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Chem 2 Heider > Ch 15 - Acids and Bases > Flashcards

Ch 15 - Acids and Bases Flashcards

1
Q

Acid Reflux

A

When HCl backs up from the stomach into the esophagus

  • painful
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2
Q

Acid

A

By one definition:

any substance that produces H+ ions in solution

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3
Q

Properties of an Acid

A
  • Sour taste
  • ability to dissolve many metals
  • ability to turn blue litmus paper red
  • ability to neutralize bases
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4
Q

Carboxylic acid

A

An acid containing the following group:

     =O(O double bonded to C) H-O-C-
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5
Q

What is carboxylic acid found in?

A

often in substances derived from living organisms

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6
Q

Examples where carboxylic acid is found

A
  • lemons
  • limes
  • malic acid
  • apples
  • grapes
  • wine
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7
Q

Properties of bases

A
  • bitter taste
  • slippery feel
  • ability to turn red litmus paper blue
  • ability to neutralize acids
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8
Q

Alkaloids

A

organic bases found in plants that are often poisonous

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9
Q

What contains alkaloids?

A
  • Coffee

- Chocolate(especially dark chocolate)

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10
Q

Why do bases feel slippery?

A

because they interact with oils in the skin to make soap like substances

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11
Q

The Arrhenius Definition:

Acid:

A

a substance that produces H+ ions in aqueous solution

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12
Q

The Arrhenius Definition:

Base:

A

a substance that produces OH- ions in aqueous solution

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13
Q

The Arrhenius Definition:

HCl -> H+(aq) + Cl-(aq)

A
  • this is an acid because H+ is produced
    - HCl is a covalent compound and does not contain ions however water ionizes it completely to form H+(aq) and Cl-(aq).
    - the H+ ions are highly reactive and bond with H2O to form H3O+
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14
Q

The Arrhenius Definition:

hydronium ion

A

H3O+

		- in water, H+ always associate with H2O to form hydronium ions
		- general form: H(H2O)n+
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15
Q

The Arrhenius Definition:

NaOH(aq) -> Na+(aq) + OH-(aq)

A
  • an ionic compound with Na+ and OH- ions

- OH- makes this a base

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16
Q

The Arrhenius Definition:

H+(aq) + OH-(aq) -> H2O(l)

A

acid + base -> water and neutralizes each other

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17
Q

Bronsted-Lowry Definition

A

the transfer of H+ ions in an acid base reaction is based on the transfer of protons

  • more widely applicable definition of acids and bases
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18
Q

Bronsted-Lowry Definition:

Acid

A

proton(H+ ion) donor

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19
Q

Bronsted-Lowry Definition:

Base

A

proton(H+ ion) acceptor

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20
Q

Bronsted-Lowry Definition:

HCl

A

an acid in solution because it donates a proton to water

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21
Q

Bronsted-Lowry Definition:

NH3

A

a base because it accepts a proton from water

- NH3(aq) + H2O(l) NH4+(aq) + OH-(aq)

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22
Q

Bronsted-Lowry Definition:

acids and bases always occur together

A
  • HCL(aq) + H2O(l) -> H3O+(aq) + Cl-(aq)
    - acid + base
  • NH3(aq) + H2O(l) NH4+(aq) + OH-(aq)
    - base + acid
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23
Q

Bronsted-Lowry Definition:

amphoteric

A

a substance that can act as an acid or base

- H2O

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24
Q

Bronsted-Lowry Definition:

Reverse: NH3(aq) + H2O(l) NH4+(aq) + OH-(aq)

to: NH4+(aq) + OH-(aq) NH3(aq) + H2O(l)

A
  • reaction is reversed
    - Acid + base
    - originally NH3 was a base but now NH4+ is the acid

creates a conjugate acid-base pair

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25
Q

Bronsted-Lowry Definition:

conjugate acid-base pair

A

two substances related to each other by transfer of a proton

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26
Q

Bronsted-Lowry Definition:

conjugate acid

A

any base to which a proton has been added

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27
Q

Bronsted-Lowry Definition:

conjugate base

A

any acid from which a proton has been removed

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28
Q

Bronsted-Lowry Definition:

How does a base become a conjugate acid?

