8. Thermochemistry

Heat Capacity

What is Heat Capacity?

Titanium metal is used as a structural material in many high-tech applications, such as in jet engines. what is the specific heat of titanium in J/(g·°C) if it takes 89.7 J to raise the temeprature of a 33.0 g block of 5.20 °C? What is the molar heat capacity of titanium J/(mol·°C)?

Understanding Heat Capacity and Specific Heat Capacity

The dissolution of CaCl(s) in water is exothermic, with ΔHsoln= - 81.3 kJ>mol. If you were to prepare a 1.00 m 
solution of CaCl2 beginning with water at 25.0 °C, what would the final temperature of the solution be in °C? Assume that the specific heats of both pure H2O and the solution are the same, 4.18 J>1K g2.

How much heat is required to warm 1.50 kg of sand from
25.0 °C to 100.0 °C?

Use standard enthalpies of formation to calculate the standard
change in enthalpy for the melting of ice. (The ΔH °f for H2O(s)
is -291.8 kJ/mol.) Use this value to calculate the mass of ice
required to cool 355 mL of a beverage from room temperature
(25.0 °C) to 0.0 °C. Assume that the specific heat capacity and
density of the beverage are the same as those of water.

GCSE Physics - Internal Energy and Specific Heat Capacity  #27

Assuming that Coca-Cola has the saem specific heat as water [4.18 J/(g C)], calculate the amount of heat in kilojoules transferred when one can (about 350 g) is cooled from 25 C to 3 C.

A house is designed to have passive solar energy features.
Brickwork incorporated into the interior of the house acts
as a heat absorber. Each brick weighs approximately 1.8 kg.
The specific heat of the brick is 0.85 J>g@K. How many bricks
must be incorporated into the interior of the house to provide
the same total heat capacity as 1.7 * 103 gal of water?

The specific heat of octane, C8H181l2, is 2.22 J>g@K. (b) Which will require more
heat, increasing the temperature of 1 mol of C8H181l2 by
a certain amount or increasing the temperature of 1 mol of
H2O1l2 by the same amount?

How much heat is required to raise a 50.0 g piece of iron from 25 °C to its melting point of 1538 °C? The specific heat capacity for iron is 0.451 J/g•°C.
(a) 34.1 kJ
(b) 168 kJ
(c) 12.1 kJ
(d) 6.78 kJ

Specific Heat Capacity  | Matter | Physics | FuseSchool

Heat Capacity Example 3

The specific heat of octane, C8H181l2, is 2.22 J>g@K. (a) How
many J of heat are needed to raise the temperature of 80.0 g
of octane from 10.0 to 25.0 °C?

50.00 g of heated metal ore is placed into an insulated beaker containing 822.5 g of water. Once the metal heats up the final temperature of the water is 32.08 ºC. If the metal gains 14.55 kJ of energy, what is the initial temperature of the water? The specific heat capacity of copper is 4.184 J / g  ºC.

A sample of copper absorbs 3.53 kJ of heat, which increases the temperature by 25 ºC, determine the mass (in kg) of the copper sample if the specific heat capacity of copper is 0.385 J / g  ºC.

Heat Capacity, Specific Heat, and Calorimetry

Instant cold packs used to treat athletic injuries contain solid NH4NO3 and a pouch of water. When the pack is squeezed, the pouch breaks and the solid dissolves, lowering the tem-perature because of the endothermic reaction NH4NO31s2 ¡ NH4NO31aq2 ∆H = +25.7 kJ What is the final temperature in a squeezed cold pack that contains 50.0 g of NH4NO3 dissolved in 125 mL of water? Assume a specific heat of 4.18 J/(g C) for the solution, an initial temperature of 25.0 °C, and no heat transfer between the cold pack and the environment.

Sodium metal is sometimes used as a cooling agent in heat-exchange units because of its releatively high molar heat capacity fo 28.2 J/(mol·°C). What is the specific heat and molar heat capacity of sodium in J/g·°C?

A kilogram of aluminum metal and a kilogram of water are each
warmed to 75 °C and placed in two identical insulated containers.
One hour later, the two containers are opened and the temperature
of each substance is measured. The aluminum has
cooled to 35 °C, while the water has cooled only to 66 °C.
Explain this difference.

snow melts on a mountaintop (snow is the system)

water vapor condenses on a cold surface (water vapor is the system)

a rock tumbles from a mountain (the rock is the system)

In both of these examples, the system undergoes a heat loss.

 An increase in the heat of a system is denoted by a positive sign.  When a system has work done on it, work is positive. 

Which of the following exchanges of energy has a positive sign for heat (relative to the system)?

Heat Capacity Example 2

What is the difference between Heat Capacity and Specific Heat Capacity?

