Nature of Matter. Matter Defined as anything that occupies space and has mass Everything we see around us is matter Soil Air Water Rocks

SECTION 2.1Nature of Matter

Matter

Defined as anything that occupies space and has mass

Everything we see around us is matter Soil Air Water Rocks Plants Animals

States of Matter Kinetic Energy—energy of motion; all

particles of matter are in constant motion and therefore have kinetic energy

The temperature of matter is an indirect measurement of the particles’ average kinetic energy

At low temperatures, the particles have very little kinetic energy and move very little

At high temperatures, the particles can have enough kinetic energy to fly about at thousands of kilometers per hour

States of Matter

Solid Has a definite shape A definite volume Very difficult to compress Particles are packed close together and

held rigidly

States of Matter

Liquid Has the shape of its container Has a definite volume Difficult to compress Particles close together but free to move

States of Matter

Gas Has no definite shape Has no definite volume Relatively easy to compress Particles far apart

Plasma A hot gas in which atoms are partially

broken down to form charged particles, or ions

States of Matter

When the substance reaches a certain temperature, called the melting point, the particles have enough energy to break loose from their rigid positions and form a liquid

When a liquid is heated to the boiling point, the particles gain enough kinetic energy to break away from each other forming a gas

Classification of Matter

Substance: a form of matter with unique properties that make it different from every other substance

Examples: Water, copper, baking soda, table salt, aluminum, etc.

Atoms: the smallest particles of an element (not the smallest particles of matter)

Elements: substances that cannot be broken down into simpler substances by ordinary chemical means

Examples: Oxygen, hydrogen, gold, iron, lead,

potassium, nitrogen, silicon (other elements listed in the periodic table of the elements)

Not all elements are found as single atoms

Molecules: groups of two or more atoms that are linked by chemical bonds to form distinct units

These two-atoms elements are called diatomic elements; other elements exist as larger molecules Examples of diatomic elements: hydrogen,

oxygen, nitrogen, flourine, chlorine, bromine, and iodine

Atomic Symbols

Memorize tables 2.2 and 2.3

Abundance of Elements 118 elements identified thus far 88 occur naturally in significant amounts More than 99% of the earth’s crust,

oceans, and atmosphere is made up of only eight elements

Oxygen-the most abundant element in the earth’s crust

Iron-the most abundant element in the earth as a whole

Hydrogen-the most abundant element in the universe

Compounds

A substance that can be decomposed into simpler substances

Examples: water (H2O) and table salt (NaCl),

A compound has properties quite different from those of the elements from which it is formed

Example: hydrogen-highly flammable; oxygen-flammable; H2O-commonly used to put out fires

Formula

A molecule of a compound is composed of different types of atoms linked together

The compound may be represented by a formula-a grouping of symbols that tells what types of atoms compose the compounds and the number of each type of atom in one molecule of the compound

Ex. H2O; CO; CO2; C6H12O6

Pure Substances

Elements and compounds Definite composition (relative

amounts of each element in a given compound are unchangeable, no matter where the compound may be found)

Homogeneous (same kind)-composed of the same kind of matter throughout the sample

Mixtures

A substance consisting of two or more pure substances that are physically mixed but not chemically combined

Each component retains its own properties

Each can be present in various proportions and can be separated from the mixture by physical means

Homogenous Mixtures

Solutions; same throughout Examples: air, vinegar, sweetened

coffee, steel

Heterogeneous Mixtures

Consists of pure substances that are incompletely mixed

Each component in a heterogeneous mixture forms a distinct phase—a homogeneous part of a system that is in contact with but physically distinct from other parts of the system

Examples:

Dalton’s Atomic Theory 1. Every element consists of tiny,

indivisible, indestructible particles called atoms

2. All of the atoms of a particular element have the same size, mass, and chemical behavior

3. Differences in properties of elements result from differences in the atoms of the elements

4. The atoms of the elements combined in a compound are combined in a definite ratio

5. A chemical reaction is the result of rearrangement, combination, or separation of atoms

The Law of Definite Composition States that the ratios of the masses

of each element in a given compound are always the same

Example on pg. 22

The Law of Multiple Proportions States that when two elements can

combine to form more than one compound, the masses of one element that combine with a fixed amount of the other element are in a ratio of small whole numbers

Example on pg. 22

Section 2.2: Properties and Changes of Matter One type of matter may be

described, identified, or distinguished from another by its properties.

