TRO3 Lecture 03

8/7/2019 TRO3 Lecture 03

1/111

Roy Kennedy

Massachusetts Bay Community College

Wellesley Hills, MA

Introductory Chemistry, 3rd Edition

Nivaldo Tro

Chapter 3

Matter and Energy

2009, Prentice Hall


2/111

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.


3/111


What Is Matter? Matter is defined as

anything that occupiesspace and has mass.

Even though it appears tobe smooth and continuous,matter is actually composed

of a lot of tiny little pieceswe call atoms andmolecules.


4/111


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 inspecific geometric

arrangements.


5/111

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 Dioxide

1. 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


6/111


Classifying Matter

by 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 doesnt change when placed in a container.

Indefinite = Takes the property of the container.


7/111


8/111


Solids

The particles in a solid are packedclose together and are fixed inposition.Although they may vibrate.

The close packing of the particlesresults in solids beingincompressible.

The inability of the particles to

move around results in solidsretaining their shape and volumewhen placed in a new containerand prevents the particles from

flowing.


9/111


Solids, Continued Some solids have their particles

arranged in an orderly geometricpatternwe call these crystalline

solids.Salt and diamonds.

Other solids have particles that donot show a regular geometricpattern over a long rangewecall these amorphous solids.Plastic and glass.


10/111


Liquids

The particles in a liquid are closely packed,but they have some ability to move around.

The close packing results in liquids beingincompressible.

The ability of the particles to move allowsliquids to take the shape of their container

and to flow. However, they dont haveenough freedom to escape and expand to fillthe container.


11/111


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.


12/111


Gases, Continued

Because there is a lot of emptyspace, 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 andtake the shape of their container,

and will flow.

l ifi i f


13/111


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.


14/111


PracticeClassify 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.


15/111



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


16/111


Classifying Matter

by 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 ofatoms or molecules is called a mixture.

Because pure substances always have only onekind of piece, all samples show the same

properties. However, because mixtures have variable

composition, different samples will showdifferent properties.


17/111


CopperA 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


18/111

BrassA MixtureType Color % Cu % Zn Density

g/cm3MP

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 deepyellow 80 20 8.67 999 74K musical instrumentsand 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


19/111


Classification of Matter

Pure Substance = All samples are made of the same pieces inthe same percentages.

Salt

Mixtures= Different samples may have the same pieces indifferent percentages. Salt water

Pure Substance

Constant Composition

Homogeneous

Mixture

Variable Composition

Matter


20/111


Pure Substances vs. MixturesPure Substances

1. All samples have the samephysical and chemicalproperties.

2. Constant composition = Allsamples have the same pieces

in the same percentages.3. Homogeneous.

4. Separate into componentsbased on chemical properties.

5. Temperature stays constant

while melting or boiling.

Mixtures

1. Different samples may showdifferent properties.

2. Variable composition =Samples made with the samepure substances may have

different percentages.3. Homogeneous or

heterogeneous.

4. Separate into componentsbased on physical

properties.5. Temperature usually

changes while melting orboiling because compositionchanges.


21/111



Pure Substances or Mixtures A homogeneous liquid whose temperature stays

constant while boiling.

Granitea rock with several visible minerals in it. A red solid that turns blue when heated and

releases water that is always 30% of the solids

mass.

A gas that when cooled and compressed, a liquid

condenses out but some gas remains.


22/111



Pure Substances or Mixtures,

Continued A homogeneous liquid whose temperature stays

constant while boiling = pure substance.

Granitea rock with several visible minerals in it= mixture.

A red solid that turns blue when heated andreleases water that is always 30% of the solids

mass = pure substance. A gas that when cooled and compressed, a liquid

condenses out but some gas remains = mixture.


23/111


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 theelements 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.


24/111


Atoms and Molecules AtomsAre submicroscopic particles that are the

unit pieces of elements.

Are the fundamental building blocks of allmatter.

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 andpatterns.


25/111


Classification of Pure Substances

1. Made of one

type of atom.

(Some elements

are found as

multi-atom

molecules in

nature.)

2. Combinetogether to make

compounds.

1. Made of one

type of

molecule, orarray of ions.

2. Molecules

contain 2 or

more differentkinds of atoms.

Elements Compounds


26/111



Elements or Compounds Chlorine, Cl

2

Table sugar, C12H

22O

11

A red solid that turns blue when heated and


mass.

