Brainscape's Knowledge GenomeTM


Top Flashcards (Certified)
All Flashcards

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
How well did you know this?
1
Not at all
2
3
4
5
Perfectly
2
Q

Acid

A

By one definition:

any substance that produces H+ ions in solution

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
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
How well did you know this?
1
Not at all
2
3
4
5
Perfectly
4
Q

Carboxylic acid

A

An acid containing the following group:

     =O(O double bonded to C) H-O-C-
How well did you know this?
1
Not at all
2
3
4
5
Perfectly
5
Q

What is carboxylic acid found in?

A

often in substances derived from living organisms

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
6
Q

Examples where carboxylic acid is found

A
  • lemons
  • limes
  • malic acid
  • apples
  • grapes
  • wine
How well did you know this?
1
Not at all
2
3
4
5
Perfectly
7
Q

Properties of bases

A
  • bitter taste
  • slippery feel
  • ability to turn red litmus paper blue
  • ability to neutralize acids
How well did you know this?
1
Not at all
2
3
4
5
Perfectly
8
Q

Alkaloids

A

organic bases found in plants that are often poisonous

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
9
Q

What contains alkaloids?

A
  • Coffee

- Chocolate(especially dark chocolate)

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
10
Q

Why do bases feel slippery?

A

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

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
11
Q

The Arrhenius Definition:

Acid:

A

a substance that produces H+ ions in aqueous solution

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
12
Q

The Arrhenius Definition:

Base:

A

a substance that produces OH- ions in aqueous solution

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
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+
How well did you know this?
1
Not at all
2
3
4
5
Perfectly
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+
How well did you know this?
1
Not at all
2
3
4
5
Perfectly
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

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
16
Q

The Arrhenius Definition:

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

A

acid + base -> water and neutralizes each other

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
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
How well did you know this?
1
Not at all
2
3
4
5
Perfectly
18
Q

Bronsted-Lowry Definition:

Acid

A

proton(H+ ion) donor

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
19
Q

Bronsted-Lowry Definition:

Base

A

proton(H+ ion) acceptor

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
20
Q

Bronsted-Lowry Definition:

HCl

A

an acid in solution because it donates a proton to water

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
21
Q

Bronsted-Lowry Definition:

NH3

A

a base because it accepts a proton from water

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

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
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
How well did you know this?
1
Not at all
2
3
4
5
Perfectly
23
Q

Bronsted-Lowry Definition:

amphoteric

A

a substance that can act as an acid or base

- H2O

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
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

