Chapter 5 Thermochemistry Ashley Warren Kings High School AP Chemistry 2013-2014 Lecture Presentation © 2012 Pearson Education, Inc.

Chapter 5

Thermochemistry

Ashley Warren Kings High School

AP Chemistry 2013-2014

Lecture Presentation

© 2012 Pearson Education, Inc.

Thermodynamics v. Thermochemistry

• The study of energy and its transformations is known as thermodynamics. – Industrial Revolution – Relationships between chemical reactions and

energy changes that involve heat = thermochemistry

Thermochemistry


Energy

• Energy is the ability to do work or transfer heat.– Energy used to cause an object

that has mass to move is called work.

– Energy used to cause the temperature of an object to rise is called heat.

Thermochemistry


Definitions: Work

• Energy used to move an object over some distance is work:

• w = F dwhere w is work, F is the force, and d is the distance over which the force is exerted.

Thermochemistry


Heat

• Energy can also be transferred as heat.

• Heat flows from warmer objects to cooler objects.

Thermochemistry


Kinetic EnergyKinetic energy is energy an object possesses by virtue of its motion:

12

Ek = mv2

Thermochemistry


Potential Energy• Potential energy is

energy an object possesses by virtue of its position or chemical composition.– It arises why????***

• The most important form of potential energy in molecules is electrostatic potential energy, Eel:

KQ1Q2

dEel =

Thermochemistry


Units of Energy

• The SI unit of energy is the joule (J):

• An older, non-SI unit is still in widespread use: the calorie (cal):

1 cal = 4.184 J

1 J = 1 kg m2

s2

Thermochemistry


Definitions: System and Surroundings

• The system includes the molecules we want to study (here, the hydrogen and oxygen molecules).– 3 types

• The surroundings are everything else (here, the cylinder and piston).

Thermochemistry

Describing and Calculating Energy Changes

• A bowler lifts a 5.4 kg (12 lb) bowling ball from ground level to a height of 1.6 m (5.2 ft) and then drops it. – a.) What happens to the potential energy of the ball as

it is raised? – b.) What quantity of work, in J, is used to raise the

ball? – c.) The ball is dropped. If all the work done in part (b)

has been converted to kinetic energy by the time the ball strikes the ground, what is the ball’s speed just before it hits the ground?


Thermochemistry


First Law of Thermodynamics• Energy is neither created nor destroyed.• In other words, the total energy of the universe is

a constant; if the system loses energy, it must be gained by the surroundings, and vice versa.

Thermochemistry


Internal EnergyThe internal energy of a system is the sum of all kinetic and potential energies of all components of the system; we call it E.

Thermochemistry


Internal EnergyBy definition, the change in internal energy, E, is the final energy of the system minus the initial energy of the system:

E = Efinal − Einitial

Thermochemistry

Energy Diagram

• The internal energy for Mg (s) and Cl2 (g) is greater than that of MgCl2 (s). Sketch an energy diagram that represents the following reaction:

• MgCl2 (s) Mg (s) + Cl2 (g)


Thermochemistry

Internal Energy• Thermodynamic quantities such as ΔE have 3

parts: – A #– A Unit -- A sign that signifies direction


Thermochemistry


Changes in Internal Energy• If E > 0, Efinal > Einitial

– Therefore, the system absorbed energy from the surroundings. (meaning that E is positive)

– This energy change is called endergonic.

Thermochemistry


Changes in Internal Energy• If E < 0, Efinal < Einitial

– Therefore, the system released energy to the surroundings.

– This energy change is called exergonic.

Thermochemistry


Changes in Internal Energy

• When energy is exchanged between the system and the surroundings, it is exchanged as either heat (q) or work (w).

• That is, E = q + w.

Thermochemistry


E, q, w, and Their Signs

Thermochemistry

Relating Heat & Work to Internal Energy

• Gases A and B are confined in a cylinder and piston arrangement and they react to form a solid product C. As the reaction occurs, the system loses 1150J of heat to the surroundings. The piston moves downward as the gases react to form the solid. As the volume of the gas decreases under the constant pressure of the atmosphere, the surroundings do 480 J of work on the system. What is the change in the internal energy of the system?


