Chemistry Exam 3

Front 1

particles

Back 1

can be counted, have mass, can have charge

Front 2

waves

Back 2

"fuzzy" (diffuse)

Front 3

classical wave energy & amplitude relationship

Back 3

energy is proportional to amplitude squared

Front 4

electric field

Back 4

perpendicular to magnetic field

Front 5

spectroscopy

Back 5

what color is something?

Front 6

light

Back 6

electromagnetic radiation

Front 7

types of radiation, increasing wavelength, decreasing frequency

Back 7

gamma rays, x-rays, UV, visible (400-750), infrared, microwave, radio

Front 8

energy is

Back 8

discontinuous, quantized (comes in photons)

Front 9

photoelectric effect

Back 9

Heinrich Hertz 1888, electrons ejected from surface of certain metals when light strikes, frequency above certain threshold frequency, # electrons proportional to intensity of light, velocity of electrons proportional to frequency of light

Front 10

E=mc^2

Back 10

gives light properties of particle

Front 11

Bohr atom

Back 11

electrons move in circular orbits, electrons can only exist in discrete orbits (orbitals), only discrete quanta of energy can be absorbed/emitted to change an electron orbit

Front 12

Heisenberg Uncertainty Principle

Back 12

you can know where something was, but not where it is (can know either position or momentum, not both, to threshold frequency)

Front 13

Balmer series for H atom emission

Back 13

light given off when system jumps between energy levels, bigger jumps=more energy=bigger frequency=bluer, all visible light ends up at n=2

Front 14

quantum numbers

Back 14

describe wave function

Front 15

# of nodes corresponds to

Back 15

energy

Front 16

principle quantum number

Back 16

n, shell (energy)

Front 17

angular momentum quantum number

Back 17

l, shape

Front 18

magnetic quantum number

Back 18

m sub-l, orientation of orbital

Front 19

m sub-s

Back 19

electron spin

Front 20

standing waves

Back 20

nodes remain at constant position

Front 21

more nodes

Back 21

higher energy

Front 22

Schrodinger

Back 22

electrons can be described as standing waves orbiting nucleus
represented electrons with mathematical equation

Front 23

l=0

Back 23

s orbitals

Front 24

l=1

Back 24

p orbitals, smallest possible is 2p, have planar nodes (orientation along cartesian x/y/z axes)

Front 25

l=2

Back 25

d orbitals, smallest possible is 3d, 4 clover leaf shapes

Front 26

l=3

Back 26

f orbitals, smallest possible is 4f, double clover leaf shape, 7 orientations

Front 27

Pauli exclusion principle

Back 27

Each orbital can contain 2 electrons. No 2 electrons can have the same set of quantum numbers.

Front 28

Aufbau process

Back 28

electrons will always populate the lowest energy state available

Front 29

Hund's rule

Back 29

populate all degenerate orbitals (same energy) with a single electron before putting 2 electrons in an orbital (exceptions: transition metals)

Front 30

transition metal electron configurations

Back 30

half-filled or completely filled subshells have extra stability

Front 31

covalent radius

Back 31

1/2 distance between bonded nuclei

Front 32

radii of cations

Back 32

always smaller as lose electrons

Front 33

radii of anions

Back 33

bigger than neutral form

Front 34

transition metal ionization

Back 34

valence electrons come out first

Front 35

isoelectronic

Back 35

same electron configuration

Front 36

ionization potential

Back 36

energy required to ionize (eject an electron) in the gas phase, increases to the right and up PT

Front 37

electron affinity

Back 37

energy released when an electron is added to atom in the gas phase, increases to the right and up PT
group 2 will not accept electrons, more negative want electrons the most

Front 38

atomic radius increases

Back 38

to the left and down PT

Front 39

valence electrons

Back 39

outermost electrons, involved in chemical bonding

Front 40

ionic compounds

Back 40

held together by electrostatic attraction of opposite charges (coulombic attraction)