A brown-red liquid that, when energy is applied toit in any form, causes only physical changes in the

material, not chemical.


27/111



Elements or Compounds, Continued Chlorine, Cl

2= element.

Table sugar, C12H

22O

11= compound.

A red solid that turns blue when heated and


mass = compound.

A brown-red liquid that, when energy is applied toit in any form, causes only physical changes in the

material, not chemical = element.


28/111


Classification of Mixtures

Mixtures are generally classified based ontheir uniformity.

Mixtures that are uniform throughout arecalled homogeneous.Also known as solutions.

Mixing is on the molecular level.

Mixtures that have regions with differentcharacteristics are called heterogeneous.


29/111


Classification of Mixtures, Continued

1. Made of

multiple

substances, but

appears to be

one substance.

2. All portions of

a sample have

the same

compositionand properties.

1. Made of

multiple

substances,

whose

presence can

be seen.

2. Portions of a

sample have

differentcomposition

and properties.

Heterogeneous Homogeneous


30/111

30

Classifying Matter


31/111


Properties Distinguish Matter

Each sample of matter is distinguished byits characteristics.

The characteristics of a substance are calledits properties.

Some properties of matter can be observeddirectly.

Other properties of matter are observedwhen it changes its composition.


32/111


Properties of Matter

Physical Properties are the characteristics of matterthat 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.


33/111


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


34/111


Some Physical Properties of Iron

Iron is a silvery solid at room temperature with a metallic taste and smoothtexture.

Iron melts at 1538 C and boils at 4428 C.

Irons density is 7.87 g/cm3.

Iron can be magnetized.

Iron conducts electricity, but not as well as most other common metals.

Irons 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 ofiron by 1C.


35/111


Some Chemical Properties

Acidity Basicity (aka alkalinity)

Causticity Corrosiveness

Reactivity Stability

Inertness Explosiveness

(In)Flammability Combustibility

Oxidizing ability Reducing ability


36/111


Some Chemical Properties of Iron

Iron is easily oxidized in

moist air to form rust.

When iron is added to

hydrochloric acid, it producesa solution of ferric chloride

and hydrogen gas.

Iron is more reactive thansilver, but less reactive than

magnesium.


37/111


PracticeDecide 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.


38/111


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.


39/111


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.


40/111


Changes in Matter, Continued

Physical ChangesChanges inthe properties of matter that donot 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 separatedback into sugar and water byevaporation.


41/111


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 andwater.

Silver combines with sulfur in the air to

make tarnish.


42/111


Is it a Physical or Chemical Change?

A physical change results in a different form of

the same substance.

The kinds of molecules dont change.

A chemical change results in one or morecompletely 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.

Ph Ch A


43/111


Phase Changes Are

Physical 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 nota simple phase change, it is a solution

process.


44/111


PracticeClassify 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.


45/111


PracticeClassify 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.


46/111


Separation of Mixtures Separate mixtures based on different

physical properties of the components.Physical change.

Centrifugation and

decantingDensity

EvaporationVolatility

ChromatographyAdherence to a surface

FiltrationState of matter (solid/liquid/gas)

DistillationBoiling point

TechniqueDifferent Physical Property


47/111


Distillation


48/111


Filtration


49/111


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.


50/111


Conservation of Mass

Total amount of matter remains constant in achemical reaction.

58 grams of butane burns in 208 grams of oxygen toform 176 grams of carbon dioxide and 90 grams of

water.butane + oxygen carbon dioxide + water

58 grams + 208 grams 176 grams + 9

266 grams =

Practice A Student Places Table Sugar and


51/111


PracticeA 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 WasProduced?

Practice A Student Places Table Sugar and


52/111


PracticeA 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 WasProduced?

Total of reactants and beaker = 144.0 g.

Conservation of mass says total of products andbeaker must be 144.0 g.

Mass of steam = 144.0 g 129.6 g = 14.4 g.


53/111


Energy

There are things that do not have mass andvolume.

These things fall into a category we call energy.

Energy is anything that has the capacity to dowork.

Although chemistry is the study of matter, matter

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

matter.


54/111


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 increaseor decrease that amount.

However, we can transfer energy from one

place in the universe to another, and we canchange its form.


