Chemistry: Atoms First 2e · Thermochemistry

Energy Basics

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In 30 seconds
Why this matters
The college version
Eli explains
Worked example
Key takeaway
Check yourself
Study tools
Sources & references

In 30 seconds

Thermochemistry is the study of the heat absorbed or released during chemical reactions and physical changes. Before any calculation — before coffee-cup calorimeters, before enthalpy diagrams — you need a clear picture of what is, how it moves, and how chemists keep score. This topic lays that foundation: the two fundamental forms of energy (kinetic and potential), the difference between heat and , the concept of a and its , the units chemists use, and the — the conservation rule that governs every thermochemical calculation.

Everything in the next three topics (calorimetry, enthalpy, and bond energies) applies these ideas. If you can identify the system, track heat and work with correct signs, and convert between joules and calories, the rest of Chapter 9 becomes routine.

Why this matters

Every reaction either absorbs or releases energy. Digestion, combustion, and dissolving salt are all thermochemical events; the basics let you predict whether a process warms or cools its surroundings.
Nutrition labels are thermochemistry. Food "Calories" are kilocalories — the energy released when food is metabolized.
Energy units appear everywhere. Fuels are rated in kJ per gram, food in Calories; chemists and engineers convert among them routinely.
The first law protects you from errors. If your calculation creates energy from nowhere, something is wrong — the law is a built-in check.
Foundation for the chapter: calorimetry measures energy as heat; enthalpy is defined using the first law; bond strengths are measured as energy changes.

The college version

Core Concepts

Kinetic and potential energy

Energy is the capacity to do work or transfer heat. It comes in two fundamental forms:

is the energy of motion: KE = 12mv². A moving car, a flowing river, and vibrating, rotating molecules all possess it.
is stored energy due to position or composition: PE = mgh for an object at height h. Chemical potential energy is stored in the positions of electrons and nuclei — in the arrangements of atoms within molecules.

A stretched spring, a raised weight, and a gasoline molecule about to react all store potential energy that can be released as kinetic energy (motion or heat) when the arrangement changes.

Thermal energy, heat, and temperature — three different ideas

Thermal energy is the total kinetic energy of all the particles in a sample — the sum of all their jiggling, rotating, and moving. It depends on how many particles there are and how fast they move.
Temperature is the average kinetic energy per particle — how fast the average particle moves. A cup of tea and a swimming pool can both be at 30 °C, but the pool holds vastly more thermal energy because it holds vastly more particles.
Heat (q) is the energy transferred between objects at different temperatures. Heat flows spontaneously from hot to cold until thermal equilibrium. Heat is not a substance objects "contain"; it is energy in transit.

System and surroundings

Thermochemistry needs a clear accounting boundary. The system is the part of the universe under study — the reacting chemicals, the gas in a cylinder, the ice cube. The surroundings are everything else — the beaker, the air, the room.

Systems come in three types:

Open system: both energy and matter are exchanged (a beaker open to the air).
Closed system: energy is exchanged but matter is not (a sealed flask).
Isolated system: neither energy nor matter is exchanged (a well-insulated thermos — the ideal, never perfectly achieved).

The first law of thermodynamics: energy is conserved

The first law states that the total energy of the universe is constant: energy can be converted or transferred but never created or destroyed. For a system:

$Δ E = q + w$

where ΔE is the change in internal energy, q is heat transferred to the system, and w is work done on the system. The sign convention is crucial:

q > 0: heat flows into the system (endothermic, system warms).
q < 0: heat flows out of the system (exothermic, system cools).
w > 0: work is done on the system (compressing a gas).
w < 0: the system does work on the surroundings (expanding gas).

In chemistry, the work is usually pressure–volume work: w = -PΔV. Expansion (ΔV > 0) gives w < 0 — the system spends energy pushing the surroundings aside.

Energy units and conversions

The SI unit of energy is the : 1 J = 1 kg m²s^-2. The calorie is the older unit, defined as the energy to raise 1 gram of water by 1 °C. The link is exact:

$1 cal = 4.184 J$

A nutritional is actually a kilocalorie: 1 Cal = 1000 cal = 4184 J. Because chemical energies are large, kilojoules (1 kJ = 10³ J) are the everyday workhorse unit.

Common Confusions

Do Not Confuse	With	Difference
Heat (q)	Temperature	Heat is energy transferred; temperature is average particle kinetic energy
Thermal energy	Temperature	Thermal energy is total particle kinetic energy (depends on amount); temperature is average per particle
Calorie (cal)	Calorie (Cal)	1 Cal = 1000 cal = 1 kcal; nutrition labels use capital-C Calorie
Energy	Power	Energy is the capacity to do work (J); power is energy per unit time (J/s = W)
q > 0	q < 0	q > 0: heat into system (endothermic); q < 0: heat out (exothermic)
Work done on system	Work done by system	On system: w > 0 (compression); by system: w < 0 (expansion)
Closed system	Isolated system	Closed exchanges energy only; isolated exchanges neither energy nor matter

Eli explains

The same idea, in plain words

Explain it like I’m 10

Energy is like money that can never disappear — it only changes form or moves to a different pocket. Kinetic energy is the money of motion: a running kid has it. Potential energy is money saved in a bank: a stretched rubber band has it, ready to be spent. Heat is money handed from a hot object to a cold one until they both feel the same. Thermochemistry counts where that money goes when chemicals react — how much is handed out, how much is taken in, and who pays.

