LECTURE № 7 Autoionization of water Hydrolysis of salts 14.03.2017

calculate pH and pOH
Be able to calculate hydrogen and hydroxide

ion concentration from pH or pOH
Hydrolysis of salts

of mass action is applied to the dissociation of water,

we have:

1 mole 1 mole 1 mole

+

molecule in 550 million naturally dissociates into OH- and H+

ions

const

negatively charged hydroxide ion OH-
A water molecule that gains a

hydrogen ion becomes a positively charged hydronium ionH3O+







Self ionization of water – the reaction in which water molecules produce ions.

Hydroxide
ion

Hydronium
ion

that a small amount of ions will form in pure

water. Some molecules of H2O will act as acids, each donating a proton to a corresponding H2O molecule that acts as a base. Thus, the proton-donating molecule becomes a hydroxide ion, OH-, while the proton-accepting molecule becomes a hydronium ion, H3O+.

hydronium
ion, H3O+

hydroxide
ion, OH-

=> H3O+(aq) + OH-(aq)
or  2 H2O(l)  => H3O+(aq) + OH-(aq)

The equilibirum expression

for the above reaction is written below and is treated mathematically like all equilibrium expressions: Kw = [H3O+][OH-]=1 x 10-14
At 25oC, the value of Kw has been determined to be 1 x 10-14.  This value, because it refers to the auto-ionization of water, has been given a special symbol, Kw, but, it is just a special case of Kc.
If one knows the concentration of either the hydronium ions or of the hydroxide ions in a water solution, the other ion concentration can be determined:

equal:
So any aqueous solution in which H+ and OH-

are equal is a neutral solution.

Not all solutions are neutral (example HCl+H2O or NaOH+H2O).

When some substances (acids, bases, salts) dissolve in water, they release hydrogen ions:

previous HCl solution – acidic solution, in which [H+] is

greater than [OH-]:

When solid sodium hydroxide dissolves in water, it forms hydroxide ions in solution:



In the above NaOH solution – basic solution, in which [H+] is less than [OH-]:

Soren Sorensen introduced the concepts of pH and pOH values:

x 10-7 has a pH less than 7.0 and is

acidic.

A solution in which [H+] is less than 1 x 10-7 has a pH greater than 7.0 and is basic.

The pH of pure water or a neutral aqueous solution is 7.0

Acidic solution: pH < 7.0 [H+] > 10-7mol/L
Neutral solution: pH = 7.0 [H+] = 10-7mol/L
Basic solution: pH > 7.0 [H+] < 10-7 mol/L

be used acid-base indicators, and for more accurate measurement -

pH meters

For more precious measuring of pH it is widely used the special tools – pH-meters, which provides assurance of measuring within the limits of ± 0,01.

that react with ions in solution) that change color depending

on the relative concentrations of H+ and OH- ions and added in small amounts to a solution so that the pH (acidity or basicity) of the solution can be determined visually.

This visually method is called a colorimetric.

The color change of different indicators occurs at different hydrogen ion concentrations, which is important for chemical analysis.

blend of several compounds that exhibits several smooth colors changes

over a pH value range from 1-14 to indicate the acidity or basicity of solutions.
Definition: A universal indicator is typically composed of water, propan-1-l, phenolphthalein sodium salt, sodium hydroxide, methyl red, bromothymol blue monosodium salt, and thymol blue monosodium salt.

comparison (qualitative)

Ways to Test pH:

organisms that are a combination of algae and fungus)

Color changes
red

paper→ blue (base)
blue paper→ red (acid)

acid

base

readout

Most accurate way of determining pH because it is

quantitative

ability to resist changes in pH.  If you add acid or base to a

buffered solution, its pH will not change significantly. Similarly, adding water to a buffer or allowing water to evaporate will not change the pH of a buffer.

A buffer is most easily prepared by dissolving an acid together with its conjugate base in the same solution:

СН3СООН + СН3СООNa

NH4OH + NH4Cl

base conjugate
acid

acid conjugate base

is an example of a typical buffer.


CH3COOH and CH3COO-

(source is the completely ionized CH3COONa) act as reservoirs of neutralizing power.

CH3COOH (aq)
ethanoate ion

hydrogen ion ethanoic acid

When an acid is added to the solution, the ethanoate ions act as a hydrogen-ion sponge.


