Analytical Chemistry and Chemical Analysis

 

Analytical chemistry is one of the chemical disciplines. Analytical chemistry is united with other chemical sciences with common chemical laws and based on studying of chemical properties of substances.

Analytical chemistry is the chemical science about

–       theoretical base of chemical analysis of substances;

–       method of detection and identification of chemical elements;

–       methods of qualitative determination of substances;

–       methods of selection (separation) of chemical elements and its compounds;

–       methods of establishing the structure of chemical compounds.

Subjects of analytical chemistry are: chemical elements and its compounds and processing of transformation of substances in run chemical reactions.

The main tool of analytical analysis is chemical reaction as a source of information about chemical composition of substances using for qualitative and quantitative analysis.

Aims of analytical chemistry are:

1. Establishing the chemical composition of analysed object (isotopic, elementary, ionic, molecular, phase) – qualitative analysis.

         Qualitative analysis consist from

–       identification – establishing of identity of researched chemical compounds with well-known substance du to compare its physical and chemical properties

–       and detection – checking the presence in analysed objects some components, impurities, functional groups etc.

2. Determination of content (amount and concentration) some components in analysed objects – quantitative analysis.

3. Determination (establishing) of structure of chemical compound – nature and number of structural elements, its bonds one to another, disposition in space.

4. Detection of heterogeneity on surface or in volume of solids, distribution of elements in layers.

5. Research process in time: establishing character, mechanism and rate of molecular regrouping.

6. Developing of present analytical methods theory, working out the new methods of analysis.

 

Analytical chemistry achieves the aims by various methods of analysis:

I.   Physical – determination of components of investigated substances without chemical reactions (destroying of sample):

1. Spectral analysis – investigation of emission and absorption spectra.

2. Fluorescence analysis – investigation of luminescence, caused action of UV-radiation.

3. Roentgen-structural analysis – using X-ray.

4. Mass-spectra analysis.

5. Densimetry – measurement of density.

II. Instrumental (physical-chemical) – based on measurement of physical parameters (properties) of substances in run of chemical reaction. This method divides on

1. Electrochemical – measurement of electrical parameters of electrochemical reactions.

2. Optical – investigation the influence of various electromagnetic radiation on substance.

3. Thermal (heating) – investigation the changes the properties of substance by heat (undergo) action.

III. Chemical – measurement of chemical bonds energy.

Chemical analysis has some steps:

1. Sampling.

2. Dissolving the sample (in water, acid or alkali).

3. Executing (running) the chemical reaction X + R ® P.

4. Measurement of definite parameter.

In accordance to analytical reaction (X + R ® P) applies three groups of chemical analysis methods:

I.   Measurement of amount (quantity) of reaction product P: mass, physical properties.

II. Measurement of amount of reagent R that interacted with determined substance: volume of solution reagent R with known concentration.

III. Registration changes of substance X acting with reagent: measurement of gas volumes.

 

IUPAC Classification of analytical methods in accordance with mass and volume of analytic sample

Method name	Mass of sample, g	Volume of sample, ml
Gramm-method	1–10	10–100
Santigramm-method	0,05–0,5	1–10
Milligramm-method	10-6–0,001	10-4–0,1
Microgramm-method	10-9–10-6	10-6–10-4
Nanogramm-method	10-12–10-9	10-10–10-7
Picogramm-method	10-12	10-10

 

Analytical Reactions and Requirements to Analytical Reactions

For identification (detection) and determination of substances the chemical reactions runs in solution or by “dry” way. These reactions always accompany the various external effects (analytical signals):

–       precipitation or dissolving of precipitate;

–       formation of coloured compound;

–       evolution of gas with specific properties (colour, odour).

“Dry” way testing (without dissolving of sample) can be make by:

1) pyrochemical methods:

–       flame test (colouring of gas torch flame),

–       making a glass (alloys with Na₂CO₃, K₂CO₃, Na₂B₄O₇, Na(NH₄)₂PO₄),

–       tempering;

2) crush (rub) sample to powder with analytical reagent;

3) microcrystalloscopic analysis – produce (receive) the specific crystals with analytical reagent and watching its with microscope (forms of crystals);

4) analysis in drops on filter paper – reaction between analysed substance and analytical reagent run on filter paper with some drops (1-2) of solutions – arise a coloured spots.

Requirements (demands) to analytical reactions:

1) reaction must run quickly, in practice – immediately;

2) reaction must accompanied with accordance (special) analytical effect;

3) reaction must be irreversible – run in one way (in one side);

4) reaction must have high specificity and have high sensitivity.

 

Description (characteristic) of analytical reactions.

At field of application in qualitative analysis the analytical reactions divide into group and individual (characteristic) reactions.

Group reactions use for selection from complex (complicated) mixes some substances. Substances with definite properties are united in special analytical groups.

This reactions use for:

a) detection the present analytical group;

b) selection this analytical group from another during systematic path (way) of analysis;

c) concentration of small amounts of substances;

d) separation groups, which prevent to analysis path.

