5.1.3 - Acid, Bases and Buffers

Question 1

define bronsted-lowry acid

Answer 1

a species that donates a proton

Question 2

define brosted-lowry base

Answer 2

a species that accepts a proton

Question 3

what happens in terms of electrons in a brosted lowry base

Answer 3

lone pair of electrons that form a dative covalent bond with H

Question 4

what is an amopheric substance

Answer 4

a substance that can act as either an acid or a base depending on the condition

Question 5

example of an amopheric substance

Answer 5

water

Question 6

monobasic acids

Answer 6

an acid that can donate 1 proton

Question 7

example of a monobasic acid

Answer 7

HCl and CH3COOH

Question 8

dibasic acids

Answer 8

acids with 2 protons to donate

Question 9

example of a dibasic acid

Answer 9

H2SO4 and H2CO3

Question 10

tribasic acids

Answer 10

acids with 3 protons to donate

Question 11

example of a tribasic acid

Answer 11

H3BO3

Question 12

ionic equation for: acid + metal

Answer 12

2H+(aq) + metal(s) –> metal ion2+(aq) + H2(g)

Question 13

ionic equation for: acid + carbonate

Answer 13

2H+(aq) + CO3 2+(aq) –> H2O(l) + CO2(g)

Question 14

ionic equation for: acid + metal oxide

Answer 14

2H+(aq) + metal oxide(s) –> metal ion2+(aq) + H2O(aq)

Question 15

ionic equation for: acid + alkali

Answer 15

H+(aq) + OH-(aq) –> H2O(l)

Question 16

what is the acid dissociation constant

Answer 16

Ka

Question 17

equation for Ka

Answer 17

Ka = [H+][A-]/[HA]

Question 18

assumptions when calulation things for a weak acid using Ka

Answer 18

1. when HA dissociates, [H] = [A]
2. [HA]eq = [HA]start

Question 19

problems with the assumption that [HA]eq = [HA]start

Answer 19

It becomes less accurate the bigger Ka becomes

Question 20

what does the size of Ka tell us about an acid

Answer 20

The larger the Ka, the stronger the acid and the higher the H+ concentration at equilibrium

Question 21

formula for pH

Answer 21

pH = -log[H+]

Question 22

how to find [H+] from pH

Answer 22

[H+] = 10^-pH

Question 23

formula for pKa

Answer 23

pKa = -log(Ka)

Question 24

how to find Ka from pKa

Answer 24

Ka = 10^-pKa

Question 25

what does 'p' imply

Answer 25

negative logarithmic scale

Question 26

what is Kw

Answer 26

ionic product of water: H2O(l) ⇌ H+(aq) + OH-(aq)

Question 27

equation for Kw

Answer 27

Kw = [H+][OH-]

Question 28

why can [H2O] be ignored when calculating Kw

Answer 28

it is constant at a given temperature

Question 29

what is Kw at room temperature (give room temperature in K)

Answer 29

1x10^-14 room temperature: 298K

Question 30

assumptions when calculating Kw

Answer 30

in a neutral substance [H+] = [OH] therefore Kw = [H+]^2

Question 31

calculating the pH of strong bases using Kw

Answer 31

use [OH-] and Kw to find [H+]

Question 32

assumptions when calculating the pH of strong bases

Answer 32

1. fully ionise in water eg. NaOH --> Na+ + OH- 2. assume [NaOH] = [Na+] = [OH-] 3. [base] = [OH-]

Question 33

define buffer

Answer 33

a solution that resists pH changes when small amounts of acid or base are added, or when diluted

Question 34

what are buffers formed from, give eg.

Answer 34

weak acid + salt of weak acid ethanoic acid + sodium ethanoate

Question 35

why do we assume all ethanoate ions come from sodium ethanoate in a buffer

Answer 35

CH3COOH ⇌ CH3COO- + H+ CH3COONa ---> CH3COO- + Na+ equilibrium for dissociation of ethanoic acid is so far left its negligable therefore [CH3COONa] = [CH3COO-]

Question 36

what happens when you add an acid to a buffer

Answer 36

pH decreases increased [H+] H+ reacts with CH3COO- as H+ is removed, equilibrium shifts left and pH increases

Question 37

what happens when you add alkali to a buffer

Answer 37

pH increases increased [OH-] OH- reacts with H+ [H+] decreases so equilibrium shifts right to increase [H+] again decreasing pH

Question 38

calculating pH of buffers

Answer 38

Ka = [H+][A-]/[HA]

Question 39

why cant we assume [H+] = [A-] when calculating the pH of buffers

Answer 39

[A-] from the base

Question 40

equation for the blood buffer system

Answer 40

H2CO3 ⇌ H+ + HCO3-

Question 41

function of the blood buffer system

Answer 41

to maintain the pH of blood between 7.35 and 7.45

Question 42

acidocis

Answer 42

when blood pH falls below 7.35 - fatigue, shock and death

Question 43

alkalosis

Answer 43

when blood pH falls above 7.45 - nausea, muscle spasms, headaches

Question 44

what is the equivalence point on a pH curve

Answer 44

when the volume of acid reacts exactly with the volume of alkali in the amounts matching the stoichiometry

Question 45

how to match an indicator to a pH curve

Answer 45

the verical section of the curve to match the pH range of the indicator