Chapter 10

Energy

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Energy and Energy Changes

• Energy: ability to do work or produce heat– Chemical, mechanical, thermal, electrical, radiant,

sound, nuclear– Potential and kinetic

• Energy may affect matter.– e.g. Raise its temperature, eventually causing a state

change, or cause a chemical change such as decomposition

• All physical changes and chemical changes involve energy changes.

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Energy and Energy Changes

• Potential Energy: energy due to composition or position

• Kinetic Energy: energy

due to motion– - ½ mv2

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Energy and Energy Changes (cont.)

• Law of Conservation of Energy: energy can be converted from one form to another, but cannot be created or destroyed

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Work and Energy

• Work: force acting over a distance– w = f • d– Work done on a system will increase the

energy of the system, whereas work done by the system will decrease the energy of the system

• State function: a property that changes independent of pathway

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Temperature and Heat

• Heat: a flow of energy due to a temperature difference

• Temperature: a measure of the random motions of the components of a substance

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Temperature and Heat (cont.)

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Exothermic vs. Endothermic

• System: that part of the universe that we wish to study

• Surroundings: everything else in the universe• Exothermic process: a process that results in the

evolution of heat- Example: when a match is struck, it is an

exothermic process because energy is produced as heat.

• Endothermic process: absorbs heat- Example: melting ice to form liquid water is an

endothermic process because the ice absorbs heat in order to melt

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Exothermic Process

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Thermodynamics

• The Law of Conservation of Energy is also known as The First Law of Thermodynamics. It can be stated as “the energy of the universe is constant.”

• Internal Energy (E) = kinetic energy + potential energy

• ΔE = q + w = change in internal energy q = heat absorbed by the system w = work done on the system

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Units of Energy

• One calorie = amount of energy needed to raise the temperature of one gram of water by 1°C– kcal = energy needed to raise the temperature of

1000 g of water 1°C

• joule – 4.184 J = 1 cal

• In nutrition, calories are capitalized.– 1 Cal = 1 kcal

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Example - Converting Calories to Joules

251Jcal 1

J 4.184 60.1cal

joules 4.184 cal 1

Convert 60.1 cal to joules.

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Energy & Temperature of Matter

• The amount the temperature of an object increases depends on the amount of heat added (q).– If you double the added heat energy the temperature

will increase twice as much.

• The amount the temperature of an object increases when heat is added depends on its mass– If you double the mass it will take twice as much heat

energy to raise the temperature the same amount.

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Specific Heat Capacity

• Specific heat (s): the amount of energy required to raise the temperature of one gram of a substance by one degree Celsius

C g

J 4.184 is water ofheat specific the,definitionBy

Amount of Heat = Specific Heat x Mass x Temperature ChangeQ = s x m x T

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Specific Heat Capacity

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Calculate the amount of heat energy (in joules) needed to raise the temperature of 7.40 g of water from 29.0°C to 46.0°C.

Example #1:

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Mass = 7.40 g

Temperature change = 46.0°C – 29.0°C = 17.0°C

Q = s • m • T

Example #1 (cont.)

J 526 C17.07.40gC g

J 4.184 Heat

Specific heat of water = 4.184 C g

J

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A 1.6 g sample of metal that appears to be gold requires 5.8 J to raise the temperature from 23°C to 41°C. Is the metal pure gold?

Example #2

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Example #2

C g

J0.20

C18 x g 1.6

J 5.8 s

C18 C23 - C41 TT m

Q s

Tms Q

Table 10.1 lists the specific heat of gold as 0.13Therefore the metal cannot be pure gold.

C g

J

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Enthalpy

• Change in enthalpy (ΔHp = qp): the amount of heat exchanged when heat exchange occurs under conditions of constant pressure

• Enthalpy is a state function

• ΔH is independent of the path taken

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• Hess’s Law: in going from a set of reactants to a set of products, the change in enthalpy is the same whether the reaction takes place in one step or in a series of steps.

Hess’s Law

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• Overall reaction: N2 + 2O2 2NO2 ΔH = 68 kJ• This reaction can be carried out in 2 steps: N2 + O2 2NO ΔH = 180 kJ 2NO + O2 2NO2 ΔH = -112 kJ

-------------------------------------------------------- N2 + 2O2 2NO2 ΔH = 68 Kj

Note: the sum of the two reactions gives the overall reaction and the same is true for the sum of the enthalpy change values.

Hess’s Law (cont.)

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Calorimetry

• The amount of heat flow transferred during a reaction is determined from temperature measurements made in a calorimeter.

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Calorimetry (cont.)

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Energy Quality & Quantity

• While the total amount or quantity of energy in the universe is constant (1st Law) the quality of energy is degraded as it is used.

Burning of petroleum:

High grade concentrated energy Low grade energy (heat)

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Fuels

• Petroleum– A fossil fuel composed mainly of hydrocarbons

• Natural gas– Consists largely of methane – Also contains ethane, propane, and butane

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Fuels (cont.)

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Fuels (cont.)

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Fuels (cont.)

• Coal– Matures geologically through stages

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Global Warming

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Global Warming (cont.)

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Energy Use and Sources

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Energy as a Driving Force

• Most processes that occur spontaneously involve an “energy spread.”– Heat flows from high to low temperature and

“spreads”

…or a “matter spread”– Salt dissolves or “spreads” in water

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Entropy

• Entropy (S) is a measure of disorder or randomness.– As a system becomes more disordered, ΔS >0

• Second Law of Thermodynamics: the entropy of the universe is always increasing.