Unit 6 Beta

Thermochemistry: the one-page sheet

6.1 Endothermic and Exothermic Processes

Every process moves energy between a system and its surroundings. Exothermic processes release energy (ΔH < 0, surroundings warm); endothermic processes absorb it (ΔH > 0, surroundings cool). At the particle level, breaking bonds or attractions absorbs energy and forming them releases it.

The system is the part you study, usually the reacting particles; the surroundings are everything else, including the water they are in.
Exothermic: energy leaves the system, the surroundings warm, ΔH is negative.
Endothermic: energy enters the system, the surroundings cool, ΔH is positive.
Breaking bonds or attractions absorbs energy; forming them releases energy. The balance sets the sign of ΔH.
Physical changes count: melting and evaporating are endothermic; freezing and condensing are exothermic.

the system must absorb energy for that step the system releases energy for that step energy leaves the system: exothermic, ΔH < 0, and the surroundings warm energy enters the system: endothermic, ΔH > 0, and the surroundings cool a rising temperature means an exothermic process, a falling one an endothermic process

system (thermodynamics): The system is the part of the universe being studied, usually the reacting particles; the surroundings are everything else, such as the water, container and air.
enthalpy: A measure of a system's energy; its change, ΔH, is the energy the system absorbs (positive) or releases (negative) as heat at constant pressure.
exothermic: Describes a process that releases energy from the system to the surroundings, so the surroundings warm and ΔH is negative.
endothermic: Describes a process that absorbs energy from the surroundings into the system, so the surroundings cool and ΔH is positive.

6.2 Energy Diagrams

An energy diagram shows reactants and products at their enthalpy levels. The vertical gap between the ends is ΔH, negative when products are lower (exothermic) and positive when higher (endothermic). The hump between them is the activation energy, which a catalyst lowers without changing ΔH.

An energy diagram plots enthalpy against the progress of the reaction.
ΔH = H(products) − H(reactants): products lower means exothermic (ΔH < 0); products higher means endothermic (ΔH > 0).
The hump gives the activation energy, the climb from the starting level to the top. It controls speed, not ΔH.
Reverse reaction: same picture read backward; ΔH changes sign and Ea(reverse) = Ea(forward) − ΔH.
A catalyst lowers the hump and leaves both end levels, and ΔH, unchanged.

reactants and products sit at definite levels on the diagram products below reactants give a negative ΔH (exothermic); products above give a positive ΔH (endothermic) a hump between the levels marks the activation energy ΔH flips sign and the activation energy is measured from the other side

energy diagram: A graph of a system's enthalpy against the progress of a reaction; the gap between the reactant and product levels is ΔH, and the hump between them gives the activation energy.

6.3 Heat Transfer and Thermal Equilibrium

When objects at different temperatures touch, collisions pass energy from the faster particles of the hotter object to the slower particles of the cooler one. The flow continues until both reach the same temperature, thermal equilibrium. In an insulated system, the energy one object loses equals the energy the other gains.

Heat is energy transferred because of a temperature difference; it always flows from warmer to cooler.
At the particle level, heat transfer happens by collisions: fast particles speed up slow ones.
Thermal equilibrium: equal temperatures, no net energy flow, though collisions continue.
In an insulated (isolated) system, energy lost by one object = energy gained by the other.
The final temperature lies between the starting temperatures.

particles at the boundary collide, the faster ones from the hot object with the slower ones from the cold object energy flows from the hotter object to the cooler one their temperatures move toward each other energy passes equally both ways, so there is no net flow: thermal equilibrium energy lost by the hot object equals energy gained by the cold one

heat transfer: The flow of energy from a warmer object to a cooler one by particle collisions; it continues until both reach the same temperature (thermal equilibrium), and energy lost by one equals energy gained by the other.

6.4 Heat Capacity and Calorimetry

Calorimetry measures the energy of a process from a temperature change. The energy gained or lost by the solution is q = mcΔT; the reaction's q has the opposite sign. Dividing by the moles that react gives ΔH per mole, with a negative sign when the solution warms.