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
25
Bronsted-Lowry Definition: conjugate acid-base pair
two substances related to each other by transfer of a proton
26
Bronsted-Lowry Definition: conjugate acid
any base to which a proton has been added
27
Bronsted-Lowry Definition: conjugate base
any acid from which a proton has been removed
28
Bronsted-Lowry Definition: How does a base become a conjugate acid?
A base accepts a proton and becomes a conjugate acid
29
Bronsted-Lowry Definition: How does an acid become a conjugate base?
An acid donates a proton and becomes a conjugate base
30
strong acid
one that completely ionizes in solution
31
weak acid
one that partially ionizes in solution
32
HA(aq) + H2O(l) H3O+ + A-(aq)
- 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
33
Six important strong acids: HCl
Hydrochloric Acid Monoprotic
34
Six important strong acids: HBr
Hydrobromic Acid Monoprotic
35
Six important strong acids: HI
Hydriodic Acid Monoprotic
36
Six important strong acids: HNO3
Nitric Acid Monoprotic
37
Six important strong acids: HClO4
Perchloric Acid
38
Six important strong acids: H2SO4
Sulfuric Acid Diprotic
39
Monoprotic
containing only one ionizable proton
40
Diprotic
containing two ionizable protons
41
in weak acids - HA(aq) + H2O(l) H3O+ + A-(aq)
- 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
42
in general, the stronger the acid, the
weaker the conjugate base and vice versa
43
Six Weak Acids: HF
Hydrofluoric Acid Monoprotic
44
Six Weak Acids: HC2H3O2
Acetic Acid Monoprotic
45
Six Weak Acids: HCHO2
Formic Acid Monoprotic
46
Six Weak Acids: H2SO3
Sulfurous Acid Diprotic
47
Six Weak Acids: H2CO3
Carbonic Acid Diprotic
48
Six Weak Acids: H3PO4
Phosphoric Acid Triprotic
49
-protic
the number of ionizable protons
50
acid ionization constant(Ka)
- 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
because water is amphoteric(can act as an acid or base) it can autoionize
H2O(l) + H2O(l) H3O+(aq) + OH-(aq)
52
ion product constant for water(Kw)(or dissociation constant for water)
- Kw = [H3O+][OH-] = [H+][OH-] - at 25 degrees C Kw = 1.0 * 10^-14 [H3O+] = [OH-] = sqrtKw = 1.0 * 10^-7 - neutral pH
53
acidic solution
contains an acid that creates additional H3O+ ions causing the [H3O+] to increase - the ion product constant still applies - 1.0*10^-14
54
basic solution
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
neutral solution
[H3O+] = [OH-] = 1.0 * 10^-7 M (at 25 C)
56
Acidic Solution
[H3O+] > [OH-]
57
Basic Solution
[H3O+] < [OH-]
58
All AQ Solns
both H3O+ and OH- are present with [H3O+][OH-] = Kw = 1.0*10^-14 at 25 C
59
pH
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
At 25 C: pH < 7
an acidic solution
61
At 25 C: pH > 7
a basic solution
62
At 25 C: pH = 7
a neutral solution
63
pH scale is logarithmic so
a change of 1pH is a 10 fold change in H3O+ concentration
64
pOH
- pOH = -log[OH-] | - same as pH scale but with respect to OH- instead of H3O+
65
At 25 C: pOH <7
Basic Solution
66
At 25 C: pOH > 7
Acidic Solution
67
At 25 C: pOH = 7
Neutral Solution
68
pH + pOH =
14.00 Always.
69
pKa = -log(Ka)
the smaller the pKa the stronger the acid
70
Two sources of H3O+
- the ionization of the acid - autoionization of water - typically it can be ignored as it is such a small amount
71
In a strong acid solution:
- 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
In a weak acid solution:
- 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
Percent Ionization
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
For weak acids: - the equilibrium H3O+ concentration increases with increasing initial concentration of the acid
- 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
What happens when you mix a strong and weak acid?
the strong acid fully ionizes to produce H3O+ and suppresses the formation of additional H3O+ which could be formed by the weak acid
76
What happens when you mix two weak acids?
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
Strong Base
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
Bases: 1A and 2A
``` 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
What is the difference between a base dissociating two OH- compared to a diprotic acid?
unlike diprotic acids which ionize in two steps bases containing two OH- ions dissociate in one step
80
Weak base
analogous to a weak acid | - the most common weak bases produce OH- by accepting a proton from water, ionizing water to form OH-
81
Base General Form with weak base example: B(aq) + H2O(l) BH+(aq) + OH-(aq)
NH3(aq) + H2O(l) NH4+(aq) + OH-(aq) | - double arrow indicates the ionization is not complete
82
Base Ionization Constant
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
The smaller the Kb the
weaker the base
84
pKb =
- log(Kb) | - like pH
85
finding the pH of a strong or weak base is
analogous to finding the H3O+ and pH of acids
86
anions tend to be either
basic or neutral solutions
87
cations tend to be
acidic or neutral solutions
88
any anion is the conjugate base of an acid
Cl- in HCL
89
every anion can act as a base BUT not every anion does act as a base
dependent on the strength of the corresponding acid
90
an anion that is the conjugate base of a weak acid is itself a weak base
F- in HF is a weak base | - F-(aq)+H2Oz(l) OH-(aq) + HF(aq)
91
an anion that is the conjugate base of a strong acid is pH-neutral(forms Solutions neither acidic or basic)
Cl- in HCl is pH neutral
92
HF(aq)+H2O(l) H3O+(aq) + F-(aq)
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
the weaker the acid the
stronger the conjugate base
94
HCl(aq) + H2O(l) -> H3O+(aq) + Cl-
a strong acid | - the reaction equilibrium lies far to the right because H and Cl- are not attracted to each other
95
Kw=(Ka)(Kb) Find the pH of a solution containing an anion
Kw = 1.0*10^-14 - find Kb for conjugate base knowing Ka acetic acid is 1.8*10^-5 - Kb = Kw/Ka = 1.0*10^-14/1.8*10^-5 = 5.6*10^-10 - can use this Kb to calculate the pH of a solution containing an anion acting as a base - 5.6*10^-10 = (X^2)/(.100-x) - do the approx. trick and cancel x in denominator(check validity at end) - x=2.4*10^-6 - 2.4*10^-6/.100 * 100% = .0024%(assumption valid!) - [OH-] = 2.4*10^-6 M - [H3O+][OH-]=Kw - [H3O+](2.4*10^-6)=1.0*10^-14 - [H3O+] = 4.2*10^-9 M - pH = -log[H3O+] = -log(4.2*10^-9) = - 8.38 = pH
96
pKa + pKb =
14
97
cations can sometimes act as
weak acids
98
3 categories of cation interaction
- cations that are counterions of strong bases - cations that are conjugate acids of weak bases - cations that are small, highly charged metals
99
cations that are counterions of strong bases
- 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
Cations That are Conjugate Acids of Weak Bases
- 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
Cations that are Small highly Charged Metals
- 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
salts in which neither the cation nor the anion acts as an acid or base form pH-neutral solutions
- 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
salts in which the cation does not act as an acid and the anion acts as a base for basic solutions
- 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
salts in which the cation acts as an acid and the anion does not act as a base form acidic solutions
- 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
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
- 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
Cation/Anion table
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
typically polyprotic acids ionize in successive steps, each with its own Ka
- H2SO3(aq) H+(aq)+HSO3-(aq) Ka1 = 1.6*10^-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
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
treat the problem as a weak acid ICE table
109
Binary acid
an acid containing hydrogen and only one other element | - H-Y
110
In a binary acid, the ease with which the hydrogen is donated(and acidic) is based on
polarity and bond strength
111
In a binary acid: Increasing electronegativity
increases acidity
112
In a binary acid: Decreasing bond strength
increases acidity
113
In a binary acid, H-Y bond must be polarized with the H on the positive end to become acidic
- 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
In a binary acid, H-Y bond strength is important
- 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
In binary acids, 6A and 7A elements become more
acidic left to right as electronegativity increases | - become more acidic down the column as bond strength decreases
116
Oxyacid
an acid containing hydrogen bonded to an oxygen atom that is bonded to another element - H-O-Y
117
Oxyacid: affected by the electronegativity of Y and the number of oxygen atoms attached to Y
- the more electronegative the greater the Ka
118
Oxyacid: the more Os bonded the
greater the electronegativity and stronger the acid
119
the lewis model is focused on
the electron pair instead of H+(Bronsted-Lowry definition)
120
Lewis Acid
electron pair acceptor
121
Lewis Base
electron pair donor
122
H+ + :NH3->[H:NH3]^+
- 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
the Lewis model significantly expands
the substances that can be considered acids without altering the substances that can be considered bases significantly
124
BF3+ :NH3 -> F3B:NH3
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)

Brainscape helps you reach your goals faster, through stronger study habits.
© 2025 Bold Learning Solutions. Terms and Conditions