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Organic Chemistry > Ch. 2: Acids and Bases > Flashcards

Ch. 2: Acids and Bases Flashcards

1
Q

What is a Bronsted-Lowry acid?

A

a species that loses a proton

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2
Q

What is a Bronsted-Lowry base?

A

a species that gains a proton

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3
Q

What is the reaction of an acid with a base called?

A

proton transfer reaction or an acid-base reaction

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4
Q

What is the conjugate acid of NH3?

A

+NH4

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5
Q

What’s the conjugate base of H2O?

A

HO-

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6
Q

Why can water act as both an acid and a base?

A

Water can behave as a base because it has a lone pair, and it can behave as an acid because it has a proton that it can lose.

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7
Q

What is acidity?

A

a measure of the tendency of a compound to lose a proton

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8
Q

What is basicity?

A

a measure of a compound’s affinity for a proton

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9
Q

Describe a strong acid and its conjugate base.

A

A strong acid has a tendency to lose a proton, has a weak conjugate base with little affinity for the proton.

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10
Q

Describe a weak acid and its conjugate base.

A

A weak acid has little tendency to lose its proton, has a strong conjugate base with a high affinity for the proton.

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11
Q

What is the equilibrium constant Keq?

A

The degree to which an acid (HA) dissociates in an acqueous solution is indicated by Keq.

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12
Q

What is the acid dissociation constant?

A

The degree to which an acid (HA) dissociates is normally determined in a dilute solution so the concentration of water remains essentially constant. Ka is the equilibrium constant multiplied by the molar concentration of water.

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13
Q

Describe pKa and its use.

A

A convenient way to describe the strength of an acid.

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14
Q

What is pH?

A

The concentration of protons in a solution.

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15
Q

What are carboxylic acids?

A

compounds with a COOH group

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16
Q

Are carboxylic acids strong or weak?

A

weak acids

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17
Q

What are alcohols?

A

compounds with an OH group. They can act as both a B-L acid and base.

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18
Q

Are alcohols stronger or weaker than carboxylic acids?

A

much weaker than carboxylic acids

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19
Q

What are amines?

A

compounds that result from replacing one or more of the hydrogens bonded to ammonia with a carbon-containing substituent. have very high pKas, and are likely to act as bases

20
Q

Explain the use of curved arrows in a equation.

A

They’re used to indicate the bonds that are broken and formed as reactants are converted into products. The arrow always points from the electron donor to the electron acceptor. In an acid-base reaction, one of the arrows is drawn from a lone pair on the base to the proton of the acid. A second arrow is drawn from the electrons that the proton shared to the atom on which they are left behind. As a result, the curved arrows let you follow the electrons to see what bond is broken and what bond is formed in the reaction.

21
Q

What is the approximate pKa value of protonated alcohols, protonated carboxylic acids, and protonated water

A

less than 0

22
Q

What is the approximate pKa value of carboxylic acids?

A

~5

23
Q

What is the approximate pKa value of protonated amines?

A

~10

24
Q

What is the approximate pKa value of alcohols and water?

A

~15

25
Q

How do you predict water will react as an acid or base?

A

Compare the pKa values of the reactants. If water is lower than the other, then it will act as the acid. If it’s higher, then it’ll act as the base in that reaction.

26
Q

In an acid-base reaction, does the equilibrium favor the formation of the strong or weak acid and why.

A

Weak acid. The reason for this is that like any chemical reaction or a process, the acid-base reactions go towards a lower energy state. A strong acid or a base means that they have a lot of energy and are very reactive while weaker acids and bases have lower energy.

27
Q

How do you calculate the equilibrium constant of an acid-base reaction?

A
28
Q

How does electronegativity affect the strength of an acid?

A

When atoms are similar in size, the strongest acid has its hydrogen attached to the most electronegative atom

29
Q

Why is a protonated alcohol (like CH3OH) more acidic than a protonated amine (like CH3NH2)?

A

because oxygen is more electronegative than nitrogen

30
Q

Rank the hybrid orbitals by electronegativity.

A

sp > sp2 > sp3

31
Q

Why does hybridization affect electronegativity?

