Chem Exam 3

Front 1

4 quantities describe the state of a gas

Back 1

1. Pressure 2. Temperature 3. Volume 4. number of moles

Front 2

Boyle's law equation

Back 2

P1V1=P2V2

Front 3

Boyle's law relationships

Back 3

Increase pressure = decrease volume

Front 4

Boyle's law constants

Back 4

Moles and temperature

Front 5

Charles's law equation

Back 5

V1/T1 = V2/T2

Front 6

Charles's law relationships

Back 6

Increase temperature = increase volume

Front 7

Charles's law constants

Back 7

Pressure and moles

Front 8

How to solve for kelvin

Back 8

C=k- 273.15

Front 9

Gay lussac's law equation

Back 9

P1/T1=P2/T2

Front 10

Gay lussac's relationships

Back 10

Increase temperature = increase pressure

Front 11

Gay lussac's constants

Back 11

Volume and moles

Front 12

Combined gas law

Back 12

(P1V1/T1n1)=(P2V2/T2n2)

Front 13

Avogadro's law equation

Back 13

V1/n1=V2/n2

Front 14

Avogadro's relationships

Back 14

Increase volume = increase number of moles

Front 15

Avogadro's law constants

Back 15

Temperature and pressure

Front 16

Ideal gas law aka universal gas law

Back 16

PV=nRT

Front 17

Universal gas constant equation

Back 17

R=PV/nT

Front 18

Universal gas constant number

Back 18

R= 0.08205

Front 19

Gas density equation

Back 19

p=PM/RT

Front 20

Gas density relationships

Back 20

Increase p= increaseP

Increase p = increase M

Increase p = decrease T

Front 21

Daltons law of partial pressures

Back 21

total pressure of a gaseous mixture is a sum of the partial pressures of each of the components gases

Front 22

Daltons law relationships

Back 22

Increase pressure =increase concentration

Front 23

Daltons law equation

Back 23

Ptotal=P1+P2+Pn

Front 24

Partial pressure equals

Back 24

Ptotal x concentration of the gas

Front 25

Concentrations of N2 O2 CO2 and H2O

Back 25

78.6%

20.9%

0.04%

0.5%

Front 26

Relative humidity definition & relationship

Back 26

- saturation of water in air

- increase air solubility =increase temperature

Front 27

Operating roomrelative humidity

Back 27

20% to 60%

Front 28

RH20% lower limit because?

Back 28

Decreases shelf life and increases chance of potential spark

Front 29

Rh 60% upper limit because?

Back 29

Increases chance for mold and mildew growth and infection and less comfortable

Front 30

Four tenets that describe ideal gases

Back 30

1. Gases consist of small particles whose volume is negligible compared to the volume of the gas

2. Gas molecules are in constant, random motion

3. Gas molecules in a sample have a range of kinetic energies, but the average kinetic energy depends on temperature

4. There are no attractive or repulsive forces between gas particles, so all collisions are elastic

Front 31

Kinetic energy relationship in gases

Back 31

Increase temperature = increase kinetic energy

Front 32

Graham's law of effusion equation

Back 32

Rate A /rate B =square root of (molar mass B/molar mass A)

Front 33

Rate of effusion relationships

Back 33

Increase molecular mass = decrease rate of effusion

Front 34

Fick's law of diffusion equation

Back 34

D =(change in P x A x S)/(d x square root of MW)

Front 35

Rate of gas diffusion relationships

Back 35

- Increase concentration or pressure gradient = increase in diffusion

-increase in solubility=increase in diffusion

-increase in cross sectional area = increase in diffusion

-increase in diffusion distance = decrease in diffusion

-increase in molecular weight = decrease in diffusion

Front 36

Real gases fail at which 2 tenets

Back 36

#1: at very high pressures, gas molecules represent a large percentage of the sample volume

#4: real gases have attractive forces between molecules, kinetic energy is needed to overcome this attractionand pull molecules away from each other

- at low temps, real gases have to use greater fraction of kinetic energy to overcome attractive forces = pressure of sample decreases

Front 37

The Vander Waals equation

Back 37

PV = nRT

Front 38

Van der Waals forces include

Back 38

Attraction and repulsion between atoms, molecules, and surfaces as well as other intermolecular forces

Front 39

Van der Waals is designed to account for what

Back 39

The intermolecular forces between gas molecules and the volume of gas molecules

Front 40

Adiabatic changes

Back 40

Occur when gases are rapidly expanded or compressed without surrounding environment

Front 41

Adiabatic changes do what to systems energy

Back 41

No increase or decrease in systems energy

Front 42

Adiabatic changes do what to temperature of system?

