#### Transcript Chapter 5 – Thermochemistry

Thermochemistry Chapter 5

First Law of Thermodynamics •states that energy is conserved •Energy that is lost by a system must be gained by the surroundings or vice versa.

Enthalpy (H)

### • accounts for heat flow in chemical changes occurring at constant pressure

Enthalpy (H)

• State function: • H=H final -H initial

Enthalpy (H)

• When D H > 0, the system has gained heat from surroundings (endothermic) • When D H < 0, the system has released heat to the surroundings (exothermic)

Enthalpy (H)

Enthalpies of Reaction

### Thermochemical equations are balanced chemical equations that show the associated enthalpy change

Enthalpies of Reaction

• D

Guidelines

• (1)

Guidelines

• (2)

D

Guidelines

• (2)

Guidelines

• (3)

D

### H for a reaction depends on the state of the reactants and products and we assume they are at constant T

Practice 5.5

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

• Given: • CH 4 (g) + 2O 2 (g)

CO 2 (g) +2H 2 O(l)

D

H = -890kJ

Enthalpies of Reaction

• D H can be determined experimentally by measuring the heat flow accompanying a reaction at constant pressure which is done by measuring Temperature changes

Enthalpies of Reaction

• Calorimetry is the measurement of heat flow • A calorimeter will measure this heat flow

Heat Capacity

• the amount of heat required to raise the temperature of an object by 1K or 1 o C.

Molar Heat Capacity

• heat capacity of one mole of a substance

Specific Heat Capacity

• heat capacity of 1g of substance

Specific Heat Capacity

Enthalpies of Reaction

• Heat = specific heat x grams of substance x D T • q=cm D T

Sample Exercise 5.5 Relating Heat, Temperature Change, and Heat Capacity (a) How much heat is needed to warm 250 g of water (about 1 cup) from 22 °C (about room temperature) to near its boiling point, 98 °C? The specific heat of water is 4.18 J/g-K. (b) What is the molar heat capacity of water?

Constant- Pressure Calorimetry

Constant- Pressure Calorimetry

• (1)Assume the calorimeter prevents loss or gain of heat from the solution to its surroundings

Constant- Pressure Calorimetry

• (2a) For an exothermic rxn, heat is lost by rxn and gained by soln so the T of soln rises

Constant- Pressure Calorimetry

• (2b) For an endothermic rxn, heat is gained by rxn and lost by soln so the T of soln goes down

Constant- Pressure Calorimetry

• (3) q soln = specific heat of solution x gram of solution x D T = -q rxn

Sample Exercise 5.6 Measuring Δ Coffee-Cup Calorimeter H Using a When a student mixes 50 mL of 1.0 50 mL of 1.0 M M HCl and NaOH in a coffee-cup calorimeter, the temperature of the resultant solution increases from 21.0 °C to 27.5 °C. Calculate the enthalpy change for the reaction in kJ/mol HCl, assuming that 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.

Bomb Calorimetry

• Constant-Volume Calorimetry • Combustion reactions are studied in bomb calorimeters

Bomb Calorimetry

Bomb Calorimetry

• The heat released in combustion is absorbed by the calorimeter contents, causing a rise in the temperature of the water

Bomb Calorimetry

• To calculate D H combustion from the measured temperature increase, we must know the heat capacity of the bomb calorimeter, Ccal

Bomb Calorimetry

• q rxn = -Ccal x D T

Sample Exercise 5.7 Measuring q rxn Using a Bomb Calorimeter • Methylhydrazine (CH6N2) is used as a liquid rocket fuel. The combustion of methylhydrazine with oxygen produces N2( • • 2 CH6N2( l ) + 5 O2( g g ), CO2( ) → 2 N2( g g ), and H2O( ) + 2 CO2( g l ): ) + 6 H2O( l ) • When 4.00 g 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.

Hess’ Law

• If a reaction is carried out in a series of steps, D Hrxn will equal the sum of D H for the individual steps. D H rxn = D H 1 + D H 2 + …

Hess’ Law

• Hess’ law provides ways to calculate energy changes that are difficult to measure directly.

Sample Exercise 5.8 Using Hess’s Law to Calculate Δ H

Hess’ Law

Hess’ Law (5.9)

Practice Exercise Hess’ Law (5.9)

Standard enthalpy of a reaction • is the enthalpy change when all reactants and products are in their standard states, D H o rxn

Standard States • Standard state is when a substance is in its pure form at 1 atm and 298K.

Standard enthalpy of formation D H o f, , reported in kj/mol, is the change in enthalpy for the reaction that forms 1 mole of the compound from its elements, with all substances in their standard state.

Standard enthalpy of a formation • D H f o of the most stable form of an element is zero because there is no formation reaction needed for an element in its standard state.

• Standard enthalpy of a formation

• Standard enthalpy of a formation (5.10)

Enthalpy of Reaction • D H o rxn m D H o f = S n D H o f (products) m are the coefficients in the S (reactants) , where n and balanced chemical equation

5.11 Enthalpy of Reaction

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

5.12 Enthalpy of Reaction

• The standard enthalpy change for the decomposition reaction of CaCO 3 is 178.1 kJ. From the values for the standard enthalpies of formation of CaO( CaCO s 3 ) and CO standard enthalpy of formation of ( s ).

2 ( g ), calculate the