55/111


Matter Possesses Energy When a piece of matter possesses energy, it can give some or all of it toanother object.

It can do workon the other object.

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


56/111


Kinds of Energy

Kinetic and Potential Potential energy is energy that is

stored. Water flows because gravity pulls it

downstream. However, the dam wont allow it to

move, so it has to store that energy.

Kinetic energy is energy of motion,or energy that is being transferredfrom one object to another. When the water flows over the dam,

some of its potential energy is convertedto kinetic energy of motion.


57/111


Some Forms of Energy Electrical

Kinetic energy associated with the flow of electrical charge.

Heat or Thermal Energy

Kinetic energy associated with molecular motion.

Light or Radiant Energy

Kinetic energy associated with energy transitions in an atom. Nuclear

Potential energy in the nucleus of atoms.

Chemical

Potential energy in the attachment of atoms or because of their position.


58/111


Converting Forms of Energy

When water flows over the dam, some of itspotential 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 topush 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 youmixed, causing it to change chemically and storingsome of the energy in the new molecules that aremade.


59/111


Using Energy

We use energy to accomplish all kinds ofprocesses, but according to the Law of

Conservation of Energy we dont 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 carmove.


60/111


61/111


Theres 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 theair rather than move your car down the

road.


62/111


Units of Energy

Calorie (cal) is the amount of energy needed toraise 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)


63/111


Energy Use

Unit Energy Required toRaise Temperatureof 1 g of Water by1C

EnergyRequired toLight 100-WBulb for 1Hour

Energy Usedby AverageU.S. Citizenin 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 3.5Convert 225 Cal to Joules


64/111

p e 3.5 Co ve 5 C o Jou es

Units and magnitude are correct.Check:7. Check.

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

Solution:5. Follow the solution map to

Solve the problem.

Solution

Map:

4. Write a Solution Map.

1 Cal = 1000 cal

1 cal = 4.184 J

Conversion

Factors:

3. Write down the appropriate

Conversion Factors.

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

225 CalGiven:1. Write down the Givenquantity and its unit.

Cal

cal1J.1844

J1041.9cal1

J.1844

Cal1

cal1000Cal252 5=

cal J

Cal1cal1000

3 sig figs

3 significant figures


65/111


Example 3.5:

A candy bar contains 225 Cal of nutritional energy. How

many joules does it contain?

Example:


66/111


p

A candy bar contains

225 Cal of nutritional

energy. How manyjoules does it contain?

Write down the given quantity and its units.

Given: 225 Cal

InformationExample:


67/111


Write down the quantity to find and/or its units.

Find: ? joules

Given: 225 Cal

p




InformationExample:


68/111


Collect needed conversion factors:

1000 cal = 1 Cal4.184 J = 1 cal

Given: 225 Cal

Find: ? J

p




Information

Gi 225 C lExample:


69/111


Write a solution map for converting the units:

Given: 225 Cal

Find: ? J

Conversion Factors:

1000 cal = 1 Cal; 4.184 J = 1 cal

Cal cal J

Cal1

cal0001

cal1

J1844.




InformationExample:


70/111


cal1

J1844

Cal1

cal0001Cal252

.

Apply the solution map:

= 941400 J

= 9.41 x 105 J

Significant figures and round:

Given: 225 Cal

Find: ? J

Conversion Factors:

1000 cal = 1 Cal; 4.184 J = 1 calSolution Map: Cal cal J




Cal1

cal1000

cal1

J1844.

InformationExample:


71/111


Check the solution:

225 Cal = 9.41 x 105 J

The units of the answer, J, are correct.

The magnitude of the answer makes sense

since joules are much smaller than Cals.

Given: 225 Cal

Find: ? J

Conversion Factors:

1000 cal = 1 Cal; 4.184 J = 1 calSolution Map: Cal cal J




Cal1

cal1000

cal1

J1844.


72/111


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 attractiveforces between molecules.

When materials undergo a physical change, theattractions between molecules change as their positionchanges, resulting in a change in the amount ofchemical potential energy.

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


73/111


Energy Changes and

Chemical Reactions Chemical reactions happen most readily when

energy is released during the reaction.

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

Energy will be released when the reactants

have more chemical potential energy than theproducts.


74/111


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.


