AP®︎ Chemistry: Unit 6 Flashcards

The capacity to do work or transfer heat.

Energy due to motion; in particles, it increases as temperature increases.

Stored energy due to position or interactions, especially in chemical bonds and attractions.

Heating, cooling, phase changes, and reactions can absorb or release energy.

Solution formation absorbs or releases energy depending on relative interaction strengths before and after dissolving.

Temperature changes show energy transfer between system and surroundings during physical or chemical changes.

Energy is conserved: it cannot be created or destroyed, only transferred or converted between forms.

The system is the part being studied, such as the reacting chemicals; the surroundings are everything else that can exchange energy with it.

A graph of energy versus reaction progress showing reactants, products, and any energy barrier.

Vertical axis is potential energy; horizontal axis is reaction or process progress.

The starting and ending energy levels shown for substances before and after a process.

The highest-energy, least stable arrangement of atoms at the top of the energy barrier.

Physical changes can be graphed like reactions, showing energy absorbed or released during phase change.

Energy absorbed or released during a phase change without a temperature change.

The energy difference between the reactants and the peak, equal to the minimum energy needed to react.

On an energy diagram, exothermic processes have products lower in enthalpy than reactants and a negative ΔH, while endothermic processes have products higher than reactants and a positive ΔH.

Melting, vaporization, and sublimation are endothermic and absorb energy; freezing, condensation, and deposition are exothermic and release energy.

A condition where two bodies can exchange energy through collisions between their particles.

Energy as heat flows from the warmer body to the cooler body.

Energy still moves both ways, but equal transfers cause no overall change.

If A is in thermal equilibrium with B, and B with C, then A with C.

Collisions transfer kinetic energy from faster particles to slower particles in thermal contact.

Thermal equilibrium is the state in which objects in thermal contact have the same temperature and average kinetic energy, so no net heat is transferred between them.

The measurement of heat transfer by tracking temperature changes in an insulated system.

q = mcΔT, where heat equals mass times specific heat times temperature change.

The heat required to raise 1 gram of a substance by 1 degree Celsius.

Heat capacity is heat per degree for an object; molar heat capacity is heat per degree per mole.

The difference between final and initial temperature, calculated as Tfinal minus Tinitial.

In an isolated calorimeter, energy released by one part equals energy absorbed by another.

Positive q means heat is absorbed; negative q means heat is released.

For equal heat and mass, a larger specific heat gives a smaller temperature change.

Set heat lost equal to heat gained, then solve q = mcΔT for c.

Systems change energy through heating or cooling, phase changes, and chemical reactions.

An insulated constant-pressure setup where measured heat equals the reaction’s enthalpy change.

A calorimetry method using solution temperature changes to identify heat flow during dissolving.

A graph of temperature versus heat added showing warming slopes and flat phase-change plateaus.

A graph of temperature versus heat removed showing cooling slopes and flat phase-change plateaus.

A pure substance’s temperature stays unchanged while added or removed energy changes its phase.

The energy per mole required to melt or vaporize a substance at constant temperature.

Opposite phase changes have equal magnitudes and opposite signs, such as vaporization and condensation.

Use q = nΔH for phase changes, with moles times the molar enthalpy.

Use q = mcΔT on slopes and q = nΔH or mH on plateaus, then sum steps.

Vaporization usually requires more energy than melting because more intermolecular attractions must be overcome.

The energy released per mole when a gas condenses or a liquid freezes.

Enthalpy change per mole of reaction as written in the balanced equation, usually in kJ/mol.

Thermal energy transferred between a system and surroundings, measured in joules or kilojoules.

Calculate reaction heat by multiplying moles of reacting substance by molar enthalpy change.

Rearrange q = nΔH to n = q/ΔH, keeping units and signs consistent.

Use the moles of the limiting reactant to determine the total heat of reaction.

Energy is exchanged until products and surroundings reach the same final temperature.

Differences in chemical potential energy appear as particle kinetic energy changes and temperature changes.

At constant pressure, the reaction enthalpy change equals the heat transferred, q.

A thermodynamic quantity representing a system’s heat content at constant pressure; enthalpy change, ΔH, is the difference between final and initial enthalpy values.

Endothermic processes absorb heat from the surroundings and have positive ΔH; exothermic processes release heat to the surroundings and have negative ΔH.

The overall enthalpy change, ΔH, is the difference between products and reactants; for a reaction at constant pressure, it is the heat absorbed or released.

Triple bonds are strongest, double bonds intermediate, and single bonds weakest, with increasing bond enthalpy.

Shorter bonds are generally stronger and have higher bond enthalpies than longer, weaker bonds.

Draw reactant and product structures, count each bond type, then apply the bond enthalpy equation.

The average energy needed to break one mole of a specific gas-phase bond.

Subtract bond energies of bonds formed from bonds broken to estimate reaction enthalpy.

Breaking bonds absorbs energy, while forming bonds releases energy; these changes alter chemical potential energy and contribute to the reaction enthalpy.

Calculate using ΣnΔHf°(products) minus ΣmΔHf°(reactants), including stoichiometric coefficients.

Multiply each substance’s formation enthalpy by its balanced-equation coefficient before summing.

Use the value matching the substance’s phase, since liquids, gases, and solids have different values.

Enthalpy change when one mole of a compound forms from elements in standard states.

Elements in their most stable standard states have standard enthalpies of formation equal to zero.

Break an overall reaction into known reactions, then manipulate and add them to match the target.

Reverse reaction: change ΔH sign; multiply coefficients: scale ΔH; add reactions: add ΔH values.

Species appearing on both sides of added steps cancel, leaving only the overall reaction.

The overall enthalpy change equals the sum of step enthalpies because energy is conserved.