General Chemistry- Acids and Bases Flashcards
1
Q
What is an Arrehnius Acid?
A
An acid will dissociate to form an excess of H+ in solution
2
Q
What is an Arrhenius base?
A
A base that will dissociate to form an excess of OH- in solution
3
Q
Arrhenius acids and bases behavior is limited to what?
A
Limited to aqueous acids and bases
4
Q
How are Arrhenius acids and bases identified?
A
Acids contain H at the beginning of their formula (HCl, HNO3, H2SO4) and bases contain OH at the end of their formula (NaOH, Ca(OH)2, Fe(OH)3)
5
Q
What is more a inclusive definition of acids and bases compared to Arrhenius?
A
A Bronsted-Lowry acid: A species that donates hydrogen ions (H+)
A Bronsted-Lowry base: a species that accepts them.
6
Q
What is the advantage of the Bronsted-Lowry definition over the Arrhenius definition?
A
It is not limited to aqueous solutions.
7
Q
What are examples of Bronsted-Lowry bases?
A
OH-, NH3, and F-
Because each has the ability to accept a hydrogen ion.
8
Q
Is water an acid? Based on the Arrhenius definition?
A
No, it does not produce an excess of H+ in solution
9
Q
Is water an acid? Based on the Bronsted-Lowry definition?
A
Yes, because it is able to donate a proton to other species.
10
Q
Which definition will most likely be used on the MCAT when dealing with acids and bases?
A
The Bronsted- Lowry
11
Q
Bronsted-Lowry acids and bases always, what? Why?
A
Occur in pairs because the definitions require the transfer of a proton from the acid to the base
12
Q
What are conjugate acid-base pairs?
A
Bronsted-Lowry acids and bases
13
Q
What is a Lewis acid?
A
An electron pair acceptor.
14
Q
What is a lewis base?
A
An electron pair donor
15
Q
What are the other names for Lewis acid base chemistry?
A
Coordinate covalent bond formation,
Complex ion formation
Nucleophile-electrophile interaction
The underlying idea is that one species pushes a lone pair to form a bond with another
16
Q
What is the difference between the Lewis definition and the Bronsted-Lowry definition?
A
The focus.
Bronsted-Lowry acids and bases, we follow the exchange of the hydrogen ion (H+), which is essentially a naked proton, In the Lewis definion, the focus of the reaction is no longer on the proton, but instead the electrons forming the coordinate covalent bond. This can be seen using curved arrows.
17
Q
Which acid-base definition is the most inclusive?
A
Lewis definition
18
Q
Can all Lewis acids and bases be Bronsted-Lowry? Or Arrhenius?
A
Lewis definition encompasses some species not included in Bronsted-Lowry, and Arrhenius. Every Arrhenius acid is a Bronsted-Lowry acid, and Every Bronsted-Lowry is a Lewis
19
Q
When are Lewis acids usually used?
A
In organic chemistry reaction, Lewis acids are usually catalysts.
20
Q
What’s an amphoteric species?
A
A species that reacts like an acid in a basic environment and like a bse in an acidic environment.
21
Q
What is the most common amphoteric species on the MCAT?
A
Water is the most common example
22
Q
When a water reacts with a base, how does it behave? Acid?
A
As an acid:
H2O + B HB + OH
In acid, it acts as a base:
HA + H2O H3O+ + A-
23
Q
When a polyvalent acid partially dissociates, what does it become?
A
Amphoteric
24
Q
What is an example of a partially dissociated conjugate base of a polyvalent acid?
A
HSO4- can either gain a proton to form H2SO4 or lose a proton to form SO4 2-
25
What are also examples of amphoteric?
The hydroxides of certain metals (Al, Zn, Pb, and Cr)
Species that can act as both oxidizing and reducing agents are often considered to be amphoteric as well because of accepting or donating electron pairs, they act as Lewis acids or bases.
26
What are complex amphoteric molecules?
Amino acids that are a zwitterion intermediate with both cationic and anionic character.
27
Acids formed from anions with the name that end with -ide change to what?
Have the prefix hydro- and ending in -ic
28
What are acids formed from oxyanions called?
Oxyacids
29
If the anion ends in -ite (less oxygen), then the acid will end, how?
