Introductory Chemistry , 3 rd Edition Nivaldo Tro

Basic Principles of Chemistry OnlineSoutheast Missouri State University

Cape Girardeau, MO

Introductory Chemistry, 3rd EditionNivaldo Tro

Chapter 3Matter and Energy

2009, Prentice Hall

Tro's "Introductory Chemistry", Chapter 3

2

In Your Room• Everything you can see,

touch, smell or taste in your room is made of matter.

• Chemists study the differences in matter and how that relates to the structure of matter.


3

What Is Matter?• Matter is defined as

anything that occupies space and has mass.

• Even though it appears to be smooth and continuous, matter is actually composed of a lot of tiny little pieces we call atoms and molecules.


4

Atoms and Molecules

• Atoms are the tiny particles that make up all matter.

• In most substances, the atoms are joined together in units called molecules.The atoms are joined in

specific geometric arrangements.

5

Structure Determines Properties• The properties of matter are determined by the atoms and molecules

that compose it.

1. Composed of one carbon atom and two oxygen atoms.

2. Colorless, odorless gas.3. Incombustible.4. Does not bind to hemoglobin.

Carbon Dioxide1. Composed of one carbon

atom and one oxygen atom.2. Colorless, odorless gas.3. Burns with a blue flame.4. Binds to hemoglobin.

Carbon Monoxide


6

Classifying Matterby Physical State

• Matter can be classified as solid, liquid, or gas based on what properties it exhibits.

State Shape Volume Compress Flow

Solid Fixed Fixed No No

Liquid Indefinite Fixed No Yes

Gas Indefinite Indefinite Yes Yes

•Fixed = Property doesn’t change when placed in a container. •Indefinite = Takes the property of the container.

7

Structure Determines Properties• The atoms or molecules have different

structures in solids, liquids, and gases − leading to different properties.


8

Solids• The particles in a solid are packed

close together and are fixed in position.Although they may vibrate.

• The close packing of the particles results in solids being incompressible.

• The inability of the particles to move around results in solids retaining their shape and volume when placed in a new container and prevents the particles from flowing.


9

Solids, Continued• Some solids have their particles

arranged in an orderly geometric pattern—we call these crystalline solids.Salt and diamonds.

• Other solids have particles that do not show a regular geometric pattern over a long range—we call these amorphous solids.Plastic and glass.


10

Liquids• The particles in a liquid are closely packed,

but they have some ability to move around. • The close packing results in liquids being

incompressible.• The ability of the particles to move allows

liquids to take the shape of their container and to flow. However, they don’t have enough freedom to escape and expand to fill the container.


11

Gases

• In the gas state, the particles have complete freedom from each other.

• The particles are constantly flying around, bumping into each other and the container.

• In the gas state, there is a lot of empty space between the particles.On average.


12

Gases, Continued• Because there is a lot of empty

space, the particles can be squeezed closer together. Therefore, gases are compressible.

• Because the particles are not held in close contact and are moving freely, gases expand to fill and take the shape of their container, and will flow.


13

Classification of Matter by Appearance

• Homogeneous = Matter that is uniform throughout. Appears to be one thing. Every piece of a sample has identical properties, though another sample

with the same components may have different properties. Solutions (homogeneous mixtures) and pure substances.

• Heterogeneous = Matter that is non-uniform throughout . Contains regions with different properties than other regions.


14

Practice—Classify the Following as Homogeneous or Heterogeneous

• Table sugar.

• A mixture of table sugar and black pepper.

• A mixture of sugar dissolved in water.

• Oil and vinegar salad dressing.


15

Practice—Classify the Following as Homogeneous or Heterogeneous,

Continued• Table sugar = homogeneous

• A mixture of table sugar and black pepper = heterogeneous

• A mixture of sugar dissolved in water = homogeneous

• Oil and vinegar salad dressing = heterogeneous


16

Classifying Matterby Composition

• Matter that is composed of only one kind of atom or molecule is called a pure substance.

• Matter that is composed of different kinds of atoms or molecules is called a mixture.

• Because pure substances always have only one kind of piece, all samples show the same properties.

• However, because mixtures have variable composition, different samples will show different properties.


