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Acids And Bases Flashcards

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1
Q

Lewis Acid and Base

A

Defines an acid as any substance that accepts a pari of electrons
Defines a base as any substance that donates a pair of electrons
Includes all acids and bases in Bronsted-Lowry definition and more
Include molecules with an incomplete octet around central atom (AlCl3 and BF3)
Include all simple cations except alkali and heavier alkaline earth metal cations
Smaller the cation and higher the charge, the more electrophilic in nature and stronger the acid strength
E.g. Fe^3+ is Lewis Acid

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2
Q

Arrhenius Acid and Base

A

Acid: substance that produces hydrogen ions (H+) in water
Base: substance that produces hydroxide ions (OH-) in water
Definition only covers aqueous solutinos

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3
Q

Bronsted-Lowry Acid and Base

A

Acid: Any substance that donates a proton (H+)
Base: Any substance that accepts a proton

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4
Q

What is a convenient way to think about acids and bases?

A

Convenient to think of an acid as H+ and a base as OH-

Aqueous solutions always contain both H+ and OH-

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5
Q

What does the extent to which an acid will increase the concentration of H+ in a solution depend on?

A

Depends on the acid’s tendency to lose or hold onto its hydrogen
An acid with a weak hold on its hydrogen can lose it easily and is considered a strong acid, while an acid with a stronger hold on its hydrogen gives it up less readily and is considered a weak acid
Three aspects:
1. Strength of the bond holding the H to the molecule
2. Polarity of the bond
3. Stability of the conjugate base

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6
Q

What is an example of polarity effecting acidity?

A

Comparing C-H in methane to H-Cl, we can see that the bond strength is nearly equal, however the polarity of the H-Cl bond is much greater
This means that the proton is more easily removed in aqueous solution and HCl is far more acidic than methane
However, the most polar does not always mean the most acidic
E.g. with H-F the bond is the most polar, however HF is the least acidic of the hydrogen halides because its conjugate base is very unstable
Comparing HCl to HF, the small size of the fluoride ion causes its negative charge to be more concentrated than that of chloride, causing the fluoride ion to be more unstable

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7
Q

What are some rules of thumb for strength of oxyacids?

A

Oxyacid: Acidic compound that contains an oxygen
Electronegative oxygen draws electrons to one side of bond with H, increasing polarity
Oxygens in conjugate base of oxyacid can share negative charge, spreading it over larger area
In similar oxyacids, molecule with most oxygens is the strongest acid
Therefore H2SO4 is much stronger acid than H2SO3

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8
Q

Rank the following oxyacids in order of decreasing acidity:

HCl, HClO2, HClO3, HClO4

A

HClO4 (Perchloric acid) > HClO3 (Chloric acid) > HClO2 (Chlorous acid) > HClO (Hypochlorous acid)

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9
Q

How can the stability of a base be determined?

A

Strength of a base can be predicted base don stability of resulting species
E.g. with NaOH, the product of dissociation is Na+, a very stable cation
NaOH is therefore a strong base that readily dissociates
Can also be thought of as the tendency to accept a proton
Protonation of OH- produced by the dissociation of NaOH stabilized the negative charge, creating the more stable H2O

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10
Q

What are some qualities of organic acids and bases?

A

Organic acids may include methanol or acetic acid (carboxylic acid). Acetic acid is much more strong, because negative charge once deprotonated is stabilized throughout both oxygen atoms
Organic bases contain nitrogen as a proton acceptor
Methyl amine is less basic than guanidine because positive charge that is formed on guanidine can be resonance stabilized by three nitrogen atoms

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11
Q

Which amino acid is more basic, lysine or arginine?

A

Arginine is more basic because it has three Nitrogen atoms to stabilize positive charge
Arginine (R): NH2. O
C NH CH2 CH2 CH2 C C OH
N+H2 +H3N
Lysine (K): NH3+. O
CH2 CH2 CH2 CH2 C OH
NH3+

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12
Q

What does a strong acid mean?

A

A strong acid has a weak hold on its hydrogen, so when dissolved in water, acid completely dissociates into H+ and conjugate base
Strong acid is stronger than H3O+
Therefore, since HCl is a strong acid, a 1M aqueous solution of HCl contains 1M of H+ and 1M of Cl-

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13
Q

What does a strong base mean?

