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Calorimetry

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Wize University Chemistry Textbook > Thermodynamics: Part 1: (0th law, 1st law, Calorimetry, Enthalpy)

Calorimetry

7 Activities

Wize High School Grade 12 Chemistry Textbook > Energy Changes

Calorimetry

6 Activities

The temperature of a 12.58 g sample of calcium carbonate (CaCO3 (s)) increases from 23.6oC to 38.2oC. If the specific heat of calcium carbonate is 0.82 J/g K, how many joules of heat are absorbed?
More Calorimetry Questions:
Calorimetry: Heat of combustion
When 2g of rocket fuel, N2H4 is burned inside of a bomb-calorimeter, the temperature of the water rises from 22oC to 29oC. If there is 0.9kg of water and the heat capacity for the bomb is 822J/oC, what is the heat of the combustion for one mole of N2H4 in the bomb calorimeter?
Calorimetry: Heat of combustion
When 2g of rocket fuel, N2H4 is burned inside of a bomb-calorimeter, the temperature of the water rises from 22oC to 29oC. If there is 0.9kg of water and the heat capacity for the bomb is 822J/oC, what is the heat of the combustion for one mole of N2H4 in the bomb calorimeter?
Heat Capacity (bomb calorimeter)
1 gram of fried chicken is placed in a bomb calorimeter, which has a heat capacity of 10 kJ/oC. Upon combustion of the fried chicken, the temperature of the calorimeter rises by 5oC. Approximately how many nutritional calories are in 1 gram of this fried chicken if 1 nutritional calorie = 1000 calories and 1 calorie = 4.184 J ? Enter answer with 2 significant digits
Calorimetry
Equal masses of copper and iron are both heated to a temperature of 100°C. Each piece of hot metal is then placed into its own beaker containing 1 L of water. If both beakers start with the same initial temperature, which beaker of water will have the higher final temperature?
Additional information: Cu (specific heat = 0.377 J/g°C); Fe (specific heat = 0.450 J/g°C)
Calorimetry: Heat requirement
How much heat is required to convert a 15.5 g ice cube at -5.00oC to liquid water at 0oC? (The specific heat of ice is 2.09 J/g K, the specific heat of water is 4.18 J/g K, and the ΔHfus for water=334 J/g).
Calorimetry: Specific Heat Capacity
Addition of 3.10 kJ of heat is required to raise the temperature of 150 g of a liquid hydrocarbon from 23.0oC to 44.0oC. What is the specific heat capacity of this hydrocarbon?
A bomb calorimeter assembly has a heat capacity of 650 J/°C. Inside the bomb calorimeter 2.5g of a fuel compound is burned in excess oxygen and the temperature of the bomb calorimeter increases by 48 °C. What is the value for qrxn amount of heat released per gram of compound burned?
Please report your answer to three significant figures in kJ/g. Do NOT include units in the answer field
Water can absorb infrared radiation at a wavelength of1.06×104 m{\small 1.06\times10^4 \text{ m}}. If a 180 g sample of water is irradiated at this wavelength, calculate the number of photons required to raise the temperature by 25.0°C(CH2O=4.184Jg°C){\small \degree \text{C} \left(C_{H_2O} =4.184\frac{J}{g\degree C} \right)}
Calorimetry: Heat of combustion
When 2g of rocket fuel, N2H4 is burned inside of a bomb-calorimeter, the temperature of the water rises from 22oC to 29oC. If there is 0.9kg of water and the heat capacity for the bomb is 822J/oC, what is the heat of the combustion for one mole of N2H4 in the bomb calorimeter?
Enthalpy: Heat of reaction
2H2(g) + O2(g) --> 2H2O(l) ΔH=-572kJ/mol
1. How much heat is produced when 72g of water gas is produced?
2. Is the reaction endothermic or exothermic?
When 1.00 kg of lead (specific heat = 0.13 J / g °C) at 125 °C is added to a quantity of water at 18.0 °C, the final temperature of the lead-water mixture is 32.0 °C. What is the volume of water present? Note: the density of H2O(l) is 1.00 g/mL and the specific heat capacity is 4.18 J / g oC.
Calorimetry: Specific Heat
A piece of metal weighing 58g was heated to 100oC and then added to 100mL of water. The water's temperature was initially 22oC and after the metal was added the temperature was 26oC. What is the specific heat of the metal?