With -ic acid
30
If the anion ends in -ate (more oxygen), then the acid will end how?
-ic acid
31
What is autoinization?
As an amphoteric compound, water can react with itself
32
Which compound can go through autoionization?
Water
| H2O (l) + H2O (l) H3O+ (aq) + OH- (aq)
33
How does water go through autoionization?
One water molecule donates a hydrogen ion to another water molecule to produce the hydronium ion (H3O+) and the hydroxide ion (OH-).
34
Is autoionization reversible?
Yes
35
What is Kw?
Kw= [H3O+][OH-] = 10^-14 at 25 degrees celsius (298k)
36
What is the concentration of each ion in pure water at equilibrium at 298k if Kw is 10^-14
The hydrogen ions and the hydroxide ions are always equal in pure water at equilibirum
37
When is the only time hydgrogen ion and hydroxide ions will be equal?
When the solution is neutral
38
What will the product of hydrogen ion concentration and hydroxide ion concentration always be?
10^-14 when the temperature of the solution is 298K
39
What is Kw?
Is an equilibirum constant
40
When will Kw change?
The temperature of the water is changed
41
Temperatures avove 298K will cause Kw to do what?
Kw will increase
42
What scales are used to measure hydrogen and hydroxide ions?
pH and pOH
43
What is a p scale?
The negative logarithm of the number of items
44
How do you calculate the pH and pOH of a solution?
```
pH= -log[H+]= log(1/[H+])
pOH= -log[OH-]= log(1/[OH-])
```
45
As pH increases, what happens?
pOH decrease by the same amount
46
If a solution is acidic, what type of ions is in excess?
hydrogen ions
47
If a solution is basic what type of ions is in excess?
Hydroxide ions
48
When determining the p scale value for a solution with a base value of a power of 10, what are examples?
If [H+]=0.001 or 10^-3, then pH=3 then pOH=11
| If Kb=1.0X10^-12 then pKb= 12
49
When the base number, when determining the p scale value of a solution, is not 10, what equation can you use? Using the form nX10^-m
-log(nx10^-m)= m - log(n)
50
When using the equation m-log(n), what does the answer mean when n is closer to 1? What about 10?
1: the closer log(n) will be to 0
10: the closer log(n) will be to 1
51
What equation can be used to simplify the equation m - log(n), to get a better estimate of the p-values?
p value =(about) m - 0.n
| 0.n represents sliding the decimal point of n one position to the left
52
What are strong acids and bases?
Acids and bases that completely dissociate into their component ions in aqueous solutions
53
What happens when a strong acid or base gives into a solution
It will not remain in solution. The dissociation goes into completion
54
When is the autoionization of water negligible and when is it important?
If the concentration of the acid or base is significantly greater than 10^-7 M, then the autoionization of water is neglible.
If the concetration of the acid or base is close to 10^-7 M, then the contribution from the autoionization of water is important.
55
What are strong acids commonly found on the MCAT?
```
HCL (hydrochloric acid)
HBr (hydrobromic acid
HI (hydroiodic acid)
H2SO4 (Sulfuric acid
HNO3 (nitric acid)
HClO4 (perchloric acid)
```
56
What are strong bases commonly found on the MCAT?
NaOH (sodium hydroxide)
KOH (potassium hydroxide)
Other soluable hydroxides of Group IA metals
57
Calculation of the pH and pOH of strong acids and bases assumes what?
Complete dissociation of the acid or base in solution
58
What type of wording is prefered when talking about acids nad bases? Why?
Concentrated and dilute instead of weak and strong
| Because they are unambiguously associated with concentrations, rather than chemical behavior.
59
What are weak acids and bases?
Those acids and bases that only partially dissociate in aqueous solutions.
60
If partially dissociated acids or bases reach equilibirum, what equation can you use?
Acid dissociation constant Ka= [H3O+][A]/ [HA]
61
The smaller the Ka, means what?
The weaker the acid and the less it will dissociate
62
How can the base dissociation constant (Kb) be calculated?
Kb= [B+][OH]/[BOH]
63
The smaller the Kb, means what?
The weaker the base, and the less it dissociates
64
Generally speaking, how can you characterize a species as a weak acid or base?