CH3COOH (aq) + OH-(aq) CH3COO-(aq) + H2O(l)
Ethanoic acid hydroxide ion ethanoate ion water

When a base is added to the solution, the ethanoic acid and the hydroxide ions react to produce water and the ethanoate ion.

biological systems, proceed best at a particular pH. If the

reaction takes place in a solution that remains at that pH throughout the reaction, the most satisfactory results will be obtained.
Many life forms thrive only in a relatively small pH range so they utilize a buffer solution to maintain a constant pH. One example of a buffer solution found in nature is blood.
They are used to calibrate the pH meter.

reality, and looking at a wider variety of 'salts', the

picture is much more complicated and a 'salt' solution may be acid, neutral or alkaline depending on the nature of the interaction of the salt ions with water.
The reaction of the salt with water, whereby the salt is dissociated and decomposed to form a weak electrolyte (weak acid or weak base) called hydrolysis ("chemical decomposition by water," 1880, formed in English from hydro- + Greek lysis "a loosening, a dissolution," from lyein "to loosen, dissolve").
Hydrolysis is the reverse of neutralization.

base. Usually, a neutral salt is formed when a strong

acid and a strong base is neutralized in the reaction: H+ + OH- = H2O
There are four possible ways of forming salts:
1) If the salt is formed from a strong base and strong acid, then the salt solution is neutral, indicating that the bonds in the salt solution will not break apart (indicating no hydrolysis occurred) and is neutral (pH=7).
2) If the salt is formed from a strong acid and weak base, the bonds in the salt solution will break apart and becomes acidic (pH<7) and hydrolyzes.

salt solution is basic (pH>7) and hydrolyzes.
4) If the salt is formed

from a weak base and weak acid, will hydrolyze, but the acidity, basicity or neutral depends on the equilibrium constants of Ka and Kb. If the Ka value is greater than the Kb value, the resulting solution will be acidic and vice versa:
If Ka(cation) > Kb(anion) the solution of the salt is acidic.
If Ka(cation) = Kb(anion) the solution of the salt is neutral.
If Ka(cation) < Kb(anion) the solution of the salt is basic.

a weak base, it will hydrolyze water and an acidic

solution will form:
NH4+ + HOH => NH4OH + H+ (pH<7)
cation weak base
If the negative ion (anion) of the salt is from a weak acid, it will hydrolyze water and a basic solution will form:
NO2- + HOH => HNO2 + OH- (pH>7)
anion weak acid
If the positive and negative ion are from a strong base and strong acid, it will not react with the water molecule, and a neutral solution will form.

first showing the dissociation of the salt,

and the second

net equation showing the production of H+ or OH – ions:

strong strong
base acid
Na+ + Cl- + HOH <=> Na+ + OH- + H+ + Cl-

HOH <=> OH- + H+

(neutral medium, pH=7)

In solution strong base and strong acid are dissociated completely. The salt solution is neutral. No hydrolysis.

weak strong
base acid
NH4+  + Br- + H2O <=> NH4OH + H+ + Br-

NH4+  + H2O <=> NH4OH + H+ 
cation
hydrolyzed (acidic medium, pH<7)

Since HBr is a strong acid it breaks up and yields H+, the salt is acidic. NH4OH is a weak base. They generally stay together, however this is actually breaks down into ammonia and water.

strong weak
base acid
Na+  + NO2-  + H2O <=> Na+ + OH- + HNO2

NO2-  + H2O <=> OH-  + HNO2
anion
hydrolyzed (basic medium, pH>7)

Since NaOH is a strong base it breaks up and yields OH-, the salt is basic. HNO2 is a weak acid (does not break up in water).

CH3COOH

weak weak
base acid

CH3COO- + NH4+ + HOH<=> NH4OH + CH3COOH
anion and cation
hydrolyses (neutral medium, pH7)

It mean that so
Ka(cation)=Kb(anion) the solution of the salt is neutral.

example is the release of energy by the molecule adenosine

triphosphate (ATP).
Hydrolysis is also plays a vital role in the breakdown of food into easily absorbed nutrients. Most of the organic compounds in food do not react readily with water, and usually a catalyst is required to allow these processes to take place.
Industry
Many industrial procedures require various substances to be hydrolyzed to create useful products. Often, however, the raw materials for these processes do not react easily with water molecules, so the reactions are helped by a variety of means, such as high pressure, high temperatures and catalysts. Laboratory hydrolysis usually requires the use of a catalyst, which is typically a strong acid or alkali.
Hydrolysis has been used for a long time in the production of soap. During this process, known as saponification, fat is hydrolyzed in a reaction with water and the strong alkali, sodium hydroxide. The reaction produces fatty acid salts, commonly known as soap.
Weathering
Hydrolysis is an important process in the weathering of rocks. Various silicate minerals, such as feldspar, undergo slow hydrolysis reactions with water, forming clay and silt, along with soluble compounds. This process is important in the formation of soils, and in making essential minerals available to plants.

