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MCAT > MCAT Gen. Chemistry > Acids and Bases > Flashcards

Acids and Bases Flashcards

Always a favorite of MCAT test-makers, acid/base chemistry can also be one of the toughest topics on the exam. Use these cards to master the Henderson-Hasselbalch equation, pH and pOH, and polyprotic titrations. For enhanced practice, think about how these concepts relate to biological systems, such as the bicarbonate buffer system.

1
Q

Define:

Brønsted-Lowry acid

A

It is any species capable of donating a proton to the solution, resulting in an increase in hydronium ion concentration and a decrease in pH.

Unless otherwise stated, this is the definition that the MCAT will use when referring to an acid. Brønsted-Lowry acids react as follows:

HA + H2O ⇒ A- + H3O+

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2
Q

Define:

Brønsted-Lowry base

A

It is any species capable of accepting a proton from the solution, resulting in an increase in hydroxide ion concentration and an increase in pH.

Unless otherwise stated, this is the definition that the MCAT will use when referring to a base. Brønsted-Lowry bases react as follows:

B + H2O ⇒ BH+ + OH-

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3
Q

Is CH3COOH an acid or a base?

A

an acid

(specifically, acetic acid)

It loses a proton to solution.

Acid dissociation equation:

CH3COOH + H2O ⇒ CH3COO- + H3O+

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4
Q

Is NH3 an acid or a base?

A

a base

(specifically, ammonia)

It accepts a proton and creates an OH- ion in solution.

Base reaction:

NH3 + H2O ⇒ NH4+ + OH-

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5
Q

Define:

polyprotic acid

A

It can donate more than one proton to a solution.

For example, H2SO4 can donate 2 protons and is therefore diprotic.

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6
Q

Define:

amphoteric

A

They are substances that can act as an acid or a base depending on the solution.

A classic example is water. Acting as a base:

H2O + HA ⇔ A- + H3O+

Acting as an acid:

H2O + B- ⇔ BH + OH-

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7
Q

What is water autoionization?

A

Two water molecules can ionize each other, as such:

2 H2O ⇔ H3O+ + OH-

Since the creation of each H3O+ from a water molecule also requires the creation of an OH-, in pure water these concentrations will always be equal.

At STP, [H3O+] = [OH-] = 10-7.

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8
Q

Define:

Kw

A

It is the ion product for the water autoionization reaction.

Remember that all K values (Ksp, Keq, Ka, Kb…) are equilibrium constants and behave fairly similarly.

2 H2O ⇔ H3O+ + OH-

Kw = [H3O+] [OH-]

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9
Q

What is the value of Kw at STP?

A

10-14

Remember, Kw = [H3O+] [OH-]. In pure water at STP, [H3O+] = [OH-] = 10-7, so Kw = 10-14. When the temperature changes, this value changes as well.

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10
Q

Define:

pH

A

It measures the acidity of a substance.

A low pH means a high concentration of H+ ions, while a high pH means there are few H+ ions.

pH can be calculated using the equation:

pH= - log [H+]

In fact, p (anything) = - log (anything).

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11
Q

At 25ºC and 1 atm, what is the pH of water?

A

7.0 and is a neutral substance

When temperature increases or decreases, pure water remains neutral, but its pH still changes due to increased or decreased dissociation.

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12
Q

At STP, what is the pH of an acidic solution?

A

below 7

As a solution becomes more acidic, its pH decreases.

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13
Q

Define:

conjugate acid-base pairs

A

They are molecules which differ via the presence or absence of a proton.

The protonated form of the molecule is the conjugate acid, while the deprotonated form is the conjugate base (using the Brønsted-Lowry acid definition)

Generic equation:

HA + H2O → A- + H3O+

HA is an acid and A- is its conjugate base. Similarly, H2O is acting as a base, with H3O+ acting as its conjugate acid.

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14
Q

What is the conjugate base of acetic acid, CH3COOH?

A

acetate ion, CH3COO-

An acid’s conjugate base is the deprotonated remainder of the molecule’s acid reaction. The general acid neutralization reaction is

HA + OH- ⇒ A- + H2O

where HA is the acid and A- is the conjugate base.

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15
Q

If base X is weaker than base Y, what must be true about the conjugate acids of each?

A

The conjugate acid XH+ will be stronger than the conjugate acid YH+.

In general, the weaker the base, the stronger its conjugate acid will be.

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16
Q

Define:

strong acid

A

It is one which dissociates completely in solution.

For a monoprotic acid (one proton per molecule), each mole of acid in solution results in one mole of protons in solution as well. For example, HCl is a classic strong monoprotic acid.

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17
Q

Define:

strong base

A

It is one which dissociates completely in solution.

