Chem Exam 4 Ch 8 9 10

Front 1

matter exists in what 3 phases

Back 1

solid liquid and gas

Front 2

the state in which a compound exists depends on what

Back 2

the relative strength of the attractive forces between particles compared to the kinetic energy of the particles

Front 3

the attractive forces between particles in gases are

Back 3

very weak compared to their kinetic energy so particles move about more freely, are farther apart and have almost no influence on one another

Front 4

the attractive forces between particles in liquids are

Back 4

stronger, pulling particles close together but still allowing them considerable freedom to move about

Front 5

the attractive forces between particles in solids are

Back 5

much stronger that their kinetic energy of the particles, so the atoms/molecules/ions are held in a specific arrangement and can only wiggle around in one place

Front 6

changes of state (phase change)

Back 6

the change or transformation of a substance from one state of matter to another

Front 7

every change of state is ___ and characterized by what

Back 7

reversible and characterized by a free energy change, ∆G

Front 8

a change of state that is spontaneous in one direction (exergonic, negative ∆G) is

Back 8

non spontaneous in the other direction (endergonic, positive ∆G)

Front 9

formula for free energy change ∆G

Back 9

∆G= ∆H (enthalpy change) - T(in Kelvins) ∆S (entropy change)

Front 10

enthalpy change ∆H

Back 10

a measure of heat absorbed or released during a given change of state.

Front 11

when heat is absorbed, ∆H is

Back 11

positive- endothermic

Front 12

when heat is released, ∆H is

Back 12

negative- exothermic

Front 13

entropy change ∆S

Back 13

a measure of the change in molecular disorder or freedom that occurs during a process

Front 14

what happens to ∆S in the melting of a solid to a liquid

Back 14

disorder increases, ∆S is positive

Front 15

what happens to ∆S in the freezing of a liquid to a solid

Back 15

disorder decreases, ∆S is negative

Front 16

melting point (mp)

Back 16

the exact temperature at which solid and liquid are in equilibrium

Front 17

boiling point (bp)

Back 17

the exact temperature at which liquid and gas are in equilibrium

Front 18

sublimation

Back 18

when a solid changes directly to a gas without going through the liquid state. ex: dry ice

Front 19

intermolecular force

Back 19

a force that acts between molecules and holds molecules close to one another

Front 20

intermolecular forces in gases

Back 20

they are negligible (small and unimportant) so gas molecules act independently of one another

Front 21

intermolecular forces in liquids and solids

Back 21

they are strong enough to hold the molecules in close contact

Front 22

the stronger the intermolecular forces in a substance...

Back 22

the more difficult it is to separate the molecules and the higher the melting and boiling points of the substance

Front 23

3 major types of intermolecular forces

Back 23

dipole-dipole, london dispersion, and hydrogen bonding

Front 24

dipole-dipole force

Back 24

the attractive force between positive and negative ends of polar molecules

Front 25

dipole-dipole forces have what kind of strength

Back 25

they are weak with strengths around 1 kcal/mol

Front 26

london-dispersion force

Back 26

the short-lived attractive force due to the constant motion of electrons within molecules

Front 27

all molecules experience what type of intermolecular force

Back 27

london-dispersion force

Front 28

strength of london-dispersion force

Back 28

they are weak with strengths around 0.5-2.5 kcal/mol but increase with molecular weight and amount of surface area available for interaction between molecules

Front 29

what intermolecular force only happens between polar molecules

Back 29

dipole-dipole forces

Front 30

hydrogen bonds

Back 30

the attraction between a hydrogen atom bonded to an electronegative O, N, or F atom and another nearby O, N or F atom

Front 31

O-H, N-H, and F-H bonds are highly

Back 31

polar

Front 32

strength of hydrogen bonds

Back 32

they are strong with strengths around 10 kcal/mol

Front 33

gases have what kind of density, and what happens to volume when placed under pressure

Back 33

low density and are easily compressed to a smaller volume when placed under pressure

Front 34

gases undergo a larger_____ when their temp is changed than do liquids or solids

Back 34

they undergo a larger expansion or contraction

Front 35

kinetic-molecular theory of gas

Back 35

a group of assumptions that explain the behavior of gases

Front 36

kinetic-molecular theory of gas about their particles

Back 36

a gas consists of many particles, either atoms or molecules, moving about at random with no attractive forces between them- this explains when different gases mix together quickly

Front 37

kinetic-molecular theory of gas about their kinetic energy

Back 37

the average kinetic energy of gas particles is proportional to the Kelvin temperature- as temp increases, gas particles have more kinetic energy and move faster

Front 38

kinetic-molecular theory of gas about space

Back 38

the amount of space occupied by the gas particles themselves is much smaller than the amount of space between particles- most of the volume taken up by gases is empty space

