ThermochemistryEnergy

Foods and Fuels

Enthalpy Enthalpies of Reaction

Calorimetry

Hess’s Law

Enthalpies of Formation

First Law of Thermodynamics

History of thermodynamics Study of energy and its

transformations Study began during industrial

revolution to develop relationships among heat, work, and fuel in steam engines.

Examines relationships between chemical reactions and energy changes that involve heat.

Nature of Energy (definitions galore!)

Define energy.

Define work.

Define heat.

Kinetic vs. Potential Energy

Objects possess kinetic energy. How can we represent kinetic energy in an equation?

How does potential energy differ from kinetic energy?

How is potential energy converted to kinetic energy?

Potential Energy under the microscope

Gravitational forces play little to no role in interactions between atoms and molecules. Predict what plays a much more important role.

**electrostatic potential energy** This energy is proportional to electrical charges

on two interacting objects, Q1 and Q2, and inversely proportional to the distance, d. Predict an equation.

Electrostatic Potential Energy (Eel)

Eel = κQ1Q2 / d

κ is a constant of proportionality: 8.99 x 109 J-m/C2

How can we make electrostatic potential energy absolute zero?

At finite separation distances for two charged particles, Eel is positive for like charges and negative for opposite charges.

Think about it…

If a cyclist is at the top of the hill (not moving) and stops at the bottom of the hill, are the potential energies and kinetic energies the same at both locations? Why?

Units of Energy SI unit for energy is the joule (J) (more often

than not, use kJ) Using the kinetic energy equation, determine what

1 J of units would be.

Energy changes associated with chemical reactions often expressed as calories (cal) Originally defined as: amount of energy require to raise

the temperature of 1 g of water from 14.5 C to 15.5 C. 1 cal = 4.184 J (exactly) Cal does not = cal.

Which is the system, which is the surroundings?

Open system: matter and energy can be exchanged with surroundings (ie: boiling water)

Closed System: can exchange energy but not matter with surroundings (ie

Isolated System: no exchange of matter or energy. (ie: coffee thermos)

Transferring Energy: Work and Heat

Define Force.

How can we represent work in an equation?

Example Problem

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. What happens to the potential energy of

the ball as it is raised? What is the quantity of work, in J, used to

raise the ball? After the ball is dropped, it gains kinetic

energy. If all the work done in part b has been conver

November 29th, 2012

DO NOW: How are potential energy and

kinetic energy related? What is the kinetic energy, in J, of

the following: An Ar Atom moving at a speed of 650

m/s A mole of Ar atoms moving at 650 m/s?

THE FIRST LAW OF THERMODYNAMICS

Energy can neither be created nor destroyed Any energy lost by a system must be

gained by the surroundings (vice versa)

Energy is conserved

Internal energy

How do we define internal energy?

Chemistry is mainly concerned with the change in energy in a system

How do we represent the change in energy in a system?

Change in Energy

Change in internal energy is denoted:∆E = Efinal – Einital

Generally we can not determine Efinal – Einital for a practical system, but we only need the ∆E to apply the law.

Compare and contrast positive and negative values of ∆E. What does the sign change mean?

Example

For the reaction:2 H2(g) + O2(g) 2 H2O(l)

Which is the initial state and which is final? Label them on the equation.

As water is created, is energy lost or gained to the surroundings?

Does this indicate a positive ∆E or a negative ∆E?

Lets tie it together a little bit.

How may a system exchange energy with its surroundings?

Using the previous slide…

How can we algebraically represent a change in internal energy?

∆E = q + w As work is done on a system from its

surroundings, w > 0. As heat is gained in the system, q > 0.

Example Problem

Calculate the change in internal energy for a process in which a system absorbs 140 J of heat from the surroundings and does 85 J of work on the surroundings.

Endothermic Vs. Exothermic

Define Endothermic.

Define Exothermic.

Is the formation of H2 gas and O2 gas from H2O considered endothermic or exothermic? Why?

State Functions

A state function is a property of a system that is determined only by the present state of the system, not the path that system took. IE:

Change in sea level Vs.

Path Driven How does this relate to ∆E?

ENTHALPY

A system of a gas confined to a container can be characterized by several important properties. Which properties are important?

How are these properties similar to ∆E?

Define Enthalpy.

Why is this equation useful?

Since ∆E involves work, we must consider the most commonly work produced by chemical or physical changes open to the atmosphere. What is this type of work?

Why is the equation or work we gave yesterday no longer useful?

