Energy and States of Matter Chapter 11 23-October

Suggested HW

11.11, 11.17, 11.21, 11.23, 11.29, 11.31, 11.33, 11.39, 11.43, 11.45, 11.51, 11.57, 11.59,

11.61, 11.63

Lecture 9

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Physical Properties of Phase

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Gases, Liquids, and Solids

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Gases • Very little interaction

between molecules • Lots of motion = lots of

kinetic energy • Low density • Large compressibility • Will expand to fill a

container

Liquids • Potential and kinetic

energies are roughly equal

• High density • Little compressibility • Definite volume and

shape

Solids • Maximum interaction

between molecules • Dominated by potential

energy • High density • Very little compressibility • Definite volume and

shape

Kinetic Theory of Matter

4

Matter is composed of tiny particles that are always in motion

Kinetic Energy – energy that is possessed by an object because of particle motion Increases with temperature

Transferrable upon contact with another particle

Potential Energy – energy that is stored because of positioning Intermolecular Forces are important examples of potential energy

Objects at high temperatures have high kinetic energy

'E = Efinal - Einitial

Changes of State

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Endothermic Process Any physical or chemical process that REQUIRES

energy

Heat (energy) is necessary to melt ice and

boil water

Changes in state are characterized by changes

in energy ('E)

'E = Efinal - Einitial

Egas > Eliquid > Esolid

Changes of State

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Exothermic Process Any physical or chemical process that PRODUCES

energy

Heat (energy) is produced when

condensation or freezing occurs

'E = Efinal - Einitial

Energy and Changes of State

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Liquid Solid + Heat Exothermic

Solid + Heat Endothermic

Liquid

Gas Liquid + Heat Exothermic

Liquid + Heat Endothermic

Gas

If a chemical or physical reaction is exothermic, the reverse reaction is endothermic

Specific Heat and Heat Capacity

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The specific heat is the amount of heat needed to raise the temperature of one gram given substance by 1˚ C.

Units Æ J/(g ˚C)

Heat Capacity is about the same but it is for a specific amount of the substance.

Units Æ J/ ˚C

Molar Heat Capacity Æ Heat Capacity per mol of substance

Units Æ J/ mol ˚C

Solids and Liquids have different Specific Heats and Heat Capacities

J = Joules Æ unit of energy

1 calorie = 4.184 J

Heat change = mass * specific heat * change in temperature q = mC'T

Sample Problem

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Solid aluminum has a specific heat of 0.908 J/(g ˚C).  

1. Calculate the heat capacity of 16 g of aluminum.

2. Calculate the Molar Heat Capacity of aluminum

Sample Problem

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1. How much heat is required to raise the temperature of 16 g of aluminum from 20 ˚C 

to 82 ˚C?

2. How much heat is needed to change the temperature of 16 g ofwater from 20 ˚C to 82 ˚C (specific heat = 4.18 J/(g ˚C))?

Sample Problem

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If a 16.0 g chunk of Al (150 ˚C) is dropped in 100 g of water at 20˚C, how much will the temperature of the water change?

Sample Problem

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If a 750 g chunk of gold (C = 0.13 J/(g ˚C) at 234 ˚C is dropped into 10 kg of water at 4˚C, what will the final temperature of the solution be?

How much energy was transferred?

Energy and Phase Changes

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Solid + Heat Liquid

Fusion Æ Heat absorbed converting solid to liquid

Energy and Phase Changes

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Liquid + Heat Gas

Vaporization Æ Heat absorbed converting liquid to gas

Changes of State and Energy

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Changes of State and Energy

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Q Æ heat absorbed or consumed C Æ specific heat in specified physical state

M Æ mass

'T Æ change in temperature 'Hfusion Æ heat of fusion(melting)

'Hvap Æ heat of vaporization (boiling)

Solid

Liquid

Gas

Changes of State and Energy

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Calculate how much energy it takes to convert 50 g of ice @ -45 ˚C to steam 

@ 170 ˚C.

Cice = 2.09 J/(g ˚C) 'Hfus = 334 J/g Cwater = 4.18 J/(g ˚C) 'Hvap = 2260 J/g Csteam = 2.03 J/(g ˚C)