AP Chemistry 6.1: Endothermic and Exothermic Processes and Unit 6 Review

Thermodynamics is the study of energy and how energy is transferred during physical and chemical processes. In chemistry, this mainly means understanding what happens to energy when substances react, change temperature, or change phase.

• This unit is heavily connected to ideas you have already learned. Intermolecular forces help explain why phase changes require different amounts of energy. Chemical bonds help explain why reactions absorb or release energy. Stoichiometry lets us connect the amount of substance reacting to the amount of heat transferred.
• The biggest idea to keep in mind throughout Unit 6 is that energy is conserved. Energy can move between a chemical system and its surroundings or change from one form to another, but it does not simply disappear.
• By the end of this unit, you should be able to describe whether a process is endothermic or exothermic, calculate heat transfer using calorimetry, interpret heating curves, calculate reaction enthalpies, manipulate reactions using Hess's law, use standard enthalpies of formation, and estimate reaction enthalpy using bond enthalpies.

Thermodynamics begins by separating the part of the universe we are studying from everything around it.

The system is the specific chemical reaction, substance, or process being studied. The surroundings are everything outside the system that can exchange energy with it.

• For example, suppose a reaction occurs inside a beaker. The reacting chemicals are the system, while the water bath, beaker, air, thermometer, and rest of the room are part of the surroundings.
• This distinction matters because when energy leaves one place, it must enter another.

If the system releases heat:

• system loses energy → surroundings gain energy

If the system absorbs heat:

• surroundings lose energy → system gains energy

Temperature measures the average kinetic energy of particles. When particles have greater average kinetic energy, the temperature is higher.

Heat, represented by q, is energy transferred because of a temperature difference.

• Heat is therefore not something an object simply “contains.” It describes energy being transferred between objects or systems.

Heat naturally transfers from:

• higher temperature → lower temperature
• until the objects reach the same temperature.

The internal energy, represented by E, is the total microscopic energy contained within a system. It includes contributions from particle motion and interactions between particles.

• We usually care about the change in internal energy:
• ΔE = Efinal − Einitial

The first law of thermodynamics expresses conservation of energy:

• ΔE = q + w
• where:
• ΔE = change in internal energy
• q = heat transferred to or from the system
• w = work done on or by the system

Using the chemistry sign convention:

• q > 0 means heat enters the system.
• q < 0 means heat leaves the system.
• w > 0 means work is done on the system.
• w < 0 means the system does work on the surroundings.
• For most AP Chemistry Unit 6 problems, the major focus is heat and enthalpy rather than difficult pressure-volume work calculations, but understanding conservation of energy is important.

The quantity enthalpy, represented by H, is especially useful for chemical reactions occurring at constant pressure.

• The change in enthalpy is:
• ΔH = Hproducts − Hreactants

At constant pressure:

• qₚ = ΔH
• This means the heat transferred during a constant-pressure process corresponds to the enthalpy change.

Most reactions performed in an open beaker occur under essentially constant atmospheric pressure, which is why enthalpy is used so often in chemistry.

An endothermic process absorbs heat from the surroundings.

For an endothermic process:

• qsystem > 0
• and, under constant pressure:
• ΔH > 0
• Because the system gains energy, its final state has a higher enthalpy than its initial state.

Examples of endothermic processes include melting, vaporization, sublimation, and many chemical reactions.

• A general energy diagram looks like:
• lower-energy reactants → higher-energy products
• Therefore:
• Hproducts > Hreactants
• and:
• ΔH = positive

If the system absorbs heat, the surroundings lose heat.

• This can cause the surroundings to become cooler.
• For example, if a reaction mixture feels cold, the reaction may be absorbing thermal energy from the surroundings.
• Be careful with the wording:
• “The reaction got cold” is not the strongest explanation.
• A better explanation is:
• The reaction was endothermic, so energy was transferred as heat from the surroundings into the system.

An exothermic process releases heat into the surroundings.

For an exothermic process:

• qsystem < 0
• and:
• ΔH < 0
• The products have lower enthalpy than the reactants.
• A general energy relationship is:
• higher-energy reactants → lower-energy products
• Therefore:
• Hproducts < Hreactants
• and:
• ΔH = negative
• Many combustion reactions are strongly exothermic.
• If the surroundings become warmer during a reaction, the reaction may have released thermal energy into them.

Energy diagrams visually compare the energy of the initial and final states.

For an endothermic process:

• products are higher than reactants.
• ΔH > 0

For an exothermic process:

• products are lower than reactants.
• ΔH < 0

Do not confuse the overall energy difference on this type of diagram with activation energy from Unit 5.

Activation energy describes the energy barrier needed for the reaction pathway.

• ΔH describes:
• energy of products − energy of reactants
• A reaction can have a very large activation energy and still be strongly exothermic.

A useful idea that appears throughout thermodynamics is:

• breaking attractive interactions requires energy
• while:
• forming attractive interactions releases energy

For chemical reactions:

• breaking bonds → requires energy
• forming bonds → releases energy

For physical processes:

• separating particles against intermolecular forces → requires energy
• allowing stronger attractions to form → releases energy

Whether the overall process is endothermic or exothermic depends on the balance of all energy absorbed and released.

Do not confuse temperature with heat.

Do not say endothermic means “cold” and exothermic means “hot.” Those are possible observations of the surroundings, but the actual definitions concern energy transfer.

