Chemical Energetics

Question 1

3 types of systems

Answer 1

• Open system
  ○ Matter and energy can be exchanged between the system and the surroundings
  § E.g. Bacteria growing in a cultured medium. In this case, bacteria can exchange oxygen and carbon dioxide with the atmosphere. Energy can also flow freely in and out of the system.
  
  • Closed system
    ○ Matter cannot be exchanged between the system and the surroundings, but energy can still be transferred
    § E.g. Hydrogen gas and oxygen gas reacting in a gas jar/syringe. Although the chemical form of the hydrogen and oxygen atoms in the system is changed in the above reaction, the system has not lost or gained mass. However, the system can exchange energy with its surroundings in the form of work and heat
  • Isolated system
    ○ Matter and energy cannot be exchanged between the system and surroundings
    § E.g. Reaction carried out in an insulated and rigid container such as Dewar Flask

Question 2

Enthalpy, H

Answer 2

• Indication of a substance’s total heat energy content at constant pressure and cannot be measured directly
  
  • However, it is possible to measure the enthalpy change, ∆H, of a chemical reaction
  • For a reaction that takes place at constant pressure, the enthalpy change of the reaction is directly proportional to the heat released or absorbed by the system during the reaction
    ○ i.e. ∆H=H(products)−H(reactants)
    ○ Unit for ∆H: kJ 〖mol〗^(−1)

Question 3

Exo VS Endo

Answer 3

Exo: Bond making, heat is given off as the reactants are converted into products. Hence, the
reaction medium feels warmer. Therefore, the enthalpy of the products is lower than the enthalpy
of the reactants, and the enthalpy change (H) is negative.

Endo: Bond breaking, the reactants absorb heat to form the products. Hence, the reaction
medium feels colder. Therefore, the enthalpy of the products is higher than the enthalpy of the
reactants, and the enthalpy change (H) is positive.

note:
- Exothermic reactions are energetically more favourable than endothermic ones because a system
with lower heat content is more stable compared to that with higher heat content.
- A reaction, which is highly exothermic, will proceed with relative ease; the more negative an H
value, the more stable is the products.

Question 4

Activation energy

Answer 4

minimum energy for a reaction to occur

Question 5

standard conditions

Answer 5

• temperature: 25 C (or 298 K)
• pressure: 1 bar (or 1 × 105 Pa)
• concentration: 1 mol dm−3 (if reaction involves aqueous solutions)
• all substances are in their most stable physical form / allotropic form
  e.g. Na(s), H2O(l), Cl2(g), Br2(l), I2(s), C(graphite)

Question 6

Standard Enthalpy Change of Reaction, Hr

Answer 6

Change in enthalpy when molar quantities of reactants, as stated in the balanced stoichiometric
equation, react together at 298 K and 1 bar.

Question 7

Standard Enthalpy Change of Formation, Hf

Answer 7

Change in enthalpy when one mole of a compound is formed from its constituent elements in their standard physical states at 298 K and 1 bar.

∆Hf
ɵ of a compound is a measure of its stability relative to its constituent element.
∆Hf
ɵ < 0, the compound is energetically more stable than its constituent elements.
∆Hf
ɵ > 0, the compound is energetically less stable than its constituent elements.
For example, the value of Hf
o
(Na2O(s)) is very large and negative, which implies that Na2O(s) is

energetically much more stable than Na(s) and O2(g).
* The enthalpy changes of formation of various compounds are usually theoretical, and the values are
normally calculated indirectly from other enthalpy changes of reactions.

Question 8

Standard Enthalpy Change of Combustion, Hc

Answer 8

Heat evolved when one mole of a substance is completely burnt in excess oxygen at 298 K and 1 bar.

Enthalpy change of combustion of carbon is applied only to the complete combustion of C(s) into CO2, and not to CO.
i.e. C(s) + 1⁄2O2(g) → CO(g) is not the equation for the complete combustion of carbon.

Question 9

Standard Enthalpy Change of Neutralisation, Hn

Answer 9

Enthalpy change when an acid and a base react to form one mole of water at 298 K and 1 bar.

Question 10

Lattice Energy, L.E. (of an ionic compound)

Answer 10

Heat evolved when one mole of the solid ionic compound is formed from its isolated gaseous ions at 298K and 1 bar.

All lattice energies are exothermic because heat energy is evolved when gaseous ions come
together to form ionic bonds.
* Lattice energies provide a measure of the strength of the ionic bonds holding the ions in their
position in a crystal lattice. Hence, the more exothermic the lattice energy, the stronger the
ionic bonding and the higher is the melting point of the ionic compound.