Back 42

Temperature may change die to change in kinetic energy per area

Front 43

Energy concentration effect

Back 43

Rapid compression of gas = concentration of kinetic energy = rapid increase in temperature proportional to decrease in volume

Ex-diesel engine

Front 44

Energy dilution aka joule Thompson effect

Back 44

Rapid expansion of gas e dilution of kinetic energy = rapid decrease in temperature proportional to increase in volume

Ex: frosting on bike pump

Front 45

Which will more likely occur with compressed gas cylinders?

Back 45

Joule Thompson effect > energy concentration

Front 46

benefits of solutions

Back 46

deliver larger more precise measured volumes

better tolerated and less reactive

some meds are solid and need to be dissolved as a liquid

Front 47

homogenous mixture

Back 47

mixture that contains one or more solutes uniformly dispersed at the molecular or ionic level throughout the solution

Front 48

solute

Back 48

the substance dissolved within the solution

Front 49

solvent

Back 49

the substance that dissolves the solute

Front 50

examples of solutions

Back 50

liquids, gases, and solids

Front 51

molarity = molar concentration

Back 51

molar concentration = moles of solute per liter of solution

Front 52

molar concentration equation

Back 52

M = mol of solute/L of solution

Front 53

molar concentration depends on

Back 53

temperate of the solution, increase temperature you decrease molarity because volume increases

Front 54

Molality = molal concentration

Back 54

moles of solute per kilogram of solvent

Front 55

molality equation

Back 55

m = mol of solute/ kg of solvent

Front 56

how does temperature effect molality

Back 56

temperature has no effect on molality because it does not affect mass

Front 57

variations on percent

Back 57

1. percent by weight to volume (%w/v)

2. percent by weight (%w/w)

3. percent by volume (%v/v)

Front 58

percent by weight to volume = %w/v (actually % by mass)

Back 58

%w/v = grams of solute/ ml solution (x 100)

Front 59

percent by weight = %w/w

Back 59

%w/w = grams of solute/g of solution (x100)

Front 60

percent by volume = %v/v

Back 60

%v/v = ml of solute /ml of solution (x100)

Front 61

parts per million equation

Back 61

ppm = gram of solute /grams of solution (x10^6)

Front 62

solubility

Back 62

amount go solute that will dissolve in a given amount of solvent at room temperature

Front 63

what happens when you add solvent to an already saturated solution?

Back 63

excess solute will crystallize as a solid, separate as a liquid, or bubble out as a gas

Front 64

miscible definition

Back 64

liquids that are soluble in each other in all proportions (ex. alcohol and water)

Front 65

immiscible definition

Back 65

liquids that are not soluble in each other in all proportions (ex. oil and water)

Front 66

like dissolves?

Back 66

like

Front 67

polar solutes dissolve in what type of solvent

Back 67

polar solvents

Front 68

nonpolar solutes are more soluble in what type of solvent

Back 68

nonpolar solvents

Front 69

organic compounds are dissolved in what type of solution

Back 69

hydrochloride (HCl)

Front 70

heat of solution (aka enthalpy of solution)

Back 70

associated energy change when a solute dissolves in a solvent

Front 71

change in Hsoln =

Back 71

energy change associated with dissolving one mole of solute in a given solvent

Front 72

endothermic

Back 72

energy flows into the system

Front 73

exothermic

Back 73

energy flows out of the system

Front 74

enthalpy (heat)

Back 74

is equal to energy when the pressure remains constant

Front 75

effect of pressure on solubility (of gaseous solutes in liquid solvents)

Back 75

increase pressure = increase solubility

Front 76

effects of pressure on solids and liquids

Back 76

solids and liquids are not very compressible = pressure has little effect

Front 77

Henry's Law

Back 77

quantitative relationship between pressure and solubility

Front 78

Henrys law equation

Back 78

S = kH(aka henrys constant) x Pgas

Front 79

henrys law relationships

Back 79

increase pressure = increase solubility

Front 80

relationship between temp and solubility (solids and liquids)

Back 80

solubility of solid and liquid solutes in liquid solvents increases with increasing temperature

Front 81

effect of temp on solubility (gas and liquid)

Back 81

solubility of gaseous solutes in liquid solvents decreases with increasing temperature due to increasing vapor pressure

Front 82

colligative properties

Back 82

physical properties/characteristics of a solution that depend on the ratio of the number of solute particle to solvent particles

Front 83

4 common colligative properties

Back 83

1. vapor pressure

2. boiling point

3. freezing point

4. osmotic pressure

Front 84

vapor pressure =

Back 84

pressure exerted by a vapor (gas) in contact with its liquid or solid form

Front 85

the most energetic molecules in a liquid has sufficient what in order to do what?