Front 39

kinetic-molecular theory of gas about collision of gas particles

Back 39

collisions of gas particles, either with another particle or with the wall of the container, are elastic- the total kinetic energy of the particles is constant

Front 40

ideal gas

Back 40

a gas that obeys all of the assumptions of the kinetic-molecular theory- no such thing as a perfectly ideal gas because all gases behave differently

Front 41

when do most real gases display nearly ideal behavior

Back 41

under normal conditions

Front 42

Pressure (P)

Back 42

the force (F) per unit are (A) pushing against a surface

Front 43

formula for pressure (P)

Back 43

P= F/A

Front 44

why is the atmospheric pressure not constant

Back 44

because it varies slightly from day to day depending on weather, and varies with altitude

Front 45

where is the density of air greatest in the atmosphere

Back 45

at the earth's surface and decreases with increasing altitude

Front 46

what is the most commonly used unit of pressure

Back 46

the millimeter of mercury (mmHg)- often called a torr

Front 47

unit of pressure given in the SI system

Back 47

the pascal (Pa)

Front 48

1 pascal (Pa)= how many mmHg

Back 48

0.007500 mmHg

Front 49

1 atm (atmosphere) = how many mmHg, psi and Pa

Back 49

1 atm= 760 mmHg= 14.7 psi= 101,325 Pa

Front 50

gas laws

Back 50

a series of laws that predict the influence of pressure (P), volume (v) and temperature (T) on any gas or mixture of gases

Front 51

which gas law is the relation between volume and pressure

Back 51

boyle's law

Front 52

boyle's law

Back 52

the volume of a gas is inversely proportional to its pressure for a fixed amount of gas at a constant temperature- volume and pressure change in opposite directions

Front 53

according to boyle's law, what happens to volume as the pressure goes up

Back 53

the volume goes down

Front 54

equations for boyle's law

Back 54

PV=K (a constant value), P1V1=P2V2, P2= (P1V1)/V2, V2= (P1V1)/ P2

Front 55

which gas law is the relation between volume and temperature

Back 55

charles law

Front 56

charles law

Back 56

the volume of a gas is directly proportional to it's kelvin temperature for a fixed amount of gas at a constant pressure

Front 57

according to charle's law, is the kelvin temperature of a gas is doubled, what happens to its volume?

Back 57

it's doubled

Front 58

equations for charles law

Back 58

V/T=K, (V1/T1)=(V2/T2)

Front 59

which gas law is the relation between pressure and temperature

Back 59

Gay-lussac's law

Front 60

Gay-lussac's law

Back 60

the pressure of a gas is directly proportional to its kelvin temperature for a fixed amount of gas at a constant volume

Front 61

according to gay-lussac's law, as temperature of a gas goes up, the pressure will

Back 61

go up

Front 62

equations for gay-lussac's law

Back 62

P/T= K, (P1/T1)=(P2/T2)

Front 63

combined gas law

Back 63

when the relationships of PV, V/T, and P/T merge together when the amount of gas is fixed

Front 64

equations for the combined gas law

Back 64

PV/T=K, (P1V1)/T1= (P2V2)/T2

Front 65

which gas law is the relation between volume and molar amount

Back 65

avogadro's law

Front 66

avogadro's law

Back 66

the volume of a gas is directly proportional to its molar amount at a constant pressure and temperarture

Front 67

according to avogadro's law, a sample that contains twice the molar amount will have ____ the volume

Back 67

twice

Front 68

equations for avogadro's law

Back 68

V/n=K, (V1/n1)=(V2/n2)

Front 69

standard temperature and pressure (STP)

Back 69

temperature of 0° C (273.15 K) and a pressure of 1 atm (760 mmHg)

Front 70

standard molar volume of any ideal gas at STP

Back 70

at STP, 1 mol of any gas (6.02 x 10^23 particles) has a volume of 22.4 L- 22.4 L/mol

Front 71

the ideal gas law

Back 71

combination of pressure (p) volume (v) temperature (T) and number of moles (n)

Front 72

equations for ideal gas law

Back 72

PV/nT=R, PV=nRT

Front 73

gas constant (R)

Back 73

the constant R in the ideal gas law, PV=nRT.

Front 74

what does the gas constant (R) depend on? what are the values of R

Back 74

depends on the units of pressure. for pressure in atm, R= 0.0821 L*atm/mol*K. for pressure in mmHg, R= 62.4 L*mmHg/mol*K

Front 75

mixtures of gases vs pure gases- behavior

Back 75

mixture of gases behave the same as pure gases and obey the same laws

Front 76

partial pressure

Back 76

the contribution of a given gas in a mixture to the total pressure

Front 77

daltons law

Back 77

the total pressure exerted by a gas mixture (P-total) is the sum of the partial pressures of the components in the mixture- Ptotal= Pgas1 + Pgas2 + Pgas3....

