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MCAT 2012 Back-up > Acids and Bases -- E&D 1/31 version > Flashcards

Acids and Bases -- E&D 1/31 version Flashcards

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Now . . on to Acids & Bases!

CLICK HERE for a video introduction.

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2
Q

Define and give a general reaction for:

Brønsted-Lowry Acid

A

A Brønsted-Lowry Acid is any species capable of donating a proton to the solution. This is the definition that the MCAT uses for any acid in general, and results in an increase in hydronium ion concentration, and hence a decrease in pH.

Example acid reaction:

H2SO4 + H2O ⇒ HSO4- + H3O+

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3
Q

Define and give a general reaction for:

Brønsted-Lowry Base

A

A Brønsted-Lowry Base is any species capable of accepting a proton from the solution. This is the definition that the MCAT uses for any general base, and results in an increase in hydroxide ion concentration, and an increase in pH.

Example base reaction:

NH3 + H2O ⇒ NH4+ + OH-

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4
Q

Is CH3COOH an acid or a base?

A

An Acid. It loses a proton to solution.

Acid Equation:

CH3COOH + H2O ⇒ CH3COO- + H3O+

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5
Q

Is NH3 an acid or a base?

A

A base. It accepts a proton (or creates an OH-) in solution.

Base Reaction:

NH3 + H2O ⇒ NH4+ + OH-

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6
Q

Define and give an example of:

A polyprotic acid

A

A polyprotic acid can donate more than one proton to a solution.

Ex: H2SO4 (can donate 2 protons)

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7
Q

Define and give an example of:

amphoteric substance

A

An amphoteric substance can act as an acid or a base depending on the solution.

The classic MCAT example is water. Acting as a base:

H20 + HA ⇔ A- + H30+

Acting as an acid:

H20 + B- ⇔ BH + OH-

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8
Q
A
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9
Q

Give the equation for:

water autoionization

A

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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10
Q

What is Kw?

What is its value at STP?

A

Kw is the ion product for the water autoionization reaction,

2 H2O ⇔ H3O+ + OH-

So Kw = [H3O+] [OH-]. In pure water at STP:

[H3O+] = [OH-] = 10-7

Kw = 10-14

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11
Q

Define and give the equation for:

pH

A

pH measures the acidity of a substance. A low pH means a high concentration of hydrogen ions. A high pH means there are few H+ ions.

pH can be calculated by the equation: pH= - log [H+]

In fact: p(anything) = - log (anything).
You only need to memorize that one general equation!

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12
Q

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

A

At STP, water has a pH of 7.0 and is a neutral substance.

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13
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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14
Q
A
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15
Q

Define:

conjugate acid-base pairs

A

Conjugate acid-base pairs are molecules which differ via the presence or absence of a proton. The protonated form of the molecule is the conjugate acid, the deprotonated form is the conjugate base (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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16
Q

What is the conjugate base of acetic acid, CH3COOH?

A

The 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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17
Q

If base X is weaker than base Y, what do you know 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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18
Q
A
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19
Q

What is the equation to calculate the pH of an acid solution?

A

pH = -log[H+ ions]

You would need to be provided with the concentration of H+ ions in order to solve. Remember, for a strong acid, the acid concentration is equal to H+ ion concentration.

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20
Q

Define and give an example of:

Strong acid

A

A strong acid 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.

Ex: HCl is a classic strong acid.

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21
Q

Define and give an example of:

Strong base

A

A strong base 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.

Ex: NaOH is a classic strong base.

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22
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.

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23
Q

Why is this the ranking of increasing acid strength?

CH4 < NH3 < H2O < HF

A

CH4 < NH3 < H2O < HF

These molecules are an increasing series going from left to right across the second row of the Periodic Table. As you travel from left to right across the Periodic Table, acidity increases.

The reason for this is polarity. Each molecule is roughly identical, chemically - a set of 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 easier to dissociate in aqueous solution, so the more electronegative the central atom, the more easily it donates protons, and the more acidic it is.

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24
Q

Place these in order of increasing acid strength:

HCl; HF; HI; HBr

A

HF < HCl < HBr < HI

Recall: acid strength is determined by how easily the substance will donate a proton into solution.

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25
Q

Why is this the order of increasing acid strength?:

HF < HCl < HBr < HI

A

HF < HCl < HBr < HI

These molecules are a series going down a column of the Periodic Table. As you travel down a column in the Periodic Table, acidity increases.

The reason for this is atomic size. 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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26
Q

Please 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)
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27
Q

Please list six common strong acids.

A
  1. HI (Hydrogen Iodide)
  2. HBr (Hydrogen bromide)
  3. HCl (Hydrogen chloride)
  4. HNO3 (Nitric Acid)
  5. HClO4 (Perchloric Acid)
  6. H2SO4 (Sulfuric Acid)
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28
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 + H20

It is possible to add acid and base and NOT create water (Lewis acid/base pairs) but there will always be a salt formed. Additionally, since the MCAT uses Bronsted-Lowry as the acid/base definition, we are safe in assuming that water is always formed as well.

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29
Q

What is the shortcut equation to calculate the approximate pH of an acid?

A

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

EX: 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.

30
Q
A
31
Q

Please 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: technically any acid that is not on the list of STRONG acids, is considered weak.

32
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)
33
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(CH4)3 = trimethyl amine)
  5. Strong (Ca(OH)2 = calcium hydroxide)
34
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)
35
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)
36
Q

Please list some common weak bases.

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

Remember - any base that’s not on the list of STRONG bases, is considered a weak base by default.