For a monobasic compound (one hydroxide per molecule), each mole of base in solution results in one mole of hydroxide ions in solution as well. For example, NaOH is a classic strong monobasic base.

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18
Q

Place these in order of increasing acid strength:

H2O, NH3, HF, CH4

A

CH4 < NH3 < H2O < HF

Acidity, in general, is a measure of how easily that substance will donate a proton into solution. Acidity of a molecule increases from left to right across a row of the periodic table.

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19
Q

Why is this the ranking of increasing acid strength?

CH4 < NH3 < H2O < HF

A

The reason for this is polarity.

These molecules are ordered from left to right across the second row of the periodic table. Traveling left to right across the table, acidity always increases.

Chemically, these acids are all hydrogen atoms bound to a central atom. As the central atom becomes more electronegative, the bonds with hydrogen become more polar. More polar bonds are generally easier to dissociate in aqueous solution. Hence, when moving from left to right, the more electronegative the central atom, the more easily it donates protons, and the more acidic it is.

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20
Q

Place these in order of increasing acid strength:

HCl, HF, HI, HBr

A

HF < HCl < HBr < HI

Recall that acid strength is determined by how easily the substance will donate a proton into solution. Acid strength increases going down a column of the periodic table.

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21
Q

Why is this the order of increasing acid strength?
HF < HCl < HBr < HI

A

This is explained by atomic size.

These molecules are ordered from top down along the periodic table. Going down a column in the table, acidity increases.

Larger atoms can carry negative charges more easily, so the I- ion is more stable than the F- ion. The more stable the conjugate base, the stronger the acid.

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22
Q

List seven common strong acids.

A
  1. HI (hydrogen iodide or hydroiodic acid)
  2. HBr (hydrogen bromide or hydrobromic acid)
  3. HCl (hydrogen chloride or hydrochloric acid)
  4. HNO3 (nitric acid)
  5. HClO4 (perchloric acid)
  6. HClO3 (chloric acid; less likely to appear on the exam)
  7. H2SO4 (sulfuric acid)

While other acids can be classified as strong, these are the ones most commonly used on the MCAT.

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23
Q

List seven common strong bases.

A
  1. NaOH (sodium hydroxide)
  2. KOH (potassium hydroxide)
  3. NH2- (amide ion)
  4. H- (hydride ion)
  5. Ca(OH)2 (calcium hydroxide)
  6. Na2O (sodium oxide)
  7. CaO (calcium oxide)

While several other bases can be classified as strong, these are the ones most commonly used on the MCAT.

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24
Q

What equation is used to calculate the pH of an acidic solution?

A

pH = -log [H+]

Remember, for a strong acid, the acid concentration is equal to the H+ ion concentration.

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25
Q

What is the product of the reaction of a strong acid and a strong base?

A

Salt and water, according to the general equation

HA + BOH ⇒ AB + H2O

It is possible to add acid and base and NOT create water (Lewis acid/base pairs) but a salt will always form. Additionally, when you use Bronsted-Lowry as the acid/base definition (like the MCAT does), it is safe to assume that water is always formed as well.

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26
Q

What shortcut equation can be used to calculate the approximate pH of an acid?

A

If [H+] = n x 10 -e, then pH = {e-1}.{10-n}

For example, If [H+] = 6.2 x 10-4, then n = 6.2, and e = 4. So, the pH can be approximated as equal to

[4-1].[10-6.2] = 3.38

The real pH is 3.21, which is within the acceptable error for the multiple-choice MCAT.

27
Q

List some common weak acids.

A
  1. HCN (hydrogen cyanide)
  2. HClO (hypochlorous acid)
  3. HNO2 (nitrous acid)
  4. HF (hydrofluoric acid)
  5. H2SO3 (sulfurous acid)
  6. H2S (hydrogen sulfide)

Remember that technically, any acid that is not on the list of strong acids is considered weak.

28
Q

Which are strong bases, and which are weak?
1. NaOH
2. Na2O
3. NH3
4. H-
5. HS-

A
  1. Strong (NaOH = sodium hydroxide)
  2. Strong (Na2O = sodium oxide)
  3. Weak (NH3 = ammonia)
  4. Strong (H- = hydride ion)
  5. Weak (HS- = hydrosulfide ion)
29
Q

Which are strong bases, and which are weak?

  1. H2O
  2. NH2-
  3. CaO
  4. N(CH3)3
  5. Ca(OH)2
A
  1. Weak (H2O = water)
  2. Strong (NH2- = amide ion)
  3. Strong (CaO = calcium oxide)
  4. Weak (N(CH3)3 = trimethyl amine)
  5. Strong (Ca(OH)2 = calcium hydroxide)
30
Q

Which are strong acids, and which are weak?