• Remember:

endothermic → system absorbs heat → ΔH positive

exothermic → system releases heat → ΔH negative

Do not confuse ΔH with activation energy.

Think from the perspective of the system:

• ENDO = energy ENters
• EXO = energy EXits

Unit 6 becomes much easier when you see that almost every topic is another way of applying energy conservation.

• Start with:
• system + surroundings
• Energy can move between them as heat.

If temperature changes:

• q = mcΔT

If a phase changes:

• q = nΔHphase

If a reaction occurs at constant pressure:

• qreaction = ΔHrxn

If multiple reactions are combined:

• Hess's law → add the ΔH values

If formation data are provided:

• ΔH°rxn = products − reactants

If bond data are provided:

• ΔH ≈ broken − formed
• All of these are different tools for answering essentially the same question:
• How much energy is transferred during this process?

Unit 2 taught you that atoms form bonds because certain bonded arrangements are more stable than separated atoms.

• Unit 6 adds the energy perspective:
• breaking bonds requires energy
• while:
• forming bonds releases energy
• This explains why a reaction's ΔH depends on the bonds that disappear and the bonds that form.

Unit 3 taught you about London dispersion forces, dipole-dipole interactions, hydrogen bonding, and ion-dipole attractions.

• Unit 6 explains what happens energetically when those attractions are overcome.
• Stronger intermolecular forces generally require more energy to separate particles.
• For example, a liquid with strong intermolecular attractions will generally require more energy to vaporize than a comparable liquid with weaker attractions.
• That can correspond to a larger:
• ΔHvap

Thermochemistry calculations frequently require stoichiometry.

• For example:
• 2H₂ + O₂ → 2H₂O ΔH = −572 kJ
• This establishes a stoichiometric relationship:
• 2 mol H₂ ↔ −572 kJ
• 1 mol O₂ ↔ −572 kJ
• 2 mol H₂O formed ↔ −572 kJ
• If only 1 mol H₂ reacts:
• 1 mol H₂ × (−572 kJ / 2 mol H₂)
• = −286 kJ
• Energy can therefore be treated as another quantity connected through the balanced equation.

This distinction is extremely important:

Thermodynamics tells us about energy differences.

Kinetics tells us about reaction speed.

• A reaction can be strongly exothermic but still happen very slowly because it has a large activation-energy barrier.
• Therefore:
• large negative ΔH ≠ automatically fast reaction

Likewise, a catalyst can make a reaction faster by lowering activation energy but does not change ΔH.

Heat vs. temperature: Heat is energy transfer. Temperature reflects average kinetic energy.

System vs. surroundings: Always identify whose energy is being discussed before choosing the sign.

Endothermic: ΔH is positive.

Exothermic: ΔH is negative.

Calorimetry: If the solution gains heat, the reaction loses heat.

Temperature change: Always use Tf − Ti.

Units: Convert J and kJ correctly.

Heating curves: Use q=mcΔT on slopes and q=nΔHphase on plateaus.

Phase changes: Temperature stays approximately constant during an ideal phase transition.

Thermochemical equations: ΔH belongs to the equation exactly as written.

Reverse reaction: Reverse the sign of ΔH.

Multiply reaction: Multiply ΔH by the same factor.

Hess's law: Make sure all unwanted species actually cancel.

Formation enthalpy: Use products − reactants.

Standard elements: ΔH°f = 0 only for an element in its standard state.

Formation equations: Must form exactly 1 mol of the compound.

Bond enthalpy: Use broken − formed.

Bond counting: Count every relevant bond and coefficient.

Bond enthalpy accuracy: Average bond enthalpies normally give estimates.

Activation energy: Do not confuse activation energy with ΔH.

Catalysts: Catalysts change activation energy, not ΔH.

Before taking a Unit 6 test, you should understand the difference between a system and surroundings, as well as the difference between heat, temperature, internal energy, and enthalpy. You should be able to use the signs of q and ΔH to determine whether a process is endothermic or exothermic and interpret an energy diagram.

You should be able to explain thermal equilibrium at the particle level and understand that energy lost by one part of an isolated system must be gained by another.

For calorimetry, you should be able to use q=mcΔT, calculate an unknown specific heat or final temperature, understand coffee-cup and bomb calorimeters, apply qreaction = −qsurroundings, and calculate ΔH per mole of reaction.

For phase changes, you should know which changes are endothermic and exothermic, use q=nΔHphase, and correctly interpret a heating curve. You should know when energy changes average kinetic energy and when it changes potential energy associated with intermolecular attractions.

You should understand that ΔHrxn depends on how a balanced reaction is written. Reversing a reaction reverses ΔH, while multiplying the equation multiplies ΔH by the same amount.

You should be able to solve Hess's law problems by reversing, multiplying, and adding equations until they match a target reaction, then applying the same operations to the ΔH values.

You should know what standard enthalpy of formation means and understand why elements in their standard states have ΔH°f = 0. You should be able to calculate:

ΔH°rxn = ΣnΔH°f(products) − ΣnΔH°f(reactants)

Finally, you should understand bond enthalpy well enough to explain why breaking bonds requires energy, forming bonds releases energy, and calculate an approximate reaction enthalpy using:

ΔHrxn ≈ Σ(bonds broken) − Σ(bonds formed)

If you can move comfortably between energy diagrams, calorimetry, heating curves, balanced equations, Hess's law, formation enthalpies, and bond enthalpies, you understand the main logic behind AP Chemistry Unit 6.