Autoionization of water Hydrolysis of salts

Август 2, 2022

Главная
Химия
Autoionization of water Hydrolysis of salts

Содержание

2. LESSON OBJECTIVES: Ionic product of water. Notion of pH Be able to calculate pH and pOH
3. WATER is a weak electrolyte and dissociated to: When the law of mass action is applied
4. (Н2О) =1,8 ⋅10-9 , it is mean that one water molecule in 550 million naturally dissociates
5. A water molecule that loses a hydrogen ion becomes a negatively charged hydroxide ion OH- A
6. Water, even pure water, has an amphiprotic nature. This means that a small amount of ions
7. Water undergoes auto-ionization according to the following equation: H2O(l) + H2O(l) => H3O+(aq) + OH-(aq) or
8. In pure water the concentration of OH- and H+ are equal: So any aqueous solution in
9. When hydrogen chloride dissolves in water, it forms hydrogen-ions: In the previous HCl solution – acidic
10. In pure water and any aqueous solution: so In 1909 Danish scientist Soren Sorensen introduced the
11. The pH scale is used to express [H+]
12. Classifying Solutions A solution in which [H+] is greater than 1 x 10-7 has a pH
14. In practice for measurement of pH water or solutions may be used acid-base indicators, and for
15. Indicators are halochromic chemical compounds (weak organic acids or bases that react with ions in solution)
16. A Universal indicator is a pH indicator composed of a blend of several compounds that exhibits
17. Indicator paper = tells the pH number (value) read by color comparison (qualitative) Ways to Test
18. Ways to Test pH Litmus paper = made from lichen (symbiotic organisms that are a combination
19. ACID-BASE INDICATORS
20. pH meter Measures amount of H+ ions in the solution Digital readout Most accurate way of
21. Buffers are solutions that have constant pH values and the ability to resist changes in pH.
22. Buffers A solution of ethanoic acid (CH3COOH) and sodium ethanoate (CH3COONa) is an example of a
23. CH3COO-(aq) + H+(aq) CH3COOH (aq) ethanoate ion hydrogen ion ethanoic acid When an acid is added
24. Buffers are important because many chemical reactions, particularly those in biological systems, proceed best at a
26. HYDROLISIS The reaction of salt takes place in the solution. In reality, and looking at a
27. A salt is formed between the reaction of an acid and a base. Usually, a neutral
28. 3) If the salt is formed from a strong base and weak acid, the salt solution
29. If the positive ion (cation) of the salt is from a weak base, it will hydrolyze
30. Salt hydrolysis can be described in two chemical equations, the first showing the dissociation of the
31. 1) SODIUM CHLORIDE NaCl NaCl + HOH NaOH + HCl strong strong base acid Na+ +
32. 2) AMMONIUM BROMIDE NH4Br NH4Br + H2O NH4OH + HBr weak strong base acid NH4+ +
33. 3) SODIUM NITRITE NaNO2 NaNO2 + H2O NaOH + HNO2 strong weak base acid Na+ +
34. 4) AMMONIUM ACETATE (CH3COO)NH4 (CH3COO)NH4 + HOH NH4OH + CH3COOH weak weak base acid CH3COO- +
36. Life Many processes that are essential to life involve hydrolysis. An example is the release of
37. GLOSSARY Acid – A compound that has a proton or protons that can dissociate in water;
38. Base – A compound that has the ability to accept a proton or protons from the
39. Conjugate acid. The form of a molecule that has a dissociable proton attached to it. Since
40. pH. A measure of the acidity of a solution. It is the negative log (base 10)
42. Скачать презентацию

Слайд 2

LESSON OBJECTIVES:
Ionic product of water. Notion of pH
Be able to calculate

pH and pOH
Be able to calculate hydrogen and hydroxide ion concentration from pH or pOH
Hydrolysis of salts

Слайд 3

WATER is a weak electrolyte and dissociated to:
When the law of

mass action is applied to the dissociation of water, we have:

1 mole 1 mole 1 mole

Слайд 4

(Н2О) =1,8 ⋅10-9 , it is mean that one water molecule

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

const

Слайд 5

A water molecule that loses a hydrogen ion becomes a 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

Слайд 6

Water, even pure water, has an amphiprotic nature. This means 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-

Слайд 7

Water undergoes auto-ionization according to the following equation:
H2O(l) + H2O(l) =>

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:

Слайд 8

In pure water the concentration of OH- and H+ are 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:

Слайд 9

When hydrogen chloride dissolves in water, it forms hydrogen-ions:
In the 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-]:

Слайд 10

In pure water and any aqueous solution:
so
In 1909 Danish scientist Soren

Sorensen introduced the concepts of pH and pOH values:

Слайд 11

The pH scale is used to express [H+]

Слайд 12

Classifying Solutions
A solution in which [H+] is greater than 1 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

Слайд 13

Слайд 14

In practice for measurement of pH water or solutions may 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.

Слайд 15

Indicators are halochromic chemical compounds (weak organic acids or bases 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.

Слайд 16

A Universal indicator is a pH indicator composed of a 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.

Слайд 17

Indicator paper = tells the pH number (value)
read by color comparison

(qualitative)

Ways to Test pH:

Слайд 18

Ways to Test pH
Litmus paper = made from lichen (symbiotic organisms

that are a combination of algae and fungus)
Color changes
red paper→ blue (base)
blue paper→ red (acid)

acid

base

Слайд 19

ACID-BASE INDICATORS

Слайд 20

pH meter
Measures amount of H+ ions in the solution
Digital readout

Most accurate way of determining pH because it is quantitative

Слайд 21

Buffers are solutions that have constant pH values and the 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

Слайд 22

Buffers
A solution of ethanoic acid (CH3COOH) and sodium ethanoate (CH3COONa) is

an example of a typical buffer.
CH3COOH and CH3COO- (source is the completely ionized CH3COONa) act as reservoirs of neutralizing power.

Слайд 23

CH3COO-(aq) + H+(aq) 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.

Слайд 24

Buffers are important because many chemical reactions, particularly those in 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.

Слайд 25

Слайд 26

HYDROLISIS
The reaction of salt takes place in the solution. In 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.

Слайд 27

A salt is formed between the reaction of an acid and a 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.

Слайд 28

3) If the salt is formed from a strong base and weak acid, the 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.

Слайд 29

If the positive ion (cation) of the salt is from 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.

Слайд 30

Salt hydrolysis can be described in two chemical equations,
the first

showing the dissociation of the salt,
and the second net equation showing the production of H+ or OH – ions:

Слайд 31

1) SODIUM CHLORIDE NaCl
NaCl + HOH <=> NaOH + HCl
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.

Слайд 32

2) AMMONIUM BROMIDE NH4Br
NH4Br + H2O <=> NH4OH + HBr
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.

Слайд 33

3) SODIUM NITRITE NaNO2
NaNO2 + H2O <=> NaOH + HNO2
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).

Слайд 34

4) AMMONIUM ACETATE (CH3COO)NH4
(CH3COO)NH4 + HOH <=> NH4OH + 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.

Слайд 35

Слайд 36

Life
Many processes that are essential to life involve hydrolysis. An 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.

Слайд 37

GLOSSARY
Acid – A compound that has a proton or protons 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-).

Слайд 38

Base – A compound that has the ability to accept 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.

Слайд 39

Conjugate acid. The form of a molecule that has a 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.

Слайд 40

pH. A measure of the acidity of a solution. It 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.