Unit 2 - Thermochemistry and Kinetics

Define thermochemistry

The study of the energy changes accompanying chemical transformations

Define enthalpy

all the energy required to create a system

List what energies make up enthalpy

• Potential energy of bonds and IMF

• Kinetic energy of particles

• Rest-mass energy (e = mc²)

• Work

What energy is mainly responsible for changes in enthalpy in a chemical reaction

Potential energy - change in potential energies of bonds and IMF

• Kinetic energies change as a RESULT, AFTER the fact

Describe enthalpy in terms of bonds broken and bonds formed

deltaH = (E _{forming bonds} + E _{breaking bonds})

• Where energy is released when forming bonds and energy is absorbed when breaking bonds

Define exothermic and endothermic

• Exothermic = release energy into surroundings (negative enthalpy)

• Endothermic = absorb energy from surroundings (positive enthalpy)

Differentiate heat and thermal energy

• thermal energy = kinetic energy driving the motion (translation, vibration, rotation) of molecules

• heat = TRANSFER of thermal energy between two materials (NOT TEMPERATURE)

Describe Hess’ Law

Enthalpy is a state function - pathway INDEPENDENT - value depends on current state and not HOW the system was made (intermediate steps)

• Hess’ Law = deltaH between reactants and products remains the same NO MATTER THE PROCESS ***so long as conditions are the same

What is a formation reaction

type of synthesis reaction where elements at their standard (SATP) states form a compound

• E.g. H₂ + 1/2O₂ = H₂O

What is standard enthalpy of formation

enthalpy required at SATP for a formation reaction

• Written as deltaH^degree_f - degree indicates at SATP

What must the formation enthalpy for an element in its SATP form be and why?

Must be 0 - no energy change involved

E.g. Mg_(s) → Mg_(s)

Draw a potential energy diagram for 2NO₂ → N₂O₄ that illustrates Hess’ Law

Define heat capacity and specific heat capacity

• heat capacity (J/C) = quantity of heat required to make temperature change by 1K

• specific heat capacity (J/gC) = heat capacity per g of a substance

***energy usually in joules

provide the equation for determining heat absorbed/released in a reaction

q=mc(T_f - T_i)

or, if heat capacity, q=c(T_f - T_i)

Describe the first and second laws of thermodynamics and how they expand our understanding of thermodynamics

1. Law of Conservation of Energy - total E in universe is constant, must be transferred or transformed but not destroyed

   1. q_system = -q_surroundings → ndeltaH = -q_surroundings

2. Heat always flows from hot → until the two materials reach thermal equilibrium

   1. so, deltaT for two substances in contact should be same, or at least T_i

List and describe the three types of systems

1. Open = matter and heat flow in and out

2. Closed = matter contained, heat flow in and out

3. Isolated = both matter and heat contained

Define spontaneous and non-spontaneous processes

1. spontaneous = occurs on its own, does not require additional force

2. non-spontaneous = requires force/energy

Relate spontaneity and potential energy

if a process goes from high to low potential energy it is spontaneous

How can there be spontaneous endothermic processes? E.g. water evaporating

Spontaneous = high to low potential energy

But enthalpy doesn’t reflect this - energy is absorbed

There must be another kind of energy at play - ENTROPY

Define entropy

Entropy (S) = degree of molecular randomness/disorder in a system

• more entropy w freedom of motion vs restricted

Describe how entropy and enthalpy are related

high entropy = low/negative enthalpy

Describe entropy in terms of probability

Entropy = a description of the number of arrangements available to a system in a given state

• More ways that a certain state can be achieved = more likely it is to occur = higher entropy

*Nature proceeds towards the state that is the most likely to exist

Define microstate + p of an arrangement + p of a given microstate

number of different ways to observe or create a certain arrangement

p of an arrangement = microstates that create an arrangement/total possible microstates

p of microstate = (1/2)^# atoms

Describe the second law of thermodynamics in terms of entropy

entropy change of the universe cannot decrease

deltaS_universe >/ 0

a system can get more ordered, but the universe must get disordered

Give an equation for entropy of the universe

deltaS_universe = deltaS_system + deltaS_surroundings

Relate entropy of surroundings to the enthalpy of system

deltaS_surroundings = (-deltaH_system/T) - temperature dependent

What are three factors to consider when qualitatively determining the entropy of a system

1. state

   1. s - g = increase in entropy

2. size

   1. big = more restrictions and more bonds, small = less

3. moles

   1. more = higher entropy

What does it mean if deltaS is positive, negative or zero

1. positive = spontaneous

2. negative = non-spontaneous

3. zero = equilibrium process, no net change

Describe gibbs free energy + equation

total usable energy in a chemical system - enthalpy AND entropy

deltaG = deltaH - TdeltaS —— all for the SYSTEM

Relate deltaG and spontaneity

negative = spontaneous and exergonic

positive = non-spontaneous and endergonic

What is deltaG at the threshold temp

0

Describe the third law of thermodynamics and standard entropy values

3rd law = as a system approaches 0K, entropy approaches a constant number (often but not always 0)

standard entropy (S_degree) = the entropy change from 0K to room/current temp (molar quantity)