  1. HCl
  2. HNO3
  3. H2SO4
  4. HCN
  5. H2S
A
  1. Strong (HCl = hydrogen chloride)
  2. Strong (HNO3 = nitric acid)
  3. Strong (H2SO4 = sulfuric acid)
  4. Weak (HCN = hydrogen cyanide)
  5. Weak (H2S = hydrogen sulfide)
31
Q

Which are strong acids, and which are weak?
1. HI
2. HBr
3. HF
4. HClO4
5. H2SO3

A
  1. Strong (HI = hydrogen iodide)
  2. Strong (HBr = hydrogen bromide)
  3. Weak (HF = hydrogen fluoride)
  4. Strong (HClO4 = perchloric acid)
  5. Weak (H2SO3 = sulfurous acid)
32
Q

List some common weak bases.

A
  1. NH3
  2. N(CH3)3
  3. NH4OH
  4. HS-
  5. H2O

Remember that technically, any base that’s not on the list of strong bases is considered weak.

33
Q

How does the presence of dissolved NH4Cl affect the solubility of NH3 in a solution?

A

It is decreased.

As NH3 dissolves in water, it partially dissociates according to

NH3 + H2O ⇒ NH4+ + OH-

Le Chatelier’s Principle says that as the concentration of an ion in solution increases, that ion’s solubility decreases. So NH3, which forms NH4+ in solution, will be less soluble in the NH4Cl solution than it would be in pure water.

34
Q

What is the Ka equation for the reaction of acetic acid and water?

A

CH3COOH + H2O ⇒ CH3COO- + H3O+

Remember that Keq equations include only components in the gaseous or aqueous phase; components in the solid or liquid phase, like H2O, are ignored.

35
Q

Define:

hydrolysis of a salt

A

When a salt is dissolved in water, two separate ions are created. If either (or both) of those ions react with water to create a change in pH, that is considered hydrolysis.

Cationic hydrolysis (NH4Cl + H2O, for example) makes a solution acidic; anionic hydrolysis (like NaCH3COO + H2O) will make the solution basic.

A salt of a strong acid and strong base will never undergo hydrolysis. The resulting solution will always be neutral.

36
Q

Define:

Ka

A

It is the acid dissociation constant for an acid in solution. The higher the Ka, the stronger the acid.

37
Q

Define:

Kh

A

It is the hydrolysis constant and quantifies how much salt can be hydrolyzed in a saturated solution.

This is essentially a proportion of how much water is available in which to dissolve the salt.

A higher Kh means that more salt can be hydrolyzed.

38
Q
  1. What is the Ka of a strong acid?
  2. What is the Ka of a weak acid?
A
  1. For a strong acid, Ka > 1.
  2. For a weak acid, Ka < 1.
39
Q

How do you calculate the Ka of a weak acid?

A

The general acid equation and Ka calculation for the acid HA is:

HA ⇔ H+ + A-
Ka = [H+][A-] / [HA]

In an MCAT problem, you would be given two of these concentrations and would solve for the third.

40
Q

Acid A is stronger than acid B. What do you know about their relative Ka values?

A

The Ka for acid A (the stronger one) will be higher than the Ka for acid B.

In general, the higher the Ka value, the stronger the acid.

41
Q

Define:

Kb

A

It is the base dissociation constant.

The higher the value of Kb, the stronger the base is in solution.The general equation for Kb is:

B- + H2O ⇔ OH- + BH
Kb = [OH-][BH] / [B-]

42
Q

If base A is weaker than base B, what do you know about their relative Kb values?

A

Base B (stronger) will have a higher Kb value.

In general, the stronger the base in solution, the higher the Kb value.

43
Q

How is pKb calculated?

A

pKb = -log (Kb)

In fact. p(anything) = - log (anything).

44
Q

If base X is stronger than base Y, what must be true about their relative pKb values?

A

Base X (stronger) will have a lower pKb value than base Y.

Recall that a stronger base will have a higher Kb value, hence a lower pKb.

45
Q

How is pKa calculated?

A

pKa = - log (Ka)

In fact: p(anything) = - log (anything).

46
Q
  1. What is the pKa of a strong acid?
  2. What is the pKa of a weak acid?
A
  1. For a strong acid, pKa < 0.
  2. For a weak acid, pKa > 0.
47
Q

Acid X is stronger than acid Y. What must be true of their pKa values?

A

Acid X (stronger) will have a lower pKa than acid Y (weaker).

Recall that the stronger the acid, the higher its Ka value will be. The higher the Ka, the lower the pKa will be.

48
Q

How many Ka values do HCl and H3PO4 possess, respectively?

A

HCl (hydrochloric acid) has 1 Ka value for the single proton that it can donate.

