Unit Overview

R1: What drives chemical reactions?

IB Chemistry HL· 5 min read 📊 n/a

1. Unit at a Glance

This unit follows a logical progression from experimental measurement of heat changes to theoretical prediction of reaction spontaneity. We start with core concepts of enthalpy and experimental calorimetry, then build up to indirect methods for calculating enthalpy when experiments are not possible.

After covering enthalpy calculation methods, we introduce entropy as a measure of disorder, then extend to Gibbs free energy (the key predictor of spontaneity) for HL students. Every concept builds on the previous to answer the unit's core question.

This unit includes the following core and AHL sub-topics:

Enthalpy change and calorimetry

Define enthalpy change and learn to measure heat transfer in reactions via experimental calorimetry.

★★⏱ 10 min

Hess's law

Calculate overall enthalpy change for reactions using enthalpy values of intermediate steps.

★★★⏱ 8 min

Entropy and reaction spontaneity

Introduce entropy as a measure of disorder and explain its role in determining spontaneity.

★★★⏱ 9 min

Bond enthalpies

Estimate reaction enthalpy from average bond enthalpies of broken and formed bonds.

★★⏱ 7 min

Energy cycles

Use energy cycles to calculate enthalpy changes of formation and combustion.

★★★⏱ 8 min

AHL: Extended enthalpy and entropy calculations

Practice advanced calculations involving standard states and entropy changes of reaction.

★★★★⏱ 12 min

AHL: Gibbs free energy and reaction spontaneity

Combine enthalpy and entropy to calculate Gibbs free energy and predict reaction spontaneity.

★★★★⏱ 11 min

2. Common Pitfalls

Wrong move:

Confusing endothermic/exothermic sign conventions for enthalpy change

Why:

Different conventions across sources often lead to flipped signs in final calculations

Correct move:

Use IB convention: $Δ H < 0$ for exothermic (heat released), $Δ H > 0$ for endothermic (heat absorbed)

Wrong move:

Flipping the bond enthalpy calculation formula

Why:

Forgetting that bond breaking is endothermic and bond forming is exothermic flips the final sign of $Δ H$

Correct move:

Always use: $Δ H = \sum (E_{bonds broken}) - \sum (E_{bonds formed})$

Wrong move:

Assuming negative enthalpy means a reaction is always spontaneous

Why:

Enthalpy alone does not determine spontaneity; entropy and temperature also play critical roles

Correct move:

Always use Gibbs free energy change ( $Δ G$ ) to confirm reaction spontaneity

3. Quick Reference Cheatsheet

Concept	Key Formula/Relationship
Calorimetry enthalpy change	$q = m c Δ T$ , $Δ H = - q / n$ (constant pressure)
Hess's Law	$Δ H_{total} = \sum Δ H_{individual steps}$
Reaction enthalpy from bond enthalpies	$Δ H = \sum E_{bonds broken} - \sum E_{bonds formed}$
Standard entropy change of reaction	$Δ S^{\circ} = \sum S_{products}^{\circ} - \sum S_{reactants}^{\circ}$
Gibbs Free Energy Change	$Δ G = Δ H - T Δ S$
Standard Gibbs Free Energy of reaction	$\Delta G^\circ = \sum \Delta G^\circ_f_{\text{products}} - \sum \Delta G^\circ_f_{\text{reactants}}$
Spontaneity condition	Reaction is spontaneous if $Δ G < 0$

What's Next

Begin this unit with the first sub-topic, which introduces core concepts of enthalpy change and experimental measurement that all later topics in this unit rely on. After you complete all sub-topics in this unit and master the concepts of reaction spontaneity, you can proceed to the first topic of the next unit on reaction kinetics.