Iu C117 Chemistry Exam 2

Front 1

bondings groups: 2

lone pairs: 0

Back 1

Electron geometry: Linear

Molecular geometry: Linear

Bond Angles:180

Front 2

bondings groups: 3

lone pairs:0

Back 2

Electron geometry: Trigonal planar

Molecular geometry: Trigonal planar

Bond Angles: 120

Front 3

bondings groups: 2

lone pairs:1

Back 3

Electron geometry:Trigonal planar

Molecular geometry: bent

Bond Angles: < 120

Front 4

bondings groups:4

lone pairs:0

Back 4

Electron geometry:tetrahedral

Molecular geometry: tetrahedral

Bond Angles: 109.5

Front 5

bondings groups: 3

lone pairs: 1

Back 5

Electron geometry: tetrahedral

Molecular geometry: trigonal planar

Bond Angles: <109.5

Front 6

bondings groups:2

lone pairs:2

Back 6

Electron geometry: tetrahedral

Molecular geometry: bent

Bond Angles: <109.5

Front 7

bondings groups:5

lone pairs:0

Back 7

Electron geometry: trigonal bipyramidal

Molecular geometry: trigonal bipyramidal

Bond Angles: 120/90

Front 8

bondings groups:4

lone pairs:1

Back 8

Electron geometry: trigonal bipyramidal

Molecular geometry: seesaw

Bond Angles:<120 & <90

Front 9

bondings groups:3

lone pairs:2

Back 9

Electron geometry:trigonal bipyramidal

Molecular geometry:t-shaped

Bond Angles: <90

Front 10

bondings groups:2

lone pairs: 3

Back 10

Electron geometry:trigonal bipyramidal

Molecular geometry:linear

Bond Angles:

Front 11

bondings groups:6

lone pairs:0

Back 11

Electron geometry:octahedral

Molecular geometry: octrahedral

Bond Angles:90

Front 12

bondings groups:5

lone pairs:1

Back 12

Electron geometry: octahedral

Molecular geometry:square pyramidal

Bond Angles: <90

Front 13

How do you determine the number of valence electrons for any given atom?

Back 13

Look at unfilled outermost shell.

Based on group number for main group elements

Front 14

How do you distinguish among ionic bond, non polar covalent bond, polar covalent bond, and metallic bond?

Back 14

Electronegativity Difference:

Nonpolar covalent: 0 -.4

Polar covalent: .4 - 2.

Ionic- 2--> +

Front 15

How do you calculate lattice energy for an ionic compound based on use of a Born-Harber Cycle?

Back 15

Look at the delta H for each of the following steps:

1. Sublimation of a metal

2. Form an atom from molecule

3. Ionization to form a cation

4. Electron affinity to form anion

5. Find lattice energy

Front 16

How do you predict which compound has the largest lattice energy around a set of ionic compounds?

Back 16

Coulomb's Law

F= q1 *q2/r

Largest LE has smallest atomic radii, bigger charge makes it larger; Smallest LE has largest atomic radii

Atomic Radius increases going down and decreases across period.

Front 17

What is the difference between electron affinity and electronegativity?

Back 17

Electronegativity: The ability of an atom in a molecule to attract electrons to itself. Electron affinity: The energy change for M(g) + e- M-(g). EA deals with isolated atoms in the gas phase.

Front 18

How do you draw Lewis Dot Structures for inorganic and organic molecules and ions?

Back 18

Draw atoms, draw valence electrons, form bonds.

*carbon doesn't want lone pairs, Free Radicals- molecules and ions with an odd number of electrons, Incomplete Octets- Boron & Beryllium, Expanded octets-Elements in the 3rd row of the periodic table an beyond often have expand octets up to 12-14 electrons

Front 19

How do you draw resonance structures for a compound or ion & predict which resonance form is preferred?

Back 19

If double or tripple bonds are present, there are probably resonance structures. Look at formal charge, want 0 or negative formal charge. If a positive charge, put it on the most electronegative atom.

Front 20

How do you calculate formal charges for all atoms of a compound or ion?

Back 20

Formal charge= Valence - # of pairs- # unpaired

Front 21

What are the exceptions to the octet rule?

