Why Study Chemistry To be better informed To

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, Na. Cl Table sugar = sucrose, C 12 H 22 O 11 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 Matter 4 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

States of Matter Gases, 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. Microscopic view of a gas. Microscopic view of a liquid. Microscopic view of a solid.

States of Matter

Properties of Matter Physical Properties = characteristics of a material Color Mass Temperature Odor Density Hardness Solubility Conductivity (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 “ Ø Color Ø Melting Temperature - a physical change of state Ø Electrical conductivity Ø Density Ø Boiling Temperature - a physical change of state Ø Solubility Ø Hardness

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

Changes in Matter Physical Changes = a Chemical Changes = changes that alter the change in a physical identity of a material, a property; does NOT change in the chemical composition or atomic change the chemical arrangement of the composition or material atomic arrangement – Wood burns in air to produce CO 2 and H 2 O of the material – Increase in temperature – Phase changes – Cutting into smaller pieces – Cooking an egg (change molecular structure of the proteins, loss of water) – Formation of rust (iron to iron oxide)

Changes in Matter: Is it Physical or Chemical? • Properties: “ The characteristics that give each substance its unique identity “ • Physical Properties: “ Properties that can be observed without changing the identity of a substance Ø Color Ø Melting Temperature - a physical change of state Ø Electrical conductivity Ø Density Ø Boiling Temperature - a physical change of state Ø Solubility Ø Hardness continue…. .

Changes in Matter (cont) • Chemical Properties: “ Properties that result in changes in the identity of one or more reactants “ ØThe rusting of iron ØHydrogen and oxygen burning to form water ØThe baking of bread ØThe absorption of oxygen by hemoglobin continue…. .

Changes in Matter (cont) 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 “ Ø Reactants are on the left side of the chemical equation Products: “ Substances formed as the result of a chemical reaction “ Ø Products are on the right side of the chemical equation Reactants Products Hydrogen + Oxygen Water

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 • Example of a Physical Change: A Melting Ice Sickle Solid Water Liquid Water continue…. .

Changes in Matter - Chemical Change • Example of a Chemical Change: The Electrolysis of Water (H 2 O) Particulate Viewpoint Oxygen Gas Hydrogen Gas Negative Electrode 2 H 2 O Positive Electrode 2 H 2 + O 2 continue…. . The Chemical Identity of Water ( H 2 O ) is changed into the elements Hydrogen ( H 2 ) and Oxygen ( O 2 ) 24

2. 7 Using Chemical Symbols (cont) Chemical Equations: “ Representations of chemical reactions by the formulas of reactants and products “ 2 C (s) + O 2 (g) 2 CO At the Macroscopic Level: “ Carbon, a solid plus oxygen gas yields carbon 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 “

Classification of Matter - Anything that occupies space and has mass (solid, liquid or gas) Heterogeneous Mixture: Non-uniform composition Physically Separable Into Pure Substances: Fixed composition; cannot be further purified Homogeneous Matter: Uniform composition Physically Separable Into Solution: Homogeneous mixture Chemically decomposable Into Compounds: Elements united in fixed ratios Combine Chemically to Elements: Cannot be subdivided by chemical or physical changes 26

The Chemical View of Matter Ø What are elements and chemical compounds made of? Ø What is the difference between a mixture and a pure substance? Ø What is the difference between a chemical and a physical process? Ø What is the basic theme of chemistry? Ø How are symbols for the elements used in formulas and equations to communicate chemical information? continue….

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 “ Ø Beach Sand, Trees, Cars, Pen, CD, Mountains, Ø Planets, Galaxies, etc • Length: 101 to 109 meters continue…. .

Macroscopic, Microscopic & Particulate Matter (cont) • Microscopic Level: “ Matter that is too small to be seen by the naked eye, but can be seen under a • microscope Ø Very small plants, individual bacteria, cellular Ø 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 “ Ø Particulate matter consists of the tiny particles Ø that make up all matter Ø Molecules, atoms, protons & electron • Length: 10 - 10 meters (1 Angstrom = 10 - 10 meters ) continue…. .

Elements - The Most Simple Kind of Matter Pure Substance: “Something that with a uniform, fixed composition at the submicroscopic level” Ø 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” Ø Atoms of different elements are different and are shown on the periodic table ØEach element has a one or two letter abbreviation v Hydrogen - H v Helium - He v Sodium - Na v Lithium - Li

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

Elements

The Periodic Table and the Elements (cont) Main Group Elements Main Group Elements Transition Metals Inner Transition Elements continue…. 34

Chemical Compounds - Atoms in Combination Chemical Compounds: “ Pure substances made of atoms of different elements combined in definite ways” Examples: Ø H 2 O Water Ø Na. Cl Sodium Chloride Ø C 2 H 6 O Ethanol Ø C 6 H 12 O 6 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: continue…. .

Using Chemical Symbols Chemical Formulas: “ Combinations of the symbols for the elements that represent the stable combinations of atoms in molecules “ Examples: Water H 2 O Carbon dioxide CO 2 Ammonia NH 3 Methane CH 4 Carbon Tetrachloride CCl 4 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 Ammonia H N H Water H O H H Methane H C H 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 “ ØA mixture of water and alcohol ØSugar dissolved in water ØGold blended with silver (18 karat gold) ØThe air we breathe - a mixture of oxygen and nitrogen Solutions: “Homogeneous mixtures, either liquid, solid or gaseous” continue…. .

Mixtures and Pure Substances (cont) Table salt is stirred into water (left), forming a homogeneous mixture called a solution (right) 41 continue…. .

Mixtures and Pure Substances (cont) Heterogeneous Sample: “ Matter that does not have a uniform appearance and composition throughout “ Ø A mixture of cooking oil and water (two phases develop) Ø Concrete (sand, rock, cement, etc) Ø A mixture of sand, sawdust, iron fillings and water continue…. .

Mixtures and Pure Substances (cont) Sand water do not mix to form a uniform mixture 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 (k. W. 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 - Energy: “ The ability to cause change or, in formal terms of physics, the ability to do work “ Potential Energy: “ Energy in storage “ Ø There is potential energy in gasoline called chemical energy Ø Chemical energy is release as heat and light when it burns Ø Chemical energy can also be released as electrical energy Kinetic Energy: “ Energy in motion “ Ø 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 (o. F) • Degrees Celsius (o. C) • 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

Substances Elements • Fundamental substances from which all things are constructed • Only one type of atom is present • Can not be broken down any further

Substances Compounds • 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 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 (DT) – Specific heat capacity (s) (a physical property unique to the substance) • To calculate heat (Q): Q = c. m. DT

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 1 o. C (or 1 K) The Units of Specific Heat Capacity: 1. J/go. C (SI) 2. cal/go. C (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/(m. DT)