A

Completely dissociates in water
Strong base is stronger than OH-, either as strong or stronger than NaOH
1M aqueous solution of NaOH contains 1M of Na+ and 1M of OH-

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14
Q

Polyprotic Acids

A

Acids that can donate more than one proton
Second proton donated by a polyprotic acid is usually so weak that its effect on the acidity of the solution is negligible
Second proton from H2SO4 is strongly acidic, except in dilute concentrations (less than 1M), has negligible effect on hydrogen concentration of H2SO4 solution
H2SO4 is so much stronger than HSO4-
Percent dissociation of acid decreases as acidity of solution increases

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15
Q

What are the strong acids that you should know?

A

Hydroiodic acid (HI), Hydrobromic acid (HBr), Hydrochloric acid (HCl), Nitric acid (HNO3), Perchloric acid (HClO4), Chloric acid (HClO3), Sulfuric acid (H2SO4)

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16
Q

What are the strong bases you should know?

A

Sodium hydroxide (NaOH), Potassium hydroxide (KOH), Amide ion (NH2-), Hydride ion (H-), Calcium hydroxide (Ca(OH)2), Sodium oxide (Na2O), Calcium oxide (CaO)

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17
Q

Weak Acids and Bases

A

Reactions of weak acids and bases do not go to completion
Only a small fraction of the reaction proceeds under normal conditions
All acids that aren’t strong acids are considered weak
E.g. Acids: Hydrofluoric acid (HF), Hydrocyanic acid (HCN), Acetic acid (CH3COOH), Water (H2O)
Bases: Ammonia (NH3), Ammonium hydroxide (NH4OH), Pyridine (C5H5N), Water (H2O)

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18
Q

How do acid dissociation and acid strength change with increasing acid concentration?

A

Percentage of dissociation of the acid decreases as you increase the concentration of acid (assuming a weak acid)
Acid strength increases as you increase the acid concentration
More hydrogen ions overall dissociate, however the percentage that dissociate goes down
More hydrogen ions in the same location, means lower pH and increases acid strength

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19
Q

What is the usual amount of dissociation for acetic acid?

A

Equilibrium of reaction strongly favors the reactants

One out of every 1000 acetic acid molecules will be deprotonated at any one time in a 1M solution of acetic acid

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20
Q

Hydrides

A

Binary compounds (compounds with only two elements) that contain hydrogen
Can be basic, acidic, or neutral
Basic hydrides are to the left on the periodic table, acidic hydrides are to the right
NaH is basic and H2S is acidic
Metal hydrides are either basic or neutral, nonmetal hydrides are acidic or neutral (Ammonia is an exception)
Acidity of nonmetal hydrides increases going down periodic table: H2O < H2S < H2Se < H2Te

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21
Q

Conjugate Acids and Bases

A

If there is an acid in a reaction, there must also be a base to accept the Hydrogen ion
Hypothetical acid-base reaction in aqueous solution:
HA + H2O -> H3O+ + A-
HA is acid, water is base, H3O+ is the conjugate acid, and A- is the conjugate base
The stronger the acid, the weaker the conjugate base
Deciding which form is called the conjugate simply depends on the direction of the reaction

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22
Q

pH

A

Measurement of hydrogen ion concentration, generally runs from 0 to 14, but more extreme values are possible
PH = -log[H+]
P(x) is a function in which, given any x, p(x) = -log(x)
A pH of 7 is neutral, higher is basic, lower is acidic
A pH of 2 has 10 times as many hydrogen ions as a solution with a pH of 3

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23
Q

How can you estimate the pH of a hydrogen ion concentration of 10^-3? 4 x 10^-3? 10^-2?