that can dissociate in water; also, when one molecule has

a proton or protons that dissociate more readily than those of another (i.e., it has a higher Ka), the first is said to be the more acidic molecule.
Acid dissociation constant: A form of the equilibrium constant for the dissociation of an acidic molecule into a proton and its conjugate base. It is abbreviated "Ka." The acid dissociation serves as a measure of how acidic the molecule is; the larger the value of Ka, the more acidic the molecule.
Acid dissociation equation. The chemical equation for the separation of an acidic molecule into a proton and its conjugate base. The general form can be written: HA + H2O   H3O+ + A-
where HA is the acidic molecule, H3O+ is the stable form of the proton, and A- is the conjugate base of the acidic molecule.
Acidic solutions – Solutions containing a higher concentration of hydronium ion (H3O+) than that found in pure water (i.e., having a pH below 7); also, when one solution has a greater concentration of hydronium than another, it is said to be the more acidic solution.
Arrhenius acid – Any molecule that can dissociate in an aqueous solution to produce a proton (H+).
Arrhenius base – Any molecule that can dissociate in an aqueous solution to produce a hydroxide ion (OH-).

a proton or protons from the surrounding solution. When one

molecule associates with a proton or protons from the surrounding solution more readily than another, the first is said to be the more basic molecule. A basic compound can also be referred to as "alkaline."
Basic solutions – solutions containing a lower concentration of hydronium ion than that found in pure water (that is, having a pH above 7). When one solution has a lower concentration of hydronium than another, it is said to be the more basic solution.
Brønsted-Lowry acid – A molecule that can dissociate in an aqueous solution to produce a proton, thus increasing the concentration of hydronium ion in the solution.
Brønsted-Lowry base – A molecule that can pick up a proton from an aqueous solution, thus decreasing the concentration of hydronium ion in the solution.
Buffer system – A weak acid and the salt of its conjugate base, which together can be used to create a buffered solution.
Buffer – An aqueous solution having the property that its pH changes very little upon the addition of an acid or a base. A buffer is formed by mixing together combinations of weak acids and weak bases.

dissociable proton attached to it. Since that proton can dissociate,

this molecule is an acid.
Conjugate base. The form of a molecule that has a proton dissociated from it. Since that proton could potentially re-associate with the molecule, it is said to be a base.
Dissociation. The process in which a molecule falls apart into two pieces, commonly used to describe when an acid loses a proton (H+) and becomes its conjugate base.
Equivalence point. The point in a titration when the number of moles of added reactant is exactly equal (or stoichiometrically proportional) to the number of moles of reactant in the sample.
Hydronium ion. The conjugate acid of water. It consists of a water molecule with an extra proton attached and has the formula H3O+.
Hydroxide ion. The conjugate base of water. It consists of a water molecule with one of the protons abstracted and has the formula OH-.
Neutralization reaction. A reaction in which an equal amount of an acid and a base are mixed together, cancelling each other out, and making the solution neutral with a pH of 7.

is the negative log (base 10) of the hydronium concentration

in molar (-log10 [H3O+]).
pH scale. A logarithmic scale of the acidity of a solution. For aqueous solutions it runs from -1.7 (most acidic) to 15.7 (most basic), though typical values lie between 0 and 14.
pH unit. One unit on the pH scale. A change of one pH unit in an aqueous solution corresponds to one order of magnitude change in the hydronium concentration.
pKa. A measure of the ease with which the proton dissociates from an acidic molecule. It is equal to the negative log (base 10) of the acid dissociation constant (-log10Ka).
Strong acids – Acids that dissociate completely in solution.
Strong bases – Bases that completely dissociate in solution, usually soluble metal hydroxides.
Weak acids – Acids that do not completely dissociate in solution.
Weak bases – Bases that do not completely dissociate in solution.

C2H5ONa + H+

поваренная соль алкоголят натрия

C2H5ONa + H2O = C2H5OH + NaOH
спирт щелочь