If the Ka is less than 1.0 or the Kb less then 1.0
65
On the MCAT, molecular (nonionic) weak bases are almost always what?
exclusively amines
66
What are conjugate acids and bases?
The acid formed when a base gains a proton
| A base formed when an acid loses a proton.
67
What equation helps you find the dissociation constant for another species if one constant is known?
Kacid x Kconjugate base= Kw = 10^-14
| Kbase x Kconjugate acid = Kw= 10^-14
68
Ka, and Kb are ________ related?
Inversely
69
A strong acid will produce what?
A very weak conjugate base (HCL will produce Cl-)
70
A strong base will produce what?
A very weak conjugate acid.
71
What is the conjugate of a strong acid or base sometimes called? Why?
Inert
| Because it is almost completely unreactive
72
What does a weak acid or base produce?
Conjugates that are also weak
73
Which equilibrium equation is ideal for buffering solutions? Why?
CO3 2- with water to produce HCO3- and OH- occurs to a greater extent-- is more thermodynamically faborable-- than the reaction of HCO3- and water to produce CO3 2- and H3O+. This fact makes this equilibirum ideal for buffering solutions as part of the bicarbonate buffer system.
74
What is one important theme for acid strength?
The effect of induction.
75
Acids that have electronegative element nearer to acidic hydrogen are what? Why?
Increase the acid strength and makes the acids stronger.
The electronegative elements pull electron density out of the bond holding the acidic proton. This weakens proton bonding and facilitates dissocation.
76
What is the most common use of acid and base dissociation constants?
To determine the concentration of one of the species in solution at equilibrium.
77
In regards to acids and bases, what might you be asked to calculate on test day?
The concentration of the hydrogen ion (or pH), the concentration of the hydroxide ion (or pOH), or the concetration of either the original acid or base.
78
What is the rule of thumb for approximations? When does it typically occur?
The approximation is valid as long as "x" is less then 5 percent of the initial concentration.
This occurs when Ka is at least 100 times smaller than the concetration of the starting solution.
79
What is a neutralization reaction?
Acids and bases may react with each other to form a salt and often water
80
What is an example of a neutralization reaction?
HA (aq) + BOH (aq) -> BA (s) + H2O (l)
81
What might the salt do in a neutralization reaction? What is it based on?
Precipitate out or remain ionized in solution
| Dependent on its solubility and the amount produced
82
In general, what happens in a neutralization reaction?
Neutralization reactions go to completion
83
What is hydrolysis?
The reverse of neutralization reaction. The salt ions react with water to give back the acid or base.
84
What are the possible 4 combinations of strong/weak acids or bases?
Strong acid+ Strong base: HCl + NaOH --> NaCL + H2O
Strong acid + Weak base: HCl + NH3 --> NH4Cl
Weak acid + strong base: HClO + NaOH --> NaClO +H2O
Weak acid + weak base: HClO + NH3 --> NH4ClO
85
What is the product of a strong acid and strong base combination?
Equimolar amounts of salt and water. The acid and base neutralize each other, so the resulting solution is neutral (pH=7), and the ions formed in the reaction will not react with water because they are inert conjugates.
86
What is the product of a strong acid and weak base combination? Why?
Also salt, but often no water will be formed because weak bases are often not hydroxides
87
What is the reverse reaction to HCl (aq) + NH3 --> NH4+ (aq) + Cl- (aq)? What is the reactions name?
NH4+ (aq) + H2O (l) --> NH3 (aq) + H3O+ (aq)
| Hydrolysis
88
For the reaction HCl (aq) + NH3 --> NH4+ (aq) + Cl- (aq), Explain what is happening in the reverse reaction.
NH4+ is the conjugate acid of a weak base and is stronger than the conjugate base of the strong acid. NH4+ will then transfer a proton to H2O to form the hydronium ion. The increase in the concentration fo the hydronium ion causes the system to shift away from autoionization, therby reducing the concentration of hydroxide ion. The concentration of the hydronium will be greater than that of the hydroxide ion at equilibrium, this will result in the pH of the solution to fall below 7.
89
What is the product of a weak acid and strong base combination?