Front 78

vapor

Back 78

the gas molecules are in equilibrium with a liquid

Front 79

vapor pressure

Back 79

the partial pressure of vapor molecules in equilibrium with a liquid

Front 80

what does vapor pressure depend on

Back 80

the temperature and the chemical identity of a liquid

Front 81

vapor pressure rises with increasing temperature until

Back 81

it becomes equal to the pressure of the atmosphere

Front 82

normal boiling point

Back 82

the boiling point at a pressure of exactly 1 atm or 760 mmHg

Front 83

viscosity

Back 83

the measure of a liquids resistance to flow

Front 84

nonpolar molecules have relatively ____ viscosities, polar molecules have relativley___ viscosities

Back 84

Low (gasoline), High (glycerin)

Front 85

surface tension

Back 85

the resistance of a liquid to spreading out and increasing its surface area

Front 86

what substance has the highest specific heat of any liquid

Back 86

water- giving it the capacity to absorb a large quantity of heat while changing only slightly in temperature

Front 87

water has what degree of heat of vaporization and what does it mean

Back 87

high heat of vaporization (540 cal/g) which means that is carries away a large amount of heat when it evaporates

Front 88

what is the most fundamental distinction between solids

Back 88

some are crystalline and some are amorphous

Front 89

crystalline solid

Back 89

solid whose particles (atoms/ions/molecules) are ridgedly held in an ordered arrangement

Front 90

crystalline solids can be categorized as...

Back 90

ionic, molecular, covalent network or metallic

Front 91

ionic solids

Back 91

those like sodium chloride, whose constituent particles are ions held together by ionic bonds

Front 92

molecular solids

Back 92

those like sucrose or ice, whose constituent particles are molecules held together by intermolecular forces

Front 93

covalent network solids

Back 93

those like diamonds or quartz (SiO2) whose atoms are linked together by covalent bonds into a giant 3-dimensional array- one very large molecule

Front 94

metallic solids

Back 94

such as silver or ion, can be viewed as vast 3-dimensional arrays of metal cations immersed in a sea of electrons that are free to move about

Front 95

amorphous solid

Back 95

a solid whose particles do not have an orderly arrangement

Front 96

how are amorphous solids usually formed?

Back 96

when liquids cool before they can achieve internal order or when their molecules are large and tangled together- glass, tar, hard candies

Front 97

how do amorphous solids differ from crystalline solids

Back 97

by the softening over a wide temperature range rather than having sharp melting points and by shattering to gives pieces with curves rather than planar faces

Front 98

heat of fusion

Back 98

the quantity of heat required to completely melt one gram of a substance once it has reached its melting point

Front 99

heat of vaporization

Back 99

the quantity of heat needed to completely vaporize one gram of a liquid once it has reached its boiling point

Front 100

a liquid with a low heat of vaporization

Back 100

evaporates rapidly and is said to be volatile

Front 101

what happens if a volatile liquid spills on your skin

Back 101

you feel a cooling effect as it evaporates because it is absorbing heat from your body

Front 102

equations for the energy needed to complete a phase change (melting and boiling)

Back 102

for melting- heat (cal or J)= Mass (g) x heat of fusion (cal or J/ g). for boiling- heat (cal or J)= Mass (g) x heat of vaporization (cal or J/g)

Front 103

mixture

Back 103

an intimate combination of two or more substances, both of which retain their chemical identities

Front 104

mixtures can be classified as two things

Back 104

heterogeneous mixtures or homogeneous mixtures

Front 105

heterogeneous mixture

Back 105

a nonuniform mixture that has regions of different composition- rocky road ice cream, granite

Front 106

homogeneous mixture

Back 106

a uniform mixture that has the same composition throughout- seawater

Front 107

homogeneous mixtures can be classified as what two things. according to what

Back 107

either solutions or colloids- according to the size of their particles

Front 108

solutions

Back 108

the most important class of homogeneous mixtures- a mixture that contains particles the size of a typical ion or small molecule- roughly 0.1-2 nm in diameter- ex: air, seawater, gasoline, wine

Front 109

colloids

Back 109

a homogeneous mixture that contains particles that range in diameter from 2 to 500 nm- ex: milk, fog, butter, pearl

Front 110

characteristics of solutions

Back 110

transparent to light, does not separate on standing, nonfilterable

Front 111

characteristics of colloids

Back 111

often murky or opaque to light, does not separate on standing, nonfilterable

Front 112

characteristics of heterogeneous mixtures

Back 112

murky or opaque to light, separates on standing, filterable, particle size is greater than 500 nm

