ACIDS, BASES, SALTS

How to get a weak acid out of its salt? E.g.hydrofluoric acid, HF,out of KF.

Answer:

Treat the salt with a strong acid, e.g. with H2SO4: KF(aq) + H2SO4(aq)

HF (aq)+KHSO4

Net ionic: F- + H+HF

Same approach to get a weak, or low-soluble base, out of the salt:

Treat the salt solution

with a strong base (alkali):

CuCl2(aq) + KOH(aq)

Cu(OH)2+ 2KCl(aq)

Net Ionic: Cu2+(aq) +2OH-(aq)

Cu(OH)2
OXYACIDSBASES:

A Unifying View

Oxyacids Bases are all elementhydroxides:

E-O-H

i.e. produced by combination of an oxide with water:

K2O+H2OK-OHK++OH-

Cl2O+H2OHO-ClH++ClO-

2 ways of E-O-H ionization:

Acid: E-O-HEO- + H+

Base: E-O-HE+ + OH-

Which way prevails? depends on the relative polarity of two bonds: +EO-vs.-OH+

The more polar will be ionized with water:

Cl & O are close in electronegativity, Cl – O is a low-polar bond,while H – Cl is highly polar bond, therefore:

H – O – Cl → H+ + ClO-

In KOH, H – O bond is less polar than ionic K O bond, therefore:KOH → K+ + OH-

PERIODIC TRENDS

in BASICITY - ACIDITY

The most active metals formbasic oxides & metal hydroxides (alkali).

Non-metals form acidic oxides & acids

Transition metals may form basic, or acidic oxides hydroxides.

Basicity-acidity depends on the oxidation number of the element:

Higher oxidation number  higher acidity

(more oxygen  higher acidity)

MnO, basic oxide

Mn(OH)2a base

Mn2O7 acidic oxide

HMnO4 permanganic acid & its salts, as KMnO4

CrO basic oxide,

Cr(OH)2 a base

CrO3 acidic oxide,
H2CrO4 Chromic acid & its salts, as Na2CrO4

Arrhenius definitions of acids bases were generalized to include non-aqueous solutions dry chemistry reactions.

BRÖNSTED-LOWRY definition:

Acid is any compound that can serve as a proton donor:

HCl(g)+H2OH3O+(aq)+Cl-(aq)

HCl(g) is proton (H+) donor to water

HCl(g)+NH3(g)NH4Cl(s)

HCl(g) is proton donor to ammonia.

No need for water or other solvent!

Brönsted BASE

is aprotonacceptor

:NH3(g) accepts a proton,

it is a base

Water is an acid when reacts with ammonia:

H2O + NH3(g)NH4+(aq) + OH-(aq)

In the reverse reaction:

NH4+(aq) is acid,OH-(aq)is base.

B-L acid-base reaction is a proton transfer.

In a neutralization reaction:

H3O+ + OH-2H2O
acid base

Anions of weak acids are Brönsted bases: they produce OH- with water:

CO32- +H2OHCO3- + OH-

this is hydrolysis reaction.

It is the reason whysalts of weakacids strong bases, when dissolved, produce

alkaline solutions:

Na2CO3 2Na+ + CO32-

CO32- +H2O HCO3- + OH-

BRÖNSTEDAcid-Base PAIRS

A CONJUGATE base is the species that remains after a Brönsted acid has given up H+:

HCN(aq)+H2OH3O+(aq)+CN-(aq)

ACID CONJUGATEBASE

A CONJUGATE acid is the species formed when a Brönsted base gains a proton:

NH3(aq)+H2ONH4+(aq)+OH-(aq)

BASE CONJUGATE

ACID

H2CO3+H2O  H3O+ + HCO3-

acid1 base2 conjugate conjugate

acid2 base1

Reactions between acids bases trend to form weaker bases acids, from stronger bases acids: HCl+NaHCO3H2CO3+NaCl

or, in net ionic form:H++HCO3-H2CO3

pH

logarithm reminder:

a=log10N means 10a=N

these are two ways to present the same statement.

(They are as similar as: 3=6/2 & 2=6/3)

pH  -log10[H+] or: [H+] = 10-pH

if [H+]= 10-3 M, pH=3, or pH 3

if [H+] = 2×10-3 M, pH 2.7

if pH 2, [H+] = 10-2 M

if pH -1, [H+] = 10 M
pH of common liquids and body fluids:

Battery electrolyte pH -1

Human gastric juice 1-2

Lemon juice 2.2-2.4

Vinaigre2.4-3.4

Carbonated drinks2-4

Black coffee3.7-4.1

Tomato juice4.0-4.4

Urine5.5-7.0

Cow’s milk6.3-6.6

Saliva6.5-7.5

Human blood7.3-7.5

Sea water7.8-8.3

Bile7.8-8.8

MonoLake10-10.5

0.1 M Sodium carbonate11.7

1 M NaOH14

pH

SELF-IONIZATION of WATER

Water is a very weak electrolyte. It is self-ionized by proton transfer:

H2O + H2O  H3O+ + OH-

In pure water:

[H3O+] = [OH-] = 10-7 mol/L

[H2O] = (1000g)/(18g/mol) =55.5mol/L

10-7<55.5, & is always the same in diluted aqueous solutions.

