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Module 4 · L1 of 13 Syllabus requirement · IQ1 ~35 min ⚡ +50 XP in Learn · +25 to complete

Enthalpy & Energy Profile Diagrams

A reusable hand warmer snaps to life and stays hot for an hour; an instant cold pack turns icy the moment you squeeze it. Same idea, opposite directions, one reaction releases energy to its surroundings while the other absorbs it. Enthalpy and energy profile diagrams are how chemists put a number and a direction on that energy, and predict which way a reaction will run.

Today's hook, A reusable hand warmer releases heat as sodium acetate crystallises, while an instant cold pack absorbs heat as ammonium nitrate dissolves. Both are just chemistry moving energy. How can a single number, the enthalpy change ΔH, capture whether a reaction warms you up or cools you down?

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Worksheets

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Four printable worksheets that build from the foundations up to exam-style questions, start at whatever level suits you.

Build Foundations & guided practice Apply Application practice Master Mastery challenge Exam Exam-style questions

Where this lesson fits

Syllabus requirement

Lesson question: how can a single number, the enthalpy change ΔH, tell you whether a reaction warms its surroundings or cools them, and how fast it is willing to go?

1Define ΔH. Enthalpy change is the heat swapped with the surroundings at constant pressure.
2Read its sign. Negative ΔH releases heat (exothermic); positive ΔH absorbs heat (endothermic).
3Draw the pathway. An energy profile diagram shows reactants, products, the transition state, activation energy and ΔH.

Quick prerequisite: reactions rearrange bonds, and breaking bonds costs energy while forming them releases it. If that is shaky, start with the Supported route later in this lesson. Syllabus reference: construct energy profile diagrams to represent and analyse enthalpy changes and activation energy (ACSCH072).

Know what matters most

Must know

ΔH = H(products) − H(reactants), measured at constant pressure
Exothermic ΔH < 0 (heat released); endothermic ΔH > 0 (heat absorbed)
The five labelled features of an energy profile diagram

Should know

Activation energy is measured from the reactant level to the peak
A thermochemical equation ties ΔH to the equation exactly as written
Scaling doubles ΔH; reversing flips its sign

Going deeper

Enthalpy is a state function, H = U + pV
Why liquid vs gaseous water changes ΔH

Warm up and recall

Recall, your gut answer first

+5 XP warm-up

You've probably held a hand warmer on a cold day, it heats up without any battery or flame. And you've seen an instant cold pack snap cold the moment you crack it. Both are just chemicals reacting. Why does one reaction heat up and the other cool down?

Before we name anything, write down what you think the reaction must be doing with energy in each case. What is the system doing with energy when it gets hot? When it gets cold?

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What you'll master, and the words for it

What you'll master

Know

Key Facts

The definition of enthalpy (H) and enthalpy change (ΔH)
The signs of ΔH for exothermic and endothermic reactions
The five labelled features of an energy profile diagram

Understand

Concepts

Why exothermic reactions have ΔH < 0
How activation energy (Ea) relates to the transition state
Why reversing an equation flips the sign of ΔH

Can Do

Skills

Classify reactions as exo/endothermic from ΔH or a diagram
Sketch and label a correct energy profile diagram
Write and manipulate thermochemical equations

Key terms

Enthalpy change (ΔH)

The heat energy exchanged at constant pressure during a reaction.

Exothermic

A reaction that releases heat to surroundings (ΔH < 0).

Endothermic

A reaction that absorbs heat from surroundings (ΔH > 0).

Activation energy

The minimum energy required for reactant collisions to be effective.

Catalyst

A substance that increases reaction rate without being consumed.

Energy profile diagram

A graph showing energy changes during a reaction pathway.

Cross-lesson links: This lesson defines the core quantity, enthalpy change (ΔH), that every subsequent lesson in Module 4 builds on. L02–L04 show you how to measure ΔH using calorimetry; L05 explores how activation energy shapes the energy profile; L06–L10 calculate ΔH from bond energies and Hess's Law; L11–L13 introduce entropy and Gibbs free energy as the two drivers that together determine whether a reaction is spontaneous.

What Is Enthalpy?

Must know core concept

Think of enthalpy (H) as the heat energy stored in a substance's bonds. At constant pressure, the way most reactions run, in an open beaker or flask, the enthalpy change (ΔH) is simply the heat swapped with the surroundings during the reaction.