Physical Properties

Properties that describe a substance’s appearance

Include: color, odor, density, hardness, solubility, taste, state, boiling point, and melting point

Can be measured without changing the identity or composition of the substance

Chemical Properties

Properties that describe how matter reacts to change into other chemically different substances having different properties

That iron rusts easily in moist air and that propane burns readily are chemical properties of these substances

Can only be determined through chemical reactions, which change the identity of a substance

Physical Changes Changes in the physical appearance of

matter that do not change the identity or chemical composition of a substance

Physical changes may often be reversed by physical processes

Change of state, conversion of heterogeneous to homogeneous matter, metal hammered into a sheet, tearing a sheet of paper, crushing a large rock into gravel, cutting your hair, etc.

No change in chemical identity results, only changes in shape or size

Chemical Changes

A change in which a substance becomes a different substance with a different composition and properties

Sodium–a shiny, soft, poisonous metal that reacts violently with water

Chlorine—a greenish, poisonous gas with an irritating odor

When sodium and chlorine react chemically, a white crystalline solid with a salty taste is formed, namely sodium chloride

Chemical Changes

Other examples: Rusting of iron The digestion of food The grilling of steaks The burning of gasoline The cooking of eggs The burning of a match

Chemical Changes

Chemical changes can be reversed only by other chemical changes Not all chemical changes can be

reversed The cooking of an egg

Chemical Changes

A number of observations can be made that will suggest that a chemical change has occurred Formation of a gas Formation of a precipitate (an insoluble

substance) Liberation or absorption of heat, light, or

some other form of energy A distinct change in color

Separation of Mixtures

Distillation—used to separate homogeneous mixtures in liquid form if the components have widely differing boiling points Simple Distillation

When a mixture has components with boiling points that are close together, fractional distillation is used

Separation of Mixtures

Fractional Crystallization Used to purify a solid containing a

relatively small amount of solid impurity Chromatography

Used to separate complex mixtures Paper Chromatography Column Chromatography Gas Chromatography

Section 2.3: Subatomic Particles Subatomic Particles

Protons, neutrons, and electrons Electrons—has a negative charge Proton—has a positive charge,

located in the nucleus Neutron—has no charge, located in

the nucleus Rest of the notes in this section will

come from the videos

Section 2.4: Atomic Number, Mass Number, Isotopes, and Ions The number of protons in the nucleus

determines the identity of an element

Atomic number—the number of protons in the nucleus Denoted by the symbol Z

All atoms of the same element have the same number of protons and hence the same atomic number, Z

Mass Number

Denoted by the symbol A The sum of the number of protons

and the number of neutrons in a nucleus A = Z + N

Nuclear particles, whether protons or neutrons, are known collectively as nucleons

The mass number is not really a mass; it is a number of nucleons and must be a whole number

Isotopes

Atoms of the same element that differ in their mass numbers

Examples: Pg. 31

Ions Atoms or molecules that have a

different number of electrons than protons

An ion with more electrons than protons is called and anion Has a negative charge

An ion with fewer electrons than protons is called a cation Has a positive charge

In neutral atoms, there are equal numbers of electrons and protons

Examples: pg. 32

Example 2.1 Application 1-3

Section 2.5: Atomic Mass

An atom’s atomic mass measured in atomic mass units approximately equals the atom’s mass number Atomic mass units denoted a u

Mass Spectrometer Used to measure atomic masses with

high precision Mass Spectrum

A record of the mass distribution of particles in a sample

Average Atomic Mass

Calculated by finding the weighted average

Example: pg. 34 Track team Example 2.2 pg. 35

Chapter 2

Homework and Groupwork will be due Friday They will be handed out at the beginning

of class on Tuesday Chapter 2 Quiz on Friday