75/111


An Exothermic Reaction

Potentialenergy

Reactants

Products

Surroundings

reaction

Amount

of energy

released


76/111


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.

d h i i


77/111


An Endothermic Reaction

Potentialenergy

Products

Reactants

Surroundings

reaction

Amount

of energy

absorbed

T t S l


78/111


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


79/111

Temperature Scales

Celsius Kelvin Fahrenheit-273C-269C

-183C

-38.9C

0C

100C

0 K4 K

90 K

234.1 K

273 K

373 K

-459 F-452F

-297F

-38F

32F

212F

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 temp25C 298 K 75F 534 R


80/111


Temperature Scales

The Fahrenheit temperature scale used as its

two reference points the freezing point of

concentrated saltwater (0 F) and averagebody temperature (96 F).

More accurate measure now sets average body

temperature at 98.6 F.

Room temperature is about 72 F.


81/111


Temperature Scales, Continued

The Celsius temperature scale used as its

two reference points the freezing point of

distilled water (0 C) and boiling point ofdistilled water (100 C).

More reproducible standards.

Most commonly used in science.

Room temperature is about 22 C.


82/111


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-FC =


83/111


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 itmeasures the actual temperature of an object.

0 K is called absolute zero. It is too cold formatter 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.

Kelvin vs Celsius


84/111


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

same as on the Celsius scale.Although technically, we dont 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 lowertemperature than on the Celsius scale.

When converting between kelvins and C, rememberthat the kelvin temperature is always the larger

number and always positive!

273CK +=

Example 3.7Convert 25 C to Kelvins


85/111


258 KRound:6. Significant figures andround.


Solve the problem.

Solution

Map:


Equation:3. Write down the appropriate

Equations.

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

-25 CGiven:1. Write down the Givenquantity and its unit. units place

units place

C K

273CK +=

K258273C)25(K =+=

K = C + 273

Example 3.8Convert 55 F to Celsius


86/111

Units and magnitude are

correct.Check:7. Check.

12.778 C = 13 CRound:6. Significant figures andround.


Solve the problem.

Solution

Map:



Equations.

CFind:2. Write down the quantityyou want to Find and unit.

55 FGiven:1. Write down the Givenquantity and its unit.

units place

and 2 sig figs

units place and 2 sig figs

F C

( )

1.8

32-FC

=

( )1.8

32-FC

=

( )C778.12

1.8

32-F55C =

=

Example 3.9Convert 310 K to Fahrenheit


87/111


98.6 F = 99 FRound:6. Significant figures andround.


Solve the problem.

Solution

Map:



Equations.

FFind:2. Write down the quantityyou want to Find and unit.

310 KGiven:1. Write down the Givenquantity and its unit.

units place

and 3 sig figs

units place and 2 sig figs

( )

1.8

32-FC

=

( ) 32C1.8F +=

( ) F6.9832C37.81F =+=

K = C + 273

FCK

C = K - 273

C37273310C ==


88/111


Example 3.9:

Convert 310 K to Fahrenheit.

Example:

Convert 310 K to Fahrenheit


89/111




Given: 310 K

Information

Given: 310 K

Example:Convert 310 K to Fahrenheit


90/111



Find: ? F

Given: 310 KConvert 310 K to Fahrenheit.

Information

Given: 310 K



91/111


Collect needed equations:

Given: 310 K

Find: ? F


( )1.8

32-FC

=

273CK +=

Information

Given: 310 K



92/111


Write a solution map:

Given: 310 K

Find: ? F

Equations:

K C F


( )1.8

32-FC

= 273CK +=

( )1.8

32-FC

=273CK +=

C273K = ( )32-FC81 =.

F23C1.8 =+

Information

Given: 310 K

Example:Convert 310 K to


93/111



= 99 F

Significant figures and round:

Given: 310 K

Find: ? F

Equations:Solution Map: KC F

Convert 310 K to

Fahrenheit.

C273K =

F23C1.8 =+C273K =

F23C1.8 =+

C273103 =

C37 =

F23371.8 =+F8.69 =

Information

Given: 310 K

Example:Convert 310 K to


94/111


Check the solution:

310 K = 99 F

The units of the answer, F, are correct.

The magnitude of the answer makes sense

since both are above, but close to, room temperature.

Given: 310 K

Find: ? F

Equations:Solution Map: KC F

Convert 310 K to

Fahrenheit.

