Why Study Chemistry?

• To be better informed

• To be a knowledgeable consumer

• To make better decisions for yourself and society

• To learn problem-solving skills

• To enhance analytical thinking

Chemistry as the “Central Science”

• Chemistry = the study of matter and the transformation it undergoes

• EVERYTHING is a CHEMICAL– Table salt = sodium chloride, NaCl

– Table sugar = sucrose, C12H22O11

– Clothes: Wool? Cotton? Polyester?– Body: lipids, Proteins, Carbohydrates, DNA/RNA– You name it– it’s a chemical!

Chemistry as the “Central Science”

• Chemistry is the driving force behind many “liberal arts”– Composition of paints? Colors?– Economies of industrial nations

• #1 commercial chemical is sulfuric acid– LOTS of uses!• All idustry involves chemical processes

– Economies of Developing Nations• Agriculture depends on chemicals as fertilizers, pesticides

– Politics and Natural Resources

The Study of Chemistry

• Chemistry is everywhere!• Matter is everywhere!• Thus, chemistry matters!

• Chemistry involves the study of matter – its properties and behavior.

• Macroscopic observations are rooted in microscopic structure.

Assignment: Chemistry in your major

• Find a current news story or historical example that demonstrates the importance of chemistry to your major– For example: chemical resource as a key

issue in a political / economic rift; wars fought over chemical resources; etc

• Write a 2 paragraph summary on issue and its relevance to your studies

Chemistry as the “Central Science”

Classification of Matter4 Physical States: solid, liquid, gas, plasma

Solid:Fixed shape and fixed volume;

Atoms tightly packed together

Classification of Matter

Liquid:No fixed shape but maintains a fixed volume

Atoms loosely packed together, slide around each other

Classification of Matter

Gas:No fixed shape or volume

Atoms not really associated with neighbors at all

Classification of Matter

Plasma:

mix of subatomic particles with not organization

(sun)

States of Matter

Microscopic view of a gas.

Microscopic view of a liquid.

Microscopic view of a solid.

States of MatterGases, liquids and solids are all made up of microscopic particles, but the behaviors of these particles differ in the three phases. The following figure

illustrates the microscopic differences.

States of Matter

State Shape Volume Compress Flow

Solid Keeps Shape

Keeps Volume

No No

Liquid Takes Shape of Container

Keeps Volume

No Yes

Gas Takes Shape of Container

Takes Volume of Container

Yes Yes

Properties of MatterPhysical Properties =

characteristics of a materialColor MassTemperatureOdorDensityHardnessSolubilityConductivity (heat or

electrical)Freezing/boiling point

Chemical Properties = describe how a material reacts with another type of matter

Ability to burn

Ability to rust / corrode

Ability to make a solution acidic or basic

Lack of ability to react with something

Properties of Matter

• physical – measured without changing substance, e.g. physical state, color, odor, density, boiling point

• chemical – describes a substance’s reactivity, e.g. flammability, corrosiveness

• extensive – depends on the amount of matter present, e.g. mass, volume

• intensive – does not depend on the amount of matter present, e.g. density, color, temperature

• Properties: “ The characteristics that give each substance its unique identity “

• Physical Properties: “ Properties that can be observed

without changing the identity of a substance “

ColorMelting Temperature - a physical change of stateElectrical conductivityDensityBoiling Temperature - a physical change of stateSolubilityHardness

Chemical Properties: “ Properties that result in changes in the identity of one or more reactants “

The rusting of ironHydrogen and oxygen burning to form waterThe baking of breadThe absorption of oxygen by hemoglobin

Changes in MatterPhysical Changes = a

change in a physical property; does NOT change the chemical composition or atomic arrangement of the material– Increase in

temperature– Phase changes– Cutting into smaller

pieces

Chemical Changes = changes that alter the identity of a material, a change in the chemical composition or atomic arrangement of the material– Wood burns in air to

produce CO2 and H2O– Cooking an egg (change

molecular structure of the proteins, loss of water)

– Formation of rust (iron to iron oxide)

• Properties: “ The characteristics that give each substance its unique identity “

• Physical Properties: “ Properties that can be observed without changing the identity of a substance

Color

Melting Temperature - a physical change of state

Electrical conductivity

Density

Boiling Temperature - a physical change of state

Solubility

Hardness

continue…..

Changes in Matter: Is it Physical or Chemical?

• Chemical Properties: “ Properties that result in changes in the identity of one or more reactants “

The rusting of iron

Hydrogen and oxygen burning to form water

The baking of bread

The absorption of oxygen by hemoglobin

continue…..