We can't measure the absolute enthalpy of any substance, but we can measure the change in enthalpy (ΔH) when a reaction occurs. ΔH is defined as the enthalpy of products minus the enthalpy of reactants:

ΔH = H(products) − H(reactants)

If products have less stored energy than reactants, the difference is released to the surroundings as heat and the reaction feels warm. If products have more stored energy, the reaction absorbs heat from the surroundings and feels cold.

Always state the sign of ΔH explicitly. A number without a sign is ambiguous and will cost marks in HSC. Write −890 kJ mol⁻¹, not 890 kJ mol⁻¹.

Common error, "the reaction has energy": Be precise. Breaking bonds requires energy and forming bonds releases it; the reaction enthalpy ΔH is the net difference between the energy absorbed to break reactant bonds and the energy released as product bonds form. ΔH describes the transfer of energy between the system and surroundings, not a quantity a substance possesses.

Going deeper: the formal definition and why constant pressure

Formally, enthalpy is a thermodynamic state function defined as H = U + pV (internal energy plus pressure × volume). Because it is a state function, ΔH depends only on the initial and final states, not the pathway taken, the idea that underlies Hess's Law in L08.

We measure ΔH at constant pressure because most reactions happen in open containers, beakers, flasks, the atmosphere, not sealed pressure vessels. At constant pressure the heat transferred equals ΔH exactly.

Tabulated standard enthalpy changes in this course are reported at 298 K (25 °C) and a standard pressure of 100 kPa, written ΔH°.

Enthalpy change (ΔH = H_products − H_reactants) is a state function measuring heat energy transferred at constant pressure; negative ΔH = exothermic (heat released to surroundings), positive ΔH = endothermic (heat absorbed). Always write the sign explicitly.

Pause, copy the highlighted definition into your book before moving on.

Mini-task: A reaction has ΔH = −286 kJ mol⁻¹. (a) Is this exothermic or endothermic? (b) What does the negative sign tell you about the enthalpy of the products compared to the reactants? Write 1–2 sentences.

Exothermic vs Endothermic Reactions

Must know core concept

We just saw that ΔH = H(products) − H(reactants) defines the direction of energy transfer. That raises a question: how do we classify reactions as releasing or absorbing heat? This card answers it → by the sign of ΔH and the relative enthalpy of products versus reactants.

Whether a reaction releases or absorbs heat depends entirely on whether the products sit at lower or higher enthalpy than the reactants.

In an exothermic reaction, energy is released to the surroundings, ΔH is negative because H(products) < H(reactants). The surroundings heat up; a thermometer in the solution rises. Examples include combustion of fuels, neutralisation of strong acids and bases, and cellular respiration.

In an endothermic reaction, energy is absorbed from the surroundings, ΔH is positive because H(products) > H(reactants). The surroundings cool down; a thermometer drops. Examples include dissolving ammonium nitrate, photosynthesis, and thermal decomposition.

Feature	Exothermic	Endothermic
ΔH sign	Negative (< 0)	Positive (> 0)
Energy flow	System → surroundings	Surroundings → system
Surroundings temperature	Increases (warms up)	Decreases (cools down)
Product energy vs reactants	Lower	Higher
Common examples	Combustion, neutralisation, respiration	Dissolving NH₄NO₃, photosynthesis, thermal decomposition

Real-World Anchor, Hand Warmers vs Cold Packs: A hand warmer contains iron powder that oxidises (exothermic, ΔH < 0, heat released to your hand). An instant cold pack contains ammonium nitrate dissolving in water (endothermic, ΔH > 0, heat absorbed from your hand). The pack feels cold because the reaction is taking heat energy from the surroundings, including your skin. You'll encounter these again in Q3 of Short Answer.

Common error, system vs surroundings language: "The exothermic reaction gets hot" is imprecise. The surroundings get hot. The system (the reaction mixture) is losing enthalpy. Keep system and surroundings language precise in your HSC answers.

Concept map of what ΔH tells us, migrated to a tree layout

Exothermic reactions (ΔH < 0) release heat to surroundings, products sit at lower enthalpy than reactants. Endothermic reactions (ΔH > 0) absorb heat, products sit higher. Examples: combustion (exothermic); dissolving NH₄NO₃ (endothermic).

Add the highlighted point to your notes before the check below.

Explain it: In your own words, explain why the surroundings of an exothermic reaction heat up. Use the terms "enthalpy", "products", "reactants", and "energy flow" in your answer.