When a weak acid reacts with a strong base, the pH of the solution at equilibrium will be within the basic range because the salt hydrolyzes, with the concurrent formation of hydroxide ions. The increase in hydroxide ion concentration will cause the system to shift away from autoionization, thereby reducing the concentration of the hydronium ion. The concentration of the hydroxide ion will be greater than that of the hydronium ion at equilibirum and result in the pH to rise above 7.
90
The pH of a solution with a weak acid and a weak base is dependent on what?
the relative strengths of the reactants.
91
If a weak acid HClO has a Ka of 3.2 x 10^-8, and a weak base has a Kb of 1.8 x 10^-5, the solution will end up being _______. Why?
Basic in aqueous solution
Because the Ka for HClO is less than the Kb for NH3. The HClO is weaker as an acid than NH3 is as a base. Thus the hydroxide ions are greater than the hydronium ions in solution
92
In biology and biochemistry, _____ reactions are often ______ reactions because _________.
Neutralization reactions are often condensation reactions because they form bonds with a small molecule as a byproduct (usually water).
93
What is an example of a biology acid and base reaction?
Peptide bonds in proteins are created from the reaction of carboxyl group (acid) and an amino group (base) while forming a water molecule. The salt in the reaction is the polypetide itself.
94
The acidity or basicity of an aqueous solution is determined by what?
The relative concentrations of acid and base equivalents.
95
What is an acid equivalent?
Equal to one mole of H+ (or more properly H3O +) ions
96
What is a base equivalent?
Equal to one mole of OH- ions.
97
What is a polyvalent?
Each mole of the acid or base liberates more than one acid or base equivalent.
98
What's another name for a polyvalent, under the bronsted-lowry definition?
Polyprotic
99
The quantity of acidic or basic capacity is directly indicated by what?
The solution's normality
100
What is a gram equivalent weight?
The mass of a compound that produces one equivalent (one mole of charge)
101
What are some common polyvalent acids?
H2SO4
H2PO4
H2CO3
102
What are some common polyvalent bases?
Al(OH)3
Ca(OH)2
Mg(OH)2
103
What is titration?
A procedure used to determine the concentration of a known reactant in a solution.
104
what are the different types of titration?
Acid-base
Oxidation-Reduction
Complexometric (metal ion)
105
What 2 types of titration does the MCAT test?
Acid-base
| Oxidation-Reduction
106
How is titration generally performed?
By adding small volumnes of a solution of known concentration (the titrant) to a known volume of solution of unknown concentration (the titrand) until completion of the reaction is achieved at the equivalence point.
107
In acid-base titration, the equivalence point is reached when?
The number of acid equivalents present in the original solution equals the number of base equivalents added.
108
What is important to note in a acid-base titration?
While a strong acid/strong base titration will have its equivalence point at pH 7, the equivalence point does not always occur at pH 7. When titrating polyprotic acids or bases there are multiple equivalence points
109
What equation is used to calculate the unknown concentration of the titrand in a acid-base titration?
NaVa = NbVb
110
What do the letters represent in the titration equation?
Na and Nb: are the acid and base nomalities
Va and Vb: are the volumes of acid and base solutions.
As long as the units are the same, it does not have to be liters.
111
What are the two common ways of determining the equivalence point in an acid-base titration?
Graphicall-plotting the pH of the unknown solution as a function of added titrant by using a pH meter
Estimated- watching the color change of an added indicator
112
What are indicators?
Weak organic acids or bases that have different colors in their protonated and deprotonated states.
113
How do indicators work?
The structural change- binding or release of protons- leads to a change in the absorption spectrum of the molecules. Indicated by a color change
114
Indicators are generally __________ and can be used in ___________.
Vibrant
| Low concentrations without significantly altering the equivalence point
115
The indicators must always be what? Otherwise what happens?
Weaker acids or bases than the acid or base being titrated. Otherwise, the indicator would be titrated first.
116
What is the point at which the indicator changes to its final color?
Not the equivalence point but the endpoint.
117
How do you select an ideal indicator?
Know the pH of the reaction at the equivalence point, whether graphically or mathematically. Once you have determined where the equivalence point is, select the indicator that has the closest pKa value to it.
118
What are the most useful combinations of acid-base reactions? What aren't?
