Thermodynamics Is it hot in here or what?. Energy Many forms and sources Thermochemistry is interested in heat exchanges Breaking bonds takes energy.

Thermodynamics

Is it hot in here or what?

Energy

• Many forms and sources• Thermochemistry is interested in heat

exchanges• Breaking bonds takes energy (endothermic)• Making bonds produces energy (exothermic)• SI Unit is the Joule (J), common unit calorie• 4.18J = 1 cal• 4.18 kJ = 1 kcal

Temperature and Energy Change

• Every substance has a unique ability to absorb and release energy.

• Called heat capacity. • Water has a very large value.• Specific Heat (cp) is the measure of how

much energy a substance will absorb or release to change its temp. 4.18J/g°C for water

Q = m • cp • ΔT

Energy = mass x spec. heat x temp. change

Ex: If 20 g of water changes its temp from 25.0C to 30.0C, what amount of energy is absorbed?

Q = (20.0g) x (4.18 J/g C) x (5.0 C)

Q = 418 J

Practice!!!!

Heat of Reaction

• Energy change that occurs during a chemical reaction

• Called enthalpy• Symbol (ΔH) change in enthalpy• Negative ΔH value = exothermic reaction• Positive ΔH value = endothermic reaction

Ene

rgy

Reaction coordinate

Reactants

Products

Overall energy change

Exothermic reaction

Endothermic reaction

Ene

rgy

Reaction coordinate

Overall energy change

Heat of Formation ∆Hf°

• The energy change associated with the formation of one mole of a compound from its elements.• Sometimes uses fractional coefficients• Ex. H2(g) + ½ O2(g) → H2O (l)

• ∆Hf° = -286 kJ/mol exothermic (comp. is stable)

Thermochemical Equation

An equation which includes the heat energy as a reactant (endothermic) or product (exothermic)

H2(g) + ½ O2(g) → H2O (l) + 286kJ– Notice that the energy is written as a + even though the

∆Hf° = -286 kJ/mol

Calculation of ΔH reaction

• Two methods – both work• Hess’ Law and a shortcut

• Learn both (use the shortcut more often)

Hess’ Law - finding ΔH

• Says that if you can write a reaction in steps that add up to the original reaction, the ΔH of the overall reaction is the sum of the ΔH of the steps. (route independent)

• Guidelines• 1. If the reaction is reversed, so is the sign of ΔH.• 2. If coefficients are multiplied, so is ΔH.

Example: find ΔH CH4(g) + 2O2(g) → CO2(g) + 2H2O(l)

Solution

1. Write ΔHf eq. for each compound

C(s) + 2H2(g) → CH4(g) ΔH = -74.8 kJ

C(s) + O2(g) → CO2(g) ΔH = -393.5 kJ

H2(g) + ½ O2(g) → H2O(l) ΔH = -285.8 kJ

solution for:CH4(g) + 2O2(g) → CO2(g) + 2H2O(l)

2. Multiply by any coefficients in the orig. eq.

2H2(g) + O2(g) → 2H2O(l) ΔH = -571.6 kJ

3. Reverse reactions to get reactants on left

CH4(g) → C(s) + 2H2(g) ΔH = +74.8 kJ

4. Add the resulting equations together……

Orig. eq.CH4(g) + 2O2(g) → CO2(g) + 2H2O(l)

CH4(g) → C(s) + 2H2(g) ΔH = +74.8 kJ

2H2(g) + O2(g) → 2H2O(l) ΔH = -571.6 kJ

C(s) + O2(g) → CO2(g) ΔH = -393.5 kJ

_____________________________________

CH4(g) + 2O2(g) → CO2(g) + 2H2O(l)

(same as orig.) so add up ΔH = -890.3 kJ

Hess’ Law

• Tedious, but it works…….

• Practice…….

ALTERNATIVE TO HESS’ LAW

• Handy shortcut to use for calculations.

• ΔH = Σ ΔHf (products) - Σ ΔHf (reactants)

• Remember to multiply values by any coefficients in the balanced equation.

• ΔH = (-393.5 kJ + 2 x -285.8 kJ) – (-74.8 kJ)• ΔH = -890.3 kJ

SPONTANEOUS REACTIONS

• Proceed without outside assistance (beyond the initial Ea) -- they just happen.

• Spontaneous chemical reactions generally occur if the products have lower PE than the reactants.

• -sometimes this is not the case ( ex---H2O(s) only forms if the temp is less than 0 C)

• -reason for this is the other driving force in nature...........

Entropy

• A measure of the disorder or randomness of a system.

• -represented by S (change in ..... ΔS)

• ΔS = Σ Sf(products) - Σ Si(reactants)

All entropies (Sf and Si) are positive, but …ΔS can be negative or positive.

• if ΔS is negative ===> Less disorder (lower entropy)• Example: crystallization of a liquid

• if ΔS is positive ===> More disorder (more entropy) – Example: evaporation of a liquid….. favored in

nature!

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Prediction of entropy change

• synthesis of compounds decreases entropy (due to bonding)

• decomposition of compounds increases entropy

• mixing a solute and solvent increases entropy

• Solid→ Liquid → Gas increases entropy• Ie. evaporation of a liquid increases entropy

So>>>>>>

• Two forces influence the direction of a spontaneous reaction: ΔH and ΔS

• when ΔH decreases; lower energy results.

• when ΔS increases; higher disorder results

These two factors compete

• Like a tug of war.

• Usually the enthalpy prevails.

4 Possible Cases

ΔH ΔS Result: - + reaction favored

+ - no reaction

- - rx if ΔH is large

+ + rx if ΔS is large

Gibbs Free Energy ΔG

• An equation called the “Gibbs equation” compares the values of ΔH and Δ S.

• ΔG = ΔH - T ΔS• if ΔG is negative; the reaction is

spontaneous• if ΔG is positive: the reaction is not

spontaneous, and will not occur.

Alternative to the Gibbs equation:

• ΔG = ΣΔGf (products) - ΣΔGf (reactants)

Collision Theory•In order to react molecules and atoms must touch each other.

•They must hit each other hard enough to react.

•Anything that increase these things will make the reaction faster.

Things that Effect Rate

TEMPERATUREHigher temperature = faster particles.More and harder collisions.Faster Reactions.CONCENTRATIONMore concentrated = molecules closer together.Collide more often.Faster reaction.

Things that Effect RatePARTICLE SIZE

Molecules can only collide at the surface.

Smaller particles = bigger surface area.

Smaller particles = faster reaction.

Smallest possible is atoms or ions.

Dissolving speeds up reactions.

Getting two solids to react with each other is slow.

Things that Effect RateCATALYSTS- substances that speed up a reaction without being used up.(enzyme).

Speeds up reaction by giving the reaction a new path.The new path has a lower activation energy.More molecules have this energy.The reaction goes faster.Inhibitor- a substance that blocks a catalyst.