Back 85

they have sufficient kinetic energy to overcome the intermolecular forces binding them into the liquid state and can escape into the gas phase

Front 86

effect of concentration on vapor pressure

Back 86

increased solute concentration = decrease vapor pressure

Front 87

Raoults Law

Back 87

qualitative relationship between vapor pressure and solute concentration

Front 88

Raoults Law equation

Back 88

P = X ( mole fraction of the substance) x P^degree (vapor pressure of the substance)

Front 89

boiling point

Back 89

the temp at which the bulk of a liquid converts to vapor at a given pressure = the temp at which the vapor pressure = atmospheric pressure

Front 90

boiling point relationships

Back 90

increase solute concentration = increase boiling point

Front 91

freezing point (aka melting point)

Back 91

temp at which solid state reversibly passes into liquid state

Front 92

freezing point relationships

Back 92

increase solute concentration = decrease freezing point

Front 93

osmosis

Back 93

diffusion of water through a semipermeable membrane

Front 94

osmotic pressure relationship with solute concentration

Back 94

increase solute concentration = increase osmotic pressure

Front 95

movement of water in osmosis goes from where to where

Back 95

water diffuses from area of high to area of low (water

or

water diffuses from area of low (solute) to high (solute)

Front 96

tonicity

Back 96

relative concentration of solutes in the osmotic system

Front 97

isotonic

Back 97

two solutions with equal concentration solute particles

Front 98

hypertonic

Back 98

solution with higher concentration solute particles

Front 99

hypotonic

Back 99

solution with lower concentration solute particle

Front 100

osmosis and diffusion explained by

Back 100

second law of thermodynamics = entropy of the universe is constantly increasing

Front 101

colloids

Back 101

one phase uniformly dispersed in a second phase

Front 102

difference between colloids and solutions

Back 102

contain large particles (not molecules or ions)

Front 103

colloid exhibit what

Back 103

tyndall effect: scatter visible light passing through them

Front 104

kinetic molecular theory of matter

Back 104

attempts to describe all states of matter and the conversion between states by

1. considering the structures of the molecules

2. how the molecules interact

Front 105

common states of matter

Back 105

solid, liquid, and gas

Front 106

kinetic energy by state of matter

Back 106

(least KE) solid

Front 107

condensed states of matter =

Back 107

solids and liquids = resist compression/not easily compressed

Front 108

Solids have

Back 108

a defined shape and volume; molecules held together by intermolecular forces

Front 109

Liquids are fluids they have

Back 109

defined volume but no defined shape, they conform to the shape of the container in which they are placed

Front 110

which has weaker forces liquids or solids?

Back 110

liquids, it allows the particles to flow past each other

Front 111

solids can do what

Back 111

vibrate in place within a limited area

Front 112

Gases = fluids =

Back 112

defined volume but no defined shape, they conform to their container (expand) in which they are placed

Front 113

intermolecular forces present in gases?

Back 113

essentially no intermolecular forces between particles, this allows particles to flow past each other

Front 114

melting

Back 114

conversion from solid to liquid

Front 115

freezing

Back 115

conversions from liquid to solid

Front 116

vaporization

Back 116

conversion of liquid to gas

Front 117

condensation

Back 117

conversions of gas to liquid

Front 118

deposition

Back 118

conversion of gas to solid (aka snow)

Front 119

sublimation

Back 119

conversion of solid to gas

Front 120

chemical bonds

Back 120

hold atoms together to form molecules

Front 121

intermolecular forces

Back 121

determine how molecules interact with each other

Front 122

Coulomb's Law

Back 122

particles (atoms, molecules, oppositely charged ions) are attracted to each other unless they get too close to each other then they repel each other

Front 123

two limiting types of chemical bonds

Back 123

ionic and covalent bonds

Front 124

octet rule

Back 124

overarching force that drives formation of chemical bonds, every atom adds, removes, or shares electrons to end up with 8 in their valence shell

Front 125

atoms with full electron shells are

Back 125

very stable (aka noble gases)

Front 126

elements with nearly full outer shells

Back 126

accept electrons and become anions (right side of the table)

Front 127

elements with emptier shells

Back 127

give away electrons and become cations (left side of the table)

Front 128

ionic compounds (aka salts)

Back 128

ions held together by ionic bonds (contain cations and anions but are electrically neutral) NOT MOLECULES!

