05 Chemical Energetics Flashcards
Standard Enthalpy Change of Reaction, ΔHr⦵
Definition
Energy change in a chemical reaction when the molar quantities of reactants stated in the chemical equation reacts under standard conditions of 1 Bar & 298K.
Notes:
1. ΔHr⦵ can be +ve/ -ve.
2. ΔHr⦵ is OFTEN used to predict stability of compounds.
3. The more -ve ΔHr⦵ is, the more STABLE the compound is aka the compound is less likely to decompose into its constituent elements.
Standard Enthalpy Change of Formation of an Element, ΔHf⦵ (element)
Definition
Based on ΔHf⦵ (substance), ΔHf⦵ (element) = 0.
Standard Enthalpy Change of Formation of a Substance, ΔHf⦵ (substance)
Definition
The energy change when 1 mole of the pure substance in a specified state is formed from its constituent elements in their standard states under standard conditions of 1 Bar & 298K.
Note: ΔHr⦵ (substance) can be +ve/-ve.
Standard Enthalpy Change of Neutralisation, ΔHneut⦵
Definition
The energy change when the Acid and Base react to form 1 mole of H2O under standard conditions of 1 Bar & 298K.
Note: ΔHneut⦵ can be +ve/-ve as it depends on the strength of the Acid and Base used.
Standard Enthalpy Change of Atomisation of an Element, ΔHatom⦵ (element)
Definition
The energy absorbed when 1 mole of gaseous atoms is formed from the element under standard conditions of 1 Bar & 298 K.
Note(s):
1. ΔHatom⦵ is always +ve as bond breaking occurs.
Standard Enthalpy Change of Atomisation of a Compound, ΔHf⦵ (compound)
Definition
The energy absorbed when 1 mole of the compound is converted to gaseous atoms under standard conditions of 1 Bar and 298K.
Note(s):
1. ΔHatom⦵ (compound) is +ve as bond breaking occurs.
Bond Dissociation Energy, BDE
Definition
The energy required to break 1 mole of a particular X-Y bond in a particular compound in gaseous state.
Note(s):
BDE is usually +ve as bond breaking occurs.
Bond Energy, BE
Definition
The average energy required to break 1 mole of the X-Y bond in the gaseous state.
Note(s):
1. BE of X-Y bond in the Data Booklet is an average value of the different X-Y BE obtained from different compounds containing the X-Y bond.
2. BE of diatomic gas, X2, BE of the X-X bond = ΔHatom⦵
(1st) Ionisation Energy, IE
Definition
The energy required to remove 1 mole of electrons from 1 mole of gaseous atoms to form 1 mole of singly positively-charged gaseous ions.
(1st) Electron Affinity, EA
Definition
The energy change when 1 mole of gaseous atoms acquires 1 mole of electrons to form 1 mole of single negatively-charged gaseous ions.
Note(s):
1st EA is usually -ve as the energy released when the nucleus attracts the additional electron > energy taken in to overcome inter-electronic repulsion.
Lattice Energy, ΔHlatt⦵/ LE
Definition
The energy released when 1 mole of the solid ionic compound is formed from its constituent gaseous ions under standard conditions of 1 Bar & 298K.
Note(s):
LE is always -ve.
Standard Enthalpy Change of Hydration, ΔHhyd⦵
Definition
The energy released when 1 mole of gaseous ions is dissolved in H2O to form 1 mole of aqueous ions under standard conditions of 1 Bar and 298K.
Note(s):
1. ΔHhyd⦵ is always -ve
Standard Enthalpy Change of Solution, ΔHsoln⦵
Definition
The energy change when 1 mole of substance is completely dissolved in a solvent to form an infinitely dilute solution under standard conditions of 1 Bar & 298K.
Note(s):
1. An infinitely dilute solution is one where the addition of solvent does not produce any further ΔH.
2. ΔHsoln⦵ can be +ve/ -ve.
3. For -ve ΔHsoln⦵, salt is likely to be soluble in H2O and vice versa.
Hess’ Law
Definition
The enthalpy change of a reaction is determined by the initial and final states of the system and is independent of the pathways taken.
OR
The enthalpy change of a chemcial reaction is the same regardless of whether the reaction takes place in one step or several steps, provided the initial and final states of the reactants and products are the same.
