The combustion of glucose is exothermic and occurs according to the following equation:

C₆H₁₂O₆ (s) + 6O₂ (g) → 6CO₂ (g) + 6H₂O (g)

Which is correct for this reaction?

Which change of state is exothermic? 

A. CO₂(s) → CO₂(g)
B. H₂O(l) → H₂O(g) 
C. NH₃(g) → NH₃(l) 
D. Fe(s) → Fe(l)

Hydrazine has been used as a rocket fuel. The propulsion reaction occurs in several stages but the overall reaction is:

N₂H₄(l) → N₂(g) + 2H₂(g)

Suggest why this fuel is suitable for use at high altitudes.

Outline why the thermochemical method would not be appropriate for 0.001 mol$$dm⁻³ hydrochloric acid and aqueous sodium hydroxide of a similar concentration.

State a technique other than a pH titration that can be used to detect the equivalence point.

What is the enthalpy of combustion, ΔH_c, of ethanol in kJ mol⁻¹?
Maximum temperature of water: 30.0°C
Initial temperature of water: 20.0°C
Mass of water in beaker: 100.0 g
Loss in mass of ethanol: 0.230 g
M_r (ethanol): 46.08
Specific heat capacity of water: 4.18 J g⁻¹ K⁻¹
q = mcΔT

A.  $-\frac{100.0\times 4.18\times (10.0\times 273)}{{\displaystyle \frac{0.230}{46.08}}\times 1000}$

B.  $-\frac{0.0230\times 4.18\times 10.0}{{\displaystyle \frac{100.0}{46.08}}\times 1000}$

C.  $-\frac{100.0\times 4.18\times 10.0}{{\displaystyle \frac{0.230}{46.08}}\times 1000}$

D.  $-\frac{100.0\times 4.18\times 10.0}{{\displaystyle \frac{0.230}{46.08}}}$

Two 100 cm³ aqueous solutions, one containing 0.010 mol NaOH and the other 0.010 mol HCl, are at the same temperature.

When the two solutions are mixed the temperature rises by y °C.

Assume the density of the final solution is 1.00 g cm⁻³.

Specific heat capacity of water = 4.18 J g⁻¹ K⁻¹

What is the enthalpy change of neutralization in kJ mol⁻¹?

A.     $\frac{200\times 4.18\times y}{1000\times 0.020}$

B.     $\frac{200\times 4.18\times y}{1000\times 0.010}$

C.     $\frac{100\times 4.18\times y}{1000\times 0.010}$

D.     $\frac{200\times 4.18\times (y+273)}{1000\times 0.010}$

Which is correct when Ba(OH)₂ reacts with NH₄Cl?

Ba(OH)₂ (s) + 2NH₄Cl (s) → BaCl₂ (aq) + 2NH₃ (g) + 2H₂O (l)       ΔH^Θ = +164 kJ mol⁻¹

Calculate the standard enthalpy change, ΔH^Θ, of step 2 using section 13 of the data booklet.

The formula q = mcΔT was used to calculate the energy released. The values used in the calculation were m = 25.00 g, c = 4.18 J g⁻¹ K⁻¹.

State an assumption made when using these values for m and c.

State one other assumption that is usually made in the calculation of the heat produced.

Predict, giving a reason, how the final enthalpy of reaction calculated from this experiment would compare with the theoretical value.

The formula q = mcΔT was used to calculate the energy released. The values used in the calculation were m = 25.00 g, c = 4.18 J g⁻¹ K⁻¹.

State an assumption made when using these values for m and c.

Heat losses would make this method less accurate than the pH probe method. Outline why the thermometric method would always give a lower, not a higher, concentration.

Estimate the time at which the powdered zinc was placed in the beaker.

State what point Y on the graph represents.

Explain why the melting points of the group 1 metals (Li → Cs) decrease down the group.

Explain how the concentration may be calculated in this way.

Determine the heat change, q, in kJ, for the neutralization reaction between ethanoic acid and sodium hydroxide.