A

We know that pH = -log[H+]
This means, for [H+] = 10^-3, we have: -log[10^-3] -> -(-3) = 3
For 4 x 10^-3 this is larger than 10^-3, but not as large as 10^-2, so probably somewhere between 2 and 3, maybe 2.4
For 10^-2, following the same pattern as for 10^-3 we have pH = 2

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24
Q

Amphoteric

A

Means that a molecule is either an acid or a base depending on their environment
E.g. water, HA + H2O -> A- + H3O+ (water acts as base), but also water can act as an acid: A- + H2O -> HA + OH-
Amino acids (carboxylic acid and amine group) can carry multiple charges depending on their environment
- At pH of 7, most AAs have protonated amine group, deprotonated carboxylic acid group (both positive and negative charges, zwitterion)

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25
Q

Organic Acids

A

Unlike inorganic acids, organic acids can be resonance stabilized
- Resonance stabilization is one factor that contributes to acidity of organic molecules, molecule can distribute charge among constituent atoms that participated in resonance
- However, generally organic molecules do not like carrying a net charge
Includes molecules like amino acids

26
Q

What is the relative acidity of these organic acids: CH3CH2OH, H2O, CH3COH, CH3CH3, CH2CH2, CH3COOH, CHCH, NH3, H2

A

CH3-CH3 (weakest acid) < CH2=CH2 < NH3 < H2 < CH=-CH < CH3COH < CH3CH2OH < H2O < CH3COOH (strongest acid)

27
Q

Autoionization of water

A

H2O + H2O -> H3O+ + OH-

Pure water reacts with itself to form hydronium and hydroxide ions

28
Q

What happens when a weak acid, HA, is added to pure water?

A

Three reactions occur simultaneously:
HA + H2O -> A- + H3O+
A- + H2O -> HA + OH-
H2O + H2O -> H3O+ + OH-
Considering Le Chatelier’s, the top reaction will shift towards the right with higher concentration of HA, only a small amount of A- will be produced however due to the weak acid. This causes the second reaction to shift to the left, reducing OH- conc.
The result is a significant increase in H3O+ and OH- will about remain the same
This pushes the third reaction to the left, reducing the hydronium and hydroxide ions
As it reaches equilibrium, concentration of H3O+ ions at equilibrium will be more than OH-

29
Q

What is the equilibrium constant for autoionizaton of water?

A

Kw = [H3O+][OH-]
At 25 deg C and 1 atm, equilibrium for this reaction is Kw = 10^-14
Addition of an acid or base to aqueous solution will change concentrations of both H3O+ and OH-, but Kw will remain at 10^-14 at 25 deg C
For example, in a solution with a pH of 2, [H3O+] = 10^-2 mol / L and [OH-] = 10^-12 mol / L
pH + pOH = pKw, OR at 25 deg C: pH + pOH = 14

30
Q

Acid Dissociation Constant

A

Ka
Equilibrium constant in water
In hypothetical acid-base reaction: HA + H2O -> H3O+ + A-
Ka = ([H3O+][A-])/[HA]
Corresponding to every Ka, there is a Kb
Kb is equilibrium constant for reaction of conjugate base with water: A- + H2O -> OH- + HA, Kb = ([OH-][HA])/[A-]

31
Q

What is the product of Ka and Kb?

A

KaKb = ([H+][A-]/[HA]) x ([OH-][HA]/[A-]) = [H+][OH-] = Kw
KaKb = Kw
Formula can also be written: pKa + pKb = pKw OR pKa + pKb = 14 at 25 dec C
Ka: Equilibrium constant for any reaction in which an acid reacts with water to produce hydroniuim ion and a conjugate base
Kb: Equilibrium constant for any reaction in which a base reacts with water to produce a hydroxide ion and a conjugate acid

32
Q

What are general rules for how big Ka and Kb are and what that means?

A

The larger the Ka and the smaller the pKa, the stronger the acid
Ka > 1 or pKa < 0 indicates a strong acid

The larger the Kb and the smaller the pKb, the stronger the base
Kb > 1, or pKb < 0 indicates a strong base

33
Q

How do you find the pH of a solution of a strong acid or base?

A

Since the strong acids or bases dissociate completely, the initial form of the acid (HA) or base (BOH) will be assumed to have a concentration of 0
Therefore, cannot compute a Ka or Kb
Instead, assume initial molarity of HA or BOH is the same as the respective molarity of H3O+ or OH- ions
0.01 mol / L solution of HCl will yield 0.01 mol / L of H3O+, which means -log[10^-2] = 2 = pH

34
Q

How do you calculate the pH of a solution of a weak acid or base?

