Chemical Energetics I Flashcards

1
Q

Enthalpy and enthalpy change

A

Enthalpy of a system is measure of energy content of the system, the higher the enthalpy, the more unstable it is

Enthalpy change is difference between quantity of heat absorbed to break bonds in reactants and that released during formation of new bonds in products at constant pressure

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2
Q

Difference between exothermic and endothermic reaction

A

Definition:

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3
Q

Difference between exothermic and endothermic reaction (Definition, temperature and enthalpy change)

A

Definition:
EXOTHERMIC is a reaction where energy is released to surroundings
ENDOTHERMIC is reaction where energy is absorbed from surroundings

Temperature Change:
EXOTHERMIC: Temperature of surroundings increase
ENDOTHERMIC: Temperature of surroundings decreases

Enthalpy Change:
EXOTHERMIC: Enthalpy change is negative, products energetically more stable than reactants
ENDOTHERMIC: Enthalpy change ie positive, products energetically less stable than reactants

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4
Q

Which reaction is more energetically feasible, does that mean they will occur?

A

Exothermic reactions are more energetically feasible and more likely to occur

Enthalpy change is indication of energetic stability and not its kinetic stability of products with respect to reactants

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5
Q

Standard conditions

A

Pressure at one bar
Temp usually at 298K
Substances involved at standard state (its most stable form at above conditions)

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6
Q

Definitions of different types of enthalpy change (TEMPLATE)

A

(Name of process) is the energy (change/ released/ required) when one mole of (substances in specified physical state) is (description of process) under standard conditions.

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7
Q

Standard enthalpy change of reaction

What does kJ mol-1

A

Energy change in a chemical reaction when molar quantities of reactants stated in the chemical equation react under standard conditions.

Related to amounts of all substances given in balanced equation

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8
Q

Standard enthalpy change of formation of a substance

How can this predict stability of compound?

A

Energy change when 1 mole of pure substance in specified state is formed from its constituent elements in their standard states under standard conditions.

The more negative the enthalpy change of a compound, the more stable the compound is relative to its constituent elements and less likely decomposition of compound back into its constituent elements

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9
Q

Standard enthalpy change of formation of an element

A

Standard enthalpy change of formation of an element in its standard state under standard conditions is zero

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10
Q

Standard enthalpy change of combustion

How can this indicate energy values of fuel?

A

Energy released when 1 mole of the substance is completely burnt in excess oxygen under standard conditions

More heat being liberated upon complete combustion, the better the fuel

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11
Q

Standard enthalpy change of neutralisation

Why is enthalpy change the same for all strong acids and strong bases?

A

Energy released when acid and base react to form 1 mole of water under standard conditions

Strong acid undergoes complete dissociation in aqueous solution, while strong base undergoes complete ionisation in aqueous solution. Reaction between acid and base in aqueous solution is the reaction between H+ and OH- to produce one mole of H2O.

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12
Q

Why is neutralisation between weak acid/ base less exothermic?

A

Weak acid and base molecules have to undergo further dissociation/ ionisation, which is endothermic. Since some energy is consumed to bring about further dissociation/ ionisation of molecules during neutralisation reaction, energy released from reaction will be less.

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13
Q

Standard enthalpy change of atomisation of element

Standard enthalpy change of atomisation

Why are they always positive?

A

Energy absorbed when 1 mole of gaseous atoms (need to specify) is formed from element (at standard state) under standard conditions

Energy absorbed when 1 mole of compound (specify) is converted to gaseous atoms under standard conditions

Always positive because energy must be absorbed to break all bonds between atoms in element/ compound during atomisation reaction

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14
Q

Bond dissociation energy

Always positive? What about reverse reaction?

Why different BDEs for each different bond?

A

Energy required to break 1 mole of that particular bond (X-Y) in a particular compound (specify) in gaseous state

Always positive because bond breaking requires energy while bond forming releases energy

Chemical environment when each bond is broken is different

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15
Q

Bond energy

Enthalpy change of atomisation of diatomic gases

A

Average energy required to break 1 mole of bond in gaseous state

For diatomic gas molecules, bond energy= 2 times of standard enthalpy change of atomisation of the gas molecule

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16
Q

First electron affinity and why it is usually negative

Second electron affinity and why it is always positive

A

Energy change when 1 mole of gaseous atoms acquires 1 mole of electrons to form 1 mole of singly negatively charged gaseous ions

Usually negative because energy released when nucleus attracts additional electron is usually larger than energy taken in to overcome inter-electronic repulsion

Energy absorbed when 1 mole of singly negatively charged gaseous ions acquires 1 mole of electrons to form 1 mole of doubly negatively charged gaseous ions

Always positive because energy required to overcome repulsion between 2 negatively charged species

17
Q

Lattice energy

How can it measure strength of ionic bonding?

Factors affecting lattice energy

Enthalpy changes present

A

Energy released when 1 mole of the solid ionic compound is formed from its constituent gaseous ions under standard conditions

More negative/ greater magnitude of lattice energy of ionic compound, the stronger the ionic bonding present in the compound

Bigger cationic/ anionic charge, more negative lattice energy
Smaller cationic/ anionic radius, more negative lattice energy

FAIL (enthalpy change of formation, enthalpy change of atomisation, ionisation energy, electron affinity, lattice energy)

18
Q

Comparison between experimental lattice energy and theoretical lattice energy

A

If experimental lattice energy values are in good agreement with theoretical values, structure is quit close to being purely ionic, and is predominantly ionic, fitting the model with lattice consisting of spherical ions with evenly distributed charge

Discrepancy shows that it is not as close to ‘purely ionic’ and ionic compounds have some covalent character due to substantial polarisation of anion by cation, experimental lattice energy values would have greater magnitude as bonding is stronger than purely ionic model

19
Q

Standard enthalpy change of hydration and its relation to charge density

Steps involved

A

Energy released when 1 mole of gaseous ion is hydrated under standard conditions

The greater the charge density, the greater the ion-dipole interaction strength, the greater the magnitude of standard enthalpy change of hydration

Separation of ions in solid ionic lattice into monatomic gaseous ions, then hydration of gaseous ions

20
Q

Standard enthalpy change of solution

A

Energy change when 1 mole of substance is completely dissolved in a solvent to form an infinitely dilute solution under standard conditions

21
Q

3 formulas required in experiments

A

q = mcΔT
q = CΔT
ΔH = -q/n (of limiting reagant)

22
Q

Steps to determine ΔH

A
  1. Find ΔT
  2. Find mass (assuming density of solution = water)
  3. Assume specific heat capacity to be that of water (if not given)
  4. Find heat absorbed/ released
  5. Write reaction equation
  6. Find amount of limiting reagent
  7. Find ΔH
23
Q

Hess’ Law

A

Enthalpy change of a reaction is determined by initial and final states of the system and is independent of the pathways taken

24
Q

Formulas derived from Hess’ Law

A

Total enthalpy change of combustion = Sum of enthalpy change of combustion of reactants - products

Total enthalpy change of formation = Sum of enthalpy change of formation of products - Sum of enthalpy change of formation of reactants