Pressure – Volume Work

Work involved in expansion or compression of gas

When pressure is constant: w = -P∆V

If a system does not change its volume during the course of a process, does it do pressure-volume work?

Enthalpy Cont…

When a change occurs at constant pressure, how can we rewrite our equation?

When ∆H is positive, that means the system has gained heat. That means this is endothermic

When ∆H is negative, that means the system has lost heat. This means it is exothermic.

Practice, Practice, Practice!

Complete the worksheet.

November 30th, 2012

Do Now: Indicate the sign of the enthalpy

change, ∆H, in these processes carried out under atmospheric pressure and indicate whether each process is endothermic or exothermic:

An ice cube melts 1 g of butane is combusted in sufficient

oxygen to give complete combustion to CO2 and H2O

Blah, Blah, Blah… so what did it all mean!

Enthalpy is ∆H= ∆E + P∆V MEANING: ∆H = (q + -P∆V) + P∆V

So what does Enthalpy really tell us? HEAT!

So how is this different than ∆E? ∆E = change in heat at constant volume ∆H = change in heat at constant P.

The difference between the two is very small. It is acceptable to use ∆H

Enthalpy of Reactions

∆H = Hfinal – Hinitial

How could we re-write this equation to represent ∆H for a reaction?

∆H = Hproducts – Hreactants

Reactions

When a numerical value is given for ∆Hrxn , a numerical value along with the equation must be given.

Suppose you are considering the reaction seen at the beginning of class, and that reaction has an enthalpy of -848. How would we write this correctly?

Thermochemical Reactions

A thermochemical equation is a balanced chemical equation that shows the associated enthalpy change but does not specify amount of chemical involved.

The coefficients in the balanced equation represent the number of moles of reactants and products

**IMPORTANT GUIDELINES**

Enthalpy is EXTENSIVE. If 1 mol of CH4 is reacted with 2 moles of O2 and

has an enthalpy of -890 J, how many J would a reaction of 2 mol CH4 and 4 mol O2 give off?

Enthalpy change for a rxn is equal in magnitude, but opposite in sign, to ∆H for the reverse reaction.

The reverse rxn of the previously mentioned rxn would have what value for ∆H?

∆H for a rxn depends on states of reactants and products

Practice

How much heat is released when 4.50 g of methane gas is burned in a constant-pressure system? ( 1 mol CH4 = -890 kJ)

Was this an endothermic or exothermic process?

Calorimetry

•The value of ∆H can be determined experimentally by measuring heat flow accompanying a reaction at constant pressure.

•Measure of heat flow = calorimetry

•Determined by measure of temperature change heat flow produces

Temperature change

Why does the temperature change felt by certain objects/materials differ?

How does molar heat capacity differ from specific heat capacity?

Specific Heat Capacity

How might we solve for specific heat capacity?

Specific heat capacity values are slightly different at different temperatures. Temperature is generally given.

How can we rearrange our equation to solve for q?

Practice!

How much heat is needed to warm 250 g of water from 22 degrees celcius to 98 degrees celcius? What is the molar heat capacity of water?

December 5th, 2012Do Now:

How much heat is needed to warm 300 g of water from 20 degrees celcius to 104 degrees celcius? What is the molar heat capacity of water?

Coffee Cup Calorimetry

•Constant Pressure Calorimetry

•Why does the coffee cup act as a constant pressure calorimeter?

•How do we measure the heat change in a coffee cup calorimeter?

•How does the q of the solution relate to the q of the water surrounding it?

Example

When a student mixes 50 mL of 1.0 M HCl and 50 mL of 1.0M 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 the calorimeter loses only a negligible quantity of heat, that the total volume of the solution is 100 mL, that its density 1.0 g/mL, and specific heat is 4.18 J/gK

BOMB Calorimetry

Bomb Calorimetry

Constant-Volume Calorimetry Studies Combustion May only use after a standardization

measurement Once you have specific heat capacity of

calorimeter, you may use:qrxn = -Ccal X ∆T

Difference between ∆E and ∆H is very small

Example

The combustion of methylhydrazine (CH6N2), a liquid rocket fuel, produces N2(g), CO2 (g), and 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.

Enthalpy of Formation

How can we define Enthalpy 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]

**BE AWARE: enthalpy of formation values are generally for 1 atom.

Hess’s Law

If a reaction is carried out in a series of steps, ∆H for the overall reaction equals the sum of the enthalpy changes for the individual steps.

Example: Calculate the heat of reaction at 1 atm and 298 K for the following reaction: 2HI (g) + F2 2HF (g) + I2 (2)

PRACTICE

Complete worksheet