Assume the specific heat capacities of the solutions and their densities are those of water.

Explain why the melting points of the group 1 metals (Li → Cs) decrease down the group whereas the melting points of the group 17 elements (F → I) increase down the group.

5.35g of solid ammonium chloride, NH₄Cl(s), was added to water to form 25.0g of solution. The maximum decrease in temperature was 14 K. What is the enthalpy change, in kJmol^-1, for this reaction? (Molar mass of NH₄Cl = 53.5gmol^-1; the specific heat capacity of the solution is 4.18 Jg^-1K^-1)

A. $\mathrm{\Delta }H=+\frac{25.0\times 4.18\times \left(14+273\right)}{0.1\times 1000}$

B. $\mathrm{\Delta }H=-\frac{25.0\times 4.18\times 14}{0.1\times 1000}$

C. $\mathrm{\Delta }H=+\frac{25.0\times 4.18\times 14}{0.1\times 1000}$

D. $\mathrm{\Delta }H=+\frac{25.0\times 4.18\times 14}{1000}$

Suggest how heat loss could be reduced.

The energy released by the reaction of one mole of hydrogen peroxide with hydroquinone is used to heat 850 cm³ of water initially at 21.8°C. Determine the highest temperature reached by the water.

Specific heat capacity of water = 4.18 kJ$$kg⁻¹$$K⁻¹.

(If you did not obtain an answer to part (i), use a value of 200.0 kJ for the energy released, although this is not the correct answer.)

Which statement is correct for this reaction?

Fe₂O₃ (s) + 3CO (g) → 2Fe (s) + 3CO₂ (g)       ΔH = −26.6 kJ

A. 13.3 kJ are released for every mole of Fe produced.

B. 26.6 kJ are absorbed for every mole of Fe produced.

C. 53.2 kJ are released for every mole of Fe produced.

D. 26.6 kJ are released for every mole of Fe produced.

Explain, in terms of the molecular structure, the critical difference in properties that makes biodiesel a more suitable liquid fuel than vegetable oil.

Which expression gives the mass, in g, of ethanol required to produce 683.5 kJ of heat upon complete combustion?

(M_r for ethanol = 46.0, $\mathrm{\Delta }{H}_c^{\theta }=-1367\text{ kJ}\text{mo}{\text{l}}^{-1}$)

A.     $\frac{683.5}{1367\times 46.0}$

B.     $\frac{1367}{683.5\times 46.0}$

C.     $\frac{683.5\times 46.0}{1367}$

D.     $\frac{1367\times 46.0}{683.5}$

State what point Y on the graph represents.

To determine the enthalpy of reaction the experiment was carried out five times. The same volume and concentration of copper(II) sulfate was used but the mass of zinc was different each time. Suggest, with a reason, if zinc or copper(II) sulfate should be in excess for each trial.

State another assumption you made in (b)(i).

The energy from burning 0.250 g of ethanol causes the temperature of 150 cm³ of water to rise by 10.5 °C. What is the enthalpy of combustion of ethanol, in kJ mol^–1?

Specific heat capacity of water: 4.18 J g^–1K^–1.

A.  $\frac{150\times 4.18\times 10.5}{{\displaystyle \frac{0.250}{46.08}}}$

B.  $\frac{150\times 4.18\times 10.5}{{\displaystyle \frac{0.250}{46.08}}\times 1000}$

C.  $\frac{150\times 4.18\times \left(273+10.5\right)}{{\displaystyle \frac{0.250}{46.08}}}$

D.  $\frac{150\times 4.18\times \left(273+10.5\right)}{{\displaystyle \frac{0.250}{46.08}}\times 1000}$

The maximum temperature used to calculate the enthalpy of reaction was chosen at a point on the extrapolated (dotted) line.

State the maximum temperature which should be used and outline one assumption made in choosing this temperature on the extrapolated line.

Maximum temperature:

Assumption:

The reaction was carried out in a calorimeter. The maximum temperature rise of the solution was 7.5 °C.