Front 113

solute

Back 113

a substance that is dissolved in a solvent

Front 114

solvent

Back 114

the substance in which another substance (solute) is dissolved- ex: seawater- the dissolved salts are the solutes and the water is the solvent

Front 115

what determines whether a substance is soluble in a given liquid

Back 115

solubility depends on the strength of the attractions between solute and solvent particles relative to the strengths of the attractions within the pure substances

Front 116

rule of thumb for predicting solubility

Back 116

"like dissolves like"- substances with similar intermolecular forces form solutions with one another, whereas substances with different intermolecular forces do not

Front 117

solvation (or hydration specifically for water)

Back 117

the clustering of solvent molecules around a dissolved solute molecule or ion

Front 118

solid hydrates

Back 118

formed when ionic compounds attract water strongly enough to to hold on to water molecules even when crystalline

Front 119

in the solid hydrate, calcium sulfate hemihydrate→ CaSO₄ · ½H₂O, what does the dot mean

Back 119

the dot indicates that for every 2 CaSO₄ formula units in the crystal there is also one water molecule present

Front 120

hygroscopic

Back 120

compounds that have the ability to pull water molecules (vapor) from the surrounding atmosphere (humid air) so they can become hydrated

Front 121

miscible

Back 121

mutually soluble in all proportions- solute will continue to dissolve in solvent no matter how much is added

Front 122

saturated solution

Back 122

a solution that contains the maximum amount of dissolved solute at equilibrium

Front 123

solubility

Back 123

the maximum amount of a substance that will dissolve in a given amount of solvent at a specified temperature- usually expressed in grams per 100 mL (g/100mL)

Front 124

temperature and solubility on solids

Back 124

temperature has a dramatic effect on solubility- the effect is different for every substance and is usually unpredictable

Front 125

supersaturated solution

Back 125

a solution that contains more than the maximum amount of dissolved solute; a nonequilibrium situation

Front 126

unlike solids, the influence of temperature on the solubility of gases is

Back 126

predictable- addition of heat decreases the solubility of most gases

Front 127

the effect of pressure on the solubility of solids, liquids and gases

Back 127

pressure has no effect on the solubility of a solid or liquid, but has a strong effect on the solubility of a gas

Front 128

Henry's law

Back 128

the solubility (or concentration) of a gas is directly proportional to the partial pressure of the gas if the temperature is constant

Front 129

according to henry's law, if the partial pressure of a gas doubles, what else doubles

Back 129

solubility

Front 130

equations for henry's law

Back 130

C/Pgas=k, (C1/P1)=(C2/P2)

Front 131

percent concentrations express

Back 131

the amount of solute in one hundred units of solution- parts per hundred (pph)

Front 132

for solid solutions, percent concentrations are usually expressed as

Back 132

mass/mass percent concentrations (m/m)%

Front 133

mass/mass percent concentrations (m/m)%

Back 133

concentration expressed as the number of grams of solute per 100 grams of solution. (m/m)% concentration=mass of solute (g)/ mass of solution (g) x 100%

Front 134

the concentration of a solution made by dissolving one liquid in another is often expressed by

Back 134

volume/volume percent concentration (v/v)%

Front 135

volume/volume percent concentration (v/v)%

Back 135

concentration expressed as the number of milliliters of solute dissolved in 100 mL of solution. (v/v)% concentration= volume of solute (mL)/volume of solution (mL) x 100%

Front 136

mass/volume percent concentration (m/v)%

Back 136

concentration expressed as the number of grams of solute per 100 mL of solution- (m/v)% concentration= mass of solute (g)/ volume of solution (mL) x 100%

Front 137

when concentrations are very small, it is more convenient to use

Back 137

parts per million (ppm) or parts per billion (ppb)

Front 138

parts per million (ppm)

Back 138

number of parts per one million (10^6) parts- ppm= mass of solute/mass of solution x 10^6 or volume of solute/volume of solution x 10^6

Front 139

parts per billion (ppb)

Back 139

number of parts per one billion (10^9) parts- ppb= mass of solute/mass of solution x 10^9 or volume of solute/volume of solution x 10^9

Front 140

molarity (M)

Back 140

concentration expressed as the number of moles of solute per liter of solution- M=moles of solute/liters of solution (v)

Front 141

how would you find the moles of a solute

Back 141

molarity (moles per liter) x volume of solution (L)

Front 142

how would you find the volume of a solution

Back 142

moles of solute/molarity

Front 143

dilution

Back 143

adding additional solvent to lower the concentration- amount of solute remains constant, only volume is changed

Front 144

dilution factor

Back 144

the ratio of the initial and final solution volumes (Vc/Vd)

Front 145

electrolyte

Back 145

a substance that produces ions and therefore conducts electricity when dissolved in water