Question 11

Factors affecting Lattice Energy

Answer 11

The magnitude of lattice energy is:
(i) directly proportional to charges on the ions
the bigger the ionic charge, the more exothermic is the lattice energy;
(ii) inversely proportional to ionic radius
the smaller the ionic radius, the more exothermic is the lattice energy.
* Since lattice energy is directly proportional to the product of the charges on the ions and
inversely proportional to sum of ionic radius, charges is the more important factor when
the two exert opposing effects on lattice energy.

Question 12

Theoretical and experimental lattice energies

Answer 12

Lattice energies can be calculated theoretically based on the assumption that ionic lattice is made up of discrete spherical ions, with evenly distributed charges (i.e. compound is purely ionic with no covalent character). The experimental values are obtained from thermochemical data using Born-
Haber cycle.

Comparison of the theoretical and experimental values for lattice energy gives an indication of the
degree of covalent character in an ionic compound.
* If the theoretical value agrees well with the experimental value (e.g. the halides of sodium), the
compound is ionic with little or no covalent character.
* If the theoretical value is significantly different from the experimental value (e.g. halides of silver),
the bonding is not purely ionic but has considerable percentage of covalent character. This
makes the bond between Ag+ and X−
ions stronger than expected (in this case).

Question 13

Standard Enthalpy Change of Atomisation, Hatom

Answer 13

Heat absorbed when one mole of gaseous atoms is formed from its element in its standard state at 298 K and 1 bar.

Hatom
is always endothermic, because energy must be absorbed to pull the atoms far
apart and to break all the bonds between them.

Question 14

Ionisation Energy, I.E.

Answer 14

First ionisation energy of an element is the energy required to remove one mole of
the outermost electron from one mole of gaseous atoms of the element to form
one mole of gaseous ions with a single positive charge.

Question 15

Electron Affinity, E.A.

Answer 15

First electron affinity of an element is the energy evolved when one mole of gaseous
atoms accepts one mole of electrons to form one mole of gaseous singly charged
negative ions.First electron affinity of an element is the energy evolved when one mole of gaseous
atoms accepts one mole of electrons to form one mole of gaseous singly charged
negative ions.

The first electron affinity of an electronegative element is usually an exothermic process. However,
the second and higher electron affinities are endothermic processes.

The 1st E.A. is exothermic due to the electrostatic forces of attraction formed between the
nucleus and the electron added.
The 2nd and successive electron affinities are endothermic because energy is required to overcome
the electrostatic repulsion between the incoming electron and the negative charge on the anion.

Question 16

Standard Enthalpy Change of Hydration, Hhyd

Answer 16

Heat evolved when one mole of the gaseous ions dissolved in water to form
hydrated ions of infinite dilution at 298 K and 1 bar.

Hhyd
is always exothermic, as heat is produced when ion-dipole interactions are formed
between the ions and water molecules.

Water is a polar molecule. The negative end of the water molecule is attracted to the cation while
the positive end is attracted to the anion. An ion-dipole interaction is formed.

The hydration energies of the ions depend on the charge and size of the ions:

ΔHhyd ∝
charge
size of the ion (charge density)

The higher the charge and the smaller the size of the ions (i.e. the larger the charge density),
the more exothermic is the hydration energy (or the larger its magnitude).

Question 17

Standard Enthalpy Change of Solution, Hsol

Answer 17

Change in enthalpy when one mole of a substance is dissolved in water to form a solution of infinite dilution at 298K and 1 bar.

• Hsol can be exothermic or endothermic.
• If Hsol

is very endothermic, then the compound is insoluble in water.

• If Hsol

is exothermic, then the compound is soluble in water.

Question 18

Bond Energy (of a covalent compound)

Answer 18

Bond energy (of dissociation) is the energy required to break one mole of covalent
bonds between two atoms in a gaseous state.

Question 19

Assumptions made when measuring the enthalpy change using calorimetric methods

Answer 19

• No heat loss to / heat gained from surroundings. The reactions occur rapid enough for a
  maximum temperature change to be achieved before the reaction mixture begins to return to room
  temperature.
• No heat transfer between the solution and the container (i.e. heat capacity of container is
  assumed to be zero).
• The density and specific heat capacity of aqueous reaction mixture are taken to be those of water,
  i.e. 1.0 g cm−3 (usually given in the question but you should know if required) and 4.18 J g−1 K
  −1

(given in the Data Booklet) respectively.

