Chem. Midterm

Front 1

Atomic Mass

Back 1

Weighted average of all isotopes of an element.

Front 2

Mass number

Back 2

Total number of protons and neutrons in a certain element.

Front 3

Atomic Number

Back 3

Number of protons in the nucleus.

Front 4

Heterogeneous Mixture

Back 4

Unevenly distributed mixture with no chemical bonding. Nonuniform in composition.

Front 5

Homogeneous Mixture

Back 5

Evenly distributed mixture with no chemical bonding. Uniform in composition.

Front 6

Mixture

Back 6

Blend of two substances, not chemically combined.

Front 7

Compound

Back 7

Blend of two or more substances that become chemically combined in proportion.

Front 8

Dalton's atomic model

Back 8

A large sphere.

Front 9

Thomson Atomic Model

Back 9

A sphere with elections sprinkled throughout, "like plum pudding."

Front 10

Rutherford's Atomic Model

Back 10

Think Jimmy Neutron. Small, dense, positively charged center.

Front 11

Bohr Atomic Model

Back 11

Looks like a target. Allows the representation of electrons switching energy levels.

Front 12

Schrodinger's Atomic Model

Back 12

Shows an electron cloud with a nucleus.

Front 13

Quantum Mechanical Model

Back 13

Determines the allowed energies an electron can have and how likely it is to find the electron in various locations.

Front 14

Formula for Density

Back 14

D=(M/V)

Front 15

Formula for Percent Error

Back 15

abs(AV-EV/AV)*100

Front 16

Accuracy vs. Precision

Back 16

Accuracy is how close a result is to an accepted value.

Precision is how often that value is repeated under the same conditions.

Front 17

Schrodinger

Back 17

Emphasis on the wave-like electron properties. Electrons are held to specific regions of space.

Front 18

Heisenberg Uncertainty Principle

Back 18

It is impossible to know both the position and velocity of an electron simultaneously

Front 19

Bohr

Back 19

Took the Quantum theory and predicted that electrons orbit at fixed distances. "Target Model."

Front 20

Rutherford

Back 20

Discovered nucleus is positively charged. The atom is mostly empty space.

Front 21

Lenard

Back 21

Wavelength determined by color. Electron ejection creates light.

Front 22

Quantum

Back 22

Discreet amount of energy needed to have light

Front 23

Plank

Back 23

Light behaves as a stream of particles.

Energy is quantized.

Each electron must say in a certain region depending on its energy level.

Front 24

Thomson

Back 24

Electrons are negative and can be found throughout the atom.

Front 25

Dalton

Back 25

Elements combine in precise ratios. The law of conservation of matter.

Front 26

Democritus of Abdera

Back 26

Atoms, alchemists, pancea

Front 27

Electronegativity

Back 27

More shells=Less pull.

Increases L to R, decreases U to D.

Front 28

Electron affinity

Back 28

Energy released when non-metals gain an electron.

Increases L to R.

Front 29

Ionization Enegery

Back 29

Follows same patterns as electro negativity.

The energy required to eject an electron. Greatly increases when crossing energy levels.

Front 30

Atomic Radius

Back 30

Decreases L to R. Increases U to D because shells are added. Electron shielding.

Front 31

Octet Rule

Back 31

Every atom strives to obtain EIGHT valence electrons.

Front 32

Metals

Back 32

Lose electrons, solids, usually positive in bonds.

Front 33

Nonmetals

Back 33

Gains electron in bond. Crystalline solids.

Front 34

Metalloids

Back 34

Semi-Metals, can act like both metals and nonmetals.

Front 35

Medeleev

Back 35

Organized periodic table of elements by mass number, not by atomic number.

Front 36

Mosely

Back 36

Organized periodic table of elements by atomic number.

Front 37

Finding Number of ELECTRONS

Back 37

Refer to the atomic number.

Front 38

Finding Number of PROTONS

Back 38

Refer to the atomic number

Front 39

Finding Number of NEUTRONS

Back 39

Subtract atomic mass from atomic number.

Front 40

Pauli Exclusion Principle

Back 40

Orbital cannot contain more than two electrons.