Front 146

the ability of a solution to conduct electricity depends on

Back 146

the concentration of ions in solution

Front 147

strong electrolyte

Back 147

substance that ionizes completely when dissolved in water ex: NaCl

Front 148

weak electrolyte

Back 148

a substance that is only partly ionized in water ex: molecular substances like acetic acid CH₃CO₂H

Front 149

nonelectrolyte

Back 149

a substance that does not produce ions when dissolved in water= ex: molecular substances like glucose

Front 150

Equivalent (Eq)

Back 150

for ions, one Eq is equal to the number of ions that carry 1 mol of charge

Front 151

gram-equivalent (g-Eq)

Back 151

for ions, the molar mass of the ion (in grams) divided by the ionic charge: 1 gram-equivalent of ion= molar mass of ion(g) / charge on ion

Front 152

how do you find the number of equivalents of a given ion per liter of solution

Back 152

Molarity of ion (M) x charge on ion

Front 153

one milliequivalent (mEq) of an ion

Back 153

1 mEq of an ion is 1/1000 of an equivalent or 1 mEq= 0.001 Eq, 1 Eq= 1000 mEq

Front 154

anion gap

Back 154

the difference in concentration of positive ions and negative ions- difference is made up by the presence of negatively charged proteins and the anions of organic acids

Front 155

colligative property

Back 155

a property of a solution that depends only on the number of dissolved particles (concentration of dissolved solute), not on their chemical identity

Front 156

4 colligative properties of solutions

Back 156

1. vapor pressure is lower for a solution than for a pure solvent 2. boiling point is higher for a solution than for a pure solvent 3. freezing point is lower for a solution than for a pure solvent 4. osmosis occurs when a solution is separated from a pure solvent by a semipermeable membrane

Front 157

why is the vapor pressure lower for a solution than for a pure solvent

Back 157

if some of the liquid (solvent) molecules at surface are replaced by other (solute) particles that don't evaporate, then the rate of evaporation of solvent molecules decreases and the vapor pressure is lower

Front 158

why is the boiling point for a solution higher than for a pure solvent

Back 158

it is a consequence of the vapor pressure lowering for a solution- because vapor pressure of a solution is lower than that of the pure solvent at a given temperature, the solution must be heated to a higher temperature for its vapor pressure to reach atmospheric pressure

Front 159

for each mole of solute particles added, what happens to the boiling point (equation)

Back 159

for each mol of solute added, the boiling point of 1 kg of water is raised by 0.51 ° C or ∆Tboiling=(0.51°C kg water/mol particles)(mol particles/kg water)

Front 160

why is the freezing point of solutions lower than freezing points of pure solvents

Back 160

because solute molecules are dispersed between solvent molecules throughout the solution, making it more difficult for solvent molecules to come together and organize into ordered crystals- lowering the freezing point

Front 161

for each mole of nonvolatile solute particles added, what happens to the freezing point (equation)

Back 161

for each mol of solute particles , the freezing point of 1 kg of water is lowered by 1.86 °C: ∆Tfreezing= (-1.86°C kg water/mol particles)(mol particles/ kg water)

Front 162

semipermeable

Back 162

a material or membrane allowing certain substances like water and other small molecules to pass through it, but they block the passage of large solute molecules or ions

Front 163

osmosis

Back 163

the passage of solvent through a semipermeable membrane separating two solutions of different concentration

Front 164

osmotic pressure (π)

Back 164

the amount of external pressure that must be applied to a solution to prevent the net movement of solvent molecules across a semipermeable membrane: π=(n/V)RT- v= solution volume n= # of moles of particles in solution R= gas constant T= absolute temperature of the solution

Front 165

what is convenient to use to describe the concentration of particles in solution

Back 165

osmolarity (osmol)

Front 166

osmolarity (osmol)

Back 166

the sum of the molarities of all dissolved particles in a solution

Front 167

isotonic

Back 167

having the same osmolarity

Front 168

hypotonic

Back 168

having an osmolarity less than the surrounding blood plasma or cells

Front 169

what is hemolysis

Back 169

a process where when red blood cells are placed in a hypotonic solution, water passes through the membrane into the cells, causing the cell to swell up and burst

Front 170

hypertonic

Back 170

having an osmolarity greater than the surrounding blood plasma or cells

Front 171

what is crenation

Back 171

a process where when red blood cells are in a hypertonic solution, water passes out of the cells into the surrounding solution, causing the cells to shrivel

Front 172

dialysis

Back 172

similar to osmosis, except that the pores in a dialysis membrane are larger than those in an osmotic membrane so that both solvent molecules and small solute particles can pass through it, but large colloidal particles cannot pass

