Matter, and more!

MatterAnything that has MASS and takes up SPACE

How is Matter classified?

1) Pure Substances

2) Mixtures

1) Pure SubstancesComposition remains the same, does

not depend on a sample = Fixed composition

Homogenous—same throughout.

Example: Compound (NaCl) or Element (Fe)

2) Mixtures 2+ substances combined together

Substances do not change their properties or name.

Able to be separated, not chemically combined.

Possess a combination of properties based on the substances present.

Types of Mixtures1) Homogenous

Uniform compositionAlso known as “true solutions”Ex. Salt-water

2) HeterogeneousNo uniform compositionCan easily see the different components

of the mixtureEx. Italian dressing

Mixture Types---More Detail

True SolutionWhat we normally think of as a “solution”HomogenousSolute/solvent completely dissolved

How are mixtures separated?Thin-Layer Chromatography

Filtration

Centrifuge

Identify whether a substance is pure/mixture and

homogenous/heterogeneous. 1) Salad

2) Kool-Aid

3) Vegetable soup

4) Ca

5) Water

Properties of Matter1) Chemical

Ability to go through changes resulting in a different substance

The substance is no longer the same, different identity

Evidence of chemical reaction: color change, precipitate forms, gas formation, and/or temperature change

Ex. Burning

2) Physical Observed or measured property Substance identity is not changed Ex. melting point, boiling point, density

Classify each change as either chemical or physical.

1)Gasoline in your engine burns as you start the car.

2)Distilled water

3)Rust on a nail

4)Glow sticks

5)Medicine crushed into a powder

The Atom

How far back does the “atom” go?

Democritus400 B.C.Called the basic unit of

matter an atom or “atomos”

Law of Conservation of Mass/Matter

Matter cannot be created or destroyed

Total mass is constant in chemical reactions.

Originated with Antoine Lavoister (1700s) Quantitative mass data of

reactants and products in mercury oxide decomposition.

Law of Definite Proportions

Proposed by Joseph Proust (late 1700s)Decompositions and research with copper

carbonate

Compound composition and properties are fixedAll compound samples have the same

compositionSame % of elements in the compoundEx. H2O

TerminologyElement– basic unit of a substance, contain only

ONE type of atom, represented by symbol.

Example: Ag, only contains Ag atoms.

Atom—smallest particle of an element that still contains element properties.Example: One atom of Au, cannot have a smaller

particle of gold and still be gold.

Compound vs. MoleculeCompounds:

more than one elementelements combined in definite proportions

Molecule:Smallest unit of a compound that still retains the

properties of the compound.

Dalton Atomic Theory1800s

Atoms make up elements.

Atoms form compounds as a whole and cannot be divided. Compounds formed from atoms joining in FIXED proportions

Dalton Atomic Theory (cont.)

All matter made of atoms

Atoms of an element have the same size, mass, etc.

Different atoms have various sizes, mass, etc.

Atoms cannot be divided, destroyed, or created.

Atoms rearrange in chemical reactions.

John Thomson1897

Cathode-Ray experiments.

Discovered the electron particle and its possible charge.

Stated electrons have a negative charge

Determined ratio between mass and charge of an electron

Early Models of the AtomThompson

Must be a balance between negative and positive charges

“Raisin-Pudding” model

Uniform distribution of positive chargePositive cloud with stationary electrons

Early Models of the AtomRutherford

How are electrons distributed in an atom?

Discovered alpha particles as 42He

Experiments with Au, Ag, and Pt foils bombarded with alpha particles

Early Models of the Atom Rutherford

Mostly empty space

Small, positive nucleus

Contained protons

Negative electrons scattered around the outside

Atomic StructureNucleus

ProtonsNeutrons

Electrons

Atomic StructureElectrons

Tiny, very light particles

Have a negative electrical charge (-)

Move around the outside of the nucleus

Atomic StructureProtons

Much larger and heavier than electrons

Protons have a positive charge (+)

Located in the nucleus of the atom 

Atomic StructureNeutrons

Large and heavy like protons

Neutrons have no electrical charge

Located in the nucleus of the atom 

Describing AtomsAtomic Number (Z) =

number of protonsIn a neutral atom, the # of

protons = the # of electrons

Mass Number (A)= the number of protons + the number of neutrons

IsotopesThe number of protons for a given atom never

changes.

The number of neutrons can change. 

Two atoms with different numbers of neutrons are called isotopes have the same atomic #have different atomic Mass #’sBehave the same chemically

Isotopes

Atomic MassWeighted average of element’s natural isotopes

Some isotopes are more abundant than others….

SO atomic mass leans towards more abundant mass

How do we calculate atomic mass?

1) Masses of Isotopes

2) Fraction of the abundance of each isotope usually a percentage

Average atomic mass = mass contributed by all isotopes Fraction of abundance (isotope mass) = mass from a

particular isotope

Example 1:

Neon has 3 natural isotopes. Ne-20 (90.51%, 19.99244 u)Ne-21 (0.27%, 20.99395 u) Ne-22 (9.22%, 21.99138 u)

What is the weighted average atomic mass for Ne?

Example 2: Two natural copper isotopes are Cu-63 (62.9298 u) and Cu-65 (64.9278 u)

If copper’s atomic mass is given as 63.546 u, what are the percent abundances of these isotopes? Which isotope is the most abundant?

Homework

pp. 66-67 #7, 31-33, 37, 39-42

Finish “Atomic Theory I” worksheet