Chem Final

Front 1

What is molecular geometry?

Back 1

general shape of a molecule. the five types are linear, trigonal planar, tetrahedral, trigonal bypiramidal, and octahedral.

Front 2

What is the VSEPR model?

Back 2

Valence-shell electron-repulsion. States that shapes of molecules and ions are predicted by the fact that electrons prefer to be kept as far away from each other.

Front 3

IF an atom has two pairs of electrons...

Back 3

Its is linear. the atoms are at 180 degrees to each other

Front 4

If an atom has three pairs of electrons..

Back 4

It is trigonal planar, and the atoms are at 120 degrees to each other. However, if there is one lone pair, the molecular geometry is bent, and the angle is slightly less than 120.

Front 5

If an atom has four pairs of electrons

Back 5

then the arrangement is tetrahedral. If there is no lone pairs, then molecular geometry is tetrahedral as well and the angles are 109.5. However, with one lone pair, the geometry is trigonal pyramidal, with two lone pairs, it is bent.

Front 6

If an atom has 5 pairs of electrons

Back 6

it is trigonal bipyramidal by arrangement. If there are no lone pairs, then the geometry is trigonal bipyramidal as well. If there is one, then it is Seesaw, if there is two, then it is Tshaped, and if there is three, it is linear. The angles are: 90 for axial, and 120 for those between the equatorial atoms.

Front 7

If an atom has 6 pairs

Back 7

it is octahedral. All angles are 90. if no lone pairs- octahedral geometry. If one lone pair- square pyramidal. If 2- sqaure planar.

Front 8

Steps of prediction using VSEPR

Back 8

1) electron dot formula from molecular formula
2) determine number of electron pairs around central atom
3) determine arrangement
4) getermine geometry.

Front 9

what affects bond angles and how?

Back 9

lone pairs- they require more space, therefore, the angles between otehr atoms are less than what predicted
multiple bonds also require more space because they contain greater number of electrons

Front 10

What is a dipole moment?

Back 10

it is a quantitative measure of the degree of charge separation in a molecule. measured in debyes. Points from a positive partial charge to negative. Linear, trigonal planar, and tetrahedrtal give 0 dipole.

Front 11

Valence Bond Theory

Back 11

attempts to explain the electron pair or covalent bond by quantum mechanics.
1) an orbital of 1 atom comes to occupy a portion of the same region of space as an orbital on the other atom.= overlap
2) no more than 2e per orbital.
orbitals bond in the direction in which they protrude, or point.

Front 12

Hyprid orbitals.

Back 12

orbitals used to explain bonding=> take combinations of other orbitals. Number of hybrid orbitals formed = number of orbitals used.

Front 13

To obtain a bonding description:

Back 13

1) lewis electron dot
2) use VSEPR
3)# of hybrid orbitals.
4)assign valence electrons to hybridized orbitals
5) overlap orbitals

Front 14

Multiple Bonding

Back 14

one hybrid orbital needed for multiple bonds and lone pairs as well.
sigma and pi bonds

Front 15

sigma vs. pi

Back 15

sigma- cylindrical shape about axis- formed when two s orbitals, or s and p overlap. along axis
PI- distribution above and below axis. formed by sideways overlap of 2 p orbitals.

Front 16

double bonds

Back 16

one pi, one sigma. one unhybridized 2p

Front 17

triple bond

Back 17

one sigma, tWO pi. two unhybridized 2P's

Front 18

Molecular Theory

Back 18

electrons put into orbitals with least energy.
binding vs. antibonding

Front 19

bonding

Back 19

where electrons are concentrated

Front 20

antibonding

Back 20

where electrons spend less time.

Front 21

order of molecular orbitals

Back 21

see book.

Front 22

Bond Order

Back 22

.5(#e bonding-#e antibonding)

Front 23

diamagnetic vs. paramagnetic

Back 23

dia- substance not attracted by magnetic field- only paired e
para- slightly attracted

Front 24

Energy involved in Ionic bonding

Back 24

can be described by Coulomb's law.
E= kQQ/r

Front 25

lattice energy

Back 25

cahnge in energy wen idonic solid spearates into isolated ions.

Front 26

Main-group ions

Back 26

losing electrons requires more energy-> IIIA elements show less tendency to form ionic compuonds.

Front 27

Ionic radius

Back 27

decreases in period, increases in group. if isoelectronic, go by proton #

Front 28

Bond Length

Back 28

distance between nucleus and point of minimum energy

Front 29

Coordinate covalent bond

Back 29

wehre both electrosn are donated by one atom:
NH3+

Front 30

Polar Covalent Bonds and electronegativity

Back 30

Measure of how able the atom is to draw bonding electrons to itself. I.E-E.A/2
Electronegativity increases in period, decreases in group.

Front 31

Which Atoms Central?

Back 31

the one that is least electronegative?