Front 41

covalent bonds

Back 41

share electrons, no charges

Front 42

Lewis structures

Back 42

symbol notation used to describe bonded/unbonded atoms

Front 43

octet rule

Back 43

most atoms are stable when there are 8 electrons in their valence shell

Front 44

electronegativity

Back 44

ability to compete for electrons in a chemical bond (max 4.0, min 0.8)

Front 45

electrons shared equally only when

Back 45

electronegativities are the same

Front 46

ionic bond difference in electronegativity

Back 46

>/= 2

Front 47

lattice energy

Back 47

energy required to completely separate a mole of solid ionic compound into individual ions in the gas phase

Front 48

lattice energy trends

Back 48

as ionic radius increases, lattice energy decreases
as charge increases, lattice energy increases
as # of ions increases, lattice energy increases

Front 49

dipole moment

Back 49

mu, measure of charge separation
units: Debye (D)

Front 50

in Lewis structures, more electronegative atoms go

Back 50

on the periphery

Front 51

formal charge

Back 51

measure of how electron rich/poor a bonded atom is if all bonds were assumed 100% covalent

Front 52

formal charge =

Back 52

# valence electrons - nonbonded electrons - 1/2 bonded electrons

Front 53

resonance structures

Back 53

single Lewis structures that in combination give accurate electron assignment for a molecule, don't change what atoms are bonded to each other, result from reorganization of electrons

Front 54

octet exceptions

Back 54

compounds with odd # of electrons, group 13 compounds (B & Al, 6 electron rule), expanded octets (never for 1st & 2nd rows)

Front 55

bond enthalpy (bond dissociation energy)

Back 55

energy required to break a bond (1 mol in gas phase) & give 1/2 electrons in bond to each atom (homolytic cleavage)

Front 56

bond enthalpies can be used to

Back 56

determine heat of reactions

Front 57

molecular dipole

Back 57

sum of bond dipoles

Front 58

can test polarity by

Back 58

putting near electric field

Front 59

valence shell electron-pair repulsion theory

Back 59

electrons that make up bonds & lone pairs repel each other and will get as far apart in space as possible

Front 60

2 electron domains molecular geometry

Back 60

linear

Front 61

3 electron domains, all atoms molecular geometry

Back 61

trigonal planar

Front 62

3 electron domains, 2 atoms & 1 electron pair molecular geometry

Back 62

bent

Front 63

4 electron domains, all atoms molecular geometry

Back 63

tetrahedral

Front 64

4 electron domains, 3 atoms & 1 electron pair molecular geometry

Back 64

pyramidal

Front 65

4 electron domains, 2 atoms & 2 electron pairs molecular geometry

Back 65

bent

Front 66

5 electron domains, all atoms molecular geometry

Back 66

trigonal bipyramidal

Front 67

5 electron domains, 4 atoms & 1 electron pair molecular geometry

Back 67

seesaw

Front 68

5 electron domains, 3 atoms & 2 electron pairs molecular geometry

Back 68

T-shape

Front 69

linear bond angles

Back 69

180

Front 70

bent (3) bond angles

Back 70

120

Front 71

tetrahedral bond angles

Back 71

109.5

Front 72

pyramidal bond angles

Back 72

107

Front 73

bent (4) bond angles

Back 73

104

Front 74

trigonal bipyramidal bond angles

Back 74

120 & 90

Front 75

lone pairs (bond angles)

Back 75

push away a little more than bonds

Front 76

5 electron domains, 2 atoms & 3 electron pairs molecular geometry

Back 76

linear

Front 77

6 electron domains, all atoms molecular geometry

Back 77

octahedral

Front 78

6 electron domains, 5 atoms & 1 electron pair molecular geometry

Back 78

square pyramidal

Front 79

6 electron pairs, 4 atoms & 2 electron pairs molecular geometry

Back 79

square planar

Front 80

octahedral bond angles

Back 80

90

Front 81

square planar electron placement

Back 81

opposite

Front 82

seesaw/T-shape electron placement

Back 82

around middle (equatorial positions)