Specific heat capacity, c, is the energy to warm 1 g by 1 °C; water's is 4.18 J/(g·°C). Heat capacity is for a whole object: C = m × c.
q = m × c × ΔT, with ΔT = T(final) − T(initial) and m the mass of everything that changes temperature.
In a calorimeter, q(reaction) = −q(solution): a warmer solution means an exothermic reaction.
ΔH per mole = q(reaction) ÷ moles that react. Watch kJ vs J and the sign.
Energy lost to the room makes the measured |ΔH| too small.

almost all the energy it releases or absorbs goes into or out of the solution q(solution) = m × c × ΔT gives the energy the solution gained (positive) or lost (negative) q(reaction) = −q(solution) ΔH per mole = q(reaction) ÷ moles of the limiting reactant (or product named) the calculated |ΔH| comes out too small or too large in a predictable direction

specific heat capacity: The energy needed to raise the temperature of 1 g of a substance by 1 °C; for liquid water, 4.18 J/(g·°C). It appears in q = mcΔT.
heat capacity: The energy needed to raise the temperature of a whole object by 1 °C, in J/°C; for one substance, mass × specific heat.
calorimetry: Measuring the energy of a process from the temperature change of a known mass of material, usually water, in an insulated container (a calorimeter); q(process) = −q(water).

6.5 Energy of Phase Changes

During a phase change, energy overcomes (or forms) attractions between particles, so the temperature stays constant. The energy is n × ΔHfus for melting and n × ΔHvap for boiling, with ΔHvap the larger. A heating curve combines sloped parts (q = mcΔT) and flat parts (q = nΔH).

The enthalpy of fusion (ΔHfus) is the energy to melt 1 mol of solid at its melting point; the enthalpy of vaporization (ΔHvap) is the energy to vaporize 1 mol of liquid.
Melting, vaporizing and subliming absorb energy (q = +nΔH); freezing and condensing release it (q = −nΔH).
ΔHvap > ΔHfus: vaporizing separates molecules almost completely.
On a heating curve, slopes use q = mcΔT and flat parts use q = nΔH; add the steps.
Stronger attractions between particles mean larger ΔHfus and ΔHvap.

its particles vibrate faster and its temperature rises until it reaches the melting point the solid melts at constant temperature; potential energy rises, kinetic energy does not the liquid warms until the boiling point the boiling plateau is longest: ΔHvap is larger than ΔHfus condensing and freezing release the same energies: q = −nΔH

enthalpy of fusion: The energy absorbed to melt 1 mol of a solid at its melting point (ΔHfus; water 6.01 kJ/mol). Freezing releases the same amount.
enthalpy of vaporization: The energy absorbed to vaporize 1 mol of a liquid at its boiling point (ΔHvap; water 40.7 kJ/mol). Condensing releases the same amount.
heating curve: A graph of temperature against energy added (or time); sloped parts show one phase warming (q = mcΔT), and flat parts show a phase change at constant temperature (q = nΔH).

6.6 Introduction to Enthalpy of Reaction

The enthalpy of reaction, ΔH, is the energy change for a reaction exactly as written, per mole of reaction. It scales with the coefficients and changes sign when the reaction is reversed. Combined with stoichiometry, it gives the energy released or absorbed by any amount of reactant.

A thermochemical equation is a balanced equation with states and the enthalpy of reaction, ΔH, per mole of reaction as written.
Multiply the equation, multiply ΔH; reverse it, change the sign of ΔH.
Energy for an amount: convert to moles, then use ΔH ÷ coefficient as a conversion factor (kJ per mol of that substance).
Base the energy on the limiting reactant.
States matter: forming H₂O(g) instead of H₂O(l) releases less energy.