When an acid is added to water, the concentration of protons it releases is so much higher than what water would produce on its own, that the latter can be ignored.

In a [HCl] = 0.1 mol/L, HCl  H+ + Cl-

HCl, as a strong electrolyte, gives

[H+] = [HCl] = 0.1 mol/L, pH = -log0.1 = 1

In a 0.01 M HNO3, HNO3  H+ + NO3-

[H+] = [HNO3] = 0.01 mol/L, pH = -log0.01 = 2

In 1.0 M HCl: HCl  H+ + Cl-

[H+]= [HCl] = 1.0mol/L,

pH = -log1 = 0 (100=1)

Water self-ionization, as a source of protons, may be neglected compared to HCl ionization.

When acid is added, [H3O+] (or [H+])increases, while [OH-] goes down.

However:

the product of both[H+] &[OH-]concentrations is a constant independent of each of these concentrations.

At room temperature:

Kw= [H+].[OH-] = 10-14M2

H2O  H+ + OH-

backward rxn will occur when H+ & OH- collide.
The collisions frequency is proportional to their concentrations, i.e.
the rate of the backward rxn:

R = k [H+][OH-]

The rate of the forward rxn:

R = k [H2O]

At equilibrium:

R = R

k [H2O] = k [H+][OH-]

[H+][OH-] is equilibrium

K = constant

[H2O]

or Kw = K[H2O] = [H+][OH-]

Since [H2O] = 55.5 always, we put it into the constant.

Kw=[H+][OH-]=10-14

Ionic product for water

In pure water: [H+]=[OH-] = 10-7 M

In a 1.0 M HCl, [H+] = 1.0 M

In a 1.0 M KOH [OH-] = 1.0M

Since Kw= [H+][OH-] = 10-14

[H+] = Kw/[OH-] = 10-14/1.0 = 10-14M

i.e. pH = -log(10-14) = 14

The range of possible [H+] is about 16 orders of magnitude wide.

Therefore, logarithmic scale is convenient.

pH = -log[H+]

Range of pH values: -1 to 15.

Neutral solutions: [H+] = 10-7M, pH 7

Acidic solutions: low pH, pH<7

Alkaline solutions: high pHpH>7

For strong acids & bases in concentrations that are exact power of 10, pH is self-evident: [HCl] = 10-2M, [H+] = 10-2M, pH 2

[KOH] = 10-3M. [OH-] = 10-3M, p[OH-]=3

where by definition: pOH = - log[OH-]

Since Kw=[H+][OH-] = 10-14

pH + pOH = 14

pH =14 - pOH = 14 - 3 = 11

When the concentration differs from integer power of 10, use calculator.

[HCl] = 3.4x10-5M. pH-? pH = -log(3.4x10-5) = 4.47

[KOH] = 0.001 M

KOH  K+ + OH- [OH-] = 0.001 M = 10-3M

pOH = 3, pH = 14-3=11 or [H+] = 10-11M

pHINDICATORS

are weak acids HIn H+ + In-

which have different colors of their neutral molecular form HIn as compared to that of anion In-

During titration of an alkali, KOH, with phenolphtalein, initial KOH solution has color of In- anion (pink), since excess of alkali shifts the equilibrium

HIn H+ + In-

When alkali is neutralized by the acid, pH decreases sharply, the color changes to that of HIn(colorless for phenolphtalein). This moment is equivalence point.

Phenolphthalein is an example of an indicator which establishes this type of equilibrium in aqueous solution:

Colourless (Acidic) / Rasberry red (Alkaline)

Transition occurs at pH9

Indicator / Acid / Base / pH range
Thymol Blue - 1st change / red / yellow / 1.2 - 2.8
Methyl Orange / red / yellow / 3.2 - 4.4
Bromocresol Green / yellow / blue / 3.8 - 5.4
Methyl Red / yellow / red / 4.8 - 6.0
Bromothymol Blue / yellow / blue / 6.0 - 7.6
Phenol Red / yellow / red / 6.8 - 8.4
Thymol Blue - 2nd change / yellow / blue / 8.0 - 9.6
Phenolphthalein / colourless / pink / 8.2 - 10.0

How pH is estimated?

Paper strips arepre-socked with indicator solutions covering the entire pH range:

Another way: use an electronic instrument: pH meter

(was invented byBeckman in California, in 1930-s, who combined the glass electrode with then new electronic amplifyer. This instrument, initially developed for CA orange juice companies to determine the quality of juice by estimating the content of ascorbic acid in it, started his multi-billion buisiness. Beckmann was a major benefactor of American research & education: Beckman institutes, etc.)