Front 129

salts

Back 129

ionic compounds - form highly organized crystalline lattice in solid state, almost all are solid at room temp and pressure

Front 130

salts in liquid

Back 130

conduct electricity or dissolve in water

Front 131

which bonds are stronger ionic or covalent?

Back 131

ionic bonds are stronger

Front 132

covalent bonds result

Back 132

from sharing one or more pairs of electrons

Front 133

valence bond theory

Back 133

the presence of electrons between the nuclei shield them from each other = reduces coulombic repulsive forces

Front 134

shared electrons are able to orbit either nuclei

Back 134

= increases the electron density between the two nuclei; this stabilizes the molecule by reducing overall energy

Front 135

decreased energy =

Back 135

increased stability when forming covalent bonds

Front 136

there is an ideal bond length or distance where

Back 136

the energy is minimum and stability is maximum

Front 137

energy increases if

Back 137

distance increases or decreases

Front 138

bond dissociation energy

Back 138

amount of energy needed to break the covalent bond

Front 139

strength of bonds vary

Back 139

stronger bonds are harder to break and are more stable

Front 140

electronegativity

Back 140

atom's propensity for pulling electrons toward itself

Front 141

most electronegative element

Back 141

Fluorine (F); the closer to fluorine an element is on the table the more electronegative it is

Front 142

non polar covalent bonds

Back 142

bond between two atoms with the same electronegativity

Front 143

non polar covalent bond electrons are

Back 143

evenly shared, and have no areas of charge, hemodiatomic molecules are always non polar (aka H2 or O2)

Front 144

polar covalent bonds

Back 144

bond between two atoms with different electronegativity

Front 145

increased differences in electronegativity =

Back 145

increasingly polar covalent bond

Front 146

the more electronegative atom =

Back 146

the more pull it has, electrons spend more time here, will have partial negative charge and other atoms will have partial positive charge

Front 147

intermolecular forces (4)

Back 147

1. dipole-dipole attraction

2. hydrogen bonding

3. London forces

4. ion-dipole interactions

Front 148

dipole-dipole attraction

Back 148

attraction between the opposite partial charges (poles) of polar molecules

Front 149

hydrogen bonding

Back 149

only possible when hydrogen is bonded to oxygen, nitrogen, or fluorine; H becomes positive charge and all partial negative charges are highly attracted to it

Front 150

which bond is very important for intermolecular forces in life?

Back 150

hydrogen bonds

Front 151

London Forces (weakest intermolecular force)

Back 151

instantaneous dipole created whenever electrons in an atom or molecule are unevenly distributed

Front 152

what increases an atoms London forces?

Back 152

size, the larger the atom the more London force occurs

Front 153

ion-dipole attraction

Back 153

occurs between ions and polar molecules

Front 154

ion-dipole attraction strength depends on

Back 154

the strength of the charges of the poles and the size of the ion

Front 155

ion-dipole attraction allows

Back 155

ionic solids to dissolve in water, NaCl for example

Front 156

polar compounds dissolve in

Back 156

polar compounds

Front 157

non polar compounds are insoluble in

Back 157

polar compounds

Front 158

non polar compounds dissolve easily in

Back 158

nonpolar compounds

Front 159

ionic compounds dissolve easily in

Back 159

polar solvents

Front 160

surface tension

Back 160

molecules near or on the surface are subject to surrounding forces and forces from below = unbalanced force, which creates "skin" on the surface of the liquid

Front 161

substances with increased intermolecular attractions have

Back 161

increased surface tensions

Front 162

if somethings density/weight/force is less than the surface tension of water

Back 162

it will not sink in water

Front 163

water is very cohesive liquid =

Back 163

water molecules have a strong attraction for other water molecules, causes water to form spherical drops

Front 164

surface tensions does what in the lungs

Back 164

causes collapse of alveoli in the absence of surfactant in the lungs of premature infants