Worked example

Example 1: Converting food energy from Calories to joules

A chocolate bar's label says it provides 230 Calories. Express this in kilocalories, calories, joules, and kilojoules.

Step 1 — Write the conversion relationships:

$1 Cal = 1 kcal = 1000 cal = 4184 J$

Step 2 — Convert Cal to kcal and cal:

$230 Cal \times 1 kcal 1 Cal = 230 kcal$

$230 kcal \times 1000 cal 1 kcal = 2.30 \times 105 cal$

Step 3 — Convert cal to J with the exact factor:

$2.30 \times 105 cal \times 4.184 J 1 cal = 9.62 \times 105 J$

Step 4 — Convert J to kJ:

$9.62 \times 105 J \times 1 kJ 1000 J = 962 kJ$

Dimensional check: each factor is written so the old unit cancels and the new unit survives. ~960 kJ is a typical candy-bar energy — enough to run for roughly 15 minutes.

Example 2: Applying the first law to a gas

A gas in a cylinder absorbs 450 J of heat while expanding, doing 175 J of work on the surroundings. Calculate ΔE.

Step 1 — Write the first law:

$Δ E = q + w$

Step 2 — Assign signs. Heat enters the system, q = +450 J. The system does work on the surroundings, w = -175 J:

$Δ E = (+450 J) + (-175 J) = +275 J$

The gas's internal energy rose by 275 J: it absorbed more energy as heat (450 J) than it spent pushing the surroundings aside (175 J). If the gas had been compressed instead, w would be positive and ΔE larger — same equation, different sign.

Example 3: Pressure–volume work during expansion

A gas expands from 2.0 L to 5.0 L against a constant external pressure of 1.5 atm. Calculate the work done by the gas in L·atm and in joules (1 L atm = 101.325 J).

Step 1 — Write the work formula:

$w = -P Δ V$

Step 2 — Compute ΔV and substitute:

$Δ V = 5.0 L - 2.0 L = 3.0 L$

$w = -(1.5 atm)(3.0 L) = -4.5 L atm$

Step 3 — Convert to joules:

$w = -4.5 L atm \times 101.325 J 1 L atm = -4.6 \times 102 J$

The negative sign means the gas spent ~460 J pushing the surroundings aside — energy it must have absorbed as heat or drawn from its own internal energy.

Key takeaways

Energy = capacity to do work or transfer heat; kinetic (motion) vs. potential (position/composition).
Temperature = average kinetic energy per particle; thermal energy = total kinetic energy of all particles; heat = energy in transit due to a temperature difference.
First law: ΔE = q + w; q > 0 heat in, q < 0 heat out; w > 0 work on system, w < 0 work by system.
Pressure–volume work: w = -PΔV.
Exact conversion: 1 cal = 4.184 J; 1 Cal = 1 kcal = 4184 J.
System types: open (energy + matter), closed (energy only), isolated (neither).
Heat flows hot → cold until thermal equilibrium.

Check yourself

6 review questions from the chapter. Try each one, then open the answer.

Distinguish kinetic energy, potential energy, and thermal energy with one example each.
Show answer
Kinetic: a moving molecule or car (12mv²). Potential: a raised weight or energy in chemical bonds (mgh, bond energy). Thermal: the total jiggling energy of all particles in a sample.
Why can a cup of coffee and a swimming pool have the same temperature but vastly different thermal energies?
Show answer
Temperature is the average kinetic energy per particle; both have particles at the same average speed. Thermal energy is the sum over all particles — the pool has far more particles, hence far more total thermal energy.
State the first law in words and equation form, and define every symbol's sign convention.
Show answer
The energy of the universe is constant; energy is converted or transferred but never created or destroyed. ΔE = q + w, with q > 0 heat into the system and w > 0 work on the system.
A reaction releases 850 J of heat to the surroundings and does no work. What are q and ΔE for the reacting system?
Show answer
Heat leaves the system, so q = -850 J; w = 0; ΔE = -850 J. The system's internal energy fell by 850 J.
Convert 1500 cal to joules and kilojoules. Convert 2.5 kJ to calories.
Show answer
1500 cal × 4.184 J cal^-1 = 6.28 × 10³ J = 6.28 kJ. 2.5 kJ × 1000 J1 kJ × 1 cal4.184 J = 6.0 × 10² cal.
Classify each as open, closed, or isolated: (a) an open cup of hot coffee; (b) a sealed, rigid flask of gas; (c) an ideal thermos.
Show answer
(a) open — steam (matter) and heat (energy) escape; (b) closed — energy can flow, matter cannot; (c) isolated — by design, neither crosses the boundary.

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Study tools & related lessonsKey vocabulary · Related

Key vocabulary

Energy: Capacity to do work or transfer heat
Kinetic energy: Energy of motion, 12mv²
Potential energy: Stored energy from position or composition
Temperature: Average kinetic energy per particle
Heat (q): Energy transferred due to a temperature difference
Work (w): Energy transferred by a force acting through a distance
System: The part of the universe under study
Surroundings: Everything outside the system
Joule (J): SI unit of energy, 1 J = 1 kg m²s^-2
Calorie (cal): Energy to warm 1 g of water by 1 °C; exactly 4.184 J
First law of thermodynamics: ΔE = q + w; energy is conserved

Sources & references

openstax.org — Chemistry Atoms First 2e

This lesson was adapted from the open educational references above; their licenses and attributions are preserved. See Copyright & Licensing.

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