Reactivity 1.1.4—The standard enthalpy change for a chemical reaction, ΔH⦵, refers to the heat transferred at constant pressure under standard conditions and states. It can be determined from the change in temperature of a pure substance. Apply the equations Q = mcΔT and ΔH = − Qn in the calculation of the enthalpy change of a reaction.

Calculate the energy released, in kJ g⁻¹, when 3.49 g of starch are completely combusted in a calorimeter, increasing the temperature of 975 g of water from 21.0 °C to 36.0 °C. Use section 1 of the data booklet.

Calculate the energy released, in kJ g⁻¹, when 3.49 g of starch are completely combusted in a calorimeter, increasing the temperature of 975 g of water from 21.0 °C to 36.0 °C. Use section 1 of the data booklet.

Calculate the energy released, in kJ g⁻¹, when 3.49 g of starch are completely combusted in a calorimeter, increasing the temperature of 975 g of water from 21.0 °C to 36.0 °C. Use section 1 of the data booklet.

Calculate the energy released, in kJ g⁻¹, when 3.49 g of starch are completely combusted in a calorimeter, increasing the temperature of 975 g of water from 21.0 °C to 36.0 °C. Use section 1 of the data booklet.

Calculate the energy released, in kJ g⁻¹, when 3.49 g of starch are completely combusted in a calorimeter, increasing the temperature of 975 g of water from 21.0 °C to 36.0 °C. Use section 1 of the data booklet.

Calculate the energy released, in kJ g⁻¹, when 3.49 g of starch are completely combusted in a calorimeter, increasing the temperature of 975 g of water from 21.0 °C to 36.0 °C. Use section 1 of the data booklet.

Calculate the energy released, in kJ g⁻¹, when 3.49 g of starch are completely combusted in a calorimeter, increasing the temperature of 975 g of water from 21.0 °C to 36.0 °C. Use section 1 of the data booklet.

Calculate the energy released, in kJ g⁻¹, when 3.49 g of starch are completely combusted in a calorimeter, increasing the temperature of 975 g of water from 21.0 °C to 36.0 °C. Use section 1 of the data booklet.

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

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

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

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

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.  $14$

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.  $14$

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.  $14$

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.  $14$

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)

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)

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)

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)

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

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

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

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.

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.

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.

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.

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.

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.

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.

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.

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.

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.

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.

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