17

Copper—A Pure Substance

• Color is brownish red.

• Shiny, malleable, and ductile.

• Excellent conductor of heat and electricity.

• Melting point = 1084.62 °C

• Density = 8.96 g/cm3 at 20 °C

Brass—A MixtureType Color % Cu % Zn Density

g/cm3

MP

°C

Tensile

Strength

psi

Uses

Gilding reddish 95 5 8.86 1066 50K pre-83 pennies,munitions, and

plaques

Commercial bronze 90 10 8.80 1043 61K door knobs andgrillwork

Jewelry bronze 87.5 12.5 8.78 1035 66K costume jewelry

Red golden 85 15 8.75 1027 70K electrical sockets,fasteners, and

eyelets

Low deep yellow

80 20 8.67 999 74K musical instruments and clock dials

Cartridge yellow 70 30 8.47 954 76K car radiator cores

Common yellow 67 33 8.42 940 70K lamp fixtures andbead chain

Muntz metal yellow 60 40 8.39 904 70K nuts and bolts


19

Classification of Matter

• Pure Substance = All samples are made of the same pieces in the same percentages. Salt

• Mixtures = Different samples may have the same pieces in different percentages. Salt water

Pure SubstanceConstant Composition

Homogeneous

MixtureVariable Composition

Matter


20

Pure Substances vs. MixturesPure Substances

1. All samples have the same physical and chemical properties.

2. Constant composition = All samples have the same pieces in the same percentages.

3. Homogeneous.4. Separate into components

based on chemical properties.

5. Temperature stays constant while melting or boiling.

Mixtures1. Different samples may show

different properties.2. Variable composition =

Samples made with the same pure substances may have different percentages.

3. Homogeneous or heterogeneous.

4. Separate into components based on physical properties.

5. Temperature usually changes while melting or boiling because composition changes.


21

Practice—Classify the Following as Pure Substances or Mixtures

• A homogeneous liquid whose temperature stays constant while boiling.

• Granite—a rock with several visible minerals in it.• A red solid that turns blue when heated and

releases water that is always 30% of the solid’s mass.

• A gas that when cooled and compressed, a liquid condenses out but some gas remains.


22

Practice—Classify the Following as Pure Substances or Mixtures,

Continued• A homogeneous liquid whose temperature stays

constant while boiling = pure substance.• Granite—a rock with several visible minerals in it

= mixture.• A red solid that turns blue when heated and

releases water that is always 30% of the solid’s mass = pure substance.

• A gas that when cooled and compressed, a liquid condenses out but some gas remains = mixture.


23

Classification of Pure Substances• Substances that cannot be broken down into simpler

substances by chemical reactions are called elements.Basic building blocks of matter.Composed of single type of atom.

Although those atoms may or may not be combined into molecules.

• Substances that can be decomposed are called compounds.Chemical combinations of elements.

Although properties of the compound are unrelated to the properties of the elements in it!

Composed of molecules that contain two or more different kinds of atoms.

All molecules of a compound are identical, so all samples of a compound behave the same way.

• Most natural pure substances are compounds.


24

Atoms and Molecules• Atoms

Are submicroscopic particles that are the unit pieces of elements.

Are the fundamental building blocks of all matter.

• MoleculesAre submicroscopic particles that are the

unit pieces of compounds. Two or more atoms attached together.

Attachments are called bonds.Attachments come in different strengths.

Molecules come in different shapes and patterns.


25

Classification of Pure Substances

1. Made of one type of atom. (Some elements are found as multi-atom molecules in nature.)

2. Combine together to make compounds.

1. Made of one type of molecule, or array of ions.

2. Molecules contain 2 or more different kinds of atoms.

Elements Compounds


26

Practice—Classify the Following as Elements or Compounds

• Chlorine, Cl2

• Table sugar, C12H22O11

• A red solid that turns blue when heated and releases water that is always 30% of the solid’s mass.

• A brown-red liquid that, when energy is applied to it in any form, causes only physical changes in the material, not chemical.


27

Practice—Classify the Following as Elements or Compounds, Continued

• Chlorine, Cl2 = element.

