1 Matter and Energy Chapter 3. Why does soda fizz when you open the bottle? Why is the sun hot? When iron rusts, what’s happening?

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Matter and Energy

Chapter 3

Why does soda fizz when you open the bottle?

Why is the sun hot? When iron rusts, what’s happening?

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Universe Classified

• Matter is the part of the universe that has mass and volume

- There are three states of matter

Solid, liquid, and gas

• Chemistry is the study of matter– The properties of different types of matter– The way matter behaves when influenced by

other matter and/or energy

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Elements and Compounds

• a pure substance

• homogeneous – same composition throughout

• Contains only one type of atom

• Can not be broken down into simpler substances

by ordinary chemical reactions (not a nuclear

reaction)

Element

Chemical Symbols of Elements

• System started by Jons Berzelius (Sweden, 1779-1848)

• 90-91 naturally occurring elements• One or two first letters of name of the

element.• Many elements names have roots from:

Latin, Greek, mythology, geography, names of scientists.

Examples:

• Americium, Am

• Einsteinium, Es

• Bromine, Br

• Helium, He

• Lead(Plumbum), Pb

• Niobium, Nb

• Iron (Ferrum), Fe

• Mendelevium, Md

Examples of other elements: O2, H2, I2

Compound – chemical combination of two or more elementsHave two properties in common with elements:

1. pure substance

2. Homogeneous – same chemical composition at all times

Have two properties that differ from elements:

1. two or more elements that are chemically combined, in a definite ratio

2. Compounds can be broken down by chemical reactions using energy:

a) decomposition - uses heat

b) electrolysis - uses electricity

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> 10 million compounds

•elements are represented by symbols•compounds are represented by chemical formulas

•chemical formula

1. symbol --> tells which elements are present in compounds

2. subscript (little # lower right) --> tells the number of atoms of each element

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When the elements sodium and chlorine combine chemically to form sodium chloride, there is a change in composition and a change in properties.

Properties of Compounds

• Sodium chloride (commonly known as table salt) is a white solid.

Distinguishing ElementsDistinguishing Elementsand Compoundsand Compounds

Compounds have different properties from their individual elements.

Breaking down Compounds

Breaking down NaCl

•Sodium is a soft gray metal.


Breaking down NaCl

Breaking down Compounds

• Chlorine is a pale yellow poisonous gas.


Classification of Matter(by composition)

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Classification of Matter

Mixtures can be classified as:

•Homogeneous Mixtures

•Heterogeneous Mixtures

P u re S ub s ta n ceC o n s tan t C o m p o s it ion

H o m o ge n e o us

M ix tu reV a ria b le C o m p o s it ion

M a tte r

• Homogeneous mixture = uniform throughout, appears to be one thing

– Also called solutions

– Examples: olive oil, salt water, lemonade, coffee, air

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The substances in the olive oil are evenly distributed throughout the mixture

Example: Stainless Steel

A homogeneous mixture of:

-Iron (Fe)

-Chromium (Cr)

-Nickel (Ni)

• Heterogeneous mixture = non-uniform, contains regions with different properties than other regions

- Examples: oil and vinegar, salad, chicken soup

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Pure Substances vs. Mixtures• Pure Substances

– All samples have the same physical and chemical properties– Constant Composition all samples have the same composition– Homogeneous– Separate into components based on chemical properties

• Mixtures– Different samples may show different properties– Variable composition– Homogeneous or Heterogeneous– Separate into components based on physical properties

• All mixtures are made of pure substances

Copyright©2004 by Houghton Mifflin Company. All rights reserved. Copyright © Houghton Mifflin Company.All rights reserved. 3–18

Figure 3.4: When table salt is stirred into water (left), a homogeneous mixture called a solution forms (right).

Copyright©2004 by Houghton Mifflin Company. All rights

reserved.

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Identity Each of the following as a Pure Substance, Homogeneous Mixture or

Heterogeneous Mixture

Gasoline

A stream with gravel on the bottom

Copper metal


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Identity Each of the following as a Pure Substance, Homogeneous Mixture or

Heterogeneous Mixture

Gasoline – a homogenous mixture

A stream with gravel on the bottom– a heterogeneous mixture

Copper metal– A pure substance (all elements are pure substances)


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Separation of Mixtures• Separate mixtures based on different

physical properties of the components

EvaporationVolatility

ChromatographyAdherence to a Surface

FiltrationState of Matter (solid/liquid/gas)

DistillationBoiling Point

TechniqueDifferent Physical Property

Example: Separate iron filings from sulfur using a magnet.

Filtration: separates a solid from a liquid in a heterogeneous mixture


Figure 3.6: Distillation of a solution consisting of salt dissolved in water.

-separate dissolved solids from a liquid in a homogeneous mixture

-uses boiling and condensation.