A

A base accepts a proton and becomes a conjugate acid

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29
Q

Bronsted-Lowry Definition:

How does an acid become a conjugate base?

A

An acid donates a proton and becomes a conjugate base

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30
Q

strong acid

A

one that completely ionizes in solution

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31
Q

weak acid

A

one that partially ionizes in solution

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32
Q

HA(aq) + H2O(l) H3O+ + A-(aq)

A
  • HA is any acid
    - A- is the conjugate base
    - if equilibrium lies far to the right then it’s a strong acid
    - if equilibrium lies to the left then it is a weak acid
    - the range is continuous but in most purposes the categories of strong and weak acid suffice
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33
Q

Six important strong acids:

HCl

A

Hydrochloric Acid

Monoprotic

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34
Q

Six important strong acids:

HBr

A

Hydrobromic Acid

Monoprotic

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35
Q

Six important strong acids:

HI

A

Hydriodic Acid

Monoprotic

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36
Q

Six important strong acids:

HNO3

A

Nitric Acid

Monoprotic

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37
Q

Six important strong acids:

HClO4

A

Perchloric Acid

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38
Q

Six important strong acids:

H2SO4

A

Sulfuric Acid

Diprotic

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39
Q

Monoprotic

A

containing only one ionizable proton

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40
Q

Diprotic

A

containing two ionizable protons

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41
Q

in weak acids

  • HA(aq) + H2O(l) H3O+ + A-(aq)
A
  • the degree to which the reaction proceeds depends on the strength of the attraction between H+ and A-
    - weak attraction between H+ and A- = forward reaction favored and the acid is strong
    - strong attraction between H+ and A- then reverse reaction favored and the acid is weak
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42
Q

in general, the stronger the acid, the

A

weaker the conjugate base and vice versa

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43
Q

Six Weak Acids:

HF

A

Hydrofluoric Acid

Monoprotic

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44
Q

Six Weak Acids:

HC2H3O2

A

Acetic Acid

Monoprotic

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45
Q

Six Weak Acids:

HCHO2

A

Formic Acid

Monoprotic

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46
Q

Six Weak Acids:

H2SO3

A

Sulfurous Acid

Diprotic

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47
Q

Six Weak Acids:

H2CO3

A

Carbonic Acid

Diprotic

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48
Q

Six Weak Acids:

H3PO4

A

Phosphoric Acid

Triprotic

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49
Q

-protic

A

the number of ionizable protons

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50
Q

acid ionization constant(Ka)

A
  • the equilibrium constant for the ionization reaction of the weak acid
    - HA(aq) + H2O(l) H3O+ + A-(aq)
    - HA(aq) H+(aq) + A-(aq)
    - Ka = ([H3O+][A-])/[HA] = ([H+][A-])/[HA]
51
Q

because water is amphoteric(can act as an acid or base) it can autoionize

A

H2O(l) + H2O(l) H3O+(aq) + OH-(aq)

52
Q

ion product constant for water(Kw)(or dissociation constant for water)

A
  • Kw = [H3O+][OH-] = [H+][OH-]
    - at 25 degrees C Kw = 1.0 * 10^-14
    [H3O+] = [OH-] = sqrtKw = 1.0 * 10^-7
    - neutral pH
53
Q

acidic solution

A

contains an acid that creates additional H3O+ ions causing the [H3O+] to increase

	- the ion product constant still applies
		- 1.0*10^-14
54
Q

basic solution

A

contains a base that creates additional OH- ions, causing the [OH-] to increase and [H3O+] to decrease

	- the ion product constant still applies
		- 1.0 * 10^-14
55
Q

neutral solution

A

[H3O+] = [OH-] = 1.0 * 10^-7 M (at 25 C)

56
Q

Acidic Solution

A

[H3O+] > [OH-]

57
Q

Basic Solution

A

[H3O+] < [OH-]