• Table sugar, C12H22O11 = compound.

• A red solid that turns blue when heated and releases water that is always 30% of the solid’s mass = compound.

• A brown-red liquid that, when energy is applied to it in any form, causes only physical changes in the material, not chemical = element.


28

Classification of Mixtures

• Mixtures are generally classified based on their uniformity.

• Mixtures that are uniform throughout are called homogeneous.Also known as solutions.Mixing is on the molecular level.

• Mixtures that have regions with different characteristics are called heterogeneous.


29

Classification of Mixtures, Continued

1. Made of multiple substances, but appears to be one substance.

2. All portions of a sample have the same composition and properties.

1. Made of multiple substances, whose presence can be seen.

2. Portions of a sample have different composition and properties.

Heterogeneous Homogeneous

30

Classifying Matter


31

Properties Distinguish Matter

• Each sample of matter is distinguished by its characteristics.

• The characteristics of a substance are called its properties.

• Some properties of matter can be observed directly.

• Other properties of matter are observed when it changes its composition.


32

Properties of Matter• Physical Properties are the characteristics of matter

that can be changed without changing its composition.Characteristics that are directly observable.

• Chemical Properties are the characteristics that determine how the composition of matter changes as a result of contact with other matter or the influence of energy. Characteristics that describe the behavior of matter.


33

Some Physical PropertiesMass Volume Density

Solid Liquid Gas

Melting point Boiling point Volatility

Taste Odor Color

Texture Shape Solubility

Electrical conductance

Thermal conductance

Magnetism

Malleability Ductility Specific heat capacity


34

Some Physical Properties of Iron• Iron is a silvery solid at room temperature with a metallic taste

and smooth texture.• Iron melts at 1538 °C and boils at 4428 °C.• Iron’s density is 7.87 g/cm3.• Iron can be magnetized.• Iron conducts electricity, but not as well as most other common

metals.• Iron’s ductility and thermal conductivity are about average for a

metal.• It requires 0.45 J of heat energy to raise the temperature of one

gram of iron by 1°C.


35

Some Chemical Properties

Acidity Basicity (aka alkalinity)

Causticity Corrosiveness

Reactivity Stability

Inertness Explosiveness

(In)Flammability Combustibility

Oxidizing ability Reducing ability


36

Some Chemical Properties of Iron

• Iron is easily oxidized in moist air to form rust.

• When iron is added to hydrochloric acid, it produces a solution of ferric chloride and hydrogen gas.

• Iron is more reactive than silver, but less reactive than magnesium.


37

Practice—Decide Whether Each of the Observations About Table Salt Is a Physical or Chemical Property

• Salt is a white, granular solid.• Salt melts at 801 °C.• Salt is stable at room temperature, it does not decompose.• 36 g of salt will dissolve in 100 g of water. • Salt solutions and molten salt conduct electricity.• When a clear, colorless solution of silver nitrate is added

to a salt solution, a white solid forms.• When electricity is passed through molten salt, a gray

metal forms at one terminal and a yellow-green gas at the other.


38

Practice − Decide Whether Each of the Observations About Table Salt Is a Physical or Chemical Property

• Salt is a white, granular solid = physical.

• Salt melts at 801 °C = physical.

• Salt is stable at room temperature, it does not decompose = chemical.

• 36 g of salt will dissolve in 100 g of water = physical.

• Salt solutions and molten salt conduct electricity = physical.

• When a clear, colorless solution of silver nitrate is added to a salt solution, a white solid forms = chemical.

• When electricity is passed through molten salt, a gray metal forms at one terminal and a yellow-green gas at the other = chemical.


39

Changes in Matter

• Changes that alter the state or appearance of the matter without altering the composition are called physical changes.

• Changes that alter the composition of the matter are called chemical changes.

During the chemical change, the atoms that are present rearrange into new molecules, but all of the original atoms are still present.


40

Changes in Matter, Continued• Physical Changes—Changes in

the properties of matter that do not effect its composition.Heating water.

Raises its temperature, but it is still water.

Evaporating butane from a lighter.Dissolving sugar in water.

Even though the sugar seems to disappear, it can easily be separated back into sugar and water by evaporation.