Energy Profile Diagrams

Must know core concept

We just saw that exo/endothermic reactions are classified by the sign of ΔH. That raises a question: can we visualise the energy pathway of a reaction, including the barrier that must be overcome? This card answers it → with the energy profile diagram, mapping enthalpy versus reaction progress.

An energy profile diagram is a graph of enthalpy versus reaction progress, it makes the invisible energy landscape of a reaction visible at a glance.

The x-axis is "reaction coordinate" (progress of reaction, not time). The y-axis is enthalpy (kJ mol⁻¹). Reactants start at one enthalpy level; products end at another. Between them is a peak, the transition state representing the activation energy (E_a), the minimum energy required to break bonds and start the reaction.

Five features to label on every diagram:

Reactants (left side, labelled with formula or "Reactants")
Products (right side, labelled)
Transition state (peak, highest point on the curve)
E_a, arrow from reactant level to the peak
ΔH, arrow from reactant level to product level

Energy profile diagrams for an exothermic reaction (left) and an endothermic reaction (right). Note that E_a is always measured from the reactant level to the peak, never from products.

Always draw E_a from the reactant level to the peak. Never from products to the peak, and never the total height of the peak from the x-axis. E_a is the energy gap between reactants and the transition state only.

Common error, E_a from zero: Many students draw E_a as the total height of the peak from the bottom of the graph. E_a is the energy difference between reactants and the transition state not the absolute height. Similarly, E_a for the reverse reaction is measured from the products level to the peak.

Going deeper: link to Module 3 and to Lesson 5

Activation energy (E_a) was introduced in Module 3 collision theory as the minimum energy for a productive collision; here it is placed precisely on a diagram. In Lesson 5 you will see how a catalyst lowers E_a without changing ΔH.

ENERGY PROFILE DIAGRAM, INTERACTIVE Interactive

Toggle exothermic/endothermic, adjust E_a and ΔH sliders, show/hide the catalysed pathway.

An energy profile diagram plots enthalpy vs reaction coordinate; the peak is the transition state, and activation energy (E_a) is the arrow from reactant level to the peak, never from zero. Five required labels: Reactants, Products, Transition state, E_a arrow, ΔH arrow.

Pause, write the highlighted definition into your book.

Match it: pair each energy-profile label on the left with its description on the right.

Reactants
Products
Transition state
Activation energy (E_a)
ΔH (enthalpy change)

The peak of the curve, the highest-energy and least stable arrangement of atoms as bonds rearrange.
The vertical gap between the product enthalpy level and the reactant enthalpy level; negative for an exothermic reaction.
The energy measured from the reactant level up to the peak, the minimum needed for a successful collision.
Plotted at the start of the reaction coordinate at their own enthalpy level, before any bonds break.
Plotted at the end of the reaction coordinate; lower than the reactants when the reaction is exothermic.

Thermochemical Equations

Should know core concept

We just saw that energy profile diagrams visualise E_a and ΔH. That raises a question: how do we communicate ΔH values precisely in written equations? This card answers it → with thermochemical equations that include state symbols and a ΔH value tied to the exact equation as written.

A thermochemical equation is a balanced chemical equation that also states the enthalpy change for the reaction as written change either the equation or the coefficients, and ΔH changes too.

Thermochemical equations must include:

State symbols (s), (l), (g), (aq), for every species
The ΔH value in kJ mol⁻¹ written after the equation
An understanding that ΔH refers to the equation exactly as written

Example:

CH₄(g) + 2O₂(g) → CO₂(g) + 2H₂O(l) ΔH = −890 kJ mol⁻¹

This means 890 kJ is released per mole of methane burned. Two rules apply universally:

Scale the equation → scale ΔH by the same factor (2 mol CH₄ burned: ΔH = −1780 kJ mol⁻¹)
Reverse the equation → flip the sign of ΔH (endothermic decomposition: ΔH = +890 kJ mol⁻¹)

State symbols are non-negotiable. Water as (l) vs (g) gives different ΔH values because of latent heat of vaporisation. In HSC, omitting state symbols loses marks.

ΔH belongs to the equation, not the substance. Students often write ΔH for a substance rather than for the reaction as written. If you change the coefficients or reverse the reaction, ΔH changes accordingly. It is a property of the equation, not the molecule.