Characteristic reactions named analytical reactions that have the individual substance nature. These reactions distinguish to selectivity.

Selective reactions give identical or alike analytical effects with small (little) number of ions (2-5).

Extreme form of selectivity is specificity. Specific reaction gives an analytical effect only with one individual substance.

For examples:  – iodine with starch – complex compound blue (navy) colour;

 – or Fe⁺³ with K₄[Fe(CN)₆] – complex compound blue (navy) colour.

 

Analytical reactions allow us to determine same quantity (amount) of substance.

Sensitivity of analytical reaction is the least amount (quantity) of substance, which can be detected with the reagent in one drop of solution (1 mm³).

The sensitivity express to next correlated values:

Limit of detection = Detected limit (m) – the least amount of substance, which present in analysed solution and which detect with the reagent. Calculate in mg. 1 mg = 0,000001 g.

Limit of concentration = Minimal concentration (C_min) – the least concentration of solution with still can be detected an analysed substance in definite (one drop) volume.

Limit of dilution (W = 1/C_min) – quantity (ml) of water solution, containing 1 g of the analysed substance, which detect with definite reaction (reagent).

         Thus, the sensitivity of analytical reaction is as more as limit of detection and limit of concentration are less.

These parameters are connected such:

m = C_min·V_min·10⁶ = V_min·10⁶ / W

Sensitivity is the most importance description of quantitative analytical reaction. Methods (modes, ways) to raise the sensitivity

1. To rise the concentration of detected substance:

–       to steam (soften by steam) of solution to small volume;

–       to extract with organic solvents to small volumes;

–       to distillate (rectify).

2. To precipitate of detected substance and dissolving the sediment in another solvent.

3. To use collector – substance, which adsorb the detected substance.

4. To mask the preventing ions (substances).

         For example, using the complex compounds for detecting Fe⁺³ and Co⁺² ions by reaction of with thiocyanate-anion:

Co⁺² + 6NH₄SCN ® [Co(SCN)₆] ^–4 + 6NH₄⁺,

        blue soution

Fe⁺³ + 6NH₄SCN ® [Fe(SCN)₆]^–3 + 6NH₄⁺

   bloody-red solution

Mixture of these cations cannot be analysed directly because Fe³⁺-complex has very colour that prevents watching the Co⁺²-complex. For masking of preventing Fe⁺³ cation to analysed solution ads ammonia fluoride, which forms strong colourless complex with iron(III) cations:

Fe⁺³ + 6NH₄F ® [FeF₆] + 6NH₄⁺

Formed fluoride complex not reacts with ammonia thiocyanate and not prevents aim reaction run. Masking (repression) is neutralisation influence of preventing agents.

 

The analysis of complex (complicated) mixes makes to next modes (ways):

I.                   Divide the mix on components (submixes) du to separation the detected substances and the preventing substances on various parts of mix (in various submixes) – systematic path (way) of analysis.

The systematic analysis – is full analysis of researched objects, which made du to separation of mix on groups (analytical groups) in definite (strong) sequence in accordance to various analytical properties of components. These separation makes until in one submix (phase) stay components, which simple detect (identify) with selective reagent.

II.                Separate and detect one component in the researched mix (without divide) with the help of (by means of) specific reactions (reagents) – fractional path (way) of analysis.

The fractional analysis – the all mix divide on identical (the same) parts. And in each part detect only one individual component.

At this path of analysis often use a masking.

Cations Classification

For selection of cations on analytical groups used group reagents. Accordance to applied group reagents all cations are divided on various systems. Cations divide to analytical groups in according with solubility of salts, formed by its.

Use of general and group reagents gave rise to creation the series of analytical cations classifications. Most widely used from them are sulphide, acid-basic and ammonia-phosphate. Analytical classifications of cations are based on chemical properties of their compounds and are associated with disposition of elements in periodic table, their structure and physico-chemical properties.

In all classifications there is a cations group, which does not have group peagent (cations of lithium, potassium, sodium, and ion of ammonium, which has the ion radius similar to the potassium ion). These are cations of the s¹ elements with electronic structure of inert gas, low electronegativity, with small radius, and small polarisation properties. Majorities of their salts are well water-soluble by reason of high tie polarity.   In periodic system they dispose in ІА-sub-group. In sulphide classification to this group is concluded a magnesium cation, which has similar lithium cation properties.

In all of classifications identical is the group of cations, which sediment by sulphate acid, ammonium carbonate, and sodium hydrogenphosphate in ammonia presence.  There are the cations of the s² elements: calcium, barium, and strontium, which are found in ІІА-sub-group of periodic system. Precipitates of their carbonates, sulphates and phosphates formed with complicated anions of oxygen-containing acids, which lightly polarize. In phosphate classification here are included cations of the s-elements – magnesium, and d-elements – iron (ІІ and ІІІ), chrome (ІІІ), manganese (ІІ), which form precipitates with three-charged phosphate-ion, and cations of the р-elements – aluminum (ІІІ) and bismuth (ІІІ), which have similarly low electronegativity.)

