Chapter 9 Acids And Bases

Front 1

An acid

Back 1

contains a hydrogen atom and dissolves in water to form a hydrogen ion H+

Front 2

a base

Back 2

contains hydroxide and dissolves in water to form -OH

Front 3

H+ and H3O+

Back 3

are sometimes used interchangeably by chemist keep in mind however that H+(aq) does not exist in aqueous solution

Front 4

A Bronsted-Lowry acid

Back 4

is a proton donor

Front 5

A Bronsted-Lowery base

Back 5

is a proton donor

Front 6

HCL is a Bronsted-lowery

Back 6

acid because it donates a proton to the solvent water

Front 7

H2O is a Bronsted- Lowery

Back 7

base because it accepts a proton from HCL

Front 8

A Bronsted-Lowery

Back 8

acid must contain a hydrogen atom

Front 9

Although HCl is a covalent molecule and a gas at room tempeture

Back 9

whin it dosolves in water it ionizes forming two ions H3O+ and Cl-
an aqueous solution of hydrogen chloride is called hydrochloric acid

Front 10

Because a Bronsted-lowery acid contains a hydrogen atom, a general Bronsted -Lowery acid is often

Back 10

written as HA A can be a single atom such as Cl or Br

Front 11

Common Bronsted-Lowery acids

Back 11

HCL hydrochloric acid
H2 SO4 sulfuric acid
HBr hydrobromic acid
HNO3 nitric acid

Front 12

A monoprotic acid

Back 12

contains one acidic proton HCl is a monoprotic acid

Front 13

A diprotic acid

Back 13

contains two acidic protons H2SO4 is a diprotic acid

Front 14

A triprotic acid

Back 14

contains three acidic protons H3PO4 is a triprotic acid

Front 15

A Bronsted-Lowry base is a proton acceptor

Back 15

a base must contain a lone pair of electrons

Front 16

Ammonia (NH3) is a Bronsted-Lowery base

Back 16

because it contains a nitrogen atom with a lone pair of electrons

Front 17

When NH3 is dissolved in water

Back 17

its N atom accepts a ptoton from H2O forming an ammoniun cation (NH4+) and Hydroxide (-OH)

Front 18

Common Bronsted-Lowery Bases

Back 18

NH3 ammonia
NaOH sodium hydroxide
Mg(OH)2 Magnesuim hydroxide
H2O Water
KOH potassium hydroxide
Ca(OH)2 calcuim hydroxide

Front 19

The Bronsted-Lowery acid donates a proton

Back 19

to the Bronsted-powery base which accepts it

Front 20

Consider for example the reaction of the general acid H-A with the general base B

Back 20

in an acid-base reaction, one bond is broken and one bond is formed

Front 21

The product formed by loss of a proton

Back 21

from acid is called its conjugate base

Front 22

the product formed by gain of a proton

Back 22

by a base is called its conjugate acid

Front 23

Two species that differ by the presence of a proton

Back 23

are called a conjugate acid-base pair

Front 24

H-Br + H2O Both ways Br- + H3O+

Back 24

H-BR loss of H+ Br- conjugate base

H2O gain of H+ conjugate acid
net charge must be the same on both sides of the equation

Front 25

When a species gains a proton (H+)

Back 25

it gains a +1 charge

Front 26

When a species loses a proton (H+)

Back 26

it effectively gains a -1 charge

Front 27

0 charge
H2O base add H+

Back 27

+1 charge
H3O+

Front 28

0 charge
H-BR acid Lose H+

Back 28

- 1 charge
Br-

Front 29

A Bronsted -Lowery acid -base reaction

Back 29

is a proton transfer reaction since it always results in the transfer of a proton from an acid to a base

Front 30

A compound that contains both a hydrogen atom and a lone pair of electrons

Back 30

can be either an acid or a base depending on the particular reaction

Front 31

H2O as a base H-O-H add H+
base

Back 31

H-O-H
H
Conjugate acid

Front 32

H2O as an acid H-O-H remove H+
acid

Back 32

H-O
conjugate base

Front 33

When a covalent acid dissolves in water, proton transfer forms H3O+ and is an anion

Back 33

This process is called dissociation acids differ in their tendency to donate a proton; that is acids differ in the extent to which they dissociate in water

Front 34

A strong acid readily donates a proton.