Distillation of Crude Oil (Refining)

• Crude Oil is a mixture of Hydrocarbons

Distillation of Crude Oil

Paper Chromatography

Chromatography separates mixtures of substances into their components.

They all have:

• a stationary phase (a solid, or a liquid supported on a solid)

•and a mobile phase (a liquid or a gas).


Objectives

• Observe and explain the difference between states of matter.

• Explain the difference between physical and chemical changes.


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Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved.

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What you observe when you look at a particular sample of matter is its properties.

– Is a solid shiny or dull?– Does a liquid flow quickly or slowly?– Is a gas odorless, or does it have a smell?

Describing Matter

Some Criteria for the Classification of Matter

• Composition (elements and ✔ compounds)

• State (solid, liquid, gas)

• Properties

3 States of Matter

• Solid

• Liquid

• Gas


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States of Matter


Figure 3.1: Liquid water takes the shape of its container.


Figure 3.11: In ice, the water molecules vibrate randomly about their positions in the solid. Their motions are represented by arrows.


Figure 3.12: Equal masses of hot water and cold water separated by a thin metal wall in an insulated box.


Figure 3.13: The H2O molecules in hot water have much greater random motions than the H2O molecules in cold water.


Figure 3.14: The water samples now have the same temperature (50°C) and have the same random motions.

Phase Changes

Melting solid liquidCondensation gas liquid

Freezing liquid solid

Deposition gas solid

Evaporation liquid gasSublimation solid gas

*Boiling: Evaporation occurring beneath the liquid’s surface.

Gallium metal has such a low melting point (30°C) that it melts from the heat of a hand.

Video - https://www.youtube.com/watch?v=wln6WSv-cro

https://www.youtube.com/watch?v=wln6WSv-cro

https://www.youtube.com/watch?v=wln6WSv-cro

Properties of Matter


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Extensive Propertiesmass - a measure of the amount of matter the object contains.

volume of an object is a measure of the space occupied by the object.–The volume of a basketball is greater than the volume of a golf ball.

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Who has a greater volume?


.

Mass and volume are both examples of extensive properties.

– An extensive property is a property that depends on the amount of matter in a sample.

Examples: mass and volume

Extensive Properties


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– An intensive property is a property that depends on the type of matter in a sample, not the amount of matter.

Examples include:

- Hardness of an object -Color- Softness -Boiling point- Absorbency -Odor

Intensive Properties

Is changing phase a physical or chemical change?

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• Physical Properties are the characteristics of matter that can be changed without changing its composition– Characteristics that are directly observable

– Examples: Color, odor, hardness, density, melting point, boiling point, state, solubility.

Substance State Color Melting

Point (C°)

Boiling Point (C°)

Density (g/cm3)

Oxygen O2 Gas Colorless -218 -183 0.0014

Mercury Hg Liquid Silvery-white

-39 357 13.5

Bromine Br2 Liquid Red-brown -7 59 3.12

Water H2O Liquid Colorless 0 100 1.00

Sodium Chloride

NaCl Solid White 801 1413 2.17

Example: Physical Properties


• Chemical Properties are the characteristics that determine how the composition of matter changes as a result of contact with other matter

- Examples: burning, decompose, reactivity, corrode, tarnish, explode, ferment


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Classify Each of the following as Physical or Chemical Properties

The boiling point of ethyl alcohol is 78°C.

Diamond is very hard.

Sugar ferments to form ethyl alcohol.


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Classify Each of the following as Physical or Chemical Properties

The boiling point of ethyl alcohol is 78°C.– Physical property – describes inherent characteristic of

alcohol – boiling point

Diamond is very hard.– Physical property – describes inherent characteristic of

diamond – hardness

Sugar ferments to form ethyl alcohol.– Chemical property – describes behavior of sugar –

forming a new substance (ethyl alcohol)


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Changes in Matter

• Physical Changes are changes to matter that do not result in a change in the composition that make that substance– State Changes – boiling, melting, condensing– Breaking, splitting, grinding, cutting

• Chemical Changes involve a change in the composition of the substance– Produce a new substance– Chemical reaction– Reactants Products

Indications of a Chemical Reaction

• Color change

• Solid forms (precipitate)

• Gas bubbles

• Odor

• Temperature change

• Fizzing


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Formation of a Precipitate

Cu(OH)2

Precipitate


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Classify Each of the following as Physical or Chemical Changes

Iron metal is melted.

Iron combines with oxygen to form rust.

Sugar ferments to form ethyl alcohol.


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Classify Each of the following as Physical or Chemical Changes

Iron is melted.– Physical change – describes a state change, but the

material is still iron

Iron combines with oxygen to form rust..– Chemical change – describes how iron and oxygen

react to make a new substance, rust

Sugar ferments to form ethyl alcohol.– Chemical change – describes how sugar forms a new

substance (ethyl alcohol)

Oxygen combines with the chemicals in wood to produce flames. Is a physical or chemical change taking place?