A thermochemical equation states ΔH for the reaction as written, with state symbols for every species. Scaling coefficients scales ΔH by the same factor; reversing the equation flips the sign of ΔH. ΔH belongs to the equation, not the substance.

Add the highlighted point to your notes before the check below.

Mini-task: The combustion of propane has ΔH = −2220 kJ mol⁻¹. Write the thermochemical equation for: (a) burning 0.5 mol of propane, and (b) the reverse reaction. Include state symbols in both.

Same science, choose your support

08b

Same science, choose your support

pick one

All three routes practise the same Must-Know idea: read ΔH and describe the energy change. Do one; try another if you want more practice.

Supported word bank + sentence frame start here if unsure

Word bank: exothermic · released · lower · surroundings

A reaction has ΔH = −92 kJ mol⁻¹. Complete the frame:

"Because ΔH is negative, the reaction is ______. Heat is ______ to the ______, so the products store ______ enthalpy than the reactants."

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Show a hint

Negative ΔH always means heat leaves the system, so the surroundings warm up and the products sit lower on an energy profile diagram.

Core independent HSC-standard the exam standard

A reaction has ΔH = +178 kJ mol⁻¹ and a forward activation energy of 190 kJ mol⁻¹. State whether it is exothermic or endothermic, describe the energy flow with the surroundings, and give the activation energy of the reverse reaction.

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Stretch reasoning in a new context push yourself

Two exothermic reactions release the same ΔH, but one fizzes instantly at room temperature while the other needs strong heating to start. Explain how this is possible using ΔH and activation energy, and sketch (in words) how the two energy profile diagrams differ.

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Worked examples

Worked examples · reveal as you go

Worked example 1 +5 XP on full reveal

Thermochemical equations. The thermochemical equation for the combustion of propane is:
C₃H₈(g) + 5O₂(g) → 3CO₂(g) + 4H₂O(l) ΔH = −2220 kJ mol⁻¹

(a) Is this reaction exothermic or endothermic? (b) How much energy is released when 2.00 mol of propane is burned? (c) Write the thermochemical equation for the reverse reaction.

Classify: check the sign of ΔH
ΔH = −2220 kJ mol⁻¹ → negative sign → exothermic. Energy is released to the surroundings.

A negative ΔH always means exothermic, products sit lower in enthalpy than reactants.

Scale ΔH for 2.00 mol propane
Energy = 2.00 × 2220 = 4440 kJ released. ΔH is per mole of equation as written, scale proportionally.

ΔH scales linearly with the number of moles. Double the fuel, double the energy released.

Write the reverse reaction
3CO₂(g) + 4H₂O(l) → C₃H₈(g) + 5O₂(g) ΔH = +2220 kJ mol⁻¹

Reversing a reaction flips the sign of ΔH. The reverse reaction is endothermic.

Final Answer
(a) Exothermic (ΔH < 0) | (b) 4440 kJ released | (c) ΔH = +2220 kJ mol⁻¹

All three parts answered concisely, this is the format HSC markers expect for thermochemical equation questions.

Worked example 2 +5 XP on full reveal

Energy profile diagrams. A reaction has ΔH = +85 kJ mol⁻¹ and a forward activation energy of 120 kJ mol⁻¹.

(a) Is the reaction exo- or endothermic? (b) What is the activation energy for the reverse reaction? (c) Describe the key features of the energy profile diagram for this reaction.

Classify from ΔH sign
ΔH = +85 kJ mol⁻¹ → positive → endothermic. Products are at higher enthalpy than reactants.

Positive ΔH = endothermic; products sit higher on the energy profile diagram.

Calculate E_a for the reverse reaction
E_a(reverse) = E_a(forward) − ΔH = 120 − 85 = 35 kJ mol⁻¹

For the reverse reaction, you start from the products (which sit 85 kJ mol⁻¹ higher than reactants) and climb to the same peak, a smaller gap of 35 kJ mol⁻¹.

Describe the energy profile diagram features
Reactants at baseline. Peak is 120 kJ mol⁻¹ above reactants (E_a = 120 kJ mol⁻¹, arrow up from reactant level to peak). Products are 85 kJ mol⁻¹ above reactants (ΔH = +85 kJ mol⁻¹, arrow pointing upward). Transition state at the peak. x-axis: "Reaction coordinate". y-axis: "Enthalpy (kJ mol⁻¹)".

Five features must be present: reactants, products, transition state, Ea arrow, ΔH arrow, each labelled.