ΔH is per mole of reaction: the mole amounts the coefficients give twice the energy is transferred, so ΔH scales with the coefficients the energy flows the other way, so ΔH changes sign q = moles × (ΔH ÷ that substance's coefficient), using the limiting reactant q(surroundings) = −q(reaction)

enthalpy of reaction: ΔH for a reaction exactly as written, in kJ per mole of reaction (the amounts given by the coefficients); a thermochemical equation lists it after the balanced equation.

6.7 Bond Enthalpies

Bond enthalpies give an estimate of ΔH: add the bond enthalpies of the bonds broken, then subtract those of the bonds formed. Breaking bonds absorbs energy and forming them releases it, so a reaction that forms stronger bonds than it breaks is exothermic.

A bond enthalpy is the energy needed to break 1 mol of a bond in gas-phase molecules; it is always positive.
ΔH ≈ Σ(bonds broken) − Σ(bonds formed). Broken first; this is not "products minus reactants".
Draw the structures and count every bond, including coefficients: 2 O₂ means 2 O=O.
Multiple bonds are stronger than single bonds between the same atoms.
Average values make the answer an estimate; it is meant for gas-phase reactions.

breaking a bond absorbs energy, its bond enthalpy (always positive) forming the bond releases the same amount of energy ΔH ≈ Σ(bonds broken) − Σ(bonds formed) more energy is released than absorbed: exothermic the result is an estimate, close to but not equal to the measured ΔH

bond enthalpy: The energy needed to break 1 mol of a given bond in gas-phase molecules, always positive; tables list averages, and ΔH ≈ Σ(bonds broken) − Σ(bonds formed).

6.8 Enthalpy of Formation

A standard enthalpy of formation is ΔH for making 1 mol of a compound from its elements in their standard states; elements in their standard states have ΔH°f = 0. Any reaction's ΔH° is the sum of ΔH°f of the products minus the sum for the reactants, each multiplied by its coefficient.

The standard enthalpy of formation, ΔH°f, is ΔH for making 1 mol of a compound from its elements in their standard states (most stable form at 1 atm, usually 25 °C).
ΔH°f = 0 for an element in its standard state: O₂(g), H₂(g), C(graphite), Br₂(l), Fe(s).
ΔH°rxn = Σ n ΔH°f(products) − Σ n ΔH°f(reactants): products minus reactants, each times its coefficient.
States matter: H₂O(l) −285.8 and H₂O(g) −241.8 kJ/mol.
Bond enthalpies: broken − formed. Formation: products − reactants. Do not mix them up.

each compound gets one tabulated number; elements in their standard states get zero taking reactants apart costs −Σ ΔH°f(reactants); building products costs +Σ ΔH°f(products) ΔH°rxn = Σ n ΔH°f(products) − Σ n ΔH°f(reactants) the result is more exact than a bond-enthalpy estimate

standard enthalpy of formation: ΔH°f: the enthalpy change for making 1 mol of a compound from its elements in their standard states (their most stable forms at 1 atm, usually 25 °C); it is zero for an element in its standard state.

6.9 Hess's Law

Because enthalpy is a state function, the ΔH of a reaction is the same whatever route connects its reactants and products. Hess's law adds known reactions, reversed or scaled as needed, to build a target reaction, and adds their ΔH values, adjusted the same way, to get its ΔH.

Enthalpy is a state function: ΔH depends only on the start and end, not the route.
Hess's law: if equations add up to the target, their ΔH values add up to the target's ΔH.
Reverse an equation, change the sign of ΔH. Multiply an equation, multiply ΔH by the same factor. Do both when needed.
Work substance by substance: put each target substance on the correct side in the correct amount, then check that everything else cancels.
Formation (6.8) is Hess's law with a route through the elements.

it is a state function they have the same overall ΔH its ΔH is the sum of their ΔH values (Hess's law) its ΔH is reversed in sign or scaled by the same factor cancel, leaving exactly the target equation

Hess's law: Because enthalpy is a state function (it depends only on the start and end states, not the route), the ΔH of a reaction equals the sum of the ΔH values of any steps that add up to it.