At least one strong species,
| Weak acid/base titrations are inaccurate and rarely performed.
119
What makes weak acid-base titration not very accurate?
The pH curve for the tration of the weak acid base lacks the sharp change that normally indicates the equivalence point. Furthermore indicators are less useful because the pH change is far more gradual
120
If using a pH meter to approximate the equivalence point, what would you do?
Locating the midpoint of the region of the curve with the steepest slope.
121
Why is the inital pH of a weak acid greater than the initial pH of a strong acid?
Weak acids do not dissociate to the same degree that strong acids do, therefore the concentration of H3O will generally be lower in equimolar solution of weak acid.
122
Why doe sthe pH meter of a strong acid/base titration show a steeper curve then the weak acid/base curve?
In the weak acid/base titration, the pH changes gradually early on in the titration and has a less sudden rise at the equivalence point.
123
What is the difference in equivalence point for a strong acid/base vs. a weak acid/base titration using a pH meter? Why?
The position of the equivalence point.
The equivalence point of a weak acid/base is above pH 7 because the reaction between the weak acid and strong base produce a weak conjugate base and even weaker conjugate acid. This produces a greater concentration of hydroxide ions then hydrogen at equilibrium
124
What is a strong acid/weak base titration curve inverse to?
The curve for a weak acid titrand and strong base tirant.
125
Why is the equivalence point in the acidic range for a strong acid/weak base titration curve?
The reaction between the weak base and strong acid will produce a weak conjugate acid and even weaker conjugate base. The stronger conjugate acid will result in an equilibrium state with a concentration of hydrogen ions greater than that of hydroxide ions.
126
What will a titration curve look like for a weak acid titrant weak base titrand combination?
Both the titrant and the tirand are weak, the initial pH is generally in the 3-11 range and will demonstrate a very shallow drop at the equivalence point. The equivalence point will be near neutral pH because the reaction is partially dissociative for both species.
127
How can you tell if a tiration curve is for polyvalent acid or base or monvalent?
The multiple equivalence points indicate it is polyvalent titration.
128
What is region I in a polyvalent titration curve?
Little acid has been added, and the predominat species is CO3 2-
129
What is region II, in a polyvalent titration curve?
More acid has been added, and the predominant species are CO3 2- and HCO3-
130
What is the flat part of the curve considered in a polyvalent titration?
The first buffer region corresponding to the pKa of HCO3 (Ka= 5.6 x 10^-11, and pKa= 10.25)
131
What is the half-equivalence point?
The center of the buffer region (the point in between regions I and II).
Where half of the given species has been protonated
132
What is region III, in a polyvalent titration curve?
Starts with the equivalence point. In the latter part of region III, the predominat species is HCO3- although some H2CO3 has formed.
133
What is region IV, in a polyvalent titration curve?
At the beginning of region IV, the acid has neutralized half of the HCO3- and now H2CO3 and HCO3- are in equal concentrations. This flat region is the second buffer region of the titration curve, corresponding to pKa of H2CO3. (Ka= 4.3 x 10^-7, pKa= 6.37
134
What is region V in a polyvalent titration curve?
Starts with the second equivalence point, A rapid change in pH is observed near the equivalence point as acid is added.
135
Titrations of acidic and basic amino acids will have what type of curves?
Curves similar to polyvalent curves, but rather than 2 equivalence points it will have 3. One corresponding to the titration of the carboxyl group and a second corresponding to the titration of the amino group. Both attached to the central carbon, as well as a third corresponding to either the acidic or basic side chain.
136
What does a buffer solution consist of?
A mixture of a weak acid and its salt (which is composed of its conjugate base and cation) or a mixture of a weak base and its salt (which is composed of its conjugate acid and anion.)
137
What are the two common buffers tested on the MCAT?
Solution of acetic acid (CH3COOH) and its salt: Sodium acetate (CH3COO-Na+)
Solution of ammonia (NH3) and its salt: ammonium chloride (NH4+Cl-)
138
The common buffer acetic acid/sodium acetate solution is what type of buffer?
Acid buffer
139
The common buffer ammonium chloride/ammonium solution is what type of buffer?
Base buffer
140
Buffer solutions have what useful property?
