Thermochemistry

Thermodynamics

Relationship and conversion between heat (Q) and work (W), depended on the laws of conservation and entropy

Thermochemistry

Thermodynamics in chemical and physical change

Energy

Ability to do work or supply heat

Potential energy

Energy due to the position

Kinetic energy

Energy due to the motion

Total energy

Potential + Kinetic energy

Central concepts

1. Energy is converted, not destroyed

2. Lower energy → More stable, favored over higher energy

Definition of system

A specific region or area, a collection of matter and energy

Definition of surrounding

Everything else than system is defined as the surrounding

Matter and energy in a system

Are always conserved, unless they are exchanged with the surrounding environment

Every process requires …

Energy

Energy divided into

1. Heat

2. Work

Heat

The energy transferred as a result of a difference in temperature between the system and the surrounding

Work

The energy transferred when an object is moved by a force

Types of work

1. Electrical work

2. PV work

Electrical work

Done by moving charged particle

PV work

Mechanical work done when the volume of the system changes in the presence of an external pressure

Open system

Exchange in energy and matter

Closed system

Only exchange in energy

Isolated system

No exchange

Q > 0

Endothermic. System receives heat

Q < 0

Exothermic. System releases heat

W > 0

Work done on the system

W < 0

Work done by the system

SI unit of energy

Joule (J)

1 J = …

1 kg.m²/s²

calorie

Energy needed to raise the temperature of 1g of water by 18^oC

1 cal = …

4,184 J

BTU

Energy needed to raise the temperature of 1lb of water by 18^oF

1 BTU = …

1055J

1 Calorie = …

1 Cal = 1000 calorie = 1kcal

Internal energy (E_{int})

Sum of potential and kinetic energies of all molecules, atoms and ions in the system

(Basically total energy but just inside the system)

Changes in internal energy (\Delta E_{int})

Total heat and total work that the system transfer

(The system exchange work and heat with the environment)

\Delta E_{int} = …

Q + W

W = …

P\Delta V

First law

In an isolated system, the total energy remains constant

Energy may be converted but not created nor destroyed

→ Total energy of universe is constant

Consequence of First law

1. In isolated system: \Delta E = \Delta E_{system} = \Delta E_{environment} = 0

2. If V = constant → \Delta E = Q

Enthalpy

Energy changes at constant pressure

→ Sum of internal energy (E) and PV work

H = …

E + PV

\Delta H = …

\Delta E + \Delta (PV)

State function

Depend only on the system’s current state and is independent of the path taken to reach that specific state

Piston-cylinder

\Delta H = Q - W + \Delta (PV)

Chemical reactions with gases cause volume expansion

W = P \Delta V → \Delta H = Q - P \Delta V + \Delta (PV)

In constant pressure

\Delta (PV) = P\Delta V → \Delta H = Q_{p} - P\Delta V + P\Delta V = Q_{p}

Pressures in reversible processes

External pressure = System pressure

\Delta H < 0

Exothermic

\Delta H > 0

Endothermic

Constant pressure involving gas

\Delta H = \Delta E + \Delta (PV) = \Delta E + \Delta nRT

\Delta H in reverse processes

\Delta H_{forward} = -\Delta H_{reverse}

R = …

1. 8.314 J/mol.K

2. 0.08206 L.atm/mol.K

Standard conditions

1. Gas, Liquid, Solid: P = 1 atm

2. Dissolved substance: C_M = 1M

3. Temperature (Not part of it) but usually, T = 298K

Phase

Matter that has the same physical and chemical

State

Physical arrangment of matters

^oC → ^oF

\ldots^{o}F=\dfrac95\ldots^{o}C+32

^oC → K

K = …^oC + 273.15

Heat capacity

Heat required to change its temperature by 1K

Heat capacity formula

\frac Q{\Delta T}J/K

Specific heat capacity

Heat required to change the temperature of 1g by 1K

Specific heat capacity formula

\frac Q{mass \times \Delta T}J/g.K

Molar heat capacity

Heat required to change the temperature of 1 mole by 1K

Molar heat capacity formula

\frac Q{mol \times \Delta T}J/mol.K

Calorimeter

Is used to measure heat released/absorbed by chemical/physical process

Constant pressure calorimetry

Heat transferred can be measured in a simple coffee-cup calorimeter

Coffee-cup calorimeter

1. Direct measuring \Delta H

2. Low precision

3. Used for most reactions

Constant volume calorimetry

Bomb calorimeter. Measure the heat of combustion reactions

Bomb calorimeter

1. Direct measuring \Delta H

2. Higher precision

3. Used for combustion reactions

Hess law

The enthalpy change of an overall process is the sum of the enthalpy changes of its individual steps

Standard enthalpy of formation

Enthalpy change for the process when one mole of a substance is formed

Reference state

\Delta H^o_f = 0kJ/mol

\Delta H^o_r formula

\sum n \Delta H^o_f (products) - \sum n \Delta H^o_f (reactants)