A

Electronegativity is a measure of the ability of an atom to pull the bonding electrons toward itself. Thus, the most electronegative atom is the one with its bonding electrons closest to the nucleus. The average distance of a 2s electron from the nucleus is less than the average distance of a 2p electron from the nucleus. Therefore, an sp hybridized atom with 50% s character is the most electronegative, an sp2 hybridized atom (33.3% s character) is next, and an sp3 hybridized atom (25% s character) is the least electronegative.

32
Q

Of electronegativity and size, what overrides the other when determining electronegativity and why?

A

As we proceed down a column in the periodic table, the atoms get larger and the stability of the anions increases even though the electronegativity of the atoms decreases. Because the stability of the bases increases going down the column, the strength of their conjugate acids increases. So when atomms are very different in size, the strongest acid has its hydrogen attached to the largest atom.

33
Q

Why are larger ions more stable than smaller ones?

A

The valence electrons of F- are in a 2sp3 orbital, the valence electrons of Cl- are in a 3sp3 orbital, those of Br- are in a 4sp3 orbital, and those of I- are in a 5sp3 orbital. The volume of space occupied by a 3sp3 orbital is significantly larger than the volume of space occupied by a 2sp3 orbital because a 3sp3 orbital extends farther from the nucleus. Because its negative charge is spread over a larger volume of space, Cl- is more stable than F-.

34
Q

What is substitution?

A

The term for replacing an atom in a compound

35
Q

What is inductive electron withdrawal?

A

Pulling electrons through sigma bonds

If we look at the conjugate base of the carboxylic acid, we see that inductive electron withdrawal decreases the electron density about the oxygen that bears the negative charge, thereby stabilizing it. And we know that stabilizing a base increases the acidity of its conjugate acid.

36
Q

What are the two factors that cause the conjugate base of a carboxylic acid to be more stable than the conjugate base of an alcohol?

A

1) inductive electron withdrawal

2) delocalized electrons

37
Q

What are delocalized electrons?

A

When a carboxylic acid loses a proton, the electrons left behind are shared by three atoms–two oxygens and a carbon. These electrons are delocalized because they belong to more than two atoms. Therefore, the negative charge is shared by both oxygens and the conjugate base is stabilized because decreasing the electron density of an atom stabilizes it.

38
Q

What are the five factors that determine acid strength?

A
  1. size: as the atom attached to the hydrogen increases in size (going down a column on the PT), the strength of the acid increases
  2. electronegativity: as the atom attached to the hydrogen increases in electronegativity, the strength of the acid increases
  3. hybridization: the electronegativity of an atom changes with hybridization as follows: sp>sp2>sp3. Because an sp carbon is the most electronegative, a hydrogen attached to an so carbon is the most acidic, and a hydrogen attached to an sp3 carbon is the least acidic
  4. inductive electron withdrawal: an electron-withdrawing group increases the strength of an acid. As the electronegativity of the electron-withdrawing group increases or as it moves closer to the acidic hydrogen, the strength of the acid increases.
  5. electron delocalization: an acid whose conjugate base has delocalized electrons is more acidic than a similar acid whose conjugate base has only localized electrons
39
Q

How is the relationship between pKa and pH represented?

A

the Henderson-Hasselbach equation

40
Q

What does the Henderson-Hasselbach equation tell us?

A
  • when the pH of a solution equals the pKa of the compound that undergoes dissociation, the concentration of the compound in its acidic form (HA) equals the concentration of the compound in its basic form (A-) (because log 1 = 0)
  • when the pH of the solution is less than the pKa of the compound, the compound exists primarily in its acidic form
  • when the pH of the solution is greater than the pKa of the compound, the compound exists primarily in its basic form
41
Q

What is a buffer solution?

A

a solution of a weak acid (HA) and its conjugate base (A-) is a buffer solution. A buffer solution maintains nearly constant pH when small amounts of acid or base are added to it, because the weak acid can give a proton to any HO- added to the solution and its conjugate base can accept any H+ that is added to the solution.

42
Q

What is a Lewis acid?

A

accepts an electron pair

43
Q

What is a Lewis base?

A

donates an electron pair

44
Q

Describe the position of equilibrium for the dissolvement of strong acid in water

A

products are favored, equilibrium lies to the right

45
Q

Describe the position of equilibrium for the dissolvement of weak acid in water

A

reactants are favored, equilibrium lies to the left

46
Q
A

Organic Chemistry (11 decks)

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