Thermochemistry


Exchange of Heat between System and Surroundings

• When heat is absorbed by the system from the surroundings, the process is endothermic.

• When heat is released by the system into the surroundings, the process is exothermic.

Thermochemistry


State Functions• We do know that the internal energy of a

system is independent of the path by which the system achieved that state.– In the system depicted in the figure below, the

water could have reached room temperature from either direction.

Thermochemistry


State Functions• Therefore, internal energy is a state function.• It depends only on the present state of the

system, not on the path by which the system arrived at that state.

• And so, E depends only on Einitial and Efinal.

Thermochemistry


State Functions

• However, q and w are not state functions.

• Whether the battery is shorted out or is discharged by running the fan, its E is the same.– But q and w are different

in the two cases.

Thermochemistry


Work

Usually in an open container the only work done is by a gas pushing on the surroundings (or by the surroundings pushing on the gas).

Thermochemistry


WorkWe can measure the work done by the gas if the reaction is done in a vessel that has been fitted with a piston:

w = −PV

Thermochemistry


Enthalpy• A system that consists of a gas confined to a container

can be characterized by several different properties. • If a process takes place at constant pressure (as the

majority of processes we study do) and the only work done is this pressure–volume work, we can account for heat flow during the process by measuring the enthalpy of the system.

• Enthalpy (denoted with symbol, H) is the internal energy plus the product of pressure and volume:

H = E + PV

Thermochemistry


Enthalpy

• When the system changes at constant pressure, the change in enthalpy, H, is

H = (E + PV)

• This can be written

H = E + PV

Thermochemistry


Enthalpy

• Since E = q + w and w = −PV, we can substitute these into the enthalpy expression:

H = E + PV

H = (q + w) − w

H = q

• So, at constant pressure, the change in enthalpy is the heat gained or lost.

Thermochemistry


Endothermicity and Exothermicity

• A process is endothermic when H is positive.

• A process is exothermic when H is negative.

Thermochemistry

Determining the Sign of ΔH. • Indicate the sign of the enthalpy change, ΔH, in the

processes carried out under atmospheric pressure and indicate whether each process is endothermic or exothermic. – a.) An ice cube melts.

– b.) 1 gram of butane (C4H10) is combusted in sufficient oxygen to give complete combustion of CO2 and H2O.

– Molten gold poured into a mold solidifies at atmospheric pressure. With the gold defined as the system, is the solidification an exothermic or endothermic process? Explain.


Thermochemistry


Enthalpy of Reaction

The change in enthalpy, H, is the enthalpy of the products minus the enthalpy of the reactants:

H = Hproducts − Hreactants

Thermochemistry


Enthalpy of ReactionThis quantity, H, is called the enthalpy of reaction, or the heat of reaction.

Thermochemistry


The Truth about EnthalpyThe following guidelines are helpful when using thermochemical equations and enthalpy diagrams:

1.Enthalpy is an extensive property.

2.H for a reaction in the forward direction is equal in size, but opposite in sign, to H for the reverse reaction.

3.H for a reaction depends on the state of the products and the state of the reactants.

Thermochemistry

Relating ΔH to Quantities of Reactants and Products

• How much heat is released when 4.50 grams of methane gas is burned in a constant pressure system?

CH4 (g) + 2 O2 (g) CO2 (g) + 2H2O (l) ΔH = -890 kJ

• Hydrogen Peroxide can decompose to water and oxygen. Calculate the quantity of heat released when 5.00 g of hydrogen peroxide decomopses at constant pressure. 2H2O2 (l) 2H2O (l) + O2 (g) ΔH = -196 kJ


Thermochemistry


Calorimetry

Since we cannot know the exact enthalpy of the reactants and products, we measure H through calorimetry, the measurement of heat flow.