Front 165

capillary rise

Back 165

caused by surface tension = tendency of a fluid to rise in a narrow diameter tube

Front 166

LaPlace's Law

Back 166

describes relationship between wall tension to pressure and the radius in cylinders and spheres

Front 167

LaPlace's Law equations

Back 167

cylinders = T = Pxr or P = T/r

spheres = 2T = Pr pr P = 2T/r

Front 168

LaPlace's Law relationships

Back 168

increased tension = increased pressure

increased radius = increased tension

increased radius = decreased pressure

Front 169

Cylinder and sphere examples

Back 169

cylinders = blood vessels and aortic aneurysms

spheres = saccular aneurysms and alveoli and cardiac chambers

Front 170

Law of Laplace of hollow viscera

Back 170

P = 2T/R where P is pressure required to keep the viscera expanded, and T is surface tension and R is radius; surfactant decreases surface tension

Front 171

surfactants

Back 171

reduce the surface tension of a liquid and increase the ability of the liquid to function as a solvent

Front 172

surfactants are commonly known as

Back 172

detergents or soaps, have a polar head and non polar tail; forming monolayers, bilayers, and micelles

Front 173

monolayers

Back 173

polar head goes into water and non polar tail sticks out of or sits on surface of water

Front 174

surfactant molecules get between the water molecules and disrupt their hydrogen bonding =

Back 174

decreased surface tension of water, aka water is less cohesive

Front 175

surfactant in lungs creates

Back 175

monolayer in the alveoli to keep from collapsing and to equilibrate the pressure needed to inflate alveoli of diff sizes

Front 176

bilayers

Back 176

tails of surfactant molecules dissolve into each other to form a double layer

Front 177

bilayer set up

Back 177

non polar tails form middle of the bilayer and polar heads form outside surfaces of bilayer; preventing polar compounds from penetrating the bilayer without carrier proteins

Front 178

Micelles

Back 178

nonpolar tails dissolve in each other forming a spherical structure; how soaps and detergent work; creates non polar microenvironment in water

Front 179

viscosity =

Back 179

measure of a fluids resistance to flow

Front 180

fluids with high viscosity

Back 180

have increased resistance to flow

Front 181

viscosity increases

Back 181

as intermolecular forces increase

Front 182

vaporization requires

Back 182

energy

Front 183

increased temp =

Back 183

increased vapor pressure

Front 184

volatility =

Back 184

tendency of a liquid to evaporate

Front 185

increased volatility =

Back 185

increased vapor pressure

Front 186

dynamic equilibrium =

Back 186

system comprised of at least 2 states when the populations of the states are constant; in a closed system rates of vaporization and condensation are equal; equilibrium is achieved

Front 187

substances with greater intermolecular forces have

Back 187

lower vapor pressure

Front 188

molecules with increased intermolecular forces are

Back 188

more tightly bond in the liquid state, cannot escape not enough kinetic energy

Front 189

volatile liquids =

Back 189

liquids that have a high vapor pressure at room temp (all our volatile anesthetics)

Front 190

if pressurized anesthetic gases flow through a closed container holding a liquid volatile anesthetic

Back 190

the anesthetic gases will mix with the gaseous portion of the volatile anesthetic and exit the vaporizer

Front 191

the temp of the vaporizer and the pressure of the anesthetic gas determine

Back 191

how much volatile anesthetic in its gas phase exits the vaporizer (this is also determined by selecting a setting on the vaporizer); each agent has different vapor pressure so must be calibrated for that specific agent or can under or over dose

Front 192

boiling point =

Back 192

the temp at which the bulk of a liquid converts to vapor at a given pressure; temp at which the vapor pressure is equal to atmospheric pressure

Front 193

the temp will never rise above boiling point

Back 193

the rest of additional heat energy transforms the liquid into the gas

Front 194

the heat energy is transformed into molecules as

Back 194

kinetic energy = molecules are able to escape liquid phase and move into gas

Front 195

increased intermolecular force =

Back 195

increased boiling point

Front 196

boiling point increases as

Back 196

pressure increases

(how pressure cooker work)

Front 197

melting point =

Back 197

temp at which a solid state reversibly passes into liquid state (also known as freezing point)

Front 198

melting point temp is the temp at which average kinetic energy is

Back 198

sufficient to overcome intermolecular forces that hold the molecules in the solid state

Front 199

increased intermolecular forces do what to melting point

Back 199

increased melting point