Final Answer
(a) Endothermic | (b) E_a(rev) = 35 kJ mol⁻¹ | (c) Products sit 85 kJ mol⁻¹ above reactants; peak is 120 kJ mol⁻¹ above reactants

Full marks in HSC require all three parts answered with correct terminology and values.

Key idea

Predict then reveal +8 XP

1 · Predict

2 · Reveal

3 · Compare

A 1 g sample of glucose is burned completely in excess oxygen. Before you calculate anything, predict whether the ΔH value will be positive or negative, and explain your reasoning.

Confidence 50%

What you'll master, and the words for it

Formula reference · this lesson

core formula

📐

Formula Reference, This Lesson

$\Delta H = H_{\text{products}} - H_{\text{reactants}}$

ΔH = enthalpy change (kJ mol⁻¹) H_products = total enthalpy of products H_reactants = total enthalpy of reactants

Exothermic: $\Delta H < 0$ (products lower than reactants, heat released) | Endothermic: $\Delta H > 0$ (products higher than reactants, heat absorbed)

"The reaction gets hot/cold"

Common errors · the 3 traps that cost marks

"The reaction gets hot/cold"

Students write "the exothermic reaction gets hot" or "the endothermic reaction gets cold."

Fix: The surroundings change temperature. The system (the chemicals reacting) loses or gains enthalpy. Precision in this language is an HSC marker requirement. Write: "the surroundings increase in temperature" not "the reaction gets hot."

"Activation energy is the height of the peak from the x-axis"

Students draw the E_a arrow starting from the x-axis (zero) rather than from the reactant enthalpy level.

Fix: E_a is the energy difference from the reactant level to the transition state only. The position of the reactant level on the y-axis is arbitrary, only the differences between levels matter.

"Reversing the equation just removes the negative sign"

Students think reversing an equation turns −890 kJ mol⁻¹ into 890 kJ mol⁻¹ (unsigned) rather than +890 kJ mol⁻¹.

Fix: Reversing flips the sign so a negative becomes positive and vice versa. ΔH = −890 becomes ΔH = +890 kJ mol⁻¹ when the equation is reversed. You must write the + sign explicitly.

Drill, then revisit

Quick-fire practice · 5 reps +2 XP per reveal

Classify each reaction as exothermic or endothermic and state the ΔH sign:
(a) Combustion of ethanol: C₂H₅OH(l) + 3O₂(g) → 2CO₂(g) + 3H₂O(l)
(b) Photosynthesis: 6CO₂(g) + 6H₂O(l) → C₆H₁₂O₆(s) + 6O₂(g)
(c) Dissolving NaOH(s) in water, the solution warms up
(d) Dissolving NH₄Cl(s) in water, the solution cools down

Student A says: "For a reaction with ΔH = −60 kJ mol⁻¹ and Ea = 100 kJ mol⁻¹, the activation energy for the reverse reaction is 100 − (−60) = 160 kJ mol⁻¹." Is this correct? What is the conceptual error if any?

Student C says: "I doubled the equation 2CH₄(g) + 4O₂(g) → 2CO₂(g) + 4H₂O(l) but kept ΔH = −890 kJ mol⁻¹ because the substances are the same." Identify the error and write the correct thermochemical equation.

Student B says: "I drew my energy profile diagram with the transition state peak touching the top of the y-axis to show it has maximum energy. The Ea arrow goes from the x-axis to the top of the peak." Identify the two errors.

The extension question: The combustion of glucose is C₆H₁₂O₆(s) + 6O₂(g) → 6CO₂(g) + 6H₂O(l), ΔH = −2803 kJ mol⁻¹. What does this tell you about the enthalpy change for photosynthesis? Use Hess's Law reasoning.

Revisit your thinking

Go back to your Think First response. Now that you've studied enthalpy and energy profile diagrams, revisit the hand warmer and the instant cold pack:

The hand warmer feels hot and the cold pack feels cold. For each, state the sign of ΔH and explain what that sign tells you about the enthalpy of the products compared with the reactants.
Sketch a labelled energy profile diagram for one exothermic and one endothermic process. Mark the reactant level, product level, transition state, E_a and the ΔH arrow on each.
What is the difference between ΔH and E_a? Why can two reactions release the same ΔH yet proceed at very different rates?