More Calorimetry Questions:
Heat of Vaporization for Hexane
Hexane is heated from 0oC to 50oC and the equilibrium vapour pressure increases from 0.06 atm to 0.53 atm.
Calculate the ΔvapH for hexanes.
Thermochemistry: Melting ice
The standard enthalpy of fusion for water, ΔHfusiono=6.01 kJ/mol.\Delta H^o_{\textrm{fusion}}=6.01\ \textrm{kJ/mol}. Calculate the energy required (in KJ) to convert 26.0 g of ice at 0oC to liquid water at 20oC.
Calorimetry: Final temperature calculation
The ΔHsoln for the process when solid sodium hydroxide (NaOH) dissolves in water is 44 kJ/mol. When a 10.0 g sample of NaOH dissolves in 250.0 g of water in a coffee-cup calorimeter, the temperature increases from 23.0oC to _______oC. Assume that the solution has the same specific heat as liquid water, i.e. 4.18 J/gK.
Calorimetry: Heat of combustion
When 2g of rocket fuel, N2H4 is burned inside of a bomb-calorimeter, the temperature of the water rises from 22oC to 29oC. If there is 0.9kg of water and the heat capacity for the bomb is 822J/oC, what is the heat of the combustion for one mole of N2H4 in the bomb calorimeter?
Thermochemistry: Enthalpy change
The enthalpy of vaporization for water, ΔHovaporization = 40.7 kJ/mol. Determine the enthalpy change (in kJ) when 54.0 mL of water is condensed from steam at 100oC.
Thermochemistry: Melting ice
The standard enthalpy of fusion for water, ΔHfusiono=6.01 kJ/mol.\Delta H^o_{\textrm{fusion}}=6.01\ \textrm{kJ/mol}. Calculate the energy required (in KJ) to convert 26.0 g of ice at 0oC to liquid water at 20oC.
Calorimetry: Heat of combustion
When 2g of rocket fuel, N2H4 is burned inside of a bomb-calorimeter, the temperature of the water rises from 22oC to 29oC. If there is 0.9kg of water and the heat capacity for the bomb is 822J/oC, what is the heat of the combustion for one mole of N2H4 in the bomb calorimeter?
Calorimetry: Heat of combustion
When 2g of rocket fuel, N2H4 is burned inside of a bomb-calorimeter, the temperature of the water rises from 22oC to 29oC. If there is 0.9kg of water and the heat capacity for the bomb is 822J/oC, what is the heat of the combustion for one mole of N2H4 in the bomb calorimeter?
Enthalpy of Formation
What is the enthalpy of the following reaction when 0.35 g of Mg2+ is produced? Use the following information:
ΔH°f[Mg(OH)2(s)] = −924.5 kJ
ΔH°f[NH4+(aq)] = −132.5 kJ
Thermodynamics: Internal energy and Enthalpy
2.3 moles of an ideal gas initially at -12oC and 16 atm underwent a complete expansion against a constant external pressure of 1atm and did 1.1kJ of work. Calculate Δ𝑈, Δ𝐻, 𝑎𝑛𝑑 𝑞 for this process.
Calorimetry: Final temperature calculation
The ΔHsoln for the process when solid sodium hydroxide (NaOH) dissolves in water is 44 kJ/mol. When a 10.0 g sample of NaOH dissolves in 250.0 g of water in a coffee-cup calorimeter, the temperature increases from 23.0oC to _______oC. Assume that the solution has the same specific heat as liquid water, i.e. 4.18 J/gK.
Calorimetry: Heat of reaction
0.23 g of ethanol (C2H5OH) was burned in a bomb calorimeter (constant
volume
) leading to a temperature change from 22.50oC to 23.81oC. If the heat capacity of the calorimeter is 5.2 kJ/K, what would be the amount of heat generated by the combustion of exactly one mole of ethanol? ___________
Calorimetry: Temperature change
1. A 56.0 g block of iron heated to 90 °C is placed in a 4.8 kg water bath at 20 °C. What will be the final temperature of the iron? heat capacity of iron c = 0.450 J /gK, heat capacity of water is 4.185 J/gK
Calorimetry: Temperature change
How much heat energy is required to raise the temperature of 35.0 g of Al (c = 0.902 J/g K) from 18 °C to 65 °C? If this much energy were instead used to raise the temperature of 2.0 L of water, by how much would the temperature of the water change? Round to the nearest oC.
Calorimetry: Heat of reaction
4. A bomb calorimeter assembly has a heat capacity of 650 J/°C. The assembly is placed in a 3.0 L water bath and 2.5g of a fuel compound is burned in excess oxygen. If the temperature of the water bath increases by 4.8 °C, what is the value of qrxn?
