Difficulty: Introductory | Prerequisites: Basic algebra, familiarity with the periodic table
This topic sits right at the start of thermochemistry, the branch of chemistry concerned with how energy moves during chemical reactions. If you have been away for a few weeks, you need to know what a mole is and have a rough sense of how chemical equations are balanced. The reason this matters: every reaction you will study from here on either releases or absorbs energy, and the language introduced here (system, surroundings, enthalpy, sign conventions) is the vocabulary the rest of the course assumes you already have.
Energy is measured in joules or calories and can change form but never appear from nothing or vanish. Chemical reactions either release heat to their surroundings (exothermic, negative ΔH) or absorb heat from them (endothermic, positive ΔH). The amount of energy involved scales directly with the quantity of reactants.
Joule (J)
The SI unit of energy. Roughly the energy needed to lift a small apple (about 100 g) one metre straight up.
In simple terms, it is the standard measuring stick for energy in science.
Kilojoule (kJ)
A larger unit of energy equal to 1,000 joules. Used because the joule is inconveniently small for most chemical reactions.
Think of it as the "kilometre" to the joule's "metre."
Calorie (cal)
The amount of energy required to raise the temperature of 1 g of water by 1 °C. 1 cal = 4.184 J.
In simple terms, it is the older, non-SI energy unit you will still see in food science and older textbooks.
Kilocalorie (kcal) / Food Calorie (Cal)
1,000 calories. This is what food labels call a "Calorie" (capital C). Fats provide about 9 kcal/g; proteins and carbohydrates each provide about 4 kcal/g.
Think of it as the nutrition label unit, confusingly spelled the same way but 1,000 times larger than the chemist's calorie.
Law of Conservation of Energy
Energy can neither be created nor destroyed; it only changes form.
In simple terms, the total energy of the universe is constant. You can move it around, but you cannot make more of it or get rid of any.
Exothermic Reaction
A reaction that releases heat to its surroundings. Products are lower in energy (more stable) than reactants. ΔH is negative.
Think of it as the reaction that warms the room.
Endothermic Reaction
A reaction that absorbs heat from its surroundings. Products are higher in energy (less stable) than reactants. ΔH is positive.
Think of it as the reaction that makes the container feel cold.
Enthalpy Change (ΔH)
The amount of energy transferred as heat at constant pressure (room temperature and pressure in this context). A negative value means heat leaves the system; a positive value means heat enters.
In simple terms, ΔH tells you whether the reaction gives off or soaks up heat, and how much.
System
The part of the universe you are studying, in chemistry usually the reaction itself.
Surroundings
Everything outside the system, essentially the rest of the universe.
The SI unit is the joule (J). Chemistry commonly uses kilojoules (kJ) because individual joules are very small at the reaction scale.
The calorie (cal) is defined by heating water: 1 cal raises 1 g of water by 1 °C. The conversion factor is 1 cal = 4.184 J.
Food labels report energy in kilocalories but print the word "Calories" (capital C). This catches students out regularly.
Fats: ~9 kcal/g
Proteins: ~4 kcal/g
Carbohydrates: ~4 kcal/g
Energy cannot be created or destroyed; it only changes form. This is the first law of thermodynamics, though at this stage the course states it as a principle rather than deriving it.
Chemical energy stored in bonds converts to heat, light, or both during a reaction.
Exothermic: heat flows from the system (the reaction) to the surroundings. The products sit at a lower energy level than the reactants, which is why the leftover energy escapes as heat.
Some exothermic reactions are gentle (iron rusting). Others are rapid (explosions, which release gas and heat at the same time).
Endothermic: heat flows from the surroundings into the system. The products are higher in energy than the reactants, so the reaction has to pull energy in.
ΔH is the enthalpy change at constant pressure (essentially room temperature and atmospheric pressure for this course).
Negative ΔH means heat leaves the system (exothermic).
Positive ΔH means heat enters the system (endothermic).
The magnitude of ΔH scales with the amount of reactant. Burning one mole of methane (CH₄) releases 890.4 kJ. Burning two moles releases twice that.
The system is the reaction you are focused on.
The surroundings are everything else (the flask, the bench, the room, the universe).
Energy moving from system to surroundings = exothermic. Energy moving from surroundings to system = endothermic.
