Principles of Chemistry Exam 1

Matter

Anything that has mass and occupies space (Not weight)

Heterogeneous

Not uniform throughout

Homogeneous

Uniform throughout, combinations of two or more substances but each still retains its chemical identity

Pure Substance

Does not have variable composition

Homogeneous Mixture

Variable composition (like a solution)

Element

Cannot be separated into simpler substances

Compound

Can be separated into two or more elements

Physical Change

Changes in appearance but not composition. Ex: water freezing

Chemical Change

Transformation into chemically different substance. Ex: H2+O2=H2O

Intensive Properties

Do not depend on amount of sample being examined

Extensive Properties

Depends on mass of sample

Peta

10^15

Tera

10^12

Giga

10^9

Mega

10^6

Kilo

10^3

Deci

10^-1

Centi

10^-2

Milli

10^-3

Micro

10^-6

Nano

10^-9

Pico

10^-12

Femto

10^-15

Atto

10^-18

Zepto

10^-21

1 Gal=

3.7854L

1in

2.54cm

1m

1.6093km

1a

10^-10 m

1lb

453.5g

1ml

1cm3

K =

C+273.15

C=

5/9(F-32)

F=

9/5C+32

Density=

mass/volume

Precision

How closely individual measurements agree with one another

Accuracy

how closely measurements agree with the correct value

Sig Figs

All nonzero digits are significant

a) Zeros between nonzero digits are always significant (105g)

b) Zeros at the beginning of a number are never significant (0.2)

c) Zeros at the end of the number are signficant if the number

contains a decimal point (0.0200g

d) Addition/Subtraction # of decimal places

e) Multiplication/Division amount of #s

Scientific Method

-collect info by observation

-find trends

-form hypothesis

-develop theory

Empedocles

-fundamental substances air, earth, fire, water

Democritus

-matter is particulate, not continuous

Dalton's Atomic Theory

1) All matter is made of atoms. Atoms are indivisible and indestructible.

2) All atoms of a given element are identical in mass and properties

3) Atoms are not created or destroyed by a chemical reaction. (law of conservation of mass)

4) Atoms form compounds with same relative number and kind of atoms (law of constant composition and law of multiple proportions)

J.J. Thompson

-discovered electrons, plum pudding model of atom, negative electrons surrounded by positive charge

Rutherford

-Gold foil experiment, atoms have a nucleus filled with protons surrounded by a cloud of electrons

Chadawick

-Neutrons

Mass #

protons+neutrons (top number)

Atomic #

# of protons or electrons (bottom number)

Atomic Weight

Average atomic mass of an element (isotope mass x fractional isotope abundance)

Periods

Horizontal Rows

Groups

Vertical columns

Stair Case Starts with

B and ends with Te

-includes Ge and Sb

Metalloids

Along the staircase, contain some metallic and nonmetallic properties

B's

Transition metals

Group 1

Alkaline metals

Group 2

Alkaline Earth metals

Molecule

At least two atoms in a unit. Ex: H20

Empirical formula

Simplified version of molecular formula. Ex: C2H6O2 becomes CH3O

Cation

Positive

Anion

Negative

How do you predict the charge on an ion?

All ions want 8e in their outer shell so they will lose or gain an electron based on their position in the periodic table

Nonmetal+nonmetal=

molecular

nonmetal+metalloid=

molecular

metal+nonmetal=

ionic

Nomenclature Anions

-monotomic anions are formed by adding "IDE"

-polyatomic anions containing O2 are formed "ATE" or "ITE"

-3 Os "ATE" add an O add a "PER" take away an O "ITE" take away 2 Os add a "HYPO"

Prefixes for Naming Binary Molecular Compounds

1Mono 5Penta 9Nona

2DI 6Hexa 10Deca

3Tri 7Hepta

4Tetra 8octa

Acid Nomenclature

-Hydrogen containing compound

-change IDE to IC add HYDRO and ACID Ex: HCl is hydrochloric acid

-ATE becomes IC ITE becomes OUS

Organic Chemistry

-Hydrogen and Carbon bonds

Organic Compounds

-Methane CH4 (with OH instead of one H becomes methanol, same for below)

-Ethane C2H6

-Propane C3H8

Polyatomic Ion Charge = -1

acetate - C2H3O2-

bicarbonate (or hydrogen carbonate) - HCO3-

bisulfate (or hydrogen sulfate) - HSO4-

chlorate - ClO3-

chlorite - ClO2-

cyanate - OCN-

cyanide - CN-

dihydrogen phosphate - H2PO4-

hydroxide - OH-

nitrate - NO3-

nitrite - NO2-

perchlorate - ClO4-

permanganate - MnO4-

thiocyanate - SCN-

acetate

C2H3O2-

bicarbonate (or hydrogen carbonate)

HCO3-

bisulfate (or hydrogen sulfate)

HSO4-

chlorate

ClO3-

chlorite

ClO2-

cyanate

OCN-

cyanide

CN-

dihydrogen phosphate

H2PO4-

hydroxide

OH-

nitrate

NO3-

nitrite

NO2-

perchlorate

ClO4-

permanganate

MnO4-

thiocyanate

SCN-

Polyatomic Ion Charge = -2

carbonate - CO32-

chromate - CrO42-

dichromate - Cr2O72-

hydrogen phosphate - HPO42-

peroxide - O22-

sulfate - SO42-

sulfite - SO32-

thiosulfate - S2O32-

carbonate

CO32-

chromate

CrO42-

dichromate

Cr2O72-

hydrogen phosphate

HPO42-

peroxide

O22-

sulfate

SO42-

sulfite

SO32-

thiosulfate

S2O32-

Polyatomic Ion Charge = -3

borate - BO33-

phosphate - PO43-

borate

BO33-

phosphate

PO43-

Combination Reaction

A+B=C

Decomposition Reaction

C=A+B

Combustion Reaction

Reactions that produce a flame adding O2 to yield CO2 and H2O

Mol

6.022*10^23 objects of anything

Molar Mass

grams per mol of a substance equal to weight in amu. Ex 1 mol O= 16g