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Interactive Tool, Enthalpy & Calorimetry Open fullscreen ↗

Use the Calorimetry Calculator. Heating 200 g of water by 5.0°C (c = 4.18 J/g·°C) requires how much heat?

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Multiple choice

+5 XP per correct · +25 XP all-correct

Pick your answer, then rate your confidence. That tells the system what to drill next.

01b

Misconceptions to fix before short answer

Wrong: An exothermic reaction has ΔH > 0 because it releases heat.

Right: Exothermic reactions have ΔH < 0 (negative) because the system loses energy to the surroundings. Endothermic reactions have ΔH > 0 (positive) because the system gains energy. The sign convention refers to the system, not the surroundings.

Wrong: "The reaction gets hot/cold."

Right: "The surroundings change temperature." The system (reaction mixture) loses or gains enthalpy, language precision is an HSC marker requirement.

Wrong: "Reversing the equation removes the negative sign."

Right: Reversing flips the sign, ΔH = −890 becomes ΔH = +890 kJ mol⁻¹. The + must be written explicitly.

Short answer

UnderstandBand 3

Q1. Distinguish between exothermic and endothermic reactions in terms of (a) the sign of ΔH, (b) the direction of energy flow, and (c) how each would appear on an energy profile diagram. Use a specific example of each. 4 MARKS

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ApplyBand 4

Q2. The thermochemical equation for the combustion of ethane is:
2C₂H₆(g) + 7O₂(g) → 4CO₂(g) + 6H₂O(l) ΔH = −3120 kJ mol⁻¹

(a) Calculate the enthalpy change when 0.500 mol of ethane is burned. (1 mark)
(b) Write the thermochemical equation for the reverse of this reaction. (1 mark)
(c) Explain why water being produced as liquid (l) rather than gas (g) affects the value of ΔH. (2 marks) 4 MARKS

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EvaluateBand 5

Q3. Real-World Application: Instant cold packs used by sports trainers contain solid ammonium nitrate (NH₄NO₃) separated from water by a thin inner bag. When the bag is cracked, NH₄NO₃ dissolves in the water and the pack becomes very cold.

(a) Is the dissolution of ammonium nitrate exothermic or endothermic? Justify using the observation described. (2 marks)
(b) Draw a labelled energy profile diagram for the dissolution of NH₄NO₃. Include: reactants, products, transition state, E_a, and ΔH arrows. Show whether products sit above or below reactants. (3 marks) 5 MARKS

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Short-answer model answers

Multiple-choice feedback appears inline as you answer. Each quick-fire drill above reveals its own answer. Below are the marker-style model answers for the three short-answer questions.

Show model answers ▼

Short Answer Model Answers

Q1 (4 marks): Exothermic: ΔH < 0; energy flows from system to surroundings; on energy profile diagram, products sit below reactants, ΔH arrow points downward; example, combustion of methane, ΔH = −890 kJ mol⁻¹ [2 marks]. Endothermic: ΔH > 0; energy flows from surroundings to system; products sit above reactants, ΔH arrow points upward; example, dissolving ammonium nitrate, ΔH > 0 [2 marks].

Q2 (4 marks): (a) Equation shows 2 mol ethane releases 3120 kJ. For 0.500 mol: ΔH = (0.500/2) × 3120 = 780 kJ released [1]. (b) 4CO₂(g) + 6H₂O(l) → 2C₂H₆(g) + 7O₂(g), ΔH = +3120 kJ mol⁻¹ (sign flips) [1]. (c) Converting water from liquid to gas requires energy (latent heat of vaporisation = 44 kJ mol⁻¹ per mol H₂O) [1]. If water were produced as gas, less energy would be released to the surroundings, ΔH would be less negative. The liquid (l) state releases the full amount including the condensation energy [1].

Q3 (5 marks): (a) Endothermic [1]. Justification: the pack becomes cold, meaning heat is being transferred from the surroundings (including the injured area) into the system, the reaction absorbs heat rather than releasing it. This means ΔH > 0 [1]. (b) Diagram requirements: reactants (solid NH₄NO₃ + liquid H₂O) at a lower enthalpy level [1]; products (NH₄⁺(aq) + NO₃⁻(aq)) at a higher enthalpy level, ΔH arrow points upward [1]; transition state at the peak above products; E_a arrow from reactant level to peak [1]. Accept: x-axis labelled "Reaction coordinate", y-axis "Enthalpy (kJ mol⁻¹)".

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