Back 34

When a strong acid dissolves in water, essentially 100% of the acid dissociates into ions

Front 35

A weak acid less readily donates a proton. When a weak acid dissolves in water

Back 35

only a small fraction of the acid dissociates into ions

Front 36

HCl (g)+ H2O(l) favors equilibrium
strong acid

H2SO4(l) + H2O (l) favors equilibrium
Strong acid

Back 36

H3O+ (aq) + Cl- conjugate base

H3O+ (aq) + HSO4 conjugate base

Front 37

Strong acids Conjugate base
HI I-
HBr Br-
HCl Cl-
H2SO4 HSO4 -
HNO3 NO3-
H3O+ H2O

Back 37

Weak Acids conjugate base
H3PO4 H2PO4-
HF F-
CH3COOH CH3COO-
H2CO3 HCO3-
NH4+ NH3
HCN -CN
H2O -OH

Front 38

CH3Cooh (l) + H2O use unequal
weak acid reaction arrows
the reactants are
favored equilibrium,

Back 38

H3O+ (aq) + CH3 COO-
conjugate base

Front 39

A strong base readily accepts a proton

Back 39

when a strong base dissolves in water essentially 100% of the base dissociates into ions

Front 40

A weak base readily accepts a proton

Back 40

when a weak base dissolves in water only a small fraction of the base forms ions

Front 41

A strong acid readily donates a proton

Back 41

forming a weak conjugate base

Front 42

A strong base readily accepts a proton

Back 42

forming a weak conjugate acid

Front 43

The position of the equilibrium

Back 43

depends upon the strength of the acids and bases

Front 44

When the stronger acid and base are the reactants on the left side

Back 44

the reaction readily occurs and the reaction proceeds to the right

Front 45

H-A + B
Strong acid strong base
a larger arrow to the reactants is favored

Back 45

A: + H-B
weaker base weaker acid
products are favored

Front 46

H--A + B
weaker acid weaker base
reactants are favored
a larger reverse arrow means that reactants are favored

Back 46

A + H---B
stronger base stronger acid

Front 47

If an acid base reaction would form the stronger acid and base

Back 47

equilibrium favors the reactants and little product forms

Front 48

Reaction HA (aq) + H2O (l)

H3O+ (aq) + A- product

Back 48

Equilibrium constant

[H30+][A-]
[HA][H2)] = K

Front 49

Acid dissociation constant Ka

Back 49

Ka = K[H2O] = [H3O+] [A-]
[HA]
the stronger the acid the larger the number for Ka

Front 50

Acid dissociation Constants (Ka) for common weak acids

Back 50

HSO4- 1.2x10 -2exponent
H3PO4 7.5 x 10 -3 exponent
HF 7.2 x 10 -4 exponent
CH3COOH 1.8x 10 -5 exponent

Front 51

Equilibrium favors the formation of the weaker acid

Back 51

that is the acid with the smaller Ka value

Front 52

One molecule of H2O donates a proton (H+)

Back 52

forming the conjugate base -OH

Front 53

one molecule of H2O accepts a proton

Back 53

forming is conjugate acid H3O+

Front 54

The product [H3O] [-OH] is a constant

Back 54

1.0x10- 14 exponent for all aqueous solutions at 25C

Front 55

In an acidic solution [H3O+] > [-OH]

Back 55

thus [H3O] .10-7 exponent

Front 56

Pure water and any solution that has an equal concentration of H3O+ and -OH ions

Back 56

1.0 x 10 -7 exponent is said to be neutral

Front 57

In a basic solution [-OH]>[H3O]

Back 57

thus {-OH] > 10 -7 exponent

Front 58

The pH scale is used to conveniently report [H3O+]

Back 58

The pH of a solution is a number generally between 1 and 14

Front 59

Acidic solution pH<7
Neutral solution pH =7
Basic solution pH > 7

Back 59

[H3O+] >1 x 10 -7 exponent
[H3O+] =1 x 10 -7 exponent
[H3O+] < 1 x 10 -7 exponent

