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Chapter 3

Matter and Energy

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CHAPTER OUTLINE

Energy & Heat Temperature Scales Specific Heat Classification of Matter Physical & Chemical Properties Physical & Chemical Changes Conservation of Mass

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ENERGY & HEAT

Energy is defined as the capacity of matter to do work.

Work is defined as the result of a force acting on a distance.

There are two types of energy:

Potential (stored)

Kinetic (moving)

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ENERGY & HEAT

Energy possesses many forms (chemical, electrical, thermal, etc.), and can be converted from one form into another.

In chemistry, energy is commonly expressed as heat.

PE is converted to

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ENERGY & HEAT

Energy is conserved. The law of conservation of energy

states that energy is neither created nor destroyed.

The total amount of energy is constant. Energy can be changed from one form

to another. Energy can be transferred from one

object to another.

The SI unit of energy is the joule (J), named after the English scientist James Joule (1818–1889).

UNITS OF ENERGY

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HEAT vs. TEMPERATURE

Heat is measured in SI units of joule or the common unit of calorie.

1 cal = 4.184 J

Although the same amount of heat is added to both containers, the temperature increases more in the container with the smaller amount of water.

Heat & temperature are NOT the same thing!

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HEAT vs. TEMPERATURE

The difference between Heat and Temperature

A form of energy associated with

particles of matter

A measure of the intensity of heat or how hot or cold a

substance is

Heat is the total energy of all

particles of matter

Temperature is the average kinetic

energy of particles of matter

TEMPERATURE

Temperature is a measure of how hot or cold a substance is.

Thermometer is an instrument used for measuring temperature, and is based on thermometric properties of matter (i.e. expansion of solids or liquids).

Three scales are used for measuring temperature.

TEMPERATURE SCALES

32 - 212

Fahrenheit

Celsius

0 - 100

Kelvin

273 - 373

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TEMPERATURESCALES

To convert from one scale to another the following relationships can be used:

K = C + 273

F = (1.8 x C) + 32

C = (F - 32) ÷ 1.8 Alternately,

F = [(C + 40) x 1.8]-40 C = [(F + 40) ÷ 1.8]-40

Example 1:

The melting point of silver is 960.8 C. What is this temperature in Kelvin?

TK = TC + 273

TK = 960.8 + 273 = 1233.8 K= 1234 K

Example 2:

Pure iron melts at 1800 K. What is this temperature in Celsius?

TC = TK - 273

TC = 1800 - 273 = 1527 C

Example 3:

On a winter day, the temperature is 5 F. What is this temperature on the Celsius scale?

TC = [(5 +40) ÷ 1.8]- 40 = -15 C

Example 4:

To make ice cream, rock salt is added to crushed ice to reach temperature of -11 C. What is this temperature in Fahrenheit?

TF = [(-11 + 40) x 1.8]- 40 = 12 F

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SPECIFICHEAT

Different materials have different capacities for storing heat.

The specific heat of a substance is the amount of heat required to change the temperature of 1 g of that substance by 1C.

Units of specific heat are:

s = J / g ºC

s = cal / g ºC

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SPECIFICHEAT

Substance (cal/gC) (J/gC)Aluminum 0.214 0.897

Copper 0.0920 0.385

Iron 0.0308 0.129

Ammonia 0.488 2.04

Ethanol 0.588 2.46

Water 1.00 4.184

Specific Heat of Some Substances

Most substances have substantially

lower specific heats compared

to water

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SPECIFICHEAT

When heated, substances with low specific heat get hot faster, while substances with high specific heat get hot at a slower rate.

When cooled, substances with low specific heat cool faster, while substances with high specific heat cool at a slower rate.

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CALCULATINGHEAT

The amount of heat lost or gained by a substance is related to three quantities:

Mass of substance

Specific heat of substance

Change in its temperature

Heat = x x

Q = m x s x T

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Example 1:

How much heat is needed to raise the temperature of 200. g of water by 10.0 C. (Specific heat of water is 4.184 J/gC)

Q = m x s x T

Q = (200. g)(4.184 J/gºC)(10.0 ºC)

Q = 8370 J or 8.37 kJ

m = 200. g

s = 4.184 J/gC

T = 10.0 C

Q = ???

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Example 2:Ethanol has a specific heat of 2.46 J/gC. When 655 J are added to a sample of ethanol, its temperature rises from 18.2 C to 32.8 C. What is the mass in grams of the ethanol sample?

m = 18.2 g

Qm = =

s x TDo o

655 J

(2.46 J/g C)(14.6 C)Q = 655 J

s = 2.46 J/gC

T = 14.6 C

m = ???

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ENERGY & NUTRITION

In the laboratory, foods are burned in a calorimeter to determine their energy. A sample of food is burned in the calorimeter, and the energy released is absorbed by water surrounding the calorimeter.

The energy of the food can be calculated from the mass of the food and the temperature increase of the water.

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Example 3:

A 2.3-g sample of butter is placed in a calorimeter containing 1900 g of water at a temperature of 17 C. After the complete combustion of the butter, the water has a temperature of 28 C. What is the energy value of butter in Cal/g?

1. Calculate heat absorbed by water

Heat absorbed by water

Heat released by butter

=

2. Calculate energy value of butter

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Example 3:

Q = m x s x T

Q = (1900 g)(1.00 cal/gºC)(11 ºC)

Q = 21000 cal = 21 Cal

m = 1900 g

s = 1.00 cal/gC

T = 11 C

Q = ???

1. Calculate heat absorbed by water

2. Calculate energy value of butter

21 Cal2.3 g = 9.1 Cal/g

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ENERGY INCHEMICAL CHANGES

In all chemical changes, matter either absorbs or releases energy.