Calculate the enthalpy change, ΔH, of the reaction, in kJ, assuming that all the heat released was absorbed by the solution. Use sections 1 and 2 of the data booklet.

A student obtained the following data to calculate $q$, using $q=mc\mathrm{\Delta }T$.

$m=20.2 \mathrm{g}\pm 0.2 \mathrm{g}$

$∆T=10°\mathrm{C}\pm 1°\mathrm{C}$

$c=4.18 \mathrm{J} {\mathrm{g}}^{-1} {\mathrm{K}}^{-1}$

What is the percentage uncertainty in the calculated value of $q$?

A.  $0.2$

B.  $1.2$

C.  $11$

D.  $14$

To determine the enthalpy of reaction the experiment was carried out five times. The same volume and concentration of copper(II) sulfate was used but the mass of zinc was different each time. Suggest, with a reason, if zinc or copper(II) sulfate should be in excess for each trial.

What is the enthalpy change, in J, when 5 g of water is heated from 10°C to 18°C?

Specific heat capacity of water: 4.18 kJ kg⁻¹ K⁻¹

A.  5 × 4.18 × 8

B.  5 × 10⁻³ × 4.18 × 8

C.  5 × 4.18 × (273 + 8)

D.  5 × 10⁻³ × 4.18 × (273 + 8)

Determine the enthalpy change of reaction, ΔH, in kJ, when 1.00 mol of gaseous hydrazine decomposes to its elements. Use bond enthalpy values in section 11 of the data booklet.

N₂H₄(g) → N₂(g) + 2H₂(g)

The reaction was carried out in a calorimeter. The maximum temperature rise of the solution was 7.5 °C.

Calculate the enthalpy change, ΔH, of the reaction, in kJ, assuming that all the heat released was absorbed by the solution. Use sections 1 and 2 of the data booklet.

Calculate the enthalpy change, ΔH, in kJ mol^–1, for the reaction between ethanoic acid and sodium hydroxide.

Iron has a relatively small specific heat capacity; the temperature of a 50 g sample rises by 44.4°C when it absorbs 1 kJ of heat energy.

Determine the specific heat capacity of iron, in J g⁻¹K⁻¹. Use section 1 of the data booklet.

Predict, giving a reason, how the final enthalpy of reaction calculated from this experiment would compare with the theoretical value.

State another assumption you made in (b)(i).

Which combination of ΔH₁, ΔH₂, and ΔH₃ would give the enthalpy of the reaction?

CS₂(l) + 3O₂(g) → CO₂(g) + 2SO₂(g)

ΔH₁  C (s) + O₂(g) → CO₂(g)
ΔH₂  S (s) + O₂(g) → SO₂(g)
ΔH₃  C (s) + 2S (s) → CS₂(l)

A.  ΔH = ΔH₁ + ΔH₂ + ΔH₃

B.  ΔH = ΔH₁ + ΔH₂ − ΔH₃

C.  ΔH = ΔH₁ + 2(ΔH₂) + ΔH₃

D.  ΔH = ΔH₁ + 2(ΔH₂) − ΔH₃

What is the heat change, in kJ, when 100.0 g of aluminium is heated from 19.0 °C to 32.0 °C?

Specific heat capacity of aluminium: 0.90 J g⁻¹K⁻¹

A.  $0.90\times 100.0\times 13.0$

B.  $0.90\times 100.0\times 286$

C.  $\frac{0.90\times 100.0\times 13.0}{1000}$

D.  $\frac{0.90\times 100.0\times 286}{1000}$

Iron has a relatively small specific heat capacity; the temperature of a 50 g sample rises by 44.4°C when it absorbs 1 kJ of heat energy.

Determine the specific heat capacity of iron, in J g⁻¹K⁻¹. Use section 1 of the data booklet.

Determine the molar enthalpy of combustion of an alkane if 8.75 × 10⁻⁴ moles are burned, raising the temperature of 20.0 g of water by 57.3 °C.