Ch. 2: Acids and Bases

Question 1

What is a Bronsted-Lowry acid?

Answer 1

a species that loses a proton

Question 2

What is a Bronsted-Lowry base?

Answer 2

a species that gains a proton

Question 3

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

Answer 3

proton transfer reaction or an acid-base reaction

Question 4

What is the conjugate acid of NH3?

Answer 4

+NH4

Question 5

What’s the conjugate base of H2O?

Answer 5

HO-

Question 6

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

Answer 6

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.

Question 7

What is acidity?

Answer 7

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

Question 8

What is basicity?

Answer 8

a measure of a compound’s affinity for a proton

Question 9

Describe a strong acid and its conjugate base.

Answer 9

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

Question 10

Describe a weak acid and its conjugate base.

Answer 10

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

Question 11

What is the equilibrium constant Keq?

Answer 11

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

Question 12

What is the acid dissociation constant?

Answer 12

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.

Question 13

Describe pKa and its use.

Answer 13

A convenient way to describe the strength of an acid.

Question 14

What is pH?

Answer 14

The concentration of protons in a solution.

Question 15

What are carboxylic acids?

Answer 15

compounds with a COOH group

Question 16

Are carboxylic acids strong or weak?

Answer 16

weak acids

Question 17

What are alcohols?

Answer 17

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

Question 18

Are alcohols stronger or weaker than carboxylic acids?

Answer 18

much weaker than carboxylic acids

Question 19

What are amines?

Answer 19

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

Question 20

Explain the use of curved arrows in a equation.

Answer 20

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.

Question 21

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

Answer 21

less than 0

Question 22

What is the approximate pKa value of carboxylic acids?

Answer 22

~5

Question 23

What is the approximate pKa value of protonated amines?

Answer 23

~10

Question 24

What is the approximate pKa value of alcohols and water?

Answer 24

~15

Question 25

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

Answer 25

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.

Question 26

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

Answer 26

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.

Question 27

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

Answer 27

Question 28

How does electronegativity affect the strength of an acid?

Answer 28

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

Question 29

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

Answer 29

because oxygen is more electronegative than nitrogen

Question 30

Rank the hybrid orbitals by electronegativity.

Answer 30

sp > sp2 > sp3

Question 31

Why does hybridization affect electronegativity?

Answer 31

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.

Question 32

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

Answer 32

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.

Question 33

Why are larger ions more stable than smaller ones?

Answer 33

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-.

Question 34

What is substitution?

Answer 34

The term for replacing an atom in a compound

Question 35

What is inductive electron withdrawal?

Answer 35

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.

Question 36

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?

Answer 36

1) inductive electron withdrawal 2) delocalized electrons

Question 37

What are delocalized electrons?

Answer 37

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.

Question 38

What are the five factors that determine acid strength?

Answer 38

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

Question 39

How is the relationship between pKa and pH represented?

Answer 39

the Henderson-Hasselbach equation

Question 40

What does the Henderson-Hasselbach equation tell us?

Answer 40

- 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

Question 41

What is a buffer solution?

Answer 41

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.

Question 42

What is a Lewis acid?

Answer 42

accepts an electron pair

Question 43

What is a Lewis base?

Answer 43

donates an electron pair

Question 44

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

Answer 44

products are favored, equilibrium lies to the right

Question 45

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

Answer 45

reactants are favored, equilibrium lies to the left