A

E.g. reaction of weak acid HCN with water
1. Write down reaction of HCN with water and then set up equilibrium equation
HCN + H2O -> H3O+ + CN-
Ka = [H3O+][CN-] / [HCN] = 6.2 x 10^-10 (Ka will be given)
2. If 0.01 moles of HCN are added to one liter of pure water, ‘x’ amount of HCN will dissociate, so ‘x’ mol / L H3O+ and ‘x’ mol / L CN- ions, and 0.01 - ‘x’ mol / L of HCN
[x][x]/[0.01 - x] = 6.2 x 10^-10
To simplify assume that x is < 5% of 0.01: x^2 / 0.01 = 6.2 x 10^-10, x ~= 2.5 x 10^-6
So pH is between 5 and 6

35
Q

Titration

A

Technique used to analyze the properties of acids and bases
Adding base to acid or acid to base in a process
- drop-by-drop mixing of an acid and a base
Performed for two reasons:
1. Find the concentration of a substance by comparing it with the known concentration of the titrant
2. Find the pKa or pKb and by extension the Ka or Kb of an acid or base
Changing pH of unknown substance as the acidic or basic titrant is added is a sigmoidal curve

36
Q

Neutralization

A

Reaction that results from the mixture of an acid and a base
Generic equation:
Acid + Base -> Water + Salt
Acid and base neutralize each other to form water from hydroxide and hydronium ion
Reaction is typically highly exothermic

37
Q

Calculate \delta G^{o} for the neutralization reaction:

HCl + NaOH -> H2O + NaCl

A

Equilibrium constant for autoionization of water is Kw = 10^-14:
H2O + H2O -> H3O+ + OH-
Reverse of Kw for this: H3O+ + OH- -> H2O, resembles neutralization
So 1/Kw is the equilibrium constant for opposite reaction
Then, to convert K to \delta G^{o}:
\delta G^{o} = -RT ln(K) = -RT ln(10^14) and using T = 298K and R = 8.314 J / mol we get \delta G^{o} = -79.9 kJ / mol where \delta G is negative because K > 1

38
Q

Equivalence Point

A

AKA stoichiometric point
Midpoint of the portion of titration curve that most nearly approximates a vertical line
For a monoprotic acid: the point in the titration when there are equal equivalents of acid and base in solution
Equivalent: amount of acid or base required to produce or consume one mole of protons
For HCl and NaOH, one to one correspondence, so equivalent point reached when the number of moles are equal
For titration of very strong monoprotic acids, equivalence point will be at pH of 7

39
Q

How does the pH affect the behavior as an acid or base?

A

Ka and pKa are intrinsic to the acid, so they are constants
The pH is environmental and therefore variable
When pH is lower than pKa, species interprets the environment as protic (full of H+ ions) and is less likely to act acidic, when pH is above the pKa a species interprets the environment as aprotic (few H+ ions) and is more likely to act acidic

40
Q

What does the titration of a weak acid with a strong base look like?

A

The key to titrations involving weak acids or bases is that the degree of protonation of the acid or base will differ according to the pH of the solution
If base is stronger than the acid, the equivalence point will be above 7
- Where at the equivalence point there is a molecule of strong base for every molecule of weak acid
The conjugate base of the weak acid at the equivalence will act as a base and so raise the pH of the equivalence point
Reverse is true if titrating a strong acid with a weak base (conjugate acid of ammonia will act as an acid and lower pH of equivalence point)

41
Q

Half Equivalence Point

A

Point where exactly one half of the acid has been neutralized by base
Concentration of the acid is equal to concentration of conjugate base
Occurs at midpoint of the section of the graph that most represents horizontal line
Section where largest amount of base or acid could be added with the least amount of change in pH
Such a solution is considered to be buffered and half equivalence point is point in titration where solution is most buffered

42
Q

How do you find the pH at the half equivalence point? At the equivalence point?

