energetics/thermochemistry

Define and distinguish heat and temperature.

Heat is a form of energy.
Temperature is a measure of the average kinetic energy of the particles

Total energy is \________________ in chemical reactions.

Total energy is conserved in chemical reactions.

Distinguish between exothermic and endothermic reactions.

exothermic: the transfer of heat from system to surroundings.
-product is more stable than reactant
-∆H is negative

endothermic: the transfer of heat from surroundings to system
-reactant is more stable than product
-∆H is positive

State the unit for enthalpy change (∆H).

The enthalpy change (∆H) for chemical reactions is indicated in kJ mol-1

Evaluate the percentage errors in a calorimetry experiment.

Some possible reasons:
-systematic error: heat loss to the surroundings
-assumptions: that solution has same density and specific heat capacity as water

List some examples of exothermic and endothermic reactions.

Exothermic:
-combustion
-neutralization

Endothermic:
-thermal decomposition
-photosynthesis

When are ∆H values expressed?

∆H values are usually expressed under standard conditions, given by ∆H°, including standard states.

Standard state refers to the normal, most pure stable state of a substance measured at 100 kPa. Temperature is not a part of the definition of standard state, but 298 K is commonly given as the temperature of interest.

What does the Hess's law state?

The total enthalpy change for a reaction is independent of the route by which the chemical change takes place.

Therefore, the enthalpy change for a reaction that is carried out in a series of steps is equal to the sum of the enthalpy changes for the individual steps.

Define average bond enthalpy.

Average bond enthalpy is the energy needed to break one mol of a bond in a gaseous molecule averaged over similar compounds.

Explain bond enthalpy change for forming and breaking bonds.

Bond-forming releases energy and is therefore exothermic, with a negative enthalpy sign.
Bond-breaking requires the absorbance of energy, so is therefore endothermic with positive enthalpy sign.

Describe and explain the bond strength in ozone relative to oxygen in its importance to the atmosphere.

The stratospheric layer of ozone functions to absorb extreme light frequencies and allow only long wavelength light to touch the earth ground.

The layer contains Oxygen and Ozone. Oxygen has a shorter bond length (with a double bond) while Ozone has a bond order of 1.5, therefore shorter light frequencies (UV-C, 242 nm) need to be absorbed to break O2 into 2 free radicals.
Ozone only needs (UV-B, 330 nm) wavelength to be broken.

The broken free radicals can then bond with O2 again to form O3. Therefore, it is the cycle of O2 and O3 being continuously formed and broken that maintains the equilibrium in the stratospheric layer.

Write out the representative equations for enthalpy/energy of hydration, ionization, atomization, electron affinity, lattice, covalent bond and solution.

Enthalpy/Energy of hydration:
X+ (g) → X+ (aq)

Ionization:
(IE1)X (g) → X+ (g) + e-
(IE2) X+ (g) → X 2+ (g) + e-

Atomization:
X (s) + Y (g) → X (g) + Y (g)

Electron Affinity:
Y (g) + e- → Y- (g)

Lattice enthalpy:
XY (s) → X+ (g) +Y− (g)

Covalent bond:
*show the breaking of one mole of a covalent bond in gaseous molecules:
such as Cl2(g) --\> 2Cl(g) or N2(g) --\> 2N(g).

Solution:
XY (s) → X+ (aq) +Y− (aq)

State the equation for enthalpy of formation under born Haber Cycle.

formation \= atomisation + IE + bond enthalpy + EA - lattice

Explain the relationship between size and charge of ions to lattice and hydration enthalpies.

greater charge \= decrease in atomic radii \= stronger force \= more energy released (exothermic) \= decrease in hydration enthalpy

greater charge \= decrease in atomic radii \= stronger force \= more energy required to break bonds \= increase in lattice enthalpy
**higher ionic charge results in a stronger electrostatic attraction between the ions, leading to a higher value of lattice enthalpy.

State the formula for finding enthalpy of solution

Δ H sol \= Δ H hyd + Δ H lat

Construction of energy cycles from hydration, lattice and solution enthalpy. For example dissolution of solid NaOH or NH4Cl in water.

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Explain what entropy is.

Entropy (S) refers to the distribution of available energy among the particles. The more ways the energy can be distributed the higher the entropy.

Entropy of gas\>liquid\>solid under same conditions.

Explain what Gibbs free energy is.

Gibbs free energy (G) relates the energy that can be obtained from a chemical reaction to the change in enthalpy (ΔH), change in entropy (ΔS), and absolute temperature (T).

G \= ΔH - TΔS
T \= ΔH/ΔS

Explain what the value of ΔG mean.

-ΔG: reaction is spontaneous
+ΔG: reaction is non-spontaneous
ΔG \= 0: reaction is in equilibrium

What are the various conditions that affect ΔG?

-sign of ΔH
-sign of ΔS
-T: temperature at which reaction takes place.

Temperature only plays in a role when:
-ΔH\>0 and ΔS\>0
reaction is ONLY spontaneous at HIGH temperatures

-ΔH

State the formula for Gibbs free energy of formation.

Σ ΔG (products) - Σ ΔG (reactant)

Explain what the enthalpy of solution indicates.

The enthalpy of solution indicates the solubility of an ionic compound.

-ΔH sol \= exothermic \= soluble
The high ΔH hyd compensates for the endothermic ΔH lat

+ΔH sol \= endothermic \= insoluble

Define lattice enthalpy.

the enthalpy change that occurs when 1 mol of solid crystal is formed from its gaseous ions.
or
the enthalpy change required to separate 1 mol of ionic compound into gaseous ions.

Define the enthalpy change.

enthalpy change when one mole of the compound is formed from its elements (in their standard states);
under standard conditions
25 °C/298 K and 1 atm/101.3 kPa/1.01×105 Pa101.3 kPa/1.01×105 Pa;