Front 41

Subshells

Back 41

1s 2s 2p 3s 3p 3d

Front 42

Naming molecular compounds

Back 42

The first element (with the least EN), the second element receives suffix of "-ide."

The number of atoms in indicated by the prefixes. The first element does not receive a "mono>"

Front 43

Molecular compound prefixes

Back 43

MONO
DI
TRI
TETRA
PENTA
HEXA
HEPTA
OCTA
NONA
DECA

Front 44

Naming ionic compounds

Back 44

Metal Cation goes first. Anion goes second and has suffix of "-ide." Numbers are insignificant.

If a metal has more than one oxidation number, then the number must be indicated by roman numerals to describe the charge.

Front 45

Examples of ionic compounds"

CaH2
CuCl2

Back 45

Calcium Hydroxide
Copper (II) Cloride

Front 46

Empirical Formula

Back 46

Lowest whole number ratio of elements in a molecule.

Front 47

Aufbau Principle

Back 47

Electrons occupy the lowest energy levels first.

Front 48

Hund's Rule

Back 48

Electrons occupy orbitals of the same energy in a way that makes the number of electrons with the same spin direction as large as possible.

Front 49

Oxidation State

Back 49

In both ionic and covalent bonds. Each atom assigned an oxidation number with indicated the number of electrons it gains or loses when it bonds. Neutral compounds equal zero.

Front 50

Dipole moment

Back 50

Measure of differences in the electronegativity between non-metals. MUST BE UNEQUAL.

Front 51

Sigma Bonds, Pi Bonds

Back 51

Single Bond/2+ Bonds

Front 52

Polar Bond

Back 52

Unequal sharing between atoms of electrons.

Front 53

Covalent Bonds

Back 53

Sharing between two non-metals.

Front 54

Molecular Shapes

Back 54

Linear
Bent
Trigonal Planar
Pyramidal
Tetrahedral
Trigonal Bipyramidal
Octahedral

Front 55

Electrostatic

Back 55

Between two charged polyatomic molecules.

Front 56

Hydrogen Bonding

Back 56

Between hydrogen atom of one molecule and N, O, Fl atom of another.

Front 57

Dipole-Dipole

Back 57

Between two polar. Strong, based on dipole-moment.

Front 58

Induced-Dipole

Back 58

Dipole induced in one molecule by presence, of a nearby dipole.

Front 59

Octet Rule

Back 59

Every atom strives to obtain EIGHT valence electrons.

Front 60

Metals

Back 60

Lose electrons, solids, usually positive in bonds.

Front 61

Nonmetals

Back 61

Gains electron in bond. Crystalline solids.

Front 62

Metalloids

Back 62

Semi-Metals, can act like both metals and nonmetals.

Front 63

Medeleev

Back 63

Organized periodic table of elements by mass number, not by atomic number.

Front 64

Mosely

Back 64

Organized periodic table of elements by atomic number.

Front 65

Finding Number of ELECTRONS

Back 65

Refer to the atomic number.

Front 66

Finding Number of PROTONS

Back 66

Refer to the atomic number

Front 67

Finding Number of NEUTRONS

Back 67

Subtract atomic mass from atomic number.

Front 68

Pauli Exclusion Principle

Back 68

Orbital cannot contain more than two electrons.

Front 69

Subshells

Back 69

1s 2s 2p 3s 3p 3d

Front 70

Naming molecular compounds

Back 70

The first element (with the least EN), the second element receives suffix of "-ide."

The number of atoms in indicated by the prefixes. The first element does not receive a "mono>"

Front 71

Molecular compound prefixes

Back 71

MONO
DI
TRI
TETRA
PENTA
HEXA
HEPTA
OCTA
NONA
DECA

Front 72

Naming ionic compounds

Back 72

Metal Cation goes first. Anion goes second and has suffix of "-ide." Numbers are insignificant.

If a metal has more than one oxidation number, then the number must be indicated by roman numerals to describe the charge.

Front 73

Examples of ionic compounds"

CaH2
CuCl2

Back 73

Calcium Hydroxide
Copper (II) Cloride