A

For the half equivalence point, can calculate using the Henderson-Hasselbalch equation: pH = pKa + log ([A-]/[HA]) = pKa + log(1) = pKa
For the equivalence point: Cannot use Hasselbalch equation because [HA] = 0
However use Ka and Kw to find the Kb, Kb = Kw / Ka = [OH-][HA]/[A-] and solve for OH- concentration (find pOH), subtract from 14 to find the pH

43
Q

Henderson-Hasselbalch Equation

A

pH = pKa + log[A-]/[HA]
Cannot typically be used to find the pH at the equivalence point, instead the pKb of the conjugate base must be used
Concentration of conjugate base at the equivalence point is equal to number of moles of acid divided by volume of acid plus volume of base used titrate

44
Q

What are the log(n) values for n = 1, 2, 3, …, 10

A 
Log(1) = 0, log(2) = 0.3, log(3) = 0.48, log(4) = 0.6, log(5) = 0.7, log(6) = 0.78
Log(7) = 0.85, log(8) = 0.9, log(9) = 0.95, log(10 = 1.0
45
Q

Titration of weak acid and weak base

A

Similar to a weak acid-strong base titration
One major difference is that the range of pH is compressed
There are no strong acids or bases, so it is impossible to reach the extreme pH values
More difficult to identify where the equivalence point lies because the change in pH is less pronounced
Equivalence point could fall either above or below a pH of 7, depending on particular acid and base used
If acid is stronger than the base (pKa of acid is lower than the pKb of the base), then equivalence point will fall at a pH below 7. If vice versa, equivalence point will be above pH of 7

46
Q

Indicator

A

Chemical that can be used to find the equivalence point in a titration
Usually a weak acid whose conjugate base is a different color
Designated as HIn, where In- represents conjugate base
To be detectable in color change, must reach 1/10 of concentration of original form
For titration of an acid with a base, add small amount of indicator to acid (so it won’t affect the pH)
At low pH, HIn form of indicator predominates, as titriation proceeds and pH increases, In- form increases and when In- conc reaches 1/10 of HIn, color change is visible
For titration of base with an acid, reverse process
The pH values of two points of color change give range of indicator

47
Q

What is the range of an indicator?

A

Can be predicted by using the Henderson-Hasselbalch equation:
pH = pKa + log([In-]/[HIn])
Lower range of color change: pH = pKa + log(1/10), pH = pKa - 1
Upper range of color change: pH = pKa + log (10/1), pH = pKa + 1
Point where indicator changes color is called the endpoint, usually will cover the equivalence point

48
Q

pH Meter

A

Concentration cell that compares the voltage difference between different concentrations of H+
Can be used to monitor the pH of a titration

49
Q

How can the Henderson-Hasselbalch equation be used to find the indicator range that will include the equivalence point?

A

Indicator never reaches its equivalence point in the titration
Indicator ions do not approach zero concentration near the color change range (just 1/10 and 10/1 ratios)
If asked to determine the indicator to use for a titration, choose indicator with pKa as close as possible to pH of titration’s equivalence point

50
Q

Polyprotic Titrations

A

Have more than one equivalence point and more than one half equivalence points
Assume first proton completely dissociates before the second proton begins to dissociate (valid if second proton comes from a much weaker acid than the first)
Every polyprotic acid will have multiple pKa values (molecule can be both conjugate base and acid)

51
Q

Normality

A

Similar to molarity of solution except that instead of being a measure of moles / liter, measure of acid or base equivalents per liter
One molar solution of H2SO4 has normality of two because for each mole of H2SO4, there are two equivalents of hydrogens that could be lost
symbol is N (equiv / L)

52
Q

What is the normality of Ba(OH)2? Of H2SO4?

A

Ba(OH)2 will either lose both hydroxides, or neither but still loses two hydroxides for 1 mole, so normality of two
- Note that only one equivalence point
H2SO4 can lose one or two H+ ions, loses two total for 1 mole, so normality of two
- Two equivalence points

53
Q

Titrations of Amino Acids

A

Amino acids are polyprotic acids
Amine group and carboxylic acid
At pH 1, alanine will mostly have amine group protonated and carboxylic acid protonated (two pKa values are 2.3 and 9.7)
As strong base is added and pH increases, carboxylic acid proton begins to dissociate until amino acid exists only in R-COO- form
As more base is added and pH continues to rise, proton on amine group also begins to dissociate until pH is high enough that amino acid is completely converted to RNH2 form

54
Q

For a polyprotic acid with two equivalence points, how much of the species is present at the 1st half equivalence point? 1st equivalence point? 2nd half-equivalence point? 2nd equivalence point?