Front 60

on the pH scale acidity is between and 6.9

Back 60

on the pH scale base is measured between 7.1 and 14

Front 61

Neutralization reaction

Back 61

an acid-base reaction that produces a salt and water as products

Front 62

The acid HA donates a proton (H+)

Back 62

to the -OH base to form h2o

Front 63

The anion A- from the acid combines with the cation M+

Back 63

from the base to form the salt MA

Front 64

HA(aq) + MOH (aq)
acid base

Back 64

H----OH(l) + MA (aq)
water salt

Front 65

A net ionic equation contains only the species involved in a reaction

Back 65

we can write an equation that contains only the species that are actually in the reaction such an equation is called a net ionic equation

Front 66

H+(aq) + Cl-(aq) + Na- (aq) + -OH
omit the spectator ion Cl- and Na+

Back 66

H+ (aq) + -OH reaction H-----OH(l)
Net ionic equation

Front 67

Whenever a strong acid and strong base react

Back 67

the net ionic equation is always the same H+ reacts with -OH to form H2O

Front 68

two products are always formed

Back 68

water and a metal salt

Front 69

A salt can form and acidic,base or nuetral

Back 69

solution depending on whether its cation and anion are derived from a strong or weak acid and base

Front 70

The cation M+

Back 70

comes from a base

Front 71

the anion A-

Back 71

comes from the acid HA

Front 72

A salt derived from a strong base and a strong acid

Back 72

forms a neutral solution (pH=7)

Front 73

The ion from the stronger acid or base

Back 73

determines whether the solution is acidic or basic

Front 74

A Salt derived from a strong base and a weak acid

Back 74

forms a basic solution (pH>7)

Front 75

A salt derived from a weak base and a strong acid

Back 75

forms an acidic solution (pH<7)

Front 76

NaHCO3 a salt derived from a strong base and a weak acid forms a basic solution

Back 76

HCO3-(aq) + H2O(l) reaction both ways
H2CO3 (aq) + -OH (aq)
hydroxide makes the solution basic
so the pH >7

Front 77

NH4Cl a salt derived from a strong acid and a weak base forms an acidic solution

Back 77

NH4+(aq) + H2O reaction both ways
NH3 (aq) + H3O+ (aq)
H3O+ makes the solution acidic
so the pH<7

Front 78

To determine the molarity of a solution

Back 78

we carry out a titration

Front 79

When the number of moles of base added equals the number of moles of acid in a flask

Back 79

the acid is nuetralized forming a salt and water

Front 80

Volume of solution NaOH M mol/l conversion factor=mole of chemical mol mol conversion= moles of HCl m(mol/L) conversion factor Molarity of solution

Back 80

Volume of NaOH
22.5 mL NaOH x 1L
1000mL =0.00225 mol NaOH
in this reaction one mole of HCl react with one mole of NaOH so the number of mole of NaOH equals the number of mole of HCl

mol = 0.00225 mol HCl x 1000 mL
L 25.0 ml solution 1L

= 0.900 M Cl

Front 81

A buffer is a solution whose

Back 81

pH changes very little when acid or base is added

Front 82

The weak acid of the buffer

Back 82

reacts with added base -OH

Front 83

the conjugate base of the buffer

Back 83

reacts with added acid H3O+

Front 84

If the pH of the blood is lower then 7.35 the blood is more acidic this condition is called

Back 84

acidosis

Front 85

If the pH of the blood is higher then 7.45 the blood is more basic this condition is called

Back 85

alkalosis

Front 86

Respiratory acidosis

Back 86

results when the body fail to eliminate adequate amounts of CO2 through the lungs

Front 87

Respiratory alkalosis

Back 87

is caused by hyperventilation it decreases CO2

Front 88

Metabolic acidosis

Back 88

results when excessive amounts of acid are produced and blood pH begins to fall

Front 89

Metabolic alkalosis

Back 89

may occur when recurrent vomiting decre4ases the amount of acid in the stomach thus causing a rise in pH