41

Changes in Matter, Continued• Chemical Changes involve a change

in the properties of matter that change its composition.A chemical reaction.Rusting is iron combining with oxygen to

make iron(III) oxide.Burning results in butane from a lighter to

be changed into carbon dioxide and water.

Silver combines with sulfur in the air to make tarnish.


42

Is it a Physical or Chemical Change?• A physical change results in a different form of

the same substance.The kinds of molecules don’t change.

• A chemical change results in one or more completely new substances.Also called chemical reactions.

The new substances have different molecules than the original substances.

You will observe different physical properties because the new substances have their own physical properties.


43

Phase Changes ArePhysical Changes

• Boiling = liquid to gas.• Melting = solid to liquid.• Subliming = solid to gas.• Freezing = liquid to solid.• Condensing = gas to liquid.• Deposition = gas to solid.• State changes require heating or cooling the substance.

Evaporation is not a simple phase change, it is a solution process.


44

Practice—Classify Each Change as Physical or Chemical

• Evaporation of rubbing alcohol.

• Sugar turning black when heated.

• An egg splitting open and spilling out.

• Sugar fermenting.

• Bubbles escaping from soda.

• Bubbles that form when hydrogen peroxide is mixed with blood.


45

Practice—Classify Each Change as Physical or Chemical, Continued

• Evaporation of rubbing alcohol = physical.• Sugar turning black when heated = chemical.• An egg splitting open and spilling out =

physical.• Sugar fermenting = chemical.• Bubbles escaping from soda = physical. • Bubbles that form when hydrogen peroxide is

mixed with blood = chemical.


46

Separation of Mixtures• Separate mixtures based on different

physical properties of the components.Physical change.

Centrifugation and

decanting

Density

EvaporationVolatility

ChromatographyAdherence to a surface

FiltrationState of matter (solid/liquid/gas)

DistillationBoiling point

TechniqueDifferent Physical Property


47

Distillation


48

Filtration


49

Law of Conservation of Mass

• Antoine Lavoisier

• “Matter is neither created nor destroyed in a chemical reaction.”

• The total amount of matter present before a chemical reaction is always the same as the total amount after.

• The total mass of all the reactants is equal to the total mass of all the products.


50

Conservation of Mass• Total amount of matter remains constant in a

chemical reaction.• 58 grams of butane burns in 208 grams of oxygen to

form 176 grams of carbon dioxide and 90 grams of water.

butane + oxygen carbon dioxide + water 58 grams + 208 grams 176 grams + 90 grams

266 grams = 266 grams


51

Practice—A Student Places Table Sugar and Sulfuric Acid into a Beaker and Gets a Total

Mass of 144.0 g. Shortly, a Reaction Starts that Produces a “Snake” of Carbon Extending from the Beaker and Steam Is Seen Escaping. If the Carbon Snake and Beaker at the End Have a

Total Mass of 129.6 g, How Much Steam Was Produced?


52

Practice—A Student Places Table Sugar and Sulfuric Acid into a Beaker and Gets a Total

Mass of 144.0 g. Shortly, a Reaction Starts that Produces a “Snake” of Carbon Extending from the Beaker and Steam Is Seen Escaping. If the Carbon Snake and Beaker at the End Have a

Total Mass of 129.6 g, How Much Steam Was Produced?

• Total of reactants and beaker = 144.0 g.• Conservation of mass says total of products and

beaker must be 144.0 g.• Mass of steam = 144.0 g − 129.6 g = 14.4 g.


53

Energy• There are things that do not have mass and

volume.• These things fall into a category we call energy.• Energy is anything that has the capacity to do

work.• Although chemistry is the study of matter, matter

is effected by energy.It can cause physical and/or chemical changes in

matter.


54

Law of Conservation of Energy

• “Energy can neither be created nor destroyed.”

• The total amount of energy in the universe is constant. There is no process that can increase or decrease that amount.

• However, we can transfer energy from one place in the universe to another, and we can change its form.


55

Matter Possesses Energy• When a piece of matter possesses

energy, it can give some or all of it to another object.It can do work on the other object.

• All chemical and physical changes result in the matter changing energy.