Higher energy systems are less stable than lower energy systems.

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ENERGY INCHEMICAL CHANGES

When energy is released during a chemical change, it is said to be exothermic.

When energy is gained during a chemical change, it is said to be endothermic.

Exothermic reactions heat up

Endothermic reactions cool down

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4.4

energy is given offenergy is absorbedhigher potential energy lower potential energy

EXOTHERMIC vs.ENDOTHERMIC

Which is exothermic and which is endothermic?

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CLASSIFICATIONOF MATTER

Matter is anything that has mass, and occupies space.

Matter can be classified by its physical state as solid, liquid or gas.

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SOLIDS

Solid particles have strong forces of attraction towards each other.

Solids are not very compressible.

Ice, diamond, quartz, and iron are examples of solid matter.

Solids are densely packed particles with definite shape and volume.

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LIQUIDS

Liquid particles have moderate forces of attraction towards each other and are mobile.

Liquids are slightly compressible. Water, gasoline, alcohol, and

mercury are all examples of liquid matter.

Liquids are loosely packed particles with definite volume but indefinite shape.

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GASES

Gas particles have little or no forces of attraction towards each other.

Gases are very compressible. Oxygen, helium, and carbon

dioxide are examples of gases.

Gases are very loosely packed particles with indefinite shape or volume.

 Since the atoms or molecules that compose gases are not in contact with one another, gases can be compressed.

GASES ARE COMPRESSIBLE

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SUMMARY OFPROPERTIES OF MATTER

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CLASSIFICATIONOF MATTER

MATTERAnything that has mass

PURE SUBSTANCE

Fixed composition & properties

MIXTURE

Variable composition & properties

Mixtures can be converted into pure substances by simple physical processes (e.g. filtration, evaporation)

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MIXTURES

HOMOGENEOUS

Uniform composition & properties

MIXTURE

Variable composition & properties

HETEROGENEOUS

Non-uniform composition & properties

Tea, Coke Ink

Salad dressing Cement

Also called solutions

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PURE SUBSTANCES

PURE SUBSTANCE

Fixed composition & properties

COMPOUNDS

2 or more elements chemically combined

ELEMENTS

Composed of one type of atom

Compounds can be converted into elements by chemical processes or reactions (e.g. electrolysis)

hydrogen, copper, gold

water, salt aspirin

Properties are unique compared to their

components

Smallest particle is a molecule

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PURE SUBSTANCES

separation of compound

through chemical methods

(electrolysis)

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CONCEPTCHECK

Classify each substance below as element, compound or mixture.

Element: only one type of atomElement: only one type of atomCompound: composition is fixedMixture: made of two or more types of substances

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MIXTURES

Mixtures are 2 or more substances physically combined together.

Mixtures possess properties similar to those of their components.

Mixtures can be separated easily by a physical process.

Two types of mixtures are possible:

homogeneousheterogeneous

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HETEROGENEOUSMIXTURES

Heterogeneous mixtures are non-uniform in their composition.

Examples include vegetable soup, cement and salad dressing.

Heterogeneous

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HOMOGENEOUSMIXTURES

Homogeneous mixtures are uniform in their composition.

Examples include gasoline, soda pop and salt solution.

Homogeneous

Homogeneous mixtures are called solutions.

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MIXTURES vs.COMPOUNDS

List 3 differences between compounds & mixtures.

Composition

Compounds have fixed composition while mixtures have varied composition

Properties

Compounds have unique properties while mixtures have blended properties

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MIXTURES vs.COMPOUNDS

List 3 differences between compounds & mixtures.

Make-up

Compounds are chemically combined (cannot be easily separated) while mixtures are physically combined (easily separated)

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PHYSICAL & CHEMICALPROPERTIES

The characteristics of a substance are called its properties.

Physical properties are those that describe the matter without changing its composition.

Examples are density, color, melting and boiling points, and electrical conductivity.

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PHYSICAL & CHEMICALPROPERTIES

The characteristics of a substance are called its properties.

Chemical properties are those that describe how matter behave in combination with other matter, and involve change in its composition.

Examples are flammability, corrosion, and reactivity with acids.

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Examples:

Identify each of the following properties as physical or chemical:

1. Oxygen is a gas

2. Helium is un-reactive

3. Water has high specific heat

4. Gasoline is flammable

5. Sodium is soft & shiny

Chemical

Physical

Physical

Chemical

Physical

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PHYSICALCHANGES

Changes in physical properties of matter that do not involve change in its composition are called physical changes.

Examples are melting, evaporation and other phase changes.

Physical changes are easily reversible.

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CHEMICALCHANGES

A change that alters the chemical composition of matter, and forms new substance is called a chemical change.

Examples are burning, rusting, and reaction with acids.

Chemical changes are not easily reversible, and are commonly called chemical reactions.

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Examples:

Identify each of the following changes as physical or chemical:

1. Cooking food

2. Mixing sugar in tea

3. Carving wood

4. Burning gas

5. Food molding

Chemical

Physical

Physical

Chemical

Chemical

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CONSERVATIONOF MASS

Similar to the law of conservation of energy, the law of the conservation of mass states that matter is neither created nor destroyed.

The total mass of substances does not change during a chemical reaction.

Mass of Reactants

= Mass of Products

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266 g product

mass products

266 g reactant →

mass reactants =

CONSERVATIONOF MASS

The number of substances and their properties may change, but the total amount of matter remains constant.

Suppose that we burn 58 g of butane in a lighter. It will react with 208 g of oxygen to form ??? g of carbon dioxide and 90 g of water.

Burning: a Chemical Change

The butane molecules react with oxygen molecules in air to form new molecules, carbon dioxide and water.

This is a chemical change.

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THE END