Front 173

according to the Arrhenius definition, what is an acid

Back 173

a substance that produces hydrogen ions, H⁺ when dissolved in water

Front 174

according to the Arrhenius definition, what is a base

Back 174

a substance that produces hydroxide ions, OH⁻ when dissolved in water

Front 175

according to the Arrhenius definition, a neutralization reaction of an acid with a base yields

Back 175

water plus a salt

Front 176

what is a salt

Back 176

an ionic compound composed of the cation from the base and the anion from the acid

Front 177

why are swedish chemist, Svante Arrhenius's definition of acids and bases limited

Back 177

because they only refer to reactions that take place in aqueous solutions

Front 178

hydronium ion

Back 178

the H₃O⁺ ion, formed when an acid reacts with water

Front 179

acids generally have what kind of taste

Back 179

sour- nearly every sour food contains an acid

Front 180

bases are present in many

Back 180

household cleaning agents- perfumed bar soap, ammonia based window cleaners, substances you put down the drain to dissolve hair and grease

Front 181

what are the 9 most common acids and bases

Back 181

1. sulfuric acid (H₂So₄) 2. hydrochloric acid (HCl) 3. phosphoric acid (H₃PO₄) 4. nitric acid (HNO₃) 5. Acetic acid (CH₃CO₂H) 6. sodium hydroxide (NaOH) 7. calcium hydroxide (Ca(OH)₂) 8. magnesium hydroxide (Mg(OH)₂) 9. Ammonia (NH₃)

Front 182

what is sulfuric acid

Back 182

(H₂So₄)- most important raw material and most manufactured chemical. used in hundreds of industrial processes including preparation of phosphate fertilizers, Most common consumer use is as the acid found in automobile batteries

Front 183

what is hydrochloric acid

Back 183

(HCl)- historically known as muriatic acid- many industrial applications including its use in metal cleaning and in the manufacture of high-frutose corn syrup. Also present as stomach acid in the digestive systems of most mammals

Front 184

what is phosphoric acid

Back 184

(H₃PO₄)- used in the manufacture of phosphate fertilizers. Also used as an additive in foods and toothpastes- the tart taste of many soft drinks is due to the presence of phosphoric acid

Front 185

what is nitric acid

Back 185

(HNO₃)- strong oxidizing agent that is used for many purposes including the manufacture of ammonium fertilizer and military explosives. when spilled on skin, it leaves a characteristic yellow coloration because of its reaction with skin proteins

Front 186

what is acetic acid

Back 186

CH₃CO₂H- the primary organic constituent of vinegar. Also occurs in all living cells and is used in many industrial processes such as the preparation of solvents, lacquers and coatings

Front 187

what is sodium hydroxide

Back 187

NaOH- also called caustic soda or lye- most commonly used of all bases. Industrially it is used in production of aluminum from its ore, production of glass and in manufacture of soap from animal fats. Often found in drain cleaners because it reacts with the fats and proteins found in grease and hair

Front 188

what is calcium hydroxide

Back 188

Ca(OH)₂- or slaked lime, its made industrially by treating lime (CaO) with water. Many applications, including its use in mortars and cements

Front 189

what is magnesium hydroxide

Back 189

Mg(OH)₂- or milk of magnesia. an additive in foods, toothpastes and many over the counter medications.

Front 190

what is Ammonia

Back 190

NH₃- used primarily as a fertilizer, but it also has many other applications including the manufacture of pharmaceuticals and explosives. dilute solution is frequently used as a glass cleaner

Front 191

bronsted-lowry acid

Back 191

a substance that can donate a hydrogen ion H⁺ to another molecule or ion- also called a proton donor

Front 192

monoprotic acids

Back 192

acids with one proton to donate, such as HCl or HNO₃

Front 193

diprotic acids

Back 193

acids that have two protons to donate, such as H₂SO₄

Front 194

triprotic acids

Back 194

acids that have three protons to donate such as H₃PO₄

Front 195

bronsted-lowry base

Back 195

a substance that can accept H⁺ ions from an acid

Front 196

an acid-base reaction is one in which a ____ is transferred

Back 196

proton

Front 197

a bronsted-lowry base can be two things

Back 197

neutral or negatively charged

Front 198

if a bronsted-lowry base is neutral, then the product has a

Back 198

positive charge after H⁺ has been added

Front 199

if the bronsted-lowry base is negatively charged, then the product

Back 199

is neutral

Front 200

conjugate acid-base pairs

Back 200

two substances whose formulas differ only by a hydrogen ion H⁺- they are found on opposite sides of the chemical equation

Front 201

conjugate base

Back 201

the substance formed by the loss of H⁺ from an acid

Front 202

conjugate acid

Back 202

the substance formed by the addition of H⁺ to a base

Front 203

the number of protons in a conjugate acid-base pair is always

Back 203

one greater than the number of protons in the base of the pair

Front 204

what acids and bases are highly corrosive (3) and how do they react

Back 204

sulfuric acid (H₂SO₄), hydrochloric acid (HCl), sodium hydroxide (NaOH). They react readily and can cause serious burns to skins

Front 205

why are some acids and bases relatively safe while others must be handled with extreme cation?