The enthalpy of fusion of water is 6.01 kJ/mol. Sunlight striking Earth's surface supplies 168 W per square meter (1 W = 1 watt = 1 J/s). (b) The specific heat capacity of ice is 2.032 J/g°C. If the initial temperature of a 1.00 square emter patch of ice is -5.0°C, what is its final temperature after being in sunlight for 12 h, assuming no phase changes and assuming that sunlight penetration uniformly to a depth of 1.00 cm?

Given 400.0 g of hot tea at 80.0 °C, what mass of ice at 0 °C must be added to obtain iced tea at 10.0 °C? The specific heat of the tea is 4.18 J>1g °C2 and ΔHfusion for ice is 
+ 6.01 kJ>mol.

Based on their given specific heat capacities which compound would show the greatest temperature change upon absorbing 25.0 J of heat?&nbsp;<img src="https://static.studychannel.pearsonprd.tech/courses/general-chemistry/thumbnails/images/7a1e0b81"alt=""/>

How much heat is required to warm 1.50 L of water from 25.0 °C
to 100.0 °C? (Assume a density of 1.0 g/mL for the water.)

Consider the data about gold metal in Exercise 5.26(b).
(b) Suppose that the same amount of heat is added to two
10.0-g blocks of metal, both initially at the same temperature.
One block is gold metal, and one is iron metal.
Which block will have the greater rise in temperature after
the addition of the heat?

An unknown mass of each substance, initially at 23.0 °C,
absorbs 1.95 * 103 J of heat. The final temperature is recorded.
Find the mass of each substance.
a. Pyrex glass (Tf = 55.4 °C)

Calculate the amount of heat required to raise the tempera- ture of 250.0 g (approximately 1 cup) of hot chocolate from 25.0 °C to 80.0 °C. Assume hot chocolate has the same spe-cific heat as water 34.18 J>1g °C24.

Instant cold packs used to ice athletic injuries on the field contain
ammonium nitrate and water separated by a thin plastic
divider. When the divider is broken, the ammonium nitrate dissolves
according to the endothermic reaction:
NH4NO3(s)¡NH4 + (aq) + NO3- (aq)
In order to measure the enthalpy change for this reaction,
1.25 g of NH4NO3 is dissolved in enough water to make
25.0 mL of solution. The initial temperature is 25.8 °C and the
final temperature (after the solid dissolves) is 21.9 °C. Calculate
the change in enthalpy for the reaction in kJ. (Use 1.0 g >
mL as the density of the solution and 4.18 J>g # °C as the specific
heat capacity.)

Imagine that you dissolve 10.0 g of a mixture of NaNO3 and KF in 100.0 g of water and find that the temperature rises by 2.22 °C. Using the following data, calculate the mass of each compound in the original mixture. Assume that the specific heat of the solution is 4.18 J>1 g °C2 
NaNO31s2 S NaNO31aq2 ΔH = + 20.4 kJ>mol 
KF1s2 S KF1aq2 ΔH = - 17.7 kJ>mol

Palmitic acid (C16H32O2) is a dietary fat found in beef and butter.
The caloric content of palmitic acid is typical of fats in general.
Which dietary substance (sugar or fat) contains more Calories
per gram? The standard enthalpy of formation of palmitic acid
is -208 kJ/mol and that of sucrose is -2226.1 kJ>mol. [Use
H2O(l ) in the balanced chemical equations because the metabolism
of these compounds produces liquid water.]

Two solid objects, A and B, are placed in boiling water and allowed
to come to the temperature of the water. Each is then
lifted out and placed in separate beakers containing 1000 g
of water at 10.0 °C. Object A increases the water temperature
by 3.50 °C; B increases the water temperature by 2.60 °C.
(a) Which object has the larger heat capacity?

The complete combustion of methane, CH41g2, to form
H2O1l2 and CO21g2 at constant pressure releases 890 kJ of
heat per mole of CH4. (b) Draw an enthalpy diagram for
the reaction.

(b) The specific heat of
aluminum is 0.9 J/1g # K2. Calculate its molar heat capacity.

(d) How many kJ of heat are needed to raise the
temperature of 5.00 kg of liquid water from 24.6 to 46.2 °C?

Ethyl chloride (C2H5Cl) boils at 12 °C. When liquid C2H5Cl under pressure is sprayed on a room-temperature (25 °C) surface in air, the surface is cooled considerably. (a) What does this observation tell us about the specific heat of C2H5Cl(g) as compared with that of C2H5Cl(l)?

(b) Calculate the energy needed for
this temperature change.