56

Kinds of EnergyKinetic and Potential

• Potential energy is energy that is stored. Water flows because gravity pulls it

downstream. However, the dam won’t allow it to

move, so it has to store that energy.• Kinetic energy is energy of motion,

or energy that is being transferred from one object to another. When the water flows over the dam,

some of its potential energy is converted to kinetic energy of motion.


57

Some Forms of Energy• Electrical

Kinetic energy associated with the flow of electrical charge.

• Heat or Thermal EnergyKinetic energy associated with molecular motion.

• Light or Radiant EnergyKinetic energy associated with energy transitions in an atom.

• NuclearPotential energy in the nucleus of atoms.

• ChemicalPotential energy in the attachment of atoms or because of their

position.


58

Converting Forms of Energy• When water flows over the dam, some of its

potential energy is converted to kinetic energy.Some of the energy is stored in the water because it is

at a higher elevation than the surroundings. • The movement of the water is kinetic energy.• Along the way, some of that energy can be used to

push a turbine to generate electricity.Electricity is one form of kinetic energy.

• The electricity can then be used in your home. For example, you can use it to heat cake batter you mixed, causing it to change chemically and storing some of the energy in the new molecules that are made.


59

Using Energy

• We use energy to accomplish all kinds of processes, but according to the Law of Conservation of Energy we don’t really use it up!

• When we use energy we are changing it from one form to another.For example, converting the chemical energy

in gasoline into mechanical energy to make your car move.


60

“Losing” Energy

• If a process was 100% efficient, we could theoretically get all the energy transformed into a useful form.

• Unfortunately we cannot get a 100% efficient process.

• The energy “lost” in the process is energy transformed into a form we cannot use.


61

There’s No Such Thing as a Free Ride

• When you drive your car, some of the chemical potential energy stored in the gasoline is released.

• Most of the energy released in the combustion of gasoline is transformed into sound or heat energy that adds energy to the air rather than move your car down the road.


62

Units of Energy

• Calorie (cal) is the amount of energy needed to raise one gram of water by 1 °C.kcal = energy needed to raise 1000 g of water 1 °C.food calories = kcals.

Energy Conversion Factors

1 calorie (cal) = 4.184 joules (J)

1 Calorie (Cal) = 1000 calories (cal)

1 kilowatt-hour (kWh) = 3.60 x 106 joules (J)


63

Energy Use

Unit

Energy Required to Raise Temperature of 1 g of Water by 1°C

Energy Required to Light 100-W Bulb for 1 Hour

Energy Used by Average U.S. Citizen in 1 Day

joule (J) 4.18 3.6 x 105 9.0 x 108

calorie (cal) 1.00 8.60 x 104 2.2 x 108

Calorie (Cal) 1.00 x 10-3 86.0 2.2 x 105

kWh 1.1 x 10-6 0.100 2.50 x 102

Example 3.5—Convert 225 Cal to Joules

Units and magnitude are correct.

Check:7. Check.

225 Cal = 9.41 x 105 JRound:6. Significant figures and round.

Solution:5. Follow the solution map to Solve the problem.

Solution Map:

4. Write a Solution Map.

1 Cal = 1000 cal1 cal = 4.184 J

Conversion Factors:

3. Write down the appropriate Conversion Factors.

? JFind:2. Write down the quantity you want to Find and unit.

225 CalGiven:1. Write down the Given quantity and its unit.

Cal

cal 1

J .1844

J 1041.9cal 1

J .1844

Cal 1

cal 1000Cal 252 5

cal J

Cal 1

cal 1000

3 sig figs

3 significant figures


72

Chemical Potential Energy• The amount of energy stored in a material is its

chemical potential energy.• The stored energy arises mainly from the

attachments between atoms in the molecules and the attractive forces between molecules.

• When materials undergo a physical change, the attractions between molecules change as their position changes, resulting in a change in the amount of chemical potential energy.

• When materials undergo a chemical change, the structures of the molecules change, resulting in a change in the amount of chemical potential energy.


73

Energy Changes andChemical Reactions

• Chemical reactions happen most readily when energy is released during the reaction.