Changes in Matter (cont)

Changes in Matter (cont)

Reactants Products

Chemical Reactions: “ Process in which one or more pure substances are converted to one or more different pure substances “

Reactants: “ Substances that undergo change in a chemical reaction “

Products: “ Substances formed as the result of a chemical reaction “

Hydrogen + Oxygen Water

Reactants are on the left side of the chemical equation

Products are on the right side of the chemical equation

Changes in Matter - Physical & Chemical

• Physical Change: “ A change that alters the physical form of matter without changing its chemical identity “

• Chemical Change: “ A change which changes the chemical identity of the substance and creates one or more new substances “

continue…..

Changes in Matter - Physical Change

A Melting Ice Sickle

Solid Water

Liquid Water

continue…..

•Example of a Physical Change:

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• Example of a Chemical Change: The Electrolysis of Water (H2O)

Hydrogen Gas

Oxygen Gas

Negative Electrode Positive Electrode

ParticulateViewpoint

The Chemical Identity of Water ( H2O ) is changed

into the elements Hydrogen ( H2 ) and Oxygen ( O2 )

2H2O 2H2 + O2continue…..

Changes in Matter - Chemical Change

2.7 Using Chemical Symbols (cont)

Chemical Equations: “ Representations of chemical reactions by the formulas of reactants and products “

2 C (s) + O2 (g) 2 CO

At the Macroscopic Level: “ Carbon, a solid plus oxygen gas yieldscarbon monoxide “

At the Particulate Level: “ Two atoms of carbon plus one diatomic molecule of oxygen yields two molecules

of carbon monoxide “

Equation Coefficients: “ Gives the relative amount of each compound involved in the chemical equation “

Balanced Chemical Equations: “ The number of each kind of atom on the reactant side must equal the number of each kind of atom on the product side “

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Classification of Matter

Matter - Anything that occupies space and has

mass (solid, liquid or gas)

Heterogeneous Mixture: Non-uniform composition

Homogeneous Matter: Uniform composition

Pure Substances: Fixed composition; cannot be

further purified

Physically Separable Into

Solution: Homogeneous

mixturePhysically Separable Into

Compounds: Elements united in

fixed ratios

Elements: Cannot be subdivided by

chemical or physical changes

Chemically decomposable Into

Combine Chemically to

The Chemical View of Matter

continue….

What are elements and chemical compounds made of?

What is the difference between a mixture and a pure substance?

What is the difference between a chemical and a physical process?

What is the basic theme of chemistry?

How are symbols for the elements used in formulas and equations to communicate chemical information?

Macroscopic, Microscopic & Particulate Matter

• Matter: - “ Anything that has mass and takes up space • (occupies volume) “• Matter can be studied on three levels:• Macroscopic Level: “ Matter that can be seen with the human eye “

Beach Sand, Trees, Cars, Pen, CD, Mountains, Planets, Galaxies, etc

• Length: 101 to 109 meters

continue…..

• Microscopic Level: “ Matter that is too small to be seen by the naked eye, but can be seen under a

• microscope Very small plants, individual bacteria, cellular structures, DNA Molecule, Semiconductors, etc

• Length: 10- 6 meters

continue…..

Macroscopic, Microscopic & Particulate Matter (cont)

• Particulate Level: “ Matter too small to be seen with even the most powerful optical microscope “

Particulate matter consists of the tiny particles that make up all matter

Molecules, atoms, protons & electron

• Length: 10 - 10 meters (1 Angstrom = 10 - 10

meters )

continue…..

Macroscopic, Microscopic & Particulate Matter (cont)

Elements - The Most Simple Kind of Matter

Pure Substance: “Something that with a uniform, fixed composition at the submicroscopic level”

Recognized by the unchanging nature of their properties

Element: “A pure substance composed of only one kind of atom”

Atom: “The smallest particle of an element”

Atoms of different elements are different and are shown on the periodic table

Each element has a one or two letter abbreviation

Hydrogen - H Helium - He Sodium - Na Lithium - Li

                                                               

Microscopic view of the atoms of the element argon (gas phase).

Microscopic view of the molecules of the element nitrogen (gas phase).

Elements

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The Periodic Table and the Elements (cont)

Transition Metals

continue….

Main GroupElements

Main Group Elements

Inner Transition Elements

Chemical Compounds - Atoms in Combination

Chemical Compounds: “ Pure substances made of atoms of different elements combined in definite ways”

Examples:

H2O Water

NaCl Sodium Chloride

C2H6O Ethanol

C6H12O6 Sugar

Chemical Compounds (cont) • Compound: “ Any pure substance that can be decomposed by a• chemical change into two or more pure substances • is a compound “ - (another definition)

• Compounds are made up of elements

• Examples of Compounds:

Water - H2O Ethanol - C2H6O

Salt - NaCl Sugar - C6H12O6

• Examples of Mixtures of Compounds:

Pepper Beer, Wine & Soda Pop Milk Cheese

continue…..

