Σ(formed)estimating ΔH for a gas-phase reaction
Percentage error(experimental − accepted) ÷ accepted × 100%comparing a lab result with an accepted value
Rate lawrate = k[A]ᵐ[B]ⁿfinding orders from initial-rate data
Key equations from the unit and when to use each one.
How much energy? (ΔH)exo (ΔH < 0) or endo (ΔH > 0)Q = mcΔT and calorimetryHess's law: add equationsΔH = ΣnΔHf(products) − ΣnΔHf(reactants)bond energies: broken − formedfuels, efficiency, technologiesHow fast? (rate)rate = change in amount ÷ timecollision theory: energy + orientationfactors: T, conc., surface area, catalystPE diagrams and activation energycatalysts and enzymes lower Eamechanisms: slowest step sets ratechoosing efficient, sustainable reaction conditions
The unit in one picture: thermochemistry tells you how much energy a reaction transfers; kinetics tells you how fast it happens. Both guide the choice of efficient, sustainable conditions.
Just for fun

You have now used Hess's law to find energy changes no one ever measured directly. Germain Hess published his law in 1840; chemists have been taking the scenic route between equations ever since.

Energy Changes in Physical, Chemical and Nuclear Processes

1

Rank the typical size of the energy change per mole: melting ice, burning methane, nuclear fission of uranium-235.

2

Is condensing steam exothermic or endothermic? Explain in terms of particles.

3

In an endothermic reaction in a beaker, does the beaker feel warm or cold? Why?

Enthalpy and Thermochemical Equations

1

For 2H₂(g) + O₂(g) → 2H₂O(l), ΔH = −571.6 kJ. What is ΔH for H₂O(l) → H₂(g) + ½O₂(g)?

2

Write the equation from the first question, 2H₂(g) + O₂(g) → 2H₂O(l), with ΔH as a heat term.

3

How much heat is released when 10.0 g of H₂ burns, according to that equation? (M of H₂ = 2.02 g/mol)

Heat Transfer and Coffee-Cup Calorimetry

1

How much heat raises the temperature of 250.0 g of water from 20.0 °C to 35.0 °C? (c = 4.18 J/(g·°C))

2

50.0 mL of 1.00 mol/L HCl and 50.0 mL of 1.00 mol/L NaOH mix in a coffee-cup calorimeter; the temperature rises 6.7 °C. Find the molar enthalpy of neutralization. (Assume 100.0 g of solution, c = 4.18 J/(g·°C).)

3

Name one assumption made in coffee-cup calorimetry.

Combustion Calorimetry, Fuel Energy and Experimental Error

1

Burning 0.500 g of a fuel heats 200.0 g of water by 12.0 °C. Find the energy per gram of fuel.

2

Why is a value measured with a simple can calorimeter usually smaller than the accepted value?

3

A student finds −1.1 × 10³ kJ/mol for a fuel with an accepted value of −1.37 × 10³ kJ/mol. Calculate the percentage error.

Hess's Law: Adding Thermochemical Equations

1

State Hess's law.

2

Given C(s) + O₂(g) → CO₂(g), ΔH = −393.5 kJ, and CO(g) + ½O₂(g) → CO₂(g), ΔH = −283.0 kJ, find ΔH for C(s) + ½O₂(g) → CO(g).

3

When you multiply an equation by 3, what happens to its ΔH?

Inquiry: Testing Hess's Law with the Combustion of Magnesium

1

Why can't the combustion of magnesium be measured directly in a coffee-cup calorimeter?

2

Which two reactions are measured in the calorimeter in this inquiry?

3

A student's result is −593 kJ/mol; the accepted value is −601.6 kJ/mol. Calculate the percentage error.

Standard Enthalpies of Formation

1

What is ΔH°f for O₂(g)? Why?

2

Use ΔH°f values (CH₄ −74.8, CO₂ −393.5, H₂O(l) −285.8 kJ/mol) to find ΔH°c of methane.

3

Write the formation equation for C₂H₅OH(l).

4

Find ΔH° for CaCO₃(s) → CaO(s) + CO₂(g). (CaCO₃ −1206.9, CaO −635.1, CO₂ −393.5 kJ/mol)

Bond Breaking, Bond Forming and Bond Energies

1

Is bond breaking endothermic or exothermic?

2

Estimate ΔH for H₂(g) + Cl₂(g) → 2HCl(g) using H–H 436, Cl–Cl 242 and H–Cl 431 kJ/mol.

3

Why do bond-energy estimates differ from values found with ΔH°f?

Energy Technologies: Efficiency and Environmental Impact

1

Which source supplied the largest share of Ontario's electricity in 2025?

2

A power plant releases 1000 MJ of chemical energy and delivers 380 MJ of electricity. Find its efficiency.

3

Give one advantage and one drawback of a hydrogen fuel cell vehicle.

Reaction Rates and How to Measure Them

1

The concentration of a reactant falls from 0.80 mol/L to 0.50 mol/L in 60 s. Find the average rate of consumption.

2

For 2H₂O₂ → 2H₂O + O₂, if H₂O₂ is consumed at 0.040 mol/(L·s), at what rate is O₂ produced?

3

Name two properties that could be measured to follow the rate of CaCO₃(s) + 2HCl(aq) → CaCl₂(aq) + H₂O(l) + CO₂(g).

Collision Theory and the Factors That Affect Reaction Rate

1

State the two conditions for an effective collision.