Resisting changes in pH when small amounts of acid or base are added.
141
How does a buffer work?
The weak acid component of the buffer thereby serves to neutralize the strong base that has been added. The resulting increase in the concentration of the acetate ion (the conjugate base) does not create nearly as large an increase in hydroxide ions as the unbuffered NaOH would. Thus, the addition of the strong base does not result in a significant increase in [OH-] and does not appreciably change the pH.
When a small amount of HCl is added to the buffer, H+ ions from the HCl react with the acetate ions to form acetic acid. acetic acid is weaker than the added hydrochloric acid, so the increased concentration of acetic acid does not significantly contribute to the hydrogen ion concentration in the solution. Because the buffer maintains the [H+] at constant values, the pH of the solution is relatively unchanged.
142
What is the most important buffers in the human body?
One of the most important buffers is the H2CO3/HCO3 conjugate pain in the plasma component of the blood, called the bicarbonate buffer system.
143
What specific compounds are respondible for keeping the pH of the blood neutral?
Carbonic acid (H2CO3) and its conjugate base, bicarbonate (hCO3) form a weak acid buffer maintaining the pH of the blood within the fairly narrow physiological range.
144
What type of waste products does the bicarbonate buffer system take care of?
CO2
145
What reaction takes place with CO2 and the bicarbonate buffer system?
CO2 (g) + H2O (l) H2CO3 (aq) H+ (aq) + HCO3- (aq)
146
The bicarbonate buffer system is tied to what system?
Respiratory system
147
What happens when metabolic acidosis occurs?
The breathing rate will increase to compensate and blow off a greater amount of carbon dioxide gas, this causes the system to shift to the left, therby reducing [H+] and buffering against dramatic and dangerous changes in blood pH.
148
What is metabolic acidosis?
Production of excess plasma H+ not caused by the resiratory system itself.
149
The bicarbonate buffer system maintains a pH around what?
7.4, which is actually slightly outside the optimal buffering capacity of the system
150
Buffers have a narrow range of what?
optimal activity (pKa +or- 1)
151
What is acidemia?
Too much acid in the blood
152
What is alkalemia?
Too much base in the blood
153
What is more common, acidemia or alkalemia?
Acidemia
154
When does the buffer system become more effective in the human body?
When acidemia becomes more severe
155
What is the Henderson-Hasselbalch equation?
Used to estimate the pH or pOH of a buffer solution
156
What's the henderson-Hasselbalch equation for a weak acid buffer solution?
pH= pKa + log [A-]/[HA]
157
What do the letters mean in the Hendersen-Hasselbalch equation for a weak acid buffer solution?
[A-]: the concentration of the conjugate base
| [HA]: concentration of the wak acid
158
What is important to note when [conjugate base] = [weak acid]?
The pH=pKa, because log(1) =0.
| This occurs at the half-equivalence points in a titration and buffering capactiy is optimal at this pH.
159
What is the Hendersen-Hasselbalch equation for a weak base buffer solution?
pOH= pKb + log [B+]/[BOH]
160
What do the letters represent in the Hendersen-Hasselbalch equation for a weak base buffer solution?
[B+]: the concentration of conjugate acid
| [BOH]: the concetration of the weak base
161
At what pOH is buffering capacity optimal?
When pOH=pKb or [conjugate acid]= [weak base]
162
The Henderson-Hasselbalch equation is really just a rearrangement of what?
```
Acid dissociation constant
Ka=[H3O+][A-]/[HA]
-log Ka= -log [H3O+][A-]/[HA]
-log Ka= -log [H3O+] - log[A-]/[HA]
pKa=pH-log[A-]/[HA]
pH=pKa + log[A-]/[HA]
```
163
Changing the ratio of the conjugate base to the acid will lead to what in a buffering system?
A chang ein the pH of the buffer system
164
What happens if you change the concentrations of the conjugate base and acid while maintaining a constant ratio?
The pH would not change, but the buffering capacity would double or triple depending on the change in concentration. In other words, addition of a small amount of acid or base to the doubled system would now cause even less diviation in the pH. The buffering capacity is generally maintained within 1 pH unit of the pKa value.
165
What is buffering capacity?
The ability to which the system can resist changes in pH.