Source: Jim Pickerell/Stone/Getty Images

Burning of MethaneCH4 +2O2 CO2 + 2H2O

l_______________l l_________________l l l

reactants products

Burning of MethaneCH4 + 2O2 CO2 + 2H2O

The Law of Conservation of Mass (Antoine Lavoisier)

• In any chemical or physical change, mass is neither created or destroyed

• Mass is CONSTANT

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Energy and Energy Changes

• Capacity to do work– chemical, mechanical, thermal, electrical,

radiant, sound, nuclear

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

state change– All physical changes and chemical changes

involve energy changes

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

difference

1. Exothermic = 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.

2. Endothermic = A process that absorbs energy.• Example: melting ice to form liquid water is an

endothermic process.

A burning match releases energy.

Source: ElektraVision/PictureQuest

Video

• Endothermic & Exothermic Reactions


https://www.youtube.com/watch?v=-lHZg0tMdVU

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Units of Energy

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

• joule (J)– 4.184 J = 1 cal

• In nutrition, calories are capitalized– 1 Cal = 1,000 cal (1 kcal = 1,000 cal)

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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

Convert 0.1 kcal to Joules


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Video Chemmatters

• Calorie vs calorie

• https://www.youtube.com/watch?v=G0O87gWv-Xk


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https://www.youtube.com/watch?v=G0O87gWv-Xk

https://www.youtube.com/watch?v=G0O87gWv-Xk

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Energy and the 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 depends on its mass (m)– 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) is the amount of energy required to raise the temperature of one gram of a substance by one Celsius degree

C gJ

4.184 is water ofheat specific the,definitionBy


Specific Heat Capacity Calculations

Amount of Heat = Specific Heat x Mass x Temperature Change

Where T = (Final temperature – initial temperature)

= (Tf - Ti)

Copyright©2004 by Houghton Mifflin Company. All rights reserved. Copyright © Houghton Mifflin Company.All rights reserved.

Q = s x m x T

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Example 1 – 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

Mass = 7.40 g

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

J 526 C17.07.40gC g

J 4.184 Heat

Specific Heat of Water = 4.184 C-g

JC g

J

Q = s x m x T

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Example 2 – 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?

C gJ

0.20C18x g 1.6

J 5.8 s

C18 C23 - C41 TT m

Q s

Tms Q

Table 3.2 lists the specific heat of gold as 0.13Therefore the metal cannot be pure gold.

C gJ

Calorimetry

• Calorimetry is the study of the heat flow that accompanies physical and/or chemical changes.

• The apparatus used to measure heat is a calorimeter.


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“loses” heat Calorimetry

m = 75 gT = 25oC

SYSTEM

Surroundings

m = 30 gT = 100oC

AgH2O

• A hot piece of silver (Ag) at 100oC is added to water that has an initial temperature of 25oC.

• Silver will lose heat, and water will gain heat.

240 g of water (initially at 20oC) are mixed with an unknown mass of iron (initially at 500oC). When thermal equilibrium is reached, the system has a final temperature of 42oC. Find the mass of the iron.

Calorimetry Problems 2 question #5

FeT = 500oCmass = ? grams

T = 20oC

mass = 240 g

-LOSE heat = GAIN heat

- [(S,Fe) (mass) (T)] = (S,H2O) (mass) (T)

- [(0.45 J/goC) (X g) (42oC - 500oC)] = (4.184 J/goC) (240 g) (42oC - 20oC)]

Drop Units: - [(0.45 J/goC) (X) (-458oC)] = (4.184J/goC) (240 g) (22oC)

(206.1 J/g) X = 22091 J

X = 107.2 g Fe

A 97 g sample of gold at 785oC is dropped into 323 g of water, which has an initial temperature of 15oC. If the final temperature of the system has a temperature of 22.1oC, what is the the specific heat of gold?

AuT = 785oC

mass = 97 g

T = 15oC

mass = 323 g

LOSE heat = GAIN heat

- [(S,Au) (mass) (T)] = (S,H2O) (mass) (T)

- [(x) (97 g) (22.1 - 785oC)] = (4.184 J/goC) (323 g) (22.1 - 15oC)]

Drop Units: - [(x)(97 g) (-762.9oC)] = (1351.4 J/oC) (7.1oC)]

-(-74001.3 goC) x = 9594.9 J

x = 9594.9 J

74001.3 goC x = SAu = 0.13 J/goC

A Coffee Cup Calorimeter

Zumdahl, Zumdahl, DeCoste, World of Chemistry 2002, page 302

Thermometer

Styrofoamcover

Styrofoamcups

Stirrer

Documents

1 Matter and Energy Chapter 3. Why does soda fizz when you open the bottle? Why is the sun hot? When iron rusts, what’s happening?