Back 21

Free Radicals- molecules and ions with an odd number of electrons in their Lewis structures

Incomplete Octets-Boron & Beryllium Expanded octets-Elements in the 3rd row of the periodic table an beyond often have expand octets up to 12-14 electrons

Front 22

s orbital

Back 22

Front 23

p orbital

Back 23

Front 24

d orbital

Back 24

Front 25

f orbital

Back 25

Front 26

sp hybrid orbital

Back 26

Front 27

sp3d hybrid orbital

Back 27

Front 28

How do you use wedges to show dimensions of molecular geometry?

Back 28

Front 29

How do you use molecular geometry to determine polarity?

Back 29

Lone pairs cause polarity, look at vector sums

Front 30

How do you tell which compound has the largest dipole moment?

Back 30

dipole moment= length of bond x charge on each atom

tripple are shorter than double which are shorter than single

Front 31

How do you determine hybridization and bonding schemes in a molecule?

Back 31

Same number of bonds

Front 32

How do you identify sigma and pi bonds and which orbital or hybridized orbitals form each bond?

Back 32

* pi bond- p orbitals overlap side by side
* sigma bond- when orbitals overlap end to end
* double bonds- has one pi bond and one sigma bond
* single bond has one sigma bond
* A double bond in the lewis model always corresponds to one sigma and pi bond in valence theory
* pi bonds are weaker than sigma bonds because the side to side orbital overlap tends to be less efficient than the end to end orbital overlap
* pi bonds are easier to break the the sigma bond
* pi bonds don’t rotate,sigma bonds do rotate

Front 33

How do you determine whether a molecule has dipole-dipole forces?

Back 33

All polar molecules have dipole-dipole forces.

Front 34

How do you interpret an electrostatic potential map of a molecule?

Back 34

The redder the area, the more electronegative it is. Size is dependent on atomic radius.

Front 35

How do you determine whether a molecule engages in hydrogen bonding?

Back 35

Hydrogen must be bonding to Oxygen, Nitrogen, or Florine. Polar bond.

Front 36

How do you predict relative boiling points of molecules?

Back 36

The more cohesive the forces, the higher the boiling point.

ion ion > ion dipole > hydrogen > dipole > dispersion

Front 37

Sublimation

Back 37

transition from solid to gas

Front 38

depostition

Back 38

gas to solid

Front 39

melting

Back 39

solid to liquid

Front 40

freezing

Back 40

liquid to solid

Front 41

vaporizing

Back 41

liquid to gas

Front 42

condensation

Back 42

gas to liquid

Front 43

Miscibility

Back 43

Like dissolves like, polar-polar, non-polar-nonpolar

Front 44

Properties of solid

Back 44

definite shape, volume

Front 45

Properties of liquid

Back 45

definite volume, indefinite shape

Front 46

Properties of a gas

Back 46

indefinite shape and volume

Front 47

How do you predict surface tension, viscosity, and volatility?

Back 47

Stronger force: more surface tension, more viscous, weaker the force more volatile.

Front 48

Heat of vaporization

Back 48

the amount of heat required to vaporize one mole of a liquid to gas

Front 49

heat of fusion

Back 49

amount of heat required to melt 1 mole of a solid

Front 50

Dynamic equilibrium

Back 50

when the rate of condensation and evaporation become equal

Front 51

Heating curves

Back 51

Heating curves aren't curvy.

Front 52

Using heat of vaporization in calcuations

Back 52

convert kJ to moles to grams of water

Front 53

Vapor pressure and its dependence on temperature.

Back 53

Temperature rises, vapor pressure rises

Front 54

How do you determine whether a solute is soluble in a solvent?

Back 54

Like dissolves like

Front 55

Predict whether a molecule would be more soluble in water or in hexane.

Back 55

Polar vs nonpolar

Front 56

Entropy increase in solution

Back 56

because the potential for randomness is increased

Front 57

What factors improve solubility in water?

Back 57

Solids- increase temp

Gases & liquids- decrease temp.

Front 58

Enthalpy change upon solution formation:

Back 58

Add the enthalpy change for each step

1. Separation of solute particles

2. Separation of solvent particles

3. mixing of the solute and solvent particles

(first 2 endothermic, last one exothermic)

*in aqueous ionic solution its always exothermic