Section 10.1: The Nature of Energy Energy - The ability to do work or produce heat.

Two types:

1. Potential energy – energy due to position or composition (i.e. energy stored in chemicals). PE = mgh Examples include water behind a dam or

energy stored in chemical bonds.

2. Kinetic energy – energy due to motion. KE = ½mv2

Examples include water flowing downhill over a turbine or gasoline exploding.

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The Nature of Energy

Law of Conservation of Energy – Energy can be converted from one form to another but can neither be created nor destroyed.

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Section 10.2: Temperature and HeatTemperature - Measure of the random motions (average kinetic energy) of the components of a substance.

Heat – Flow of energy due to a temperature difference.

Remember: All particles are in motion. As the temperature increases, the thermal energy or vigor of their motion increases.

Heat is the transfer of this thermal energy from one object to another.

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Figure 10.2: Equal masses of hot and cold water.

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Figure 10.3: H2O molecules in hot and cold water.

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Figure 10.4: H2O molecules in same temperature water.

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CCCTT

Tcold

initial

hot

initial

final50

2

1090

2

Change in temperature (hot) = Tf – Ti =Thot = 50. C – 90. C = -40. C

Change in temperature (cold) = Tf – Ti =Tcold = 50. C – 10. C = 40. C

T =Tf – Ti

In science we use the Greek letter Delta ( ) to mean “change in.” It is ALWAYS the final value minus the initial value.

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Section 10.3: Exothermic and Endothermic ProcessesThe universe is divided into two halves: the system and the surroundings.

The system is the part we are concerned with. We define its boundaries.

The surroundings are the rest.

Every reaction has an energy change associated with it.

Energy is stored in bonds between atoms.

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Exothermic reactions release energy to the surroundings

Energy exits the system.Endothermic reactions absorb energy from the surroundings.

Energy enters the system.

Examples:A match feels hot because energy is exiting the system (the match) and entering the surroundings (your skin, for example).

An ice cube feels cold for the same reasons, but energy flows in the opposite directions.

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ThermodynamicsThermodynamic quantities always consist of two parts:

A number giving the magnitude of the change.

A sign indicating the direction of the flow (from the system’s point of view).

For an endothermic process, E >0 (the system is gaining energy).

For an exothermic process, E <0 (the system is losing energy).

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Section 10.5: Measuring Energy ChangesDifferent materials respond differently to being heated.

Units are needed to explore differences:

calorie (cal) – The amount of energy (heat) required to raise the temperature of one gram of water by one Celsius degree.

joule (J) = SI unit; 1 calorie = 4.184 joules

In chemistry calories are written with a small “c”; calories on food labels are 1000 of these (or one kilocalorie). We distinguish these with a capital “C”

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Measuring Energy Changes Practice Conversions:

1) Express 34.8 calories of energy in units of joules.

2) Express 47.3 J of energy in units of cal.

Specific heat capacity (s) – The amount of energy required to change the temperature of one gram of a substance by one Celsius degree.

Each substance has its own unique specific heat (units are J/g C).

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Measuring Energy Changes

Substance Specific Heat (J/g C)

Water (l) 4.184

Water (s) 2.03

Water (g) 2.0

Aluminum (s) 0.89

Iron (s) 0.45

Mercury (l) 0.14

The Specific Heat Capacities of Some Common Substances

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Measuring Energy ChangesWe need the following equation to calculate energy:

Q = s x m x TWhere:

Q = energy (heat) required

s = specific heat capacity

m = mass of the sample (in grams)

T = Change in temperature (in C)

This equation always applies when a substance is being heated (or cooled) and no change of state occurs.

Measuring Energy Changes: Practice Problems

REMEMBER T = Change in temperature (in C) =

Tfinal - Tinitial1) A 2.6 –g sample of a metal requires 15.6 J of

Energy to change its temperature from 21 C to 34 C. Use specific heat values from your notes to identify this metal.

2) A sample of gold requires 3.1 J of energy to change its temperature from 19 C to 27 C. What is the mass of this sample of gold?

3) Calculate the amount of energy required (in calories) to heat 145-g of water from 22.3 C to 75.0 C.

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Practice Problems

1. If it takes 526J of energy to warm 7.40g of water by 17oC, how much energy would be needed to warm 7.4 g of water by 55oC?

2. If 72.4 kJ of heat is applied to a 952 g block of metal, the temperature increases by 10.7oC. Calculate the specific heat capacity of the metal in J/goC.

If T is negative, reaction is exothermic. Energy is leaving the system as heat.

If T is positive, reaction is endothermic. Energy is entering the system as heat.

Chemists like to know how much energy is released or absorbed. We use a special energy function called ENTHALPY (H).

Section 10.6: Thermochemistry (Enthalpy) Enthalpy – (at constant pressure) the change in enthalpy

equals the energy flow as heat.

ΔHp = heat (change in enthalpy = heat)

The “p” refers to constant pressure

For endothermic reactions, ΔH>0.

For exothermic reactions, ΔH<0. Calorimeter – A device used to determine the heat

associated with a chemical reaction.

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Figure 10.6: A coffee-cup calorimeter.

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Thermochemistry (Enthalpy) Sample Problem:

When one mole of methane CH4 is burned at constant pressure, 890 kJ of energy is released as heat. Calculate ΔH for a process in which a 5.8-g sample of methane is burned at constant pressure.

qp =(Qp)= ΔH = -890 kJ/mol CH4 (negative because heat is released/exothermic)

5.8-g CH4 = 0.36 mol CH4

(0.36 mol)(-890kJ/mol) = -320 kJ

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Thermochemistry (Enthalpy) Practice Problems:

1) When one mole of sulfur dioxide reacts with excess oxygen to form sulfur trioxide at constant pressure, 198.2 kJ of energy is released as heat. Calculate ΔH for a process in which a 12.8-g sample of sulfur dioxide reacts with excess oxygen at constant pressure.

2) When one mole of hydrogen gas reacts with excess oxygen to form water at constant pressure, 241.8 kJ of energy is released as heat. Calculate ΔH for a process in which a 15.0-g sample of hydrogen gas reacts with excess oxygen at constant pressure.

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Practice Problem

The reaction that occurs in the heat packs used to treat sports injuries is

4Fe(s) + 3O2(g) 2Fe2O3 (s) ΔH=-1652kJ

1. Exothermic or Endothermic reaction?

2. How much heat is released when 1.00g of Fe(s) is reacted with excess O2(g)?

We are not covering Hess’s Law 10.7!!!!

Figure 10.7: Energy sources used in the United States.

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Figure 10.8: The earth’s atmosphere.

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Figure 10.9: The atmospheric CO2 concentration over the past 1000 years.

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