Chapter 17/Part of 18 Concepts

Question 1

Define conjugate acid and conjugate base and cite examples of conjugate pairs.

Answer 1

Conjugate acid: the substance formed from the base after the addition of an H+

Conjugate base: the substance formed from the acid after the loss of an H+

Examples:
HCL (acid) –> Cl- (conjugate base)
NH3 (base) –> NH4+ (conjugate acid)

Question 2

Give in your own words definitions of acids/bases: Brønsted and Lewis.

Answer 2

Bronsted Acid: a proton (H+) donor
Bronsted Base: a proton (H+) acceptor

Lewis Acid: an electron pair acceptor
Lewis Base: an electron pair donor

Question 3

Rank species of each of the following acids based upon relative strength: hydrohalic,
oxoacids, and carboxylic.

Answer 3

easier to lose H+, stronger acid
higher charge = stronger bond to H+ = weaker acid
shorter bonds are stronger so H+ harder to remove
less electronegative = stronger bond = weaker acid

Question 4

Use the pH scale to classify a solution as being acidic, basic or neutral.

Answer 4

pH < 7 –> acidic
pH = 7 –> neutral
pH > 7 –> basic

Question 5

Identify an acid or base as being strong or weak.

Answer 5

Strong acids: HCl, H2SO4, HNO3, HCLO4, HBR, HI

Strong bases: soluble hydroxides (NaOH, KOH, LiOH, Ba(OH)2)

Weak acids: H2CO3, H3PO4, CH3COOH, HF, HNO2, HCN, carboxylic acids

Weak bases: ammonia, amines, insoluble/slightly soluble hydroxides

Question 6

Describe the common ion function in buffers

Answer 6

a common ion prevents the weak acid/base from ionizing as much as it usually does and keeps the pH from changing

Question 7

Describe how a buffer works when absorbing external acids or bases.

Answer 7

when OH- is added, or H+ is reduced, the equilibrium will shift towards the conjugate base and the OH- will react with the H+ to form water

when H+ is added, the equilibrium will shift toward the undissociated acid