Thermochemistry


Heat Capacity and Specific HeatThe amount of energy required to raise the temperature of a substance by 1 K (1C) is its heat capacity.

Thermochemistry


Heat Capacity and Specific Heat

We define specific heat capacity (or simply specific heat) as the amount of energy required to raise the temperature of 1 g of a substance by 1 K (or 1 C).

Thermochemistry


Heat Capacity and Specific Heat

Specific heat, then, is

Specific heat =heat transferred

mass temperature change

s =q

m T

Thermochemistry

Relating Heat, Temperature Change, and Heat Capacity

• How much heat is needed to warm 250 grams of water from 22 degrees Celsius to 98 degrees Celsius?

• What is the molar heat capacity of water? • Assume the specific heat of rocks is 0.82 J/g – K.

Calculate the quantity of heat absorbed by the 50.0 kg of rocks if their temperature increases by 12.0 degrees Celsius.

• What temperature change would these rocks undergo if they emitted 450 kJ of heat?


Thermochemistry


Constant Pressure Calorimetry

By carrying out a reaction in aqueous solution in a simple calorimeter such as this one, one can indirectly measure the heat change for the system by measuring the heat change for the water in the calorimeter.

Thermochemistry


Constant Pressure Calorimetry

Because the specific heat for water is well known (4.184 J/g-K), we can measure H for the reaction with this equation:

q = m s T

qsoln = -qrxn

Thermochemistry

Measuring ΔH• When a student mixes 50 mL of 1.0 M HCl and 50

mL of 1.0 M NaOH in a coffee cup calorimeter, the temperature of the resultant solution increases from 21.0° Celsius to 27.5° Celsius. Calculate the enthalpy change for the reaction in kJ/mol HCl, assuming the calorimeter loses only a negligible quantity of heat, that the total volume of the solution is 100 mL, that its density is 1.0 g/mL, and that its specific heat is 4.18 J/g – K.


Thermochemistry

Measuring ΔH #2• When 50.0 mL of 0.100 M AgNO3 and 50.0 mL of

0.100 M HCl are mixed in a constant pressure calorimeter, the temperature of the mixture increases from 22.30° Celsius to 23.11° Celsius. Calculate ΔH for this reaction in kJ/mol AgNO3, assuming that the combined solution has a mass of 100.0 grams and a specifi heat of 4.18 J/g-°C.


Thermochemistry


Bomb Calorimetry• Reactions can be

carried out in a sealed “bomb” such as this one.

• The heat absorbed (or released) by the water is a very good approximation of the enthalpy change for the reaction.

• qrxn = -Ccal x ΔT

Thermochemistry

Bomb Calorimetry

• The combustion of methylhydrazine (CH6N2), a liquid rocket fuel, produces N2 (g), CO2 (g), and H2O (l). When 4.00 grams of methylhydrazine is combusted in a bomb calorimeter, the temperature of the calorimeter increases from 25.00° C to 39.50°C. In a separate experiment the heat capacity of the calorimeter is measured to be 7.794 kJ/°C. Calculate the heat of reaction for the combustion of a mole of CH6N2.


Thermochemistry


Bomb Calorimetry

• Because the volume in the bomb calorimeter is constant, what is measured is really the change in internal energy, E, not H.

• For most reactions, the difference is very small.

Thermochemistry


Hess’s Law

H is well known for many reactions, and it is inconvenient to measure H for every reaction in which we are interested.

• However, we can estimate H using published H values and the properties of enthalpy.

Thermochemistry


Hess’s Law

Hess’s law states that “[i]f a reaction is carried out in a series of steps, H for the overall reaction will be equal to the sum of the enthalpy changes for the individual steps.”

Thermochemistry

Hess’s Law Examples


Thermochemistry


Hess’s Law

Because H is a state function, the total enthalpy change depends only on the initial state of the reactants and the final state of the products.

Thermochemistry


Enthalpies of Formation

An enthalpy of formation, Hf, is defined as the enthalpy change for the reaction in which a compound is made from its constituent elements in their elemental forms.