Common Hess’ Law Equations
Given ΔHf⦵ data
ΔHr⦵ = ΣΔHf⦵ (products) - ΣΔHf⦵ (reactants)
Mnemonic: Fair Price Resells
Common Hess’ Law Equations
Given ΔHc⦵ data
ΔHr⦵ = ΣΔHc⦵ (reactants) - ΣΔHc⦵ (products)
Mnemonic: Crazy Ripe Potatoes
Common Hess’ Law Equations
Given ΔHhyd⦵ & LE
ΔHsol⦵ = ΣΔHhyd⦵ - LE
Common Hess’ Law Equations
Given BE
ΔHr⦵ = BE(reactants) - BE(products) = BB - BF
The Born-Haber Cycle
Normally represented by an Energy-Level Diagram
F - ΔHf⦵
A - ΔHatom⦵
B - BE
I - IE
E - EA
L - LE
Standard Enthalpy Change of Combustion, ΔHc ⦵
Definition
The energy released when 1 mole of the substance is completely burnt in excess O2 under standard conditions of 1 Bar & 298 K.
(2nd) Ionisation Energy, IE
Definition
The energy required to remove 1 mole of electrons from 1 mole of singly positively-charged ions to form 1 mole of doubly positively-charged ions.
(2nd) Electron Affinity, EA
Definition
The energy absorbed when 1 mole of singly negatively-charged ions acquires 1 mole of electrons to form 1 mole of doubly negatively-charged gaseous ions.
Note(s):
2nd EA is always +ve as energy is required to overcome the repulsion between 2 -ve-charged species.
Enthalpy, H
Definition
A measure of the energy content of the system. The higher the energy content of a system, the more unstable it is.
Enthalpy Change, ΔH
The change in energy content of a process in a system at constant pressure.
Formula: ΔH = ΣH(products)- ΣH(reactants)
Units: kJ mol-1
Spontaneous Process
Definition
A process that once started will continue without any external assistance.
Entropy, S
Definition
A measure of the disorder of matter and energy in the system. The more ways matter in the system can be arranged and energy in the system can be dispersed, the more disorder the system is and the larger its entropy.
AKA
A measure of the disorder of the system. The more disordered the system is, the larger its entropy.
Entropy Change, ΔS
Definition
A measure of the change in entropy in a system.
ΔS < 0 = decrease in S
ΔS > 0 = increase in S
Entropy, S
Factor 1: Change in Temperature, T
- T increases
- Boltzmann energy distribution curve broadens
- More possible energy states a particle can adopt at a higher temperature
- S increases
Entropy, S
Factor 2: Change in State
Change in State
- (s) to (ℓ)
- Order of (s) destroyed
- Particles in (ℓ) > randomly arranged
- Disorder increases
- S increases
- (ℓ) to (g)
- (g) is most disordered
- large increase in volume
- ΔS from (ℓ) to (g) > ΔS from (s) to (ℓ)
Entropy, S
Factor 3: Change in no. of Particles
- no. of Particles increases
- more ways to arrange particles
- more ways to distribute energy within the system
- greater S
S increases significantly in processes that increase mol of (g) particle
Entropy, S
Factor 4: Mixing of Particles
- Gas, (g), Particles
- increase in volume occupied
- more ways for (g) particles to arrange in a larger volume
- S increases
b. Liquid (ℓ), Particles
- (ℓ) of similar polarity are mixed
- molecules are randomly mixed
- more ways for (ℓ), particles to arrange in a larger volume
- S of mixture > S of liquids separately
c. Dissolution of an Ionic Solid in H2O
- Previously rigid ions can move about freely during the disruption of (s)
- S increases
- H2O molecules are arranged orderly around the ions during the hydration of (s)
- S decreases
- Weigh the ΔS of Disruption and Hydration to determine ΔS
a. & b. normally increases S while c. can cause an increase and decrease
Gibbs Free Energy Change, ΔG
Formula
ΔG = ΔH - TΔS
T = temperature
ΔS = entropy change
ΔH = enthalpy change
Unit(s:)
ΔG: kJ mol-1
ΔH: kJ mol-1
T: K
ΔS: J mol-1
* Convert the unit of ΔS to kJ mol-1
Standard Gibbs Free Energy Change, ΔG⦵
Formula
ΔG⦵ = ΔH⦵ - TΔS⦵
assuming no variation of ΔH⦵ & ΔS⦵ with T
Relationship between ΔG & Spontaneity of a reaction
ΔG < 0: Exergonic reaction
- Reaction Spontaneous/ Thermodynamically Feasible
ΔG > 0: Endogonic reaction
- Reaction not Spontaneous/ Thermodynamically Feasible
ΔG = 0: System is at equilibrium
ΔG does not indicate Kinetic Feasibility which is dependent on Activation Energy, EA
Effects of Temperature on ΔG
- ΔH < 0, ΔS > 0: ΔG < 0 at all T
- ΔH > 0, ΔS > 0: ΔG < 0 at high T
- ΔH < 0, ΔS < 0: ΔG < 0 at low T
- ΔH > 0, ΔS < 0: ΔG > 0 at all T
Determining ΔG of a Reaction
- ΔH & ΔS >/< 0
- -TΔS >/< 0
- l -TΔS l < l ΔH l
