Chem 400 Exam 3

Front 1

Coulomb's Law

Back 1

The electrostatic force between two charges separated by a distance is FΩ(QeQn)/r^2

Front 2

Effective Nuclear Charge

Back 2

Qneff ~ Qn - (# of shielding electrons)

Front 3

Qe

Back 3

-1

Front 4

Qn

Back 4

Charge of the elements nucleus

Front 5

Atomic Radius Trend

Back 5

Increases as you down and left.

Front 6

Explanation of Atomic Radius Trend

Back 6

Valence shell decreases/stays the same.

Core increases/stays the same

Qneff decreases/stays the same

Qneff decrease had bigger affect than valence shell decrease(left)

When Qneff stays so does valence, core increases size(top to bottom)

Front 7

Ionization Energy

Back 7

The energy required to remove the outermost (highest energy) electron from an atom in the gaseous state

Front 8

Ionization Energy Trend

Back 8

Increases moving right and up.

Front 9

Ionization Energy Trend Explanation

Back 9

Atomic Radius decreases

Qneff increases(right)/constant(up)

Front 10

Third Ionization Energy

Back 10

Very large increase in IE because the atom is smaller (r smaller) and Qneff is larger

Front 11

Electron Affinity

Back 11

The energy change when an electron is added to an atom in the gaseous state.

Front 12

Negative Electron Affinity

Back 12

The ion is more stable than the atom.

Front 13

Positive Electron Affinity

Back 13

The ion is unstable and doesn't form spontaneously.

Front 14

Negative Electron Affinity Tend

Back 14

An electron is added into a partially filled subshell.

Front 15

Positive Electron Affinity Trend

Back 15

An electron is added to an element that has filed subshells.

Front 16

Anion Atomic Size

Back 16

Larger than there atoms

Front 17

Cation Atomic Size

Back 17

Smaller than there atom.

Front 18

Ionic Bond

Back 18

One or more electrons is permanently transferred from one atom to another.

Front 19

Lattice Energy

Back 19

The energy reqired to create a mole of ionic solid from its constituent ions in the gas phase. Inferred from the force of attraction

Front 20

Force of Attraction

Back 20

FΩ(Q1Q2)/r^2

Front 21

Comparing Lattice Energies

Back 21

The size affects r, the charge affects Q1,Q2

Front 22

Properties of Ionic Materials

Back 22

Very high melting points

Hard but brittle

Conducts electricity: never in solid, always in aqueous solution

Front 23

Covalent Charge

Back 23

When 2 nonmetal atoms share electrons.

Front 24

Isoelectric

Back 24

Has the same number of electrons

Front 25

Electronegativity

Back 25

The ability of an atom to attract the electrons in a chemical bond.

Front 26

Electronegativity Trend

Back 26

Increases right and up, same reasons as ionization energy.

Front 27

Rules for Drawing Electron Dot Formulas

Back 27

1. Calculate the total number of valence electrons

2. Draw the skeleton structure

3. Distribute the remaining v.e. amongst the outer atoms to complete there octets

4. Place any remaining v.e. on the central atom

Front 28

Delocalized Bonding

Back 28

The electrons in the double bond are spread over two bonds

Front 29

Resonance Form

Back 29

Each possible lewis structure connected by double-headed arrows.

Same atoms arranged in the same way but with double or triple bonds in different places.

Front 30

Bond Order

Back 30

Indicates the extent of multiple bonding.

Single bond: bond order=1

Triple bond: bond order= 3

Front 31

Polyatomic Ion

Back 31

Molecule with a charge

Front 32

Formal Charge

Back 32

(Original # of valence electrons) - ( # of electrons in lone pairs) - (# of bonds)

Front 33

Linear

Back 33

2 electron groups

Front 34

Trigonal Planar

Back 34

3 electron groups

Front 35

Tetrahedral

Back 35

4 electron groups

Front 36

Trigonal Bipyramidal

Back 36

5 electron groups

Front 37

Octahedral

Back 37

6 electron groups

Front 38

AX2E

Back 38

Bent

Front 39

A3E

Back 39

Trigonal Pyramidal

Front 40

AX2E2

Back 40

Bent

Front 41

AX3E2

Back 41

T shaped

Front 42

AX4E

Back 42

See saw

Front 43

AX2E3

Back 43

Linear

Front 44

AX5E

Back 44

Square pyramidal

Front 45

AX4E2

Back 45

Square planar

Front 46

Hybridization

Back 46

The blending of occupied atomic orbitals to make the same number of hybrid orbitals

Front 47

Valence Bond Theory Steps

Back 47

1. Draw the lewis structure

2. Calculate the formal charge on the atom toy are interested in

3. Draw ground state orbital diagram

4. Promote electrons

5. Keep one p-orbital unhybridized for each pi-bond attached

6. Hybridizevall remaining occupied orbitals in the valence shell

Front 48

Rotational Isomers

Back 48

Molecules with the same atoms, but in different arrangements.

Front 49

Phase Diagram

Back 49

Shows the phase (state) of a substance as a function of pressure and temperature.