• Molecules with lots of chemical potential energy are less stable than ones with less chemical potential energy.

• Energy will be released when the reactants have more chemical potential energy than the products.


74

Exothermic Processes

• When a change results in the release of energy it is called an exothermic process.

• An exothermic chemical reaction occurs when the reactants have more chemical potential energy than the products.

• The excess energy is released into the surrounding materials, adding energy to them.Often the surrounding materials get hotter from the

energy released by the reaction.


75

An Exothermic Reaction

Pot

enti

al e

nerg

y

Reactants

Products

Surroundings

reaction

Amount of energy released


76

Endothermic Processes

• When a change requires the absorption of energy it is called an endothermic process.

• An endothermic chemical reaction occurs when the products have more chemical potential energy than the reactants.

• The required energy is absorbed from the surrounding materials, taking energy from them.Often the surrounding materials get colder due to the

energy being removed by the reaction.


77

An Endothermic Reaction

Pot

enti

al e

nerg

y

Products

Reactants

Surroundings

reaction

Amount of energy absorbed


78

Temperature Scales

• Fahrenheit scale, °F.Used in the U.S.

• Celsius scale, °C.Used in all other countries.A Celsius degree is 1.8

times larger than a Fahrenheit degree.

• Kelvin scale, K.Absolute scale.

Temperature Scales

Celsius Kelvin Fahrenheit-273°C-269°C

-183°C

-38.9°C

0°C

100°C

0 K4 K

90 K

234.1 K

273 K

373 K

-459 °F-452°F

-297°F

-38°F

32°F

212°F

Absolute zero

BP helium

Boiling point oxygen

Boiling point mercury

Melting point ice

Boiling point water

0 R7 R

162 R

421 R

459 R

671 R

Rankine

Room temp25°C 298 K 75°F 534 R


80

Temperature Scales

• The Fahrenheit temperature scale used as its two reference points the freezing point of concentrated saltwater (0 °F) and average body temperature (96 °F).More accurate measure now sets average body

temperature at 98.6 °F.

• Room temperature is about 72 °F.


81

Temperature Scales, Continued

• The Celsius temperature scale used as its two reference points the freezing point of distilled water (0 °C) and boiling point of distilled water (100 °C).More reproducible standards.Most commonly used in science.

• Room temperature is about 22 °C.


82

Fahrenheit vs. Celsius

• A Celsius degree is 1.8 times larger than a Fahrenheit degree.

• The standard used for 0° on the Fahrenheit scale is a lower temperature than the standard used for 0° on the Celsius scale.

1.8

32-F C


83

The Kelvin Temperature Scale• Both the Celsius and Fahrenheit scales have

negative numbers.Yet, real physical things are always positive amounts!

• The Kelvin scale is an absolute scale, meaning it measures the actual temperature of an object.

• 0 K is called absolute zero. It is too cold for matter to exist because all molecular motion would stop.0 K = -273 °C = -459 °F.Absolute zero is a theoretical value obtained by

following patterns mathematically.


84

Kelvin vs. Celsius• The size of a “degree” on the Kelvin scale is the

same as on the Celsius scale.Although technically, we don’t call the divisions on the

Kelvin scale degrees; we call them kelvins!That makes 1 K 1.8 times larger than 1 °F.

• The 0 standard on the Kelvin scale is a much lower temperature than on the Celsius scale.

• When converting between kelvins and °C, remember that the kelvin temperature is always the larger number and always positive!

273C K

Example 3.7—Convert –25 °C to Kelvins


Check:7. Check.

258 KRound:6. Significant figures and round.


Solution Map:


Equation:3. Write down the appropriate Equations.

KFind:2. Write down the quantity you want to Find and unit.

-25 °CGiven:1. Write down the Given quantity and its unit. units place

units place

° C K

273C K

K 258273C) 25(K

K = ° C + 273

Example 3.8—Convert 55° F to Celsius


Check:7. Check.

12.778 °C = 13 °CRound:6. Significant figures and round.


Solution Map:



° CFind:2. Write down the quantity you want to Find and unit.