Back 205

answer lies in how easily they produce the active ions for an acid (H⁺) or base (OH⁻)

Front 206

strong acid

Back 206

an acid that gives up H⁺ easily and is essentially 100% dissociated in water

Front 207

dissociation

Back 207

the splitting apart of an acid in water to give H⁺ and an anion

Front 208

weak acid

Back 208

an acid that gives up H⁺ with difficulty and is less than 100% dissociated in water

Front 209

weak base

Back 209

a base that has only a slight affinity for H⁺ and holds it weakly

Front 210

strong base

Back 210

a base that has a high affinity for H⁺ and holds it tightly

Front 211

the stronger the acid, the ____ its conjugate base; the weaker the acid, the _____ the conjugate base

Back 211

the stronger the acid, the weaker its conjugate base; the weaker the acid, the stronger the conjugate base

Front 212

a strong acid and a weak base cause the reverse reaction to

Back 212

occur at a lesser extent

Front 213

a weak acid and a strong base causes the reverse reaction to

Back 213

occur more readily

Front 214

an acid-base proton transfer equilibrium always favors the

Back 214

reaction of the stronger acid with the stronger base and formation of the weaker acid and base

Front 215

in a contest for protons, the _____ base always wins

Back 215

the stronger base

Front 216

acid dissociation constant (Ka)

Back 216

the equilibrium constant for the dissociation of an acid (HA), equal to [H⁺][A⁻]/[HA]

Front 217

important points for Ka values (4)

Back 217

1. strong acids have Ka values much greater than 1 because dissociation is favored 2. weak acids have Ka values much less than 1 because dissociation is not favored 3. donation of each successive H⁺ from a polyprotic acid is more difficult than the one before it, so Ka values become successively lower 4. most organic acids, which contain the CO₂H group, have Ka values near 10⁻⁵

Front 218

in bronsted-lowry definition, water can as

Back 218

both an acid and a base

Front 219

when in contact with a base, water reacts as

Back 219

a bronsted-lowry acid and donates a protein to the base

Front 220

when in contact with an acid, water reacts as

Back 220

a bronsted-lowry base and accepts H⁺ from the acid

Front 221

amphoteric

Back 221

a substance that can react as either an acid or a base, such as water

Front 222

what happens to water when no other acids or bases are present

Back 222

one water molecule acts as an acid while another water molecule acts as a base, reacting to form hydronium (H₃O⁺) and hydroxide (OH⁻) ions

Front 223

ion-product constant for water (Kw)

Back 223

the product of the H₃O⁺ and OH⁻ molar concentrations in water or any aqueous solution: Kw=K[H₂O][H₂O]= [H₃O⁺][OH⁻]= 1.00 x 10⁻¹⁴ at (25° C)

Front 224

solutions are identified as..... according to what

Back 224

acidic, neutral or basic (alkaline), according to their value of their H₃O⁺ and OH⁻ concentrations

Front 225

concentration of acidic solution

Back 225

[H₃O⁺]> 10⁻⁷ M and [OH⁻]< 10⁻⁷ M

Front 226

concentration of neutral solution

Back 226

[H₃O⁺]= 10⁻⁷ M and [OH⁻]= 10⁻⁷ M

Front 227

concentration of basic solution

Back 227

[H₃O⁺]< 10⁻⁷ M and [OH⁻]> 10⁻⁷ M

Front 228

the ______ of an aqueous solution is a number usually between 0 and 14 that indicates the H₃O⁺ concentration of a solution

Back 228

the pH

Front 229

pfunction

Back 229

mathematically defined as the negative common logarithm of some variable

Front 230

pH

Back 230

a measure of the acid strength of a solution; the negative common logarithm of the H₃O⁺ concentration

Front 231

pH for an acidic solution

Back 231

pH<7, [H₃O⁺]> 1 x 10⁻⁷ M

Front 232

pH for a neutral solution

Back 232

pH= 7, [H₃O⁺]= 1 x 10⁻⁷ M

Front 233

pH for a basic solution

Back 233

pH>7, [H₃O⁺]< 1 x 10⁻⁷

Front 234

converting from pH to [H₃O⁺] requires finding the

Back 234

antilogarithm of the negative pH

Front 235

converting from [H₃O⁺] to pH requires finding the

Back 235

logarithm

Front 236

significant figures for an antilogarithm

Back 236

contains the same number of sig figs as the original number has to the right of the decimal point