What is the amount of heat released per gram of compound burned?
Calorimetry: Heat capacity
5. A fuel compound is burned in a bomb calorimeter, releasing 458 kJ of heat. The temperature of the calorimeter assembly and surrounding 2.5 L water bath increases from 23.5 °C to 28.6 °C. What is the heat capacity of the calorimeter assembly?
Calorimetry
When 1.00 kg of lead (specific heat =0.13 J g-1°C-1) at 125oC is added to a quantity of water at 18.0°C, the final temperature of the lead-water mixture is 32.0°C. What is the volume of water present?
Note: the density of H2O(l) is 1.00g mL-1 and the specific heat capacity is 4.18J g-1 oC-1
A 2.00 mole sample of a monoatomic ideal gas is initially at 1.00 L and 1.00 atm. Calculate the total q, w, ΔE{\small \Delta\text{E}}and ΔH{\small \Delta \text{H}}when the final conditions are 0.500 L and 5.00 atm along the pathway where the volume is compressed first.
[Do NOT include units in your answers]
Thermochemistry: Energy required to melt ice
You are lost on an expedition and you need some water to drink. There is ice around at an ambient temperature of -22oC. Using the thermodynamic information below, how much energy is required to take 1 kg of ice and warm it up to water at 12oC.
cp,s(ice)=2.01 JK−1mol−1ΔHfuso=6.01 kJmol−1cp,s(water)=4.18 JK−1mol−1ΔHfuso=44.00 kJmol−1\begin{array}{lr} c_{p,s}(ice) = 2.01\ J K^{-1} mol^{-1} & \Delta H^o_{fus} = 6.01\ kJ mol^{-1} \\ c_{p,s}(water) = 4.18\ J K^{-1} mol^{-1} & \Delta H^o_{fus} = 44.00\ kJ mol^{-1} \end{array}
Heat of Vaporization for Hexane
Hexane is heated from 0oC to 50oC and the equilibrium vapour pressure increases from 0.06 atm to 0.53 atm.
a) Calculate the ΔvapH for hexanes.
b) The human nose can detect hexane at a partial pressure of 0.005 atm. Estimate the temperature at which we can no longer smell hexane.
Copper, Silver, water ,Gold, hydrogensulfide and have specific heats of 0.385, 0.235, 4.18, 0.129, and 1.342 Jg-1oC-1 respectively. Which of the following would need smallest quantity of heat to change the temperature of 5400kg by 65oC
An alkane’s vapour pressure at 25.00 °C is 6.00 kPa. What is its vapour pressure at 80.00 °C if the enthalpy of vaporization is 43.0 kJ/mol? (Give your answer in kPa)
Sucrose, C12H22O11, can be metabolized for energy in the body according to the reaction: C12H22O11 (s) + 12 O2 (g) →\rightarrow 12 CO2 (g) + 11 H2O (l) ΔHo = -5.65x103 kJ
How many grams of sucrose must be burned to raise the temperature of 50.0 kg of water by 1.00oC? The specific heat capacity of water is 4.184 J g-1 K-1.
Calorimetry: Heat of reaction
The temperature of a 12.58 g sample of calcium carbonate (CaCO3(s)) increases from 23.6oC to 38.2oC. If the specific heat capacity of calcium carbonate is 0.82 J/g K. How many joules of heat are absorbed?
Calorimetry: Energy of vaporization of water
How much energy is required to bring 1.00 g of liquid water from 35 °C to 165 °C water vapour? The enthalpy of vaporization of water is +40.7 kJ/mol, the specific heat capacity of liquid water is 4.18 J / g °C, and the specific heat capacity of water vapour is 1.87 J/g °C.
From the following reaction:
2C6H12O6(s)+2O2(g)→12CO2(g)+4H2O(l)ΔHrxn=−1680kJ2C_6H_{12} O_{6(s)}+2O_{2(g)} \to 12CO_{2(g)}+4H_2O_{(l)} \hspace{20pt} \Delta H_{rxn}=-1680 kJ
When an unknown quantity of sugar (C6H12O6) is burned, it was determined that the internal energy of the system decreased by 960 kJ and that 15.7 kJ of work was done on the system. How many grams of sugar were present at the beginning of this specific reaction?
A chemical compound has a molecular weight of 89.05 g/mol. 1.400 grams of this compound underwent complete combustion under constant pressure conditions in a calorimeter with a heat capacity of 2.980 × 103 J/ °C. The temperature went up by 11.95 °C. Calculate the standard heat of combustion of the compound.
Calorimetry: Heat capacity of calorimeter
The combustion of 1.00 mol of glucose (C6H12O6) gives off 2.82x103 kJ of heat (i.e. ΔHcomb = 2.82x103 kJ mol-1). If 1.80 g glucose (molar mass is 180) is burned in a bomb calorimeter and the temperature of the calorimeter increases from 19.00oC to 23.50oC, what is the heat capacity of the calorimeter (kJ/oC)?
