Energy

The ability to do work.

Not homework—real work

Types of energy

• Radiant energy: E=h(later)

• Thermal energy: q=msT

q=nHv (or mHv)

OR q=nHf (or mHf)

• Chemical energy: Enthalpy (H)

• and others of less interest to a chemist

Symbols and units

• E :Energy, in J h :Planck’s Constant• frequency, /sq :heat, in J• m :mass, in g T :change in temp.oC• n :# of moles H :enthalpy, J/mol

• Hv :enthalpy of vaporization J/mol or J/g

• Hf :enthalpy of fusion J/mol or J/g

• s :specific heat capacity, in j/goC

Law of conservation of energy

• Energy is conserved.

Law of conservation of energy

• Energy is conserved.

• Energy is neither created nor destroyed in any chemical or physical process, though it may be changed to another form.

Heat is not temperature.

• Heat cannot be measured directly, only by its affect on matter.

• If you add heat to a sample, it may…

Heat is not temperature.

• Heat cannot be measured directly, only by its affect on matter.

• If you add heat to a sample, it may…

a) warm up. q=msT

b) melt q=nHf (or mHf)

c) boil q=nHv (or mHv)

d) expand (tough to calculate, don’t bother)

Heat is not temperature.

• Heat cannot be measured directly, only by its affect on matter.

• If you add heat to a sample, it may…

a) warm up. q=msT

b) melt q=nHf (or mHf)

c) boil q=nHv (or mHv)

d) expand (tough to calculate, don’t bother)

All endothermic physical

processes!

Heat is not temperature.

• Heat cannot be measured directly, only by its affect on matter.

• If you add heat to a sample, it may…

a) warm up. q=msT

b) melt q=nHf (or mHf)

c) boil q=nHv (or mHv)

d) expand (tough to calculate, don’t bother)

Can you reverse these?

Consider the Universe

Consider the Universe

• Dude.

Consider the Universe

System

Surroundings

Consider the Universe

System

Surroundings

A system may be open, closed, or isolated

Open system

System

Surroundings

A system may be open, closed, or isolated

Matter

Energy

Closed system

System

Surroundings

A system may be open, closed, or isolated

Energy

Isolated system

System

Surroundings

A system may be open, closed, or isolated

Flow of Energy

System

Surroundings

A process may be endothermic or exothermic

Endothermic process

System

Surroundings

Energy

A process may be endothermic or exothermic

Exothermic process

System

Surroundings

A process may be endothermic or exothermic

Energy

For chemical and physical processes!

System

Surroundings

A process may be endothermic or exothermic

State functions

• Some values assigned to a system are state functions —depending only on the state of the system

• Other values are not state functions—they depend on how the system got there

State functions

• Some values assigned to a system are state functions —depending only on the state of the system

Ex: Energy, P, V, T

• Other values are not state functions—they depend on how the system got there

Ex: Work

Change

• A change in a state function is defined as the difference between the pre- and post- states.

• E=Efinal-Einitial H=Hfinal-Hinitial

• P=Pfinal-Pinitial V=Vfinal-Vinitial

• etc.

Heat

• Heat energy is measured in Joules (like any energy)

• Heat can be added to a system or lost by a system

Chemical Energy

• Chemical energy (enthalpy) is stored in bonds.

Chemical Energy

• Chemical energy (enthalpy) is stored in bonds.

• Forming bonds releases energy

• Breaking bonds requires energy

Chemical Energy

• Chemical energy (enthalpy) is stored in bonds.

• Forming bonds is exothermic

• Breaking bonds is endothermic

Chemical Work.

• When a gas expands against a pressure, it does work on its surroundings.

• When a gas is compressed, its surroundings do work on it.

Warning:

• Dr. Chang reverses the sign on work.

• Do not use his notation.

• For everyone else in the world…

Law of conservation of energy

• Energy is conserved.

• Therefore:

• The change of a system’s energy is equal to the heat it gains, minus the work it does

Law of conservation of energy

• Energy is conserved.

• Therefore:

• The change of a system’s energy is equal to the heat it gains, minus the work it does

E=q-w

Law of conservation of energy

• Energy is conserved.

• Therefore:

• If no work is done, the change in energy is indicated only by the heat. In a chemical reaction, this is the enthalpy

Law of conservation of energy

• Energy is conserved.