Using Chemical Symbols

Chemical Formulas: “ Combinations of the symbols for the elements that represent the stable combinations of atoms

in molecules “

Examples:

Water H2O

Carbon dioxide CO2

Ammonia NH3

Methane CH4

Carbon Tetrachloride CCl4

Subscripts: “ Indicate the relative numbers of atoms of each kind “

Using Chemical Symbols (cont)

Structural Formulas: “ Formulas that show the connections between atoms in molecules “

H N H

H

Ammonia

H O HWater

H C H

H

Methane

H

Microscopic view of the molecules of the compound water (gas phase). Oxygen atoms are red and hydrogen atoms are white.

Mixtures and Pure Substances

• Homogeneous Sample: “ Matter that has a uniform appearance and • composition throughout “

A mixture of water and alcohol

Sugar dissolved in water

Gold blended with silver (18 karat gold)

The air we breathe - a mixture of oxygen and nitrogen

continue…..Solutions: “Homogeneous mixtures, either liquid, solid or gaseous”

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Table salt is stirred into water (left), forming a homogeneous mixture called a solution (right)

Mixtures and Pure Substances (cont)

continue…..

Heterogeneous Sample: “ Matter that does not have a uniform appearance and composition throughout “

A mixture of cooking oil and water (two phases develop)

Concrete (sand, rock, cement, etc)

A mixture of sand, sawdust, iron fillings and water

Mixtures and Pure Substances (cont)

continue…..

Sand and water do not mix to form a uniform mixture

Mixtures and Pure Substances (cont)

continue…..

Mixtures:Homogeneous • Same composition

throughout sample• Ex- milk, tea,

others?

Heterogeneous• Different samples of

the same mixture have different compositions

• Ex- air in the room others?

Microscopic view of a gaseous mixture containing two elements (argon and nitrogen) and a compound (water).

Classification of Matter

Substances vs Mixtures

Substance – has a definite or fixed

composition– Composition does not

vary from sample to sample

Mixture– Has a varied

composition– Each individual

component can be separated by physical means

– Ex: salt and pepper, sugar in water, sea water

Energy• The “fuel” of the universe• The capacity of something to do work

– chemical, mechanical, thermal, electrical, radiant, sound, nuclear

• The SI unit of energy is the Joule (J)– Other common units are

• Calories (cal)• Kilowatt-hour (kW.hr)

• Types of energy:– Potential– Kinetic– Heat

• Energy cannot be created nor destroyed (but it does change from one type to another!)

Changes in Matter - EnergyEnergy: “ The ability to cause change or, in formal terms of physics,

the ability to do work “

Potential Energy: “ Energy in storage “

There is potential energy in gasoline called chemical energy

Chemical energy is release as heat and light when it burns

Chemical energy can also be released as electrical energy

Kinetic Energy: “ Energy in motion “

Examples are - Muscle in movement, a rocket in flight, inflation of a car air bag during collision

Heat & Temperature

• Temperature is _____.– how hot or cold something is (a physical property)– related to the average (kinetic) energy of the substance (not

the total energy)– Measured in units of

• Degrees Fahrenheit (oF)• Degrees Celsius (oC)• Kelvin (K)

• Heat is energy that _____.– flows from hot objects to cold objects– is absorbed/released by an object resulting in its change in

temperature

• Heat absorbed/released is measured by changes in temperature

SubstancesElements• Fundamental

substances from which all things are constructed

• Only one type of atom is present

• Can not be broken down any further

SubstancesCompounds• Substances made up

of two or more elements in distinct ratios

• Molecules: smallest characteristic part of a compound; composed of a distinct and unique arrangement of elements

Temperature Scales

• Fahrenheit Scale, °F– Water’s freezing point = 32°F, boiling point = 212°F

• Celsius Scale, °C– Temperature unit larger than the Fahrenheit– Water’s freezing point = 0°C, boiling point = 100°C

• Kelvin Scale, K– Temperature unit same size as Celsius– Water’s freezing point = 273 K, boiling point = 373 K

Temperature of ice water and boiling water.

Heat• Heat is the flow of energy due to a temperature difference

– Heat flows from higher temperature to lower temperature

• Heat is transferred due to “collisions” between atoms/molecules of different kinetic energy

• When produced by friction, heat is mechanical energy that is irretrievably removed from a system

• Processes involving Heat:1. Exothermic = A process that releases heat energy.

• 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.

Heat (cont.)• The heat energy absorbed by an object is

proportional to:– The mass of the object (m)

– The change in temperature the object undergoes (T)

– Specific heat capacity (s) (a physical property unique to the substance)

• To calculate heat (Q):Q = c . m . T

Specific Heat Capacity (c)• The amount of heat energy (in J or Cal) required to increase

the temperature of 1 gram of a substance by 1oC (or 1K)

• The Units of Specific Heat Capacity:1. J/goC (SI)2. cal/goC (metric & more useful in the lab)

• Specific Heat Capacity is a unique physical property of different substances– Metals have low specific heat capacity– Non-metals have higher specific heat capacity– Water has an unusually large specific heat capacity

c = Q/(mT)