Thermochemistry


Standard Enthalpies of FormationStandard enthalpies of formation, Hf°, are measured under standard conditions (25 °C and 1.00 atm pressure).

Thermochemistry

Enthalpies of Formation• For which of these reactions at 25°C does the

enthalpy change represent a standard enthalpy of formation? For each that does not, what changes are needed to make it an equation whose H is an enthalpy of formation?– 2 Na (s) + ½ O2 (g) Na2O (s)

– 2 K (l) + Cl2 (g) 2KCl (s)

– C6H12O6 (s) 6C (diamond) + 6H2 (g) + 3 O2 (g)


Thermochemistry


Calculation of H

• Consider the reaction above• Imagine this as occurring

in three steps:

C3H8(g) 3C(graphite) + 4H2(g)

3C(graphite) + 3O2(g) 3CO2(g)

4H2(g) + 2O2(g) 4H2O(l)

C3H8(g) + 5O2(g) 3CO2(g) + 4H2O(l)

Thermochemistry


C3H8(g) 3C(graphite) + 4H2(g)

3C(graphite) + 3O2(g) 3CO2(g)

4H2(g) + 2O2(g) 4H2O(l)

C3H8(g) + 5O2(g) 3CO2(g) + 4H2O(l)

Calculation of H

• The sum of these equations is

C3H8(g) + 5O2(g) 3CO2(g) + 4H2O(l)

Thermochemistry


Calculation of H

We can use Hess’s law in this way:

H = nHf,products – mHf°,reactants

where n and m are the stoichiometric coefficients.

Thermochemistry


H = [3(−393.5 kJ) + 4(−285.8 kJ)] – [1(−103.85 kJ) + 5(0 kJ)]= [(−1180.5 kJ) + (−1143.2 kJ)] – [(−103.85 kJ) + (0 kJ)]= (−2323.7 kJ) – (−103.85 kJ) = −2219.9 kJ

Calculation of H

C3H8(g) + 5O2(g) 3CO2(g) + 4H2O(l)

Thermochemistry

Guidelines to Remember for this Calculation

• Decomposition – the reverse formation reaction; use the negative version of the standard enthalpy of formation.

• Stoichiometric Coefficients – if there are 3 moles of CO2, then we must multiply the standard enthalpy of formation by 3.


Thermochemistry

Enthalpy of Reaction from Enthalpies of Formation

• Calculate the standard enthalpy change for the combustion of 1 mol of benzene, C6H6 (l) to CO2 (g) and H2O (l).

• Compare the quantity of heat produced by combustion of 1.00 g propane with that produced by 1.00 g benzene.


Thermochemistry

Example #2• The standard enthalpy change for the reaction

CaCO3 (s) CaO (s) + CO2 (g) is 178.1 kJ. Use Table 5.3 to calculate the standard enthalpy of formation of CaCO3 (s).


Thermochemistry


Energy in FoodsMost of the fuel in the food we eat comes from carbohydrates and fats.

Thermochemistry

Fuels• A 28 g serving of a popular breakfast cereal served

with 120 mL of skim milk provides 8 g protein, 26 g carbs, and 2 g fat. Using the average fuel values of these substances, estimate the fuel value (caloric content) of this serving.

• A person of average weights uses about 100 Cal/mi when running or jogging. How many servings of this cereal provide the fuel value requirements to run 3 mi?


Thermochemistry


Energy in Fuels

The vast majority of the energy consumed in this country comes from fossil fuels.

Thermochemistry

Chapter 5 Homework Problems

• Conceptual Problems: – 3, 4, 6, 8, 9

• Problems:– 16, 24, 26, 28, 32, 34, 37, 39, 41, 44, 47,

50, 53, 58, 62, 64, 73, 75, 78, 95


Chapter 5 Thermochemistry Ashley Warren Kings High School AP Chemistry 2013-2014 Lecture Presentation © 2012 Pearson Education, Inc.

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