• Therefore:

• If no work is done, the change in energy is indicated only by the heat. In a chemical reaction, this is the enthalpy

H=q

Law of conservation of energy

• Energy is conserved.

• Therefore:

• If no work is done, the change in energy is indicated only by the heat. In a chemical reaction, this is the enthalpy

H=q

Work is done when…

• A gas is produced

• A gas is used up

Work is done when…

• A gas is produced

The system does work on the surroundings

• A gas is used up

The surroundings do work on the system

Work is done when…

• A gas is produced

The system does work on the surroundings

The system loses energy

• A gas is used up

The surroundings do work on the system

The system gains energy

PV=nRT

• (PV) is work.

• It could be P(V) (at constant pressure) or

(P)V (at constant volume).

PV=nRT

• (PV) is work.

• It could be P(V) (at constant pressure) or

(P)V (at constant volume).

In either case:

(PV)=(n)RT

The change in the number of moles of gas

What is the change in energy?

1) If q=-115J and w=-10.J?2) If a system gains 200.J and does 12 J of

work?3) If a system gains 124 J of heat as it

expands from .028m3 to .043m3 at 88,000Pa?

4) If a system loses 58 J of heat as it is compressed from 12 L to 7 L at 1.5 atm?

5) If a reaction gives off 125 kJ as it produces 2.9 moles of gas at standard conditions?

How about some stoichiometry?

• The oxidation of copper releases 155.2 kJ/mol.

• How much heat is produced from the oxidation of 15 g Cu?

How about some stoichiometry?

• The oxidation of copper releases 155.2 kJ/mol.

• How much heat is produced from the oxidation of 15 g Cu?

• How much work is done if this occurs at 300k?

Calorimetry

• --the measurement of heat.

Calorimetry

• --the measurement of heat.

• If one thing gains heat…

Calorimetry

• --the measurement of heat.

• If one thing gains heat…

…something else lost it.

• If 75 g of a metal at 96oC is placed in 58 g of water at 21oC and the final temperature reaches 35oC, what is the specific heat capacity of the metal?

Step 1

• How much heat did the water gain?

Step 1

• How much heat did the water gain?

q=msT

Mass of water, in grams

Specific heat of water, 4.18 J/goC

Change in the temperature of water, in oC

Step 2

• How much heat did the metal lose?

Step 2

• How much heat did the metal lose?

• Heat lost = - heat gained

• qlost=-qgained

Step 3

• What is the specific heat capacity of the metal?

Step 3

• What is the specific heat capacity of the metal?

s=q/mT

Mass of metal, in grams

Specific heat of metal, in J/goC

Change in the temperature of metal, in oC

Heat lost by metal

• If 75 g of a metal at 96oC is placed in 58 g of water at 21oC and the final temperature reaches 35oC, what is the specific heat capacity of the metal?

.74 J/goC

Other problems

• Calculate H from calorimetry data.

• Identify and use heats of vaporization, fusion, combustion, solution, formation.

• Combine known H data using Hess’s Law.

• 175.0 g pure H2O at 10.0°C is placed in a calorimeter. 9.80 g pure H2SO4 (at 10.0°C) is added, and the temperature rises to 19.8°C.

(a) What mass increased in temperature?

(b) What was T?

(c) Write the chemical reaction.

(d) What do we assume about the specific heat of the mixture?

(e) Calculate the energy change for the reaction.

(f) What state function does this represent?

(g) Calculate the energy change in kJ/mole.

ELEMENTAL HEATS OF FUSION

• 1 H .0585 kJ/mol• 2 He 0.021 kJ/mol• 3 Li 3 kJ/mol • 4 Be 11.71 kJ/mol• 5 B 22.6 kJ/mol • 6 C ? kJ/mol• 7 N 0.36 kJ/mol • 8 O 0.222 kJ/mol• 9 F 0.26 kJ/mol • 10 Ne 0.34 kJ/mol

ELEMENTAL HEATS OF FUSION

• 2 He 0.021 kJ/mol• 1 H .0585 kJ/mol• 8 O 0.222 kJ/mol• 9 F 0.26 kJ/mol • 10 Ne 0.34 kJ/mol• 7 N 0.36 kJ/mol• 3 Li 3 kJ/mol • 4 Be 11.71 kJ/mol• 5 B 22.6 kJ/mol • 6 C ? kJ/mol

ELEMENTAL HEATS OF FUSION (kJ/mol)