A

1st half-equivalence point: 50% species 1, 50% species 2
1st equivalence point: 100% species 2
2nd half-equivalence point: 50% species 2, 50 species 3
2nd equivalence point: 100% species 3

55
Q

Isoelectric Point

A

For example with an amino acid, the point at which nearly all amino acids exist in the zwitterionic form (first equivalence point for all uncharged AAs)
Amino acid has one positive and one negative charge (so amino acid is neutral)
For alanine: pI = (pKa1 + pKa2) / 2 = 6.0
Charged AAs are triprotic acids
Isoelectric point is average of first two pKa values for acidic amino acids, and average of second and third pKa values for basic AAs

56
Q

Salts

A

Ionic compounds that dissociate in water
Dissociation of a salt often creates acidic or basic conditions
PH of a salt solution can be predicted qualitatively by comparing the conjugate of the respective ions
Conjugate of salt cation: species that remains after removal of proton, if salt cation has no protons to donate, conjugate is formed by addition of OH
Conjugate of salt anion results from addition of a proton
If conjugates are both strong, salt solution is neutral
If one is strong and other is weak, pH of salt solution favors strong conjugate

57
Q

What are examples of determining pH of salt solutions?

A

For conjugate of NaCl, conjugates are strong base NaOH and strong acid HCl, so NaCl produces neutral solution
For NH4NO3 separates into NH4+ and NO3-, where conjugates are weak base NH3 and strong acid HNO3, so NH4NO3 is weakly acidic

58
Q

Which cations act as weak Lewis acids?

A

All cations except those of the alkali metals and heavier alkaline earth metals (Ca2+, Sr2+, Ba2+) act as Lewis acids in aqueous solution

59
Q

If you have 0.05 M solution of sodium acetate (NaCH3COO) and Ka of CH3COOH is 1.8 x 10^-5, find the pH.

A

Cation Na+ -> NaOH is strong base, so Na+ is ~ neutral
Acetate anion is conjugate base of weak acetic acid -> weak or strong base, so Acetate anion is basic
CH3COO- + H2O -> CH3COOH + OH-, where Kb computed by Kw = KaKb, Kb = Kw/Ka, 10^-14/1.8 x 10^-5 ~= 5.6 x 10^-10 = Kb
Sodium salts soluble, so if add 0.05 M of sodium acetate, there will be 0.05 M of acetate anion
Therefore, to compute pH, can find [OH-] and then convert to pH
We know that [CH3COOH][OH-] / [CH3COO-] = 5.6 x 10^-10
For every CH3COO- converted to CH3COOH, there will be a hydroxide created, so if conc of CH3COO- is x then x = [OH-], and [CH3COO-] is 0.05 - x
Can use: x^2 / 0.05 -x ~= x^2 / 0.05 = 5.6 x 10^-10, where x ~= 5.3 x 10^-6, so pOH is 5.3 and pH is 8.7 (14 - 5.3 = 8.7)

60
Q

Buffers

A

Combinations of acids and salts that are used to keep pH of solution within a certain range
Start with an acid whose pKa is close to pH at which we want to buffer solution
Next, mix equal amounts of that acid with conjugate base
Concentration of buffer solution should greatly exceed concentration of outside acid or base that could affect pH of solution
E.g. to maintain pH of solution at ~ 4.75, buffer created using CH3COOH (pKa = 4.75) and acetate salt like NaCH3COO

61
Q

How can the power of a buffer be explained using the Henderson-Hasselbalch?

A

H-H equation: pH = pKa + log[A-]/[HA]
Equal amounts of buffer acid and conjugate base, so pH = pKa
Buffer concentration much higher, so when any additional other acid or base added, will not change pH much
E.g. 1 L buffered solution created with 1 M carbonic acid & sodium bicarbonate, where pKa of carbonic acid = 6.37
Add 0.01 moles of HCl which is strong acid, and will run to completion until all H+ lost from HCl: pH = 6.37 + log[1 - 0.01]/[1 + 0.01] ~= 6.36, so very little change

MCAT Chemistry (8 decks)

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