58
Q

All AQ Solns

A

both H3O+ and OH- are present with [H3O+][OH-] = Kw = 1.0*10^-14 at 25 C

59
Q

pH

A

the negative log of the hydronium ion concentration

	- pH = -log[H3O+]
	- pH = -log(1.0*10^-3)
	- pH = -(-3.00)
	- pH = 3.00
	- pH goes to the total number of significant digits for rounding
		- log 1.00*10^-3 = 3.000
60
Q

At 25 C:

pH < 7

A

an acidic solution

61
Q

At 25 C:

pH > 7

A

a basic solution

62
Q

At 25 C:

pH = 7

A

a neutral solution

63
Q

pH scale is logarithmic so

A

a change of 1pH is a 10 fold change in H3O+ concentration

64
Q

pOH

A
  • pOH = -log[OH-]

- same as pH scale but with respect to OH- instead of H3O+

65
Q

At 25 C:

pOH <7

A

Basic Solution

66
Q

At 25 C:

pOH > 7

A

Acidic Solution

67
Q

At 25 C:

pOH = 7

A

Neutral Solution

68
Q

pH + pOH =

A

14.00

Always.

69
Q

pKa = -log(Ka)

A

the smaller the pKa the stronger the acid

70
Q

Two sources of H3O+

A
  • the ionization of the acid
    - autoionization of water
    - typically it can be ignored as it is such a small amount
71
Q

In a strong acid solution:

A
  • the concentration of H3O+ is equal to the concentration of the strong acid
    - 0.10M HCl has [H3O+] of 0.10
    - pH = -log(0.10) = 1.00
    - very acidic!
72
Q

In a weak acid solution:

A
  • in weak acid solution
    - the concentration of H3O+ is NOT equal to the concentration of the weak acid
    - the acid only partially ionizes here
    - ICE tables are fun
73
Q

Percent Ionization

A

of a weak acid, the ration of the ionized acid concentration to the initial acid concentration, multiplied by 100%

	- % ionization = (concentration ionized acid/initial concentration of acid)(100%)
		- ([H3O+]equil/[HA]init)(100%)
		- ((9.6*10^-3M)/0.200M))(100%) = 4.8%
74
Q

For weak acids:

  • the equilibrium H3O+ concentration increases with increasing initial concentration of the acid
A
  • as the concentration of a weak acid solution increases so too does the H3O+ concentration(not linear the H3O+ increases more slowly)
    - the percent ionization decreases with increasing concentration of the acid
    - the more dilute a solution the greater the percent ionization(less acid so higher %)
75
Q

What happens when you mix a strong and weak acid?

A

the strong acid fully ionizes to produce H3O+ and suppresses the formation of additional H3O+ which could be formed by the weak acid

76
Q

What happens when you mix two weak acids?

A

you can assume the weaker of the two is negated if the concentrations of the two are similar or if the stronger acid is greater than the weaker

77
Q

Strong Base

A

a base that completely dissociates in solution

	- NaOH(aq) -> Na+(aq) + OH-(aq)
	- 1.0 M NaOH solution contains [OH-] = 1.0 M
		- like a strong acid
78
Q

Bases:

1A and 2A

A 
typically 1A(highly soluble) and 2A(slightly soluble and produce M(OH)2 = 2mol OH- per mole of base) metal hydroxides
			-  Sr(OH)2(aq) -> Sr^2+(aq) + 2OH-(aq)
79
Q

What is the difference between a base dissociating two OH- compared to a diprotic acid?

A

unlike diprotic acids which ionize in two steps bases containing two OH- ions dissociate in one step

80
Q

Weak base

A

analogous to a weak acid

- the most common weak bases produce OH- by accepting a proton from water, ionizing water to form OH-

81
Q

Base General Form with weak base example:

B(aq) + H2O(l) BH+(aq) + OH-(aq)

A

NH3(aq) + H2O(l) NH4+(aq) + OH-(aq)

- double arrow indicates the ionization is not complete

82
Q

Base Ionization Constant

A

the extent of ionization of a weak base is quantified by Kb

	- B(aq) + H2O(l)  BH+(aq) + OH-(aq)
	- Kb = ([BH+][OH-])/[B]
83
Q

The smaller the Kb the

A

weaker the base

84
Q

pKb =

A
  • log(Kb)