Calorimetry: Heat requirement
How much heat is required to convert a 15.5 g ice cube at -5.00oC to liquid water at 0oC? (The specific heat of ice is 2.09 J/g K, the specific heat of water is 4.18 J/g K, and the ΔHfus for water=334 J/g).
Thermochemistry: Melting ice
The standard enthalpy of fusion for water, ΔHfusiono=6.01 kJ/mol.\Delta H^o_{\textrm{fusion}}=6.01\ \textrm{kJ/mol}. Calculate the energy required (in KJ) to convert 26.0 g of ice at 0oC to liquid water at 20oC.
Enthalpy: Heat of vaporization
The enthalpy of vaporization for water, ΔHovaporization = 40.7 kJ/mol. Determine the enthalpy change (in kJ) when 54.0 mL of water is condensed from steam at 100oC.
Calorimetry: Using Heat Qualitatively
Consider the following four samples, all initially at room temperature:
The glass cubes are cooled to the same low temperature, and each is placed into one of the containers of room temperature gas.
(a) If Container A had a lower final temperature at thermal equilibrium than Container B, which gas has the lower specific heat capacity?
Calorimetry
From the following reaction:
2C6H12O6(s)+2O2(g)→12CO2(g)+4H2O(l)ΔHrxn=−1680kJ/mol⋅rxn2C_6H_{12}O_6(s)+2O_2(g)→12CO_2(g)+4H_2O(l)\qquad \qquad ΔH_{rxn}=-1680kJ/mol\cdot rxn
When an unknown quantity of sugar (C6H12O6) is burned, it was determined that the internal energy of the system decreased by 960 kJ and that 15.7 kJ of work was done on the system. How many grams of sugar were present at the beginning of this specific reaction?
Thermochemistry: Calorimetry
Consider the following balanced chemical equation:
2CℓF3(g)+2O2(g)→Cℓ2O(g)+3F2O(g)∆rH°=???2 CℓF_3(g) + 2 O_2(g) → Cℓ_2O (g) + 3 F_2O (g) \qquad ∆_rH° = ???
Substance△fH° @ 25°C (kJ/mol)CℓF3(g)??Cℓ2O(g)80.3F2O(g)–21.7F2(g)0\def\arraystretch{1.5} \begin{array}{|c|c|} \hline \rm Substance & △_fH° \ @ \ 25°C \ (kJ/mol) \\ \hline CℓF_3 (g) & ?? \\ \hline Cℓ_2O (g) & 80.3 \\ \hline F_2O (g) & – 21.7 \\ \hline F_2 (g) & 0 \\ \hline \end{array}
Thermochemistry: Calorimetry
Ethene gas (C2H4) reacts with fluorine gas according to the following balanced equation:
C2H4(g)+6F2(g)→2CF4(g)+4HF(g)ΔrH°=−3012 kJ/mol⋅rxn @ 25°CC_2H_4 (g) + 6 F_2 (g) → 2 CF_4 (g) + 4 HF (g) \qquadΔ_rH° = −3012\ kJ/mol·rxn\ @ \ 25°C
Substance△fH° @ 25°C(kJ/mol)C2H4(g)??F2(g)0CF4(g)–933HF(g)–273.3\def\arraystretch{1.5} \begin{array}{|c|c|} \hline \rm{Substance} & △_fH° \ @ \ 25°C(kJ/mol) \\ \hline C_2H_4 (g) & ?? \\ \hline F_2 (g) & 0 \\ \hline CF_4 (g) & – 933 \\ \hline HF (g) & – 273.3 \\ \hline \end{array}
Calculate ΔHof of SO2(g) given the following: (answer in kJ)
S(s)+1.5O2(g)→SO3(g),ΔH=−395.8kJS_{(s)}+1.5O_{2(g)}\to SO_{3(g)}, \Delta H=-395.8kJ
2SO2(g)+O2(g)→2SO3(g),ΔH=−198.2kJ2SO_{2(g)}+O_{2(g)}\to 2SO_{3(g)}, \Delta H=-198.2kJ
The standard enthalpy of formation for the amino acid glycine C2H5NO2(s) is –528 kJ/mol at 25°C. Which reaction below has an enthalpy change equal to this value?
Calorimetry
Equal masses of copper and iron are both heated to a temperature of 100°C. Each piece of hot metal is then placed into its own beaker containing 1 L of water. If both beakers start with the same initial temperature, which beaker of water will have the higher final temperature?
Additional information: Cu (specific heat = 0.377 J/g°C); Fe (specific heat = 0.450 J/g°C)
Thermochemistry Calculations
The standard enthalpy of fusion for water, ΔHofusion = 6.01 kJ/mol. Calculate the energy required (in kJ) to convert 26.0 g of ice at 0oC to liquid water at 20oC.
Thermochemistry: Enthalpy change
The enthalpy of vaporization for water, ΔHovaporization = 40.7 kJ/mol. Determine the enthalpy change (in kJ) when 54.0 mL of water is condensed from steam at 100oC.
Bomb Calorimetry
A fuel compound was burned in oxygen in a bomb calorimeter with a combined heat capacity of 56.5 k] / K (Bomb assembly and water). The temperature of the surrounding water bath increased from 23.5 to 34.6 °C as a result. Calculate ΔEsys\Delta E_{sys}for this process? Consider any change in pressure due to gaseous products to be negligible.