Front 32

Oxoacids

Back 32

substances in which O atoms are bonded to a central atom, with one or more H atoms usally bonded to O.
PS. Metals forms basis, Nonmentals form Acids.

Front 33

Exceptions to the OCtet rule

Back 33

atoms that have d orbitals
1)calculate total # of e for molecule
2)write skeleton structure
3)distribute electorns to the atoms surrounding the central
4)distribute the remaining electrons as pairs to the central atom. If there are fewer than 8 electrons, suggests that multiple bond is present.

Front 34

Formal charge

Back 34

addition equals charge on the molecule

Front 35

Building up principle

Back 35

do the triangle all the way up to f

Front 36

Periodic law

Back 36

when elements are arranged according to atomic number, physical and chemical properties vary perioducally

Front 37

effective nuclear charge

Back 37

positive charge electron experiences from nucleus.

Front 38

Ionization energy

Back 38

minimum energy needed to remove the highest energy electron from enutral atom. increases with increasing atomic number. It is harder to remove electrons if one must break the noble gas core.

Front 39

Electron Affinity

Back 39

energy change for the process of adding an electron to a neutral atom in the gaseous state to form a negative ion. If large, then the neutral atom is stable. increases in period

Front 40

Basic Oxide

Back 40

an oxide that reacts with acids- metals

Front 41

Acidic Oxide

Back 41

oxide that reacts with bases- nonmetals

Front 42

Speed of light

Back 42

3*10^8m/s. c=v(frequency)* lambda

Front 43

PLank's Quantization of Energy

Back 43

E=h*frequency
Energy of a photon

Front 44

Bohr

Back 44

E=-Rh/n^2- energy at a principle quantum number
hv=-deltaE=-(Ef-Ei)
hv=Rh(1/nf^2-1/ni^2)

Front 45

de BroLIE

Back 45

wawelength(lambda)=h/mv

Front 46

Uncertainty Principle

Back 46

(delta x)(delta p)>= h/4pi
uncertainty in position
p=momentum=mv- can interchange

Front 47

quantum numbers

Back 47

n=principal energy level
l=angular momentum- 0,1,2,3. n-1
ml=magnetic Q#- from -to+l
ms=spin. +/- 1/2

Front 48

Ionization energy

Back 48

increases from first ionization energy to second, etc.. getting closer to nucleus

Front 49

Electrolyte

Back 49

dissolves in water- a strong one exists in solution almost entirely as ion. a weak one- doesnt dissolve readily.

Front 50

Nonelectrolyte

Back 50

doesnt dissolve readily

Front 51

Solubility rules.

Back 51

group IA and ammonium are soluble
Acetates and nitrates are soluble
Most chlorides, bromides, and iodides anre soluble. excepts those with Ag, Hg, and Pb.
Most sulfates are soluble- except Ca, Sr, and Ba, and Ag, Hg, Pb
Most carbonates are insoluble. except group IA and ammonium
Most Phosphates, except IA and nh4+
MOst sulfides are insuluble, except as above
most hydroxides are insoluble, except for above, Ax, Sr, Ba

Front 52

Types of chemical reactions

Back 52

precipitation, acid base, oxidation reduction

Front 53

Acid

Back 53

species that donates a proton (H+) to another species.

Front 54

Base

Back 54

one that accepts proton

Front 55

Strong Acids

Back 55

HCLO4
H2SO4
HI
Hbr
HCl
HNO3

Front 56

Strong Base

Back 56

LiOH
NaOH
KOH
Ca(OH)2
Sr(OH)2
Ba(OH)2

Front 57

prefixes 1-10

Back 57

mono, di, tri, tetra, penta, hexa, hepta, octa, nona, deca.

Front 58

Acid naming

Back 58

if ate-ic
if ite-ous

Front 59

Carbonate Ion

Back 59

CO3(2-),

Front 60

Nitrite

Back 60

NO2-

Front 61

Nitrate

Back 61

NO3-

Front 62

Phostate

Back 62

PO4(3-)

Front 63

Sulfite

Back 63

SO3, 2-

Front 64

Sufate

Back 64

SO4, 2-

Front 65

Hypochlorite

Back 65

ClO-

Front 66

Chlorite

Back 66

ClO2-

Front 67

Chlorate

Back 67

ClO3-

Front 68

Perchlorate

Back 68

ClO4-

Front 69

Ammonium

Back 69

NH4+

Front 70

Cyanide

Back 70

CN-

Front 71

Carbonate

Back 71

CO3,2-

Front 72

Bicarbonate

Back 72

HCO3,-

Front 73

Acetate

Back 73

C2H3O2, 2-

Front 74

Oxalate

Back 74

C2O4,2-

Front 75

Chromate

Back 75

CrO4,2-

Front 76

Dichromate

Back 76

Cr2O7, 2-

Front 77

Permanganate

Back 77

MnO4-