- like pH

85
Q

finding the pH of a strong or weak base is

A

analogous to finding the H3O+ and pH of acids

86
Q

anions tend to be either

A

basic or neutral solutions

87
Q

cations tend to be

A

acidic or neutral solutions

88
Q

any anion is the conjugate base of an acid

A

Cl- in HCL

89
Q

every anion can act as a base BUT not every anion does act as a base

A

dependent on the strength of the corresponding acid

90
Q

an anion that is the conjugate base of a weak acid is itself a weak base

A

F- in HF is a weak base

- F-(aq)+H2Oz(l) OH-(aq) + HF(aq)

91
Q

an anion that is the conjugate base of a strong acid is pH-neutral(forms Solutions neither acidic or basic)

A

Cl- in HCl is pH neutral

92
Q

HF(aq)+H2O(l) H3O+(aq) + F-(aq)

A

hydrofloruic is a weak acid

- equilibrium lies to the left because F- and H+ have a strong affinity for each other keeping their bond together

93
Q

the weaker the acid the

A

stronger the conjugate base

94
Q

HCl(aq) + H2O(l) -> H3O+(aq) + Cl-

A

a strong acid

- the reaction equilibrium lies far to the right because H and Cl- are not attracted to each other

95
Q

Kw=(Ka)(Kb)

Find the pH of a solution containing an anion

A

Kw = 1.010^-14
- find Kb for conjugate base knowing Ka acetic acid is 1.810^-5
- Kb = Kw/Ka
= 1.010^-14/1.810^-5 = 5.610^-10
- can use this Kb to calculate the pH of a solution containing an anion acting as a base
- 5.610^-10 = (X^2)/(.100-x)
- do the approx. trick and cancel x in denominator(check validity at end)
- x=2.410^-6
- 2.410^-6/.100 * 100% = .0024%(assumption valid!)
- [OH-] = 2.410^-6 M
- [H3O+][OH-]=Kw
- H3O+=1.010^-14
- [H3O+] = 4.210^-9 M
- pH = -log[H3O+] = -log(4.210^-9) =
- 8.38 = pH

96
Q

pKa + pKb =

97
Q

cations can sometimes act as

A

weak acids

98
Q

3 categories of cation interaction

A
  • cations that are counterions of strong bases
    - cations that are conjugate acids of weak bases
    - cations that are small, highly charged metals
99
Q

cations that are counterions of strong bases

A
  • counterions do not contribute to acidity or basicity of a solution
    - strong bases typically are composed of a counterion and OH
    - NaOH, Ca(OH)2
    - counterions of strong bases are themselves pH neutral
100
Q

Cations That are Conjugate Acids of Weak Bases

A
  • a cation ion can be formed from any nonionic weak base by adding a proton(H+)
    - will be the conjugate acid of the base
    - NH4+ is conjugate acid of NH3(weak base)
    - In general a cation that is the conjugate acid of a weak base is a weak acid
101
Q

Cations that are Small highly Charged Metals

A
  • form weakly acidic solutions
    - Al^3+(aq) + 6H2O(l) -> Al(H2O)6^3+(aq)
    - the hydrated form of the ion acts as a Bronsted-Lowry Acid
    - neither alkali metal or alkaline earth metal cations ionize water in this way but many others do
    - the smaller and more highly charged the cation the more acidic the behavior
102
Q

salts in which neither the cation nor the anion acts as an acid or base form pH-neutral solutions

A
  • a salt with the cation as a counterion of a strong base and the anion as a conjugate base of a strong acid will be neutral
    - NaCl, Ca(NO3)2, KBr
    - cations pH neutral
    - anionis are conjugate bases of strong acids
103
Q

salts in which the cation does not act as an acid and the anion acts as a base for basic solutions