Question 20

Spontaneous change

Answer 20

• In the absence of any barrier (such as activation energy) takes place naturally in the direction stated (i.e. Has a natural tendency to occur.
  
  • For e.g. Ice melts when places at room temperature, a shiny nail left outdoors will eventually rust, and if you touch a hot object, heat is transferred to you
  • For all these processes, they all occur without any outside intervention, and thus they are said to be spontaneous.

Question 21

Concept of entropy, S

Answer 21

• The entropy of a system increases when the particles or energy in the system becomes more random in its arrangement.
  
  • The greater the number of ways energy or particles can be distributed in the system, the higher the entropy.
  • A system that has a greater degree of disorder/randomness is said to have higher entropy

Question 22

Entropy change, ∆S

Answer 22

• Entropy is a measure of the disorder in a system
  
  • The more disordered a system, the larger its entropy

Question 23

4 Factors affecting entropy of a chemical system

Answer 23

• Change in temperature
  ○ Increase in temperature –> increase in entropy (or result in entropy change to be positive, ∆S>0)
  ○ Because, increase in temperature –> increase in average kinetic energy of the particles in the system –> particles in the system vibrate more vigorously (in solids) or moving faster (in liquids and gases) –> due to the broadening of the Boltzman distribution and increase in number of ways that the energy can be distributed among the particles in the system –> hence the degree of disorder increases and entropy change to be positive
  • Change in phase
    ○ Melting of a solid to a liquid will lead to an increase in entropy or result in entropy change to be positive. Cuz, the orderly arrangement of particles in the solid lattice is destroyed
    ○ Boiling or vaporising of a liquid to a gas will lead to an even greater increase in entropy or result in entropy change to be large and positive.
    Cuz, gaseous state has a much greater degree of disorder compared to solids or liquids, as the gaseous particles move at very high speeds and occupies a larger volume
    ○ Hence, disorder of the particles increases from solid to liquid to gas as there is an increase in the number of ways to arrange the particles. Thus, entropy increases during phase changes from solid to liquid and from liquid to gas
  • Change in number of particles
    ○ A change that involves an increase in the number of gas particles in the system will lead to an increase in number of ways to arrange the larger number of particles.
    ○ Hence, the degree of disorder increases and entropy change to be positive
  • Mixing of particles
    ○ Mixing of particles will lead to an increase in entropy or result in entropy change to be positive.
    ○ E.g. Of mixing,
    § Solute dissolve in solvent
    § Pure gases diffuse into each other
    § Miscible liquids poured together
    ○ There is no increase in total molecules of the system. However, there is an increase in number of ways to arrange the particles in the system after mixing. Hence, the degree of disorder increases and entropy change to be positive

Question 24

Gibbs Free Energy

Answer 24

Go Home To Study
G = H - T(S)
T: kelvin under 298K and 1 bar
H: KJ mol-1
S: KJ mol-1 K-1

G<0 : spontaneous in the forward direction, and is
said to be thermodynamically feasible. products are formed

G>0: non-spontaneous in the forward direction
spontaneous in the reverse direction. reactants remained

G=0: at equilibrium, the reaction is a reversible
reaction, as both the forward and backward
reactions are spontaneous and can take place
at the same time.
Phase change is an example of reactions at
equilibrium. products and reactants exist
at the same time.

Question 25

Limitations in using G
to Predict Spontaneity of a Reaction

Answer 25

The use of G
to predict spontaneity of reactions is limited to systems at standard conditions as the values of H and S vary slightly with changes in temperature. However, it is possible to
predict the spontaneity of a reaction at various temperatures, by calculating an estimated value of G
with the assumption that H and S are relatively constant over a small temperature range.

Question 26

Factor Affecting Entropy

Answer 26

Dissolving an Ionic Solid
* In solid NaCl, the ions are in a highly ordered, crystalline state.
* When the solid dissolves, the crystal structure of the solid NaCl is disrupted. The Na+ and Cl
−
ions

are free to move about in the water resulting in a more disordered system.
* However, from the perspective of water molecules, the water molecules are held around the ions
as water of hydration. These water molecules that were originally free to move has become more
ordered.
* The disordering process of solute particles is usually more dominant. Thus, the overall effect is an
increase in disorder and increase in the entropy of the system for dissolving most salts in water.

Question 27

Ellingham Diagrams

Answer 27

Ellingham diagram is a graphical representation to illustrate the variation of change in Gibbs Free
Energy with temperature.