2

Explain why increasing the temperature increases the rate.

3

Explain why powdered zinc reacts faster with acid than a single piece.

Activation Energy and Potential Energy Diagrams

1

On a PE diagram, the reactants are at 50 kJ, the peak at 130 kJ and the products at 20 kJ. Find Ea(fwd), ΔH and Ea(rev).

2

What is at the peak of a PE diagram?

3

For an endothermic reaction, which is larger: Ea(fwd) or Ea(rev)?

Catalysts and Enzymes

1

How does a catalyst increase the rate of a reaction?

2

Does a catalyst change ΔH? Explain.

3

Why does an enzyme stop working if it is heated too much?

Inquiry: Investigating Factors That Affect Reaction Rate

1

A tablet stops fizzing after 40 s. Calculate the rate as 1/t.

2

In a surface-area inquiry, name the independent and dependent variables.

3

Suggest one improvement to an experiment where the end point is judged by eye.

Reaction Mechanisms, Rate-Determining Steps and Rate Laws

1

Mechanism: Step 1: NO + NO → N₂O₂ (fast); Step 2: N₂O₂ + O₂ → 2NO₂ (slow). Write the overall equation and name the intermediate.

2

For rate = k[A][B]², what happens to the rate if [B] is tripled?

3

Which step in a mechanism controls the overall rate?

Making Reactions More Efficient: Conditions, Catalysts and Sustainability

1

Calculate the atom economy of C₂H₄ + H₂O → C₂H₅OH.

2

Why is a condensing furnace more efficient?

3

Give one way a catalyst can make an industrial process more sustainable.

Check your understanding

Quiz

1.

Which process has the largest energy change per mole?

2.

For N₂(g) + 3H₂(g) → 2NH₃(g), ΔH = −92.2 kJ. What is ΔH for 4NH₃(g) → 2N₂(g) + 6H₂(g)?

3.

A 150.0 g sample of water absorbs 6.27 kJ. By how much does its temperature rise? (c = 4.18 J/(g·°C))

4.

Given A → B, ΔH = −50 kJ, and B → C, ΔH = +20 kJ, what is ΔH for A → C?

5.

Using ΔH°f values (C₂H₄ +52.4, C₂H₆ −84.7 kJ/mol), ΔH° for C₂H₄(g) + H₂(g) → C₂H₆(g) is

6.

In a reaction, the energy needed to break bonds is greater than the energy released when new bonds form. The reaction is

7.

Which change does NOT increase the rate of a reaction between zinc and hydrochloric acid?

8.

On a PE diagram, the reactants are at 40 kJ, the activated complex at 140 kJ and the products at 90 kJ. Ea(rev) is

9.

A catalyst increases the rate of a reaction because it

10.

Experiments show rate = k[NO₂]² for NO₂ + CO → NO + CO₂. This tells you that

For parents and teachers

Use the review to find what to revisit:

  • Confusing the size of physical, chemical and nuclear energy changes, or exo/endo from the particle view → Lesson 1 (Energy Changes in Physical, Chemical and Nuclear Processes).
  • Sign errors when reversing equations, or not scaling ΔH with moles → Lesson 2 (Enthalpy and Thermochemical Equations).
  • Q = mcΔT slips, wrong mass, or forgetting the sign of ΔH from a temperature rise → Lesson 3 (Heat Transfer and Coffee-Cup Calorimetry).
  • Fuel energy per gram or per mole, percentage error, or explaining errors → Lesson 4 (Combustion Calorimetry, Fuel Energy and Experimental Error).
  • Not reversing or multiplying equations correctly when adding them → Lesson 5 (Hess's Law).
  • Setting up the magnesium Hess's law inquiry → Lesson 6 (Inquiry: Testing Hess's Law with the Combustion of Magnesium).
  • Writing products minus reactants backwards, forgetting coefficients, or giving elements a nonzero ΔH°f → Lesson 7 (Standard Enthalpies of Formation).
  • Doing formed minus broken, or miscounting bonds → Lesson 8 (Bond Energies).
  • Efficiency calculations or weighing the impacts of energy technologies → Lesson 9 (Energy Technologies).
  • Calculating average rates, or relating rates of different substances → Lesson 10 (Reaction Rates and How to Measure Them).
  • Explaining rate factors without collision theory → Lesson 11 (Collision Theory).
  • Reading Ea(fwd), Ea(rev) and ΔH from PE diagrams → Lesson 12 (Activation Energy and Potential Energy Diagrams).
  • Thinking catalysts change ΔH, or confusing enzymes with reactants → Lesson 13 (Catalysts and Enzymes).
  • Weak hypotheses, uncontrolled variables or vague error analysis → Lesson 14 (Inquiry: Investigating Factors That Affect Reaction Rate).
  • Intermediates versus catalysts, or using equation coefficients as orders → Lesson 15 (Reaction Mechanisms, Rate-Determining Steps and Rate Laws).
  • Atom economy, condensing furnaces, or choosing efficient conditions → Lesson 16 (Making Reactions More Efficient).

Looking ahead: the Chemical Equilibrium unit studies reactions that run in both directions, and uses the energy and rate ideas from this unit to explain how temperature, pressure, concentration and catalysts affect the position of equilibrium (for example in the Haber process). The Electrochemistry unit returns to energy, turning chemical energy directly into electrical energy in cells and batteries.