A
  • the cation is the counterion of a strong base and the anion is a conjugate base of a weak acid for basic solutions
    - NaF, Ca(C2H3O2)2, KNO2
    - Cations pH-neutral
    - anions are conjugate bases of weak acids
104
Q

salts in which the cation acts as an acid and the anion does not act as a base form acidic solutions

A
  • the cation is either the conjugate acid of a weak base or a small highly charged metal ioin and the anion is the conjugate base of a strong acid for acidic solutions
    - FeCl3, Al(NO3)3, NH4Br
    - cations are conjugate acids of weak bases or small highly charged metal ions
    - anions are conjugate bases of strong acids
105
Q

Salts in which the cation acts as an acid and the anion acts as a base form solutions in which the pH depends on the relative strengths of the acid and the base

A
  • cation is either a conjugate acid of a weak base or a small highly charged metal ion and the anion is the conjugate base of a weak acid
    - pH depends on the relative strengths of the acid and base
    - FeF3, Al(C2H3O2)3,NH4NO2
    - cations are conjugate acids of weak bases or small highly charged metal ions
    - anions are conjugate bases of weak acids
106
Q

Cation/Anion table

A

Anion
Conjugate Base of strong Acid CB of weak Acid
Conjugate acid of weak base Acidic Depends
Cation Small, highly charged metal ion Acidic Depends
Counterion of strong base Neutral Basic

107
Q

typically polyprotic acids ionize in successive steps, each with its own Ka

A
  • H2SO3(aq) H+(aq)+HSO3-(aq) Ka1 = 1.610^-2
    - HSO3-(aq) H+(aq) + SO3^2-(aq) Ka2 = 6.4    10^-8
    - notice the second step Ka is always smaller than the first
108
Q

when calculating the Ka of polyprotic we typically calculate the first step and ignore latter steps as their contribution is overridden by the step 1 reaction

A

treat the problem as a weak acid ICE table

109
Q

Binary acid

A

an acid containing hydrogen and only one other element

- H-Y

110
Q

In a binary acid, the ease with which the hydrogen is donated(and acidic) is based on

A

polarity and bond strength

111
Q

In a binary acid:

Increasing electronegativity

A

increases acidity

112
Q

In a binary acid:

Decreasing bond strength

A

increases acidity

113
Q

In a binary acid,

H-Y bond must be polarized with the H on the positive end to become acidic

A
  • must be lost as a H+ and a partial + charge facilitates this
    - H-Li is not acidic because H is the negative end
    - H-C is nonpolar and not acidic
    - H-F is acidic with H+ and F-
114
Q

In a binary acid,

H-Y bond strength is important

A
  • a stronger bond will be less acidic(cant donate the H+ if stuck bonded)
    - weaker bond energy can be a stronger acid
    - H-CL has 431 kJ/mol while H-F has 565kJ/mol
115
Q

In binary acids,

6A and 7A elements become more

A

acidic left to right as electronegativity increases

- become more acidic down the column as bond strength decreases

116
Q

Oxyacid

A

an acid containing hydrogen bonded to an oxygen atom that is bonded to another element

  • H-O-Y
117
Q

Oxyacid:

affected by the electronegativity of Y and the number of oxygen atoms attached to Y

A
  • the more electronegative the greater the Ka
118
Q

Oxyacid:

the more Os bonded the

A

greater the electronegativity and stronger the acid

119
Q

the lewis model is focused on

A

the electron pair instead of H+(Bronsted-Lowry definition)

120
Q

Lewis Acid

A

electron pair acceptor

121
Q

Lewis Base

A

electron pair donor

122
Q

H+ + :NH3->[H:NH3]^+

A
  • bronsted-lowry model says ammonia accepts a proton and thus acts as a base
    - lewis model says ammonia donates an electron pair to act as a base
123
Q

the Lewis model significantly expands

A

the substances that can be considered acids without altering the substances that can be considered bases significantly

124
Q

BF3+ :NH3 -> F3B:NH3

A

a Lewis acid has an empty orbital(or can rearrange electrons to create an empty orbital) that can accept an electron pair

  • many small highly charged metal ions can act as Lewis acids

Chem 2 Heider (9 decks)

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