Front 237

significant figures for a logarithm

Back 237

contains the same number of digits to the right of the decimal point as the number of sig figs in the original number

Front 238

what is the simplest but least accurate method of measuring the pH of a solution

Back 238

using an acid-base indicator

Front 239

acid-base indicator

Back 239

a dye that changes color depending on the pH of a solution

Front 240

universal indicator

Back 240

makes pH determination easy, test kits that contain a mixture of indicators to give approximate pH measurements in the range 2-10

Front 241

buffer

Back 241

a combination of substances that act together to prevent a drastic change in pH; usually a weak acid and its conjugate base

Front 242

Henderson-Hasselbalch equation

Back 242

the logarithmic form of the Ka equation for a weak acid, used in applications involving buffer solutions: pH=pKa - log([HA]/[A⁻]) or pH= pKa + log([A⁻]/[HA])

Front 243

what conditions do most effective buffers meet (3)

Back 243

1. the pKa for the weak acid should be close to the desired pH of the buffer solution 2. the ratio of [HA] to [A⁻] should be close to 1, so that neither additional acid nor additional base changes the pH of the solution dramitically 3. the molar amounts of [HA] and [A⁻] in the buffer should be approximately 10 times greater than the molar amounts of either acid or base you expect to add so that the ratio [A⁻]/[HA] does not undergo a large change

Front 244

the pH of body fluids is maintained by how many buffer systems

Back 244

3

Front 245

what two buffer systems depend on weak acid conjugate base interactions exactly like those of the acetate buffer system

Back 245

the carbon acid- bicarbonate system (H₂CO₃-HCO₃) and the dihydrogen phosphate-hydrogen phosphate system (H₂PO₄-HPO₄²⁻)

Front 246

equivalent of acid

Back 246

amounts of acid that contains 1 mole of H⁺ ions

Front 247

1 gram-equivalent of acid

Back 247

= molar mass of acid (grams)/ number of H⁺ ions per formula unit

Front 248

equivalent of base

Back 248

amount of base that contains 1 mole of OH⁻ ions

Front 249

1 gram-equivalent of base

Back 249

= molar mass of base (grams)/ number of OH⁻ ions per formula unit

Front 250

1 equivalent of any acid does what to any base

Back 250

1 equivalent of any acid neutralizes one equivalent of any base

Front 251

normality (N)

Back 251

a measure of acid (or base) concentration expressed as the number of acid (or base) equivalents per liter of solution: N= equivalents of acid or base/ liters of solution

Front 252

the values of molarity (M) and normality (N) are the same for _____ acids, but are not the same for _____ acids

Back 252

M and N are the same for monoprotic acids, but are not the same for diprotic or triprotic acids

Front 253

normality of acid

Back 253

= (molarity of acid) X (the number of H⁺ ions produced per formula unit)

Front 254

normality of base

Back 254

= (molarity of base) X (the number of OH⁻ ions produced per formula unit)

Front 255

what are the most common kinds of bronsted-lowry acid-base reactions (3)

Back 255

1. those of an acid with hydroxide ion 2. an acid with bicarbonate or carbonate ion 3. an acid with ammonia or a related nitrogen containing compound

Front 256

reaction of an acid with hydroxide ion

Back 256

1 equivalent of an acid reacts with 1 equivalent of a metal hydroxide to yield water and a salt in a neutralization reaction

Front 257

reaction of an acid with bicarbonate and carbonate ions

Back 257

bicarbonate reacts with acid by accepting H⁺ to yield carbonic acid, H₂CO₃. Carbonate accepts 2 protons in reaction with an acid

Front 258

reaction of an acid with ammonia

Back 258

acid reacts with ammonia to yield ammonium salts

Front 259

titration

Back 259

a procedure for determining the total acid or base concentration of a solution

Front 260

when titration involves a neutralization reaction in which 1 mole of acid reacts with 1 mole of base, then

Back 260

(M-acid X V-acid)= (M-base X V-base)

Front 261

when coefficients for the acid and base in a balanced neutralization reaction are not the same, we use

Back 261

equivalents of acid and base instead of moles and normality instead of molarity: (Eq)acid= (Eq)base or (N-acid X V-acid)= (N-base X V-base)

Front 262

salt solutions can be... depending on what?

Back 262

neutral, acidic or basic, depending on the ions present because some ions react with water to produce H₃O⁺ and some ions react with water to produce OH⁻

Front 263

general rule for predicting acidity or basicity of a salt solution

Back 263

the stronger partner from which the salt is formed dominates- a salt formed from a strong acid and weak base yields an acidic solution, and vise versa. a salt formed from a strong acid and a strong base yields a neutral solution