C273K = F23C1.8 =+


95/111


PracticeConvert 0 F into Kelvin


96/111


PracticeConvert 0 F into Kelvin,

Continued

C = 0.556(F-32)

C = 0.556(0-32)

C = -18 C

K = C + 273

K = (-18) + 273

K = 255 K

Energy and the Temperature of Matter


97/111


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 asmuch heat energy to raise the temperature the

same amount.

Heat Capacity


98/111

98

p y 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/gC or J/gC.

Waters specific heat = 4.184 J/gC for liquid.

Or 1.000 cal/gC.

It is less for ice and steam.

Specific Heat Capacity


99/111


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 materials 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 doesnt matter how much material you have. It can be used to identify the type of matter.

Waters high specific heat is the reason it is such a goodcooling agent. It absorbs a lot of heat for a relatively small mass.

Specific Heat Capacities


100/111


Subst

Al i


101/111


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 thetemperature changed.

Amount of Heat = Mass x Heat Capacity x Temperature Changeq = m x Cx T

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

2 5 T 25 0 CGi1 W it d th Gi


102/111

( ) ( ) ( )J557.4

C25.0-29.90.372g2.5Cg

J

=

=

q

q


4.557 J = 4.6 JRound:6. Significant figures andround.


Solve the problem.

Solution

Map:



Equations.

q, JFind:

2. Write down the quantityyou want to Find and unit.

m = 2.5 g, T1 = 25.0 C,

T2= 29.9 C, C= 0.372 J/gC

Given:1. Write down the Given

quantity and its unit.

2 significant figures

m, C, T q

TCmq =

TCmq =


103/111


Example 3.10:

Gallium is a solid metal at room temperature, but melts at

29.9 C. If you hold gallium in your hand, it melts from

body heat. How much heat must 2.5 g of gallium absorb

from your hand to raise its temperature from 25.0 C to

29.9 C? The heat capacity of gallium is 0.372 J/gC.

Example:

How much heat must 2.5 g of


104/111


gallium absorb from your hand to

raise its temperature from 25.0 C

to 29.9 C? The heat capacity ofgallium is 0.372 J/gC.


Given: mass of Ga = 2.5 g

starting temp. = 25.0 C

final temp. = 29.9 C

spec. heat of Ga = 0.372 J/gC

Example:


Information

Given: m = 2.5 g; Ti = 25.0 C;


105/111



Find: amount of heat in joules




G g; i C;

Tf= 29.9 C; C= 0.372 J/gC

Example:


Information

Given: m = 2.5 g; Ti = 25.0 C;


106/111


Collect needed equations:

TCmq =




Tf= 29.9 C; C= 0.372 J/gC

Find: q (J)

Information

Given: m = 2.5 g; Ti = 25.0 C;

Example:



107/111


Write a solution map:

C, m, T q

TCmq =

Tf= 29.9 C; C= 0.372 J/gC

Find: q (J)

Equation: q = m C T




Information

Given: m = 2.5 g; Ti = 25.0 C;

Example:



108/111



q = 4.6 J Significant figures and round:

Tf= 29.9 C; C= 0.372 J/gC

Find: q (J)

Equation: q = m C T

Solution Map: m, C, Tq




TCmq

=( ) ( )C25.0C29.9

Cg

J0.372g2.5

=q

( ) ( )C9.4Cg

J

0.372g2.5

=q = 4.557 J

Information

Given: m = 2.5 g; Ti = 25.0 C;

Example:



109/111


Check the solution:q = 4.6 J

The units of the answer, J, are correct.

The magnitude of the answer makes sensesince the temperature change, mass, and specific heat are small.

Tf= 29.9 C; C= 0.372 J/gC

Find: q (J)

Equation: q = m C T

Solution Map: m, C, Tq




PracticeCalculate the Amount of Heat Released

h f l f


110/111


When 7.40 g of Water Cools from 49 to 29 C

PracticeCalculate the Amount of Heat ReleasedWhen 7.40 g of Water Cools from 49 to 29 C,


111/111

Continued

q = m Cs TCs = 4.18 J/g C (Table 3.4)

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

TCmq s =

( ) ( ) ( )J106.2J64.816

C02-4.18g7.40

2Cg

J

==

=

=

TCmq s

( )

C02-

C94-C29

12

==

=T

TTT

Cs m, T q

TRO3 Lecture 03

Documents