55 °FGiven:1. Write down the Given quantity and its unit.

units placeand 2 sig figs

units place and 2 sig figs

° F ° C

1.8

32-F C

1.8

32-F C

C 778.12

1.8

32-F 55 C

Example 3.9—Convert 310 K to Fahrenheit


Check:7. Check.

98.6 °F = 99 °FRound:6. Significant figures and round.


Solution Map:



°FFind:2. Write down the quantity you want to Find and unit.

310 KGiven:1. Write down the Given quantity and its unit.

units placeand 3 sig figs

units place and 2 sig figs

1.8

32-F C

32C1.8 F

F 6.9832C 37.81 F

K = °C + 273

°F°CK

°C = K - 273

C 37273310 C


95

Practice—Convert 0 °F into Kelvin


96

Practice—Convert 0 °F into Kelvin, Continued

°C = 0.556(°F-32)°C = 0.556(0-32)

°C = -18 °C

K = °C + 273K = (-18) + 273

K = 255 K


97

Energy and the Temperature of Matter

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

temperature will increase twice as much.

• The amount the temperature of an object increases depending on its mass.If you double the mass, it will take twice as

much heat energy to raise the temperature the same amount.

98

Heat Capacity• Heat capacity is the amount of heat a substance

must absorb to raise its temperature by 1 °C.cal/°C or J/°C.Metals have low heat capacities; insulators have

high heat capacities.• Specific heat = heat capacity of 1 gram of the

substance.cal/g°C or J/g°C.Water’s specific heat = 4.184 J/g°C for liquid.

Or 1.000 cal/g°C.It is less for ice and steam.


99

Specific Heat Capacity• Specific heat is the amount of energy required to raise the

temperature of one gram of a substance by 1 °C. • The larger a material’s specific heat is, the more energy it

takes to raise its temperature a given amount.• Like density, specific heat is a property of the type of

matter. It doesn’t matter how much material you have. It can be used to identify the type of matter.

• Water’s high specific heat is the reason it is such a good cooling agent. It absorbs a lot of heat for a relatively small mass.


100

Specific Heat CapacitiesSubstance Specific Heat

J/g°C Aluminum 0.903

Carbon (dia) 0.508

Carbon (gra) 0.708

Copper 0.385

Gold 0.128

Iron 0.449

Lead 0.128

Silver 0.235

Ethanol 2.42

Water (l) 4.184

Water (s) 2.03

Water (g) 2.02


101

Heat Gain or Loss by an Object

• The amount of heat energy gained or lost by an object depends on 3 factors: how much material there is, what the material is, and how much the temperature changed.

Amount of Heat = Mass x Heat Capacity x Temperature Change

q = m x C x T

J 557.4

C25.0-29.9 0.372g 2.5 Cg

J

q

q

Example 3.10—Calculate Amount of Heat Needed to Raise Temperature of 2.5 g Ga from 25.0 to 29.9 °C


Check:7. Check.

4.557 J = 4.6 JRound:6. Significant figures and round.


Solution Map:



q, JFind:2. Write down the quantity you want to Find and unit.

m = 2.5 g, T1 = 25.0 °C, T2= 29.9 °C, C = 0.372 J/g°C

Given:1. Write down the Given quantity and its unit.

2 significant figures

m, C, T q

TCmq

TCmq


110

Practice—Calculate the Amount of Heat Released When 7.40 g of Water Cools from 49° to 29 °C

Practice—Calculate the Amount of Heat Released When 7.40 g of Water Cools from 49° to 29 °C,

Continued

q = m ∙ Cs ∙ T

Cs = 4.18 J/gC (Table 3.4)

The unit and sign are correct.

T1 = 49 °C, T2 = 29 °C, m = 7.40 g

q, J

Check:• Check.

Solution:• Follow the concept plan to solve the problem.

Solution Map:

Relationships:

• Strategize

Given:

Find:

• Sort Information

TC mq s Δ

J 106.2J 64.816

C 02- 4.18g 7.40

Δ

2

Cg

J

TCmq s

C 02-

C94 - C 29

12

T

TTT

Cs m, T q

Introductory Chemistry , 3 rd Edition Nivaldo Tro

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Transcript of Introductory Chemistry , 3 rd Edition Nivaldo Tro