37
Q

Answer and explain:

The solubility of NH3 in a NH4Cl solution will be _______ than the solubility in pure water.

A

Lower.

As NH3 dissolves in water, it partially dissociates according to

NH3 + H2O ⇒ NH4+ + OH-

Le Chatelier’s Principle says that as 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.

38
Q
A
39
Q

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

CH3COOH + H2O ⇒
CH3COO- + H3O+

A

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

40
Q

Define and give the common types for:

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 (e.g. NH4Cl + H2O) makes a solution acidic; anionic hydrolysis (e.g. 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.

41
Q

Define and give the general equation for:

Ka

A

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

42
Q

Define:

Kh

A

Khis the hydrolysis constant, and quantifies how much salt can be hydrolysed in a saturated solution. This is essentially a proportion of how much water there is available to dissolve the salt in.

A higher Kh means that more salt can be hydrolysed.

43
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
44
Q

How do you calculate the Ka of a weak acid?

A

from the general equation:

HA ⇔ H+ + A-
Ka = [H+][A-] / [HA]

In an MCAT problem, you would be given at least two of these concentrations in order to solve.

45
Q

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

A

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.

46
Q

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

A

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

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

47
Q

Define and give the equation for:

Kb

A

Kb is the base dissociation constant. The higher the value of Kb, the stronger the base is in solution.

from the general equation:

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

48
Q

If base X is stronger than base Y, what do you know 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

49
Q

Give the equation for:

pKb

A

pKb = -log (Kb)

In fact: p(anything) = - log (anything).
You only need to memorize that one general equation!

50
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.
51
Q

Give the equation for:

pKa

A

pKa = - log (Ka)

In fact: p(anything) = - log (anything).
You only need to memorize that one general equation!

52
Q

If base A is weaker than base B, what do you know about the 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.

53
Q

How many Ka values does each of these acids possess?

HCl, H2SO4, H3PO4

A

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

HCl + H2O ⇒ Cl - + H3O+

H2SO4 (sulfuric acid) has 2 Ka values for the two protons it can donate.

H2SO4 + H2O ⇒ HSO4 - + H3O+
HSO4- + H2O ⇒ SO42 - + H3O+

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

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

54
Q

Acid X is stronger than acid Y, what is true of their pKa values?

A

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

Recall: the stronger the acid: the higher the Ka value will be. The higher the Ka: the lower the pKa will be.

55
Q
A
56
Q

Will having a higher concentration of buffer ions (HA and A-) in solution make a difference to the pH response when adding an acid or base?

Why or why not?

A

Yes.

Having more buffer ions will keep the pH from changing as significantly, since the buffer ions already in solution will combine with the new acid or base added and reduce their effect.

57
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.

58
Q

Give the following equation:

Henderson-Hasselbalch

A

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

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

59
Q

Please name some common buffer solutions and their pH ranges.

A

Acetic acid (pH 3.7-5.6)

Bicarbonate (pH 7-10)

Sodium citrate (pH 3-5)

60
Q
A
61
Q

What is the purpose of a titration experiment?

A

The purpose of a titration experiment is to discover the concentration of an unknown acid or base, by neutralizing it with a measured quantity of a base or acid whose concentration is known.

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

62
Q

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

Ex: CH3COOH + NaOH ⇒
NaCH3COO + H2O

A

More than 7, slightly basic.

The equivalence point is a point at which the amount of equivanents of acid and equivalents of base from the analyte and titrant are equal. Since the example given results in a weak base (acetate ion, CH3COO-) being formed, the pH at the equivalence point is above 7.

63
Q

What are the requirements for an indicator?
Name several common ones.

A

An indicator is a molecule that must change color visibly in a set pH range, usually a range of about 2 pH units. To select an indicator for a specific titration, find one whose pKa is 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)

64
Q

Define and give an example of:

indicator

(as it relates to titrations)

A

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

Ex: Phenolphthalein goes from colorless to fuchsia between pH 8.3-10
Ex: Methyl red goes from red to green to yellow between pH 4.4-5.2-6.2

65
Q

Explain and give the pH value for:

the equivalence point in a strong acid/strong base titration

A

The equivalence point is a point at which the amount of equivanents 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 one of the titrant and analyte is stronger, however, then the pH at the equivalence point will favor the stronger species.

66
Q

Define and give the common use for:

analyte

A

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

A titrant (known concentration) is used to bring the unknown (analyte) solution up to a standard pH (usually neutral 7). In this way, the concentration of the analyte can be calculated.

67
Q

Define and give an example of the use for:

titrant

A

A titrant is a strong acid or base with a known concentration.

In a titration experiment it 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.

68
Q

Define:

equivalence point

A

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

[H+ ions]original =

[OH- ions]added

In a titration, this is where all of the unknown (analyte) has been completely neutralized by the known (titrant), giving a neutral pH. At the equavalence point,

VA NA = VB NB

69
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-]

70
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, so you need the second equivalence point, which is Point D.

71
Q

WELCOME!

Since this may be your 1st time in Brainscape, here’s how it works:

  • We ask a question
  • You try to answer in your head
  • You flip / reveal the answer
  • Rate how well you knew it (1-5).

Q: What method does Brainscape use?

A

CBR: Confidence-based Repetition

We used both cognitive science (e.g ‘spaced-repetition’) and technology to create the most efficient flashcard study method available.

You will how well you knew each answer (before you saw it!)–and this rating will optimize future repetition and overall retention.

Rate this card a “5”–or even DELETE it if you don’t want it see even every once in a while (pull-down menu)!

Now . . on to Acids & Bases!

CLICK HERE for a video introduction.

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