Calorimetry Examples: How to Find Heat and Specific Heat Capacity

We said in Section 9.1 that the potential energy of water at the top of a dam or waterfall is converted into heat when the water dashes against rocks at the bottom. The potential energy of the water at the top is equal to EP = mgh, where m is the mass of the water, g is the acceleration of the falling water due to gravity 1g = 9.81 m>s22, and h is the height of the water. Assuming that all the energy is converted to heat, calculate the temperature rise of the water in degrees Celsius after falling over California’s Yosemite Falls, a distance of 739 m. The specific heat of water is 4.18 J/(g·K).

Heat Capacity Example 1

GCSE Science Revision Physics "Specific Heat Capacity"

What is the difference between heat capacity and specific heat?

Assume that the kinetic energy of a 1400-kg car moving at 115 km/h (Problem 9.46) could be converted entirely into heat. What amount of water could be heated from 20 °C to 50 °C by the car’s energy? 4.184 J are required to heat 1 g of water by 1 °C.

(c) What is the heat capacity of 370 g of
liquid water?

Consider the data about gold metal in Exercise 5.26(b). 
(c) What is the molar heat capacity 
of Au(s)?

Assume that the kinetic energy of a 1400 kg car moving at 115 km/h (Problem 1.78) is converted entirely into heat. How many calories of heat are released, and what amount of water in liters could be heated from 20.0 °C to 50.0 °C by the car’s energy? (One calorie raises the temperature of 1 mL of water by 1 °C)

When 12.5 g of NH4NO3 is dissolved in 150.0 g of water of 25.0 °C in a coffee cup calorimeter, the final temperature of the solution of 19.7 °C. Assume that the specific heat of the solution is the same as that of water, 4.18 J/(g•°C). What is the ΔH per mol of NH4NO3? (LO 9.10)
NH4NO3 (s) → NH4+ (aq) + NO3−(aq) ΔH = ?
(a) +3.60 kJ
(b) +23.0 kJ
(c) +21.3 kJ
(d) −3.60 kJ

Suppose that 25 g of each substance is initially at 27.0 °C. What 
is the final temperature of each substance upon absorbing 
2.35 kJ of heat? 
c. aluminum

LP gas burns according to the exothermic reaction:
C3H8( g) + 5 O2( g)¡3 CO2( g) + 4 H2O( g)
ΔH °rxn = -2044 kJ
What mass of LP gas is necessary to heat 1.5 L of water from room
temperature (25.0 °C) to boiling (100.0 °C)? Assume that during
heating, 15% of the heat emitted by the LP gas combustion goes
to heat the water. The rest is lost as heat to the surroundings.

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6. Chemical Quantities & Aqueous Reactions

7. Gases

9. Quantum Mechanics

10. Periodic Properties of the Elements

11. Bonding & Molecular Structure

12. Molecular Shapes & Valence Bond Theory

13. Liquids, Solids & Intermolecular Forces

14. Solutions

15. Chemical Kinetics

16. Chemical Equilibrium

17. Acid and Base Equilibrium

18. Aqueous Equilibrium

19. Chemical Thermodynamics

20. Electrochemistry

21. Nuclear Chemistry

22. Organic Chemistry

23. Chemistry of the Nonmetals

24. Transition Metals and Coordination Compounds

[BOOK CHEM] Clutch Version

Hess's Law

Hess's Law involves the use of partial reactions in the determination of the overall enthalpy of reaction.

Nature of Energy

Thermochemistry is the study of matter and energy associated with chemical reactions or physical changes.&nbsp;

Endothermic & Exothermic Reactions

Endothermic and Exothermic Reactions involve the absorbing and releasing of thermal energy respectively.&nbsp;&nbsp;

Heat Capacity is the amount of heat required to change the temperature of a substance.&nbsp;

Thermochemical Equations

Thermochemical Equations&nbsp;involve a balanced chemical equation with a given enthalpy value.

Internal Energy

The Internal Energy of a system is the total energy from all forms of kinetic and potential energy.&nbsp;

Kinetic & Potential Energy

Mechanical Energy is the energy associated with an object's velocity (Kinetic Energy) and its position (Potential Energy).&nbsp;

Formation Equations

Formation Equations involve the standard states of elements combining to form 1 mole of product.

Enthalpy of Formation

The Enthalpy of Formation for an element is a key component in determining the enthalpy of reaction.&nbsp;

Thermal Equilibrium

Thermal Equilibrium involves two substances that are in physical contact reaching the same final temperature over time.&nbsp;

Constant-Pressure Calorimetry

Constant-Pressure Calorimetry uses a coffee cup calorimeter to determine the heat transfers occurring in a liquid solution.

Constant-Volume Calorimetry

Constant-Volume Calorimetry uses a bomb calorimeter and a combustion reaction to determine its enthalpy of reaction.&nbsp;

First Law of Thermodynamics

The First Law of Thermodynamics states that energy cannot be created nor destroyed, but instead is transferred.&nbsp;