• 26 Fe 13.8• 27 Co 16.19• 28 Ni 17.2• 29 Cu 13.14• 30 Zn 7.38• 47 Ag 11.3• 48 Cd 6.07• 49 In 3.26• 50 Sn 7.2

ELEMENTAL HEATS OF FUSION (kJ/mol)

• 49 In 3.26• 48 Cd 6.07• 50 Sn 7.2• 30 Zn 7.38• 47 Ag 11.3 • 29 Cu 13.14• 26 Fe 13.8• 27 Co 16.19• 28 Ni 17.2

• 53.1 g of H2O is a liquid at 0 °C. How much heat is removed to freeze it (at 0 °C)?

• Calculate the energy needed to boil 50.0 g of water at its boiling point of 100 °C.

• How much energy is required to lower the temperature of a bottle of drinking water (280g) from 22 0C to 4 0C?

• How much heat is needed to raise the temperature of 15 grams of copper by 19° C?

• A 50. g substance loses 223 J of heat energy when the temperature falls from 35°C to 20°C.  What is the specific heat of the substance?

• 53.1 g of H2O is a liquid at 0 °C. How much heat is removed to freeze it (at 0 °C)?

• Calculate the energy needed to boil 50.0 g of water at its boiling point of 100 °C.

• How much energy is required to lower the temperature of a bottle of drinking water (280g) from 22 0C to 4 0C?

• How much heat is needed to raise the temperature of 15 grams of copper by 19° C?

• A 50. g substance loses 223 J of heat energy when the temperature falls from 35°C to 20°C.  What is the specific heat of the substance?

(-)17.7kJ

113kJ

(-)21kJ

110 J

.30J/goC

• How much copper (at its melting point) can be melted by the addition of 2500 J?

• How much heat is required to warm 10. grams of ice from –100C to 00C?

• A 200. gram sample of aluminum, initially at 00C, absorbs 660 J of heat. What is the final temperature of the aluminum sample?

• A 100-gram sample of material requires 1220 J of heat energy to raise its temperature from 200C to 80 0C. What is the specific heat capacity of the material?

• What is the change in temperature (T) if an 89g sample of copper absorbs 678J?

• What would the initial temperature be if a 1.5 mol sample of iron had a final temperature of 55°C and acquired 1111J of heat?

• How much heat is required to heat 45.0 g of H2O at 25.0 °C and then boil it all away?

• How much heat is given off when 11.3 g H2 is burned?(Hcombustion= -286kJ/mol forH2)

• How much H2 must be burned to supply 1000. kJ (=1.000 MJ) of heat?

Write a thermochemical equation to represent the following heats

• Hvap = 28.0 kJ/mol for CS2

• Hformation =- 986.6 kJ/mol for Ca(OH)2

• Hsolution =-154.3kJ/mol for ozone (O3)

• Hfusion = 5.72kJ/mol for CH3COCH3

• Hcombustion = -5471kJ/mol for C8H18

What thermochemical constant is represented?

Al(s, at 20oC) + 25.3 J Al(s, at 21oC)

C6H6+7/2 O26CO2+3H2O+3268 kJ

Ar (s,at-190oC)+1.2 kJ Ar (l,at-190oC)C2H5OH (l, at 78oC) +43.5kJC2H5OH (g, at 78oC)

NH4NO3(s)+25.7 kJ NH4NO3(aq)

2Fe+3/2 O2Fe2O3(s) + 822.1 kJ

From the following data,

CH4+2O2CO2+2H2O(l)   Ho= -890 kJ/mol

H2O(l) H2O(g)         Ho = 44 kJ/mol

(at 298K)

• Calculate the enthalpy of the reaction

• CH4+2O2CO2+2H2O (g)   Ho = ?

Tired of waiting an hour for the coals to get ready? Try LOX!

A grill is prepared using one 9.1 kg (20 pound) bag of charcoal (primarily carbon) and two 3.8L

(1 gallon) buckets of liquid oxygen:

1. Write a balanced chemical equation.

2. Determine the limiting reagent. (DLOX =1142 kg/m3)

3. Calculate the mass of each product and remaining reactant.

4. Calculate the amount of heat released.

5. If this heat is used to warm a 500 gallon tank of water at 15 C, what is the final temperature?

6. If this heat is used to warm 40 kg of ice at 200 K, what is the final temperature? Sice=2.1 J/goC