HCl + H2O ⇒ Cl - + H3O+

H3PO4 (phosphoric acid) has 3 Ka values for the three protons that it can donate.

H3PO4 + H2O ⇒ H2PO4- + H3O+
H2PO4- + H2O ⇒ HPO42- + H3O+
HPO42- + H2O ⇒ PO43- + H3O+

49
Q

How does higher concentration of HA and A- ions affect a buffer’s pH response when adding an acid or base to the solution?

A

The buffer’s pH response will be decreased; in other words, its buffering ability will be stronger.

Having more buffer ions keeps the pH from changing as significantly, since the buffer ions already in solution will neutralize the added acid or base and reduce its effect.

50
Q

What is the optimal pH for a buffer solution?

A

A buffer works most effectively when pH = pKa of the weak acid (or when pOH = pKb of the weak base) that is taking part in the buffer.

51
Q

What is the Henderson-Hasselbalch equation?

A

It relates concentrations of weak acid and conjugate base in a buffer.

pH = pKa + log([A-] / [HA])

pOH = pKb + log([BH] / [B-])

52
Q

The pKa of a weak acid is 6.3. What pH will indicate that there are equal concentrations of A- and HA?

A

The concentrations will be equal at pH = 6.3.

According to Henderson-Hasselbach, pH = pKa + log ( [A-]/[HA] ). When [A-] = [HA], the final term becomes log (1) = 0. At this point, then, pH = pKa = 6.3.

53
Q

Name 3 common buffer solutions tested on the MCAT and their effective pH ranges.

A
  1. Acetic acid (pH 3.7-5.6)
  2. Bicarbonate (pH 7-10)
  3. Sodium citrate (pH 3-5)
54
Q

What is the purpose of a titration experiment?

A

It is to discover the concentration of an unknown acid or base by neutralizing it with a measured quantity of a base or acid solution with known concentration.

In titrations, the known solution is the titrant and the unknown solution is the analyte.

55
Q

Define:

equivalence point

A

It is the point of a titration at which every molecule of acid has been neutralized by a molecule of base.

[H+ ions]original = [OH- ions]added

In a titration, this occurs when all of the unknown, or analyte, has been completely neutralized by the known, or titrant, giving a neutral pH.

56
Q

Define:

analyte

A

It is an acid or base whose concentration is determined by a titration.

A titrant with known concentration is used to bring the unknown (analyte) solution to the pH of the equivalence point for that particular titration. In this way, the concentration of the analyte can be calculated.

57
Q

Define:

titrant

A

It is a strong acid or base with a known concentration.

In a titration experiment the titrant is used to neutralize, or bring to a known pH, an unknown (analyte) solution. In this way, the concentration of the unknown can be calculated.

58
Q

Where will the equivalence point fall in a strong acid/strong base titration, and what does it signify?

A

It is a point at which the amount of equivalents of acid and equivalents of base from the analyte and titrant are equal. At the equivalence point,

VA * NA = VB * NB

For a strong acid/strong base combination, this happens at a pH of 7. If either the titrant or analyte is stronger, however, then the pH at the equivalence point will favor that stronger species.

59
Q

If a weak acid is titrated with a strong base, what is the pH at the equivalence point?

For example, CH3COOH + NaOH ⇒ NaCH3COO + H2O.

A

More than 7 (slightly basic)

The equivalence point is a point at which the amount of equivalents of acid and equivalents of base from the analyte and titrant are equal.

In the example, a weak base (acetate ion, CH3COO-) is being formed, hence the pH at the equivalence point will be above 7.

60
Q

Define:

indicator

A

In a titration, it changes the color of the solution to indicate that the equivalence point pH has been reached.

For example, phenolphthalein goes from colorless to fuchsia between pH 8.3-10. Methyl red goes from red to green to yellow between pH 4.4-5.2-6.2

61
Q

List the requirements for an indicator and name several common examples.

A

It is a molecule that must change color visibly in a set pH range, usually a range of about 2 pH units. The proper indicator for a specific titration has a pKa roughly equal to the pH of the titration’s equivalence point.

Some common indicators include:

  • Methyl red (pH range 4.4-6.2)
  • Thymol blue (8.0-9.6)
  • Azolitmin (4.5-8.3)
62
Q

In the below titration of the acid H2A, at which point does [H2A] = [HA-]?

A

Point A

Point A is the first half-equivalence point, where one-half as many equivalents of base have been added as acid molecules that were in the solution to begin with, so one half of the acid molecules have been neutralized and [H2A] = [HA-].

63
Q

In the below titration of the acid H2A, what is the earliest point where the solution is entirely A2-?

A

Point D

A2- production is not complete until 2 full equivalents of base have been added. This must be at the second equivalence point, which is Point D.

MCAT Gen. Chemistry (10 decks)

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