Front 50

Triple Point

Back 50

The point at which all three phases are in equilibrium (all present in the same place at the same time)

Front 51

Critical Point

Back 51

The end of the gas-liquid phase transition line

Front 52

Supercritical Fluid

Back 52

No distinction between liquid and gas.

Front 53

Attractive Forces

Back 53

Holding together particles that are weakly attached. Not much energy is required to separate particles. They break apart and reattach often under normal conditions.

Front 54

Bond

Back 54

Ionic and covalent attractions between atoms or ions.

Front 55

Chemical Bonds

Back 55

Ionic and covalent bonds

Front 56

Intermolecular Forces

Back 56

Attractive forces between molecules.

Front 57

Temporary Dipoles

Back 57

Form due to random motion of 'electron clouds'

Front 58

London Dispersion Forces

Back 58

The attraction between partial charges (temporary dipole and an induced dipole)

Front 59

Dipole-Dipole Forces

Back 59

Attraction between permanent dipoles.

Front 60

H-Bonds

Back 60

The strong electrostatic attraction between a very Ω+ H-atom and a 'concentrated' lone pair of electrons (Ω-) on an O, N, or F-atom.

Front 61

Dipole-Induced Dipole Forces

Back 61

A polar molecule (permanent dipole) can induce a temporary dipolein a nonpolar atom/molecule.

Front 62

Ion-Induced Dipole Forces

Back 62

An ion induces a temporary dipole in a nonpolar atom/molecule.

Front 63

Ion-Dipole Forces

Back 63

Attractive force between an ion and a molecule with a permanent dipole.

Front 64

Vapor Pressure

Back 64

The pressure of the gaseous vapor above a liquid.

Front 65

Boiling

Back 65

When the vapor pressure is equal to the atmospheric pressure.

Front 66

Surface Tension

Back 66

The energy used to create new surface area (J/m^2). The energy needed to seperate the molecules.

Front 67

Capillary

Back 67

When a liquid rises up into a narrow space. Occurs when there is a stronger attraction between liquid molecules and the sides of the glad tube than between the lewis molecules themselves.

Front 68

Viscosity

Back 68

How 'thick' or 'sticky' a liquid is.

Front 69

Solution

Back 69

Solute dissolved in a solvent.

Front 70

Molarity

Back 70

(Moles of solute)/(liters of solution)

Front 71

Molality

Back 71

(Moles of solute)/(kg of solution)

Front 72

Mass percent

Back 72

(Mass solute)/((mass solute)+(mass solvent)) × 100%

Front 73

Volume Percent

Back 73

(Volume of solute)/(volume of solution) × 100%

Front 74

Mole Fraction

Back 74

(Moles of A)/(total moles in mixture)

Front 75

^Hsolvent

Back 75

The heat required to seperate the solvent molecules so that the solute can be added - endothermic

Front 76

^Hsolute

Back 76

The heat required to separate the solute molecules so that they can be dispersed evenly throughout the solvent - endothermic

Front 77

^Hmix

Back 77

The heat produced by the attraction between the solvent and solute molecules/ions (bond formation) - exothermic

Front 78

^Hsolution

Back 78

The overall enthalpy change when a solute is dissolved in a solvent - tge solute molecules are separated and dispersed homogeneously throughout the solvent.

Front 79

Exothermic

Back 79

Enthalpy decreases

Front 80

Endothermic

Back 80

Enthalpy increases

Front 81

Gibbs Free Energy

Back 81

Thermodynamic potential

Front 82

Entropy

Back 82

The something else that turns H until G includes a measure of how disordered the system is.

Front 83

Why Endothermic Reactions Happen

Back 83

The product (the solution) is more disordered than the reactants (the solid and the solvent) - The system's entropy increases. Increase in entropy makes a system a little more stable. This added stability can compensate for the loss of stability that accompanies an endothermic process.

Front 84

Polar Solvent

Back 84

A liquid composed of polar molecule.

Front 85

Nonpolar Solvent

Back 85

A liquid composed of nonpolar molecules.

Front 86

Henry's Law

Back 86

Gas solubility = kH × Pgas

((gas solubility 1)/(gas solubility2)) = ((P1)/(P2))

Front 87

Solubility

Back 87

The maximum amount of solute that can be dissolved in a solvent.

Front 88

Colligative Properties of a Solution

Back 88

Vapor pressure lowering

Freezing point depression

Boiling point elevation

Osmotic pressure

Front 89

Van't Hoff Factor

Back 89

The number of moles of particles produced by 1 mole of solute.

Front 90

Van't Hoff Factor Equation

Back 90

i = (moles of particles in solution)/(moles of solute in solution)

Front 91

Raoult's Law

Back 91

When a solute, B, is added to a solvent, A, the vapor pressure of the solvent can be found using the mole fraction of the solvent, XA, in the solution;

PA=P°A×XA

PA - the new vapor pressure of the solvent in the solution

P°A - the vapor pressure of the pure solvent.

Front 92

Freezing Point Depression

Back 92

^Tf=i×Kf×cm

^Tf - the difference between the freezing point of the pure solvent and that of the mixture.

i - van't Hoff factor

cm - molal concentration of the solute (mol/kg)

Front 93

Boiling Point Elevation

Back 93

^Tf=i×Kb×cm

Kb - the boiling poit elevation constant.