Unit conversions
1 kJ = 1,000 J
1 cal = 4.184 J
1 kcal = 1,000 cal = 4,184 J
Energy content of macronutrients
Fats: 9 kcal/g
Proteins: 4 kcal/g
Carbohydrates: 4 kcal/g
Methane combustion (benchmark example)
CH₄ + 2 O₂ → CO₂ + 2 H₂O, ΔH = −890.4 kJ/mol
The negative sign tells you this is exothermic. Scaling is linear: two moles of CH₄ gives ΔH = −1,780.8 kJ.
The exothermic/endothermic distinction is how engineers size heat exchangers and cooling systems. A car radiator, for instance, exists because the combustion reaction in the engine is violently exothermic and that heat has to go somewhere. Cold packs used for sports injuries work in the opposite direction: the ammonium nitrate dissolving inside is endothermic, pulling heat from your skin.
"Calories on food labels are the same as calories in chemistry." They are not. A food "Calorie" (capital C) is a kilocalorie, 1,000 times larger than the chemist's calorie (lowercase c).
"Exothermic reactions are always explosions." Explosions are one dramatic type, but many exothermic reactions are slow and gentle (like iron rusting or your body metabolising food).
"A negative ΔH means the reaction is losing energy to nowhere." The energy is not lost. It transfers to the surroundings. Conservation of energy still holds.
"Endothermic reactions don't happen on their own." Some do, given the right conditions. An ice pack activates spontaneously once the inner pouch breaks, even though the dissolving step is endothermic.
⚠️ Unit conversion between joules, kilojoules, calories, and kilocalories is a near-certain exam question. Know 1 cal = 4.184 J cold.
⚠️ The sign of ΔH is the single most tested concept here. Negative = exothermic (heat out). Positive = endothermic (heat in). Mixing these up is the most common mark lost.
⚠️ Scaling ΔH with moles: if the question doubles the reactant, you double ΔH. Simple, but students forget under time pressure.
⚠️ System vs. surroundings and which direction heat flows. Exams like to phrase this as "the beaker feels warm" (exothermic, heat went to surroundings) or "the beaker feels cold" (endothermic, heat left surroundings).
True or false: 1 food Calorie = 1 cal.
False. 1 food Calorie = 1 kcal = 1,000 cal.
Fill in the blank: A reaction with ΔH = −200 kJ is ____.
Exothermic.
True or false: In an endothermic reaction, the surroundings gain heat.
False. The surroundings lose heat (it flows into the system).
If burning 1 mol of CH₄ releases 890.4 kJ, how much is released by 0.5 mol?
445.2 kJ.
True or false: Energy can be created in a nuclear reaction.
False. Energy changes form; it is neither created nor destroyed.
Q: Convert 250 cal to joules.
A: 250 × 4.184 = 1,046 J (or about 1.046 kJ).
Q: A reaction has ΔH = +45 kJ. Is it exothermic or endothermic? Does the system absorb or release heat?
A: Endothermic. The system absorbs heat from the surroundings.
Q: Burning 1 mol of methane releases 890.4 kJ. How much energy is released when 3 mol of methane burn?
A: 3 × 890.4 = 2,671.2 kJ.
Q: A student touches a beaker during a reaction and finds it is warm. Is the reaction exothermic or endothermic? Explain using "system" and "surroundings."
A: Exothermic. Heat has left the system (the reaction) and entered the surroundings (the beaker and the student's hand), making the beaker feel warm.
Q: A food label says a serving contains 150 Calories. Express this in kilocalories and in kilojoules.
A: 150 Calories = 150 kcal. In kilojoules: 150 × 4.184 = 627.6 kJ.
This material leads directly into Hess's Law, where you will combine multiple ΔH values to find the enthalpy change of a reaction you cannot measure directly. The system/surroundings framework also underpins calorimetry (measuring heat with a calorimeter), which is typically the next lab exercise. Later in the course, you will revisit energy when studying bond energies and Gibbs free energy, both of which build on the sign conventions and definitions introduced here.
Thermochemistry, enthalpy, ΔH, delta H, exothermic reaction, endothermic reaction, joule, kilojoule, calorie, kilocalorie, food Calorie, conservation of energy, first law of thermodynamics, system and surroundings, heat of reaction, heat transfer, methane combustion, energy units, calorimetry, Hess's Law, bond energy, chemical energy, heat release, heat absorption, CHEM 101, general chemistry energy