All classifications also include a group of cations, which form precipitates with НСІ: silver (I), mercury (I), and lead (ІІ). First two are the d-elements and lead is the р-element.

From cations of other groups can be picked out the ampholytic cations of the р– and d-elements, which have amphoterric properties and disposed bias of periodic table – zinc (ІІ), aluminium (ІІІ), teen (ІІ, IV), arsenic (ІІІ, V), chrome (ІІІ). They are found identical groups of analytic classifications. The ampholytes inherent small electronrgativity, high polarising properties and their compounds is capable to dependence on conditions to display oneself both base and acids.

Sameness disposition attitude in analytic groups has cations giving the complexes with ammonia. There are cations of the d-elements – nickel (ІІ), cobalt (ІІ), cadmium (ІІ), mercury (ІІ), copper (ІІ). High ability to complex compounds formation intrinsic explains by acceptor properties of unfilled in d-orbitals.

Types of Analytical Classifications of Cations

 

Anions Classification 

p-Elements and some d-elements (chrome, manganese) form anions. High ability to anions formation have the p-elements, disposed in right top quadrant of the periodic table. On the strength of that the р-elements have a variable oxidation degree, they are capable to form various acids and acids force increases with increasing of element oxidation degree. 

For oxidising-reducing properties the anions divide on anions-oxidisers with high oxidation degree (nitrate-anion), anions-reducers with lower oxidation degree (chlorides, bromides, iodides) and neutral anions, which not display nor reducing no oxidising properties (carbonate-, sulphate-, phosphate-anions). Oxidising-reducing properties of some anions can change (sulphite-, nitrite-anions) dependency on reaction conditions.

The analytical classification of anions is based on formation of insoluble in water precipitate with barium and silver salts. In accordance to this classification all anions divide on three groups:

–       the first group of anions forms precipitate with barium salt: sulphate-, sulphite, carbonate-, phosphate-, thiosulphate-, oxalate-, tetraborate-, iodate-, arsenate-, arsenite-, fluoride-, tartrate-, citrate-ions;

–       the second group of anions form insoluble in water and nitrate acid precipitates with silver salt: chlorides, bromides, iodides, thiocyanates, cyanides, benzoates;

–       the third group of anions not form insoluble compounds with barium and silver salts: nitrate-, nitrite-, acetate- bromate-, perchlorate-, salicylate-ions.

Majority of anions detect by fractional method, that’s why the group reagents use only for separation of anions groups, that exclude necessity to search in solution the anions of given group in case of negative reaction with group reagents.

 

 

 

 

 

 

 

 

 

 

 

Scheme of Fractional Analysis of Complex Mixtures

 

 

 

 

 

 

 

 

 

 

 

 

 

 

Scheme of Systematic Path of Complex Mixtures Analysis

 

 

 

 

 

 

 

 

Reagent B

 

 

 

Scheme of Analysis of Group j

 

 

 

 

 

 

 

                                                                

 

 

 

 

 

 

 

 

Law of mass action and its application to various types

The laws of mass action have universal importance in chemistry. The law of mass action is a reaction that states that the values of the equilibrium – constant expression K_c are constant for a particular reaction at a given temperature, whatever equilibrium concentrations are substitute.

aA + bB « cC + dD              K_c =

 

Getting the maximum amount of product from a reaction depends on the proper selection of reaction conditions. By changing these conditions, we can increase or decrease the yield of product. We might change the yield by:

1. Changing concentrations by removing products or adding reactants to the reaction vessel.

2. Changing the partial pressure of gaseous reactants and products.

3. Changing the temperature.

The equilibrium-constant expression is defined in terms of the balanced chemical equation. All analytical reactions, as a rule, run in solutions. For solutions we caot change the pressure. Sometimes we might heat or freeze the reaction vessel. But, in general, all reactions (processes) occur at isothermal condition. Therefore, we may use the equilibrium-constant expression in term of concentrations for both types of equilibrium:

I.   A homogeneous equilibrium is an equilibrium that involves reactants and products in a single phase (in solution, particle):

–      solutions of electrolytes;

–      protolytic equilibrium (hydrolysis, buffer systems);

–      complex compounds;

–      red-ox systems.

II. A heterogeneous equilibrium is an equilibrium involving reactants and products in more than one phase:

         a) liquid–solid systems:

–      saturated solution–precipitate (sediment);

–      colloids;

b) liquid–liquid system:

–      extraction.

In analytical chemistry law of mass action use for calculation of:

1) equilibrium ions concentration of dissociated weak electrolyte;

2) equilibrium concentration of reactants and products of chemical-analytical process;

3) equilibrium concentration of hydrogen and hydroxide ions and ionisation degree of electrolytes solutions;

4) equilibrium concentration of hydrogen and hydroxide ions in buffers and solutions of hydrolysed salts;

5) equilibrium concentration of cations and anions and solubility of electrolytes;

6) equilibrium concentration of ions of oxidant and reduce agent in red-ox reactions;

7) equilibrium concentration of ions in complex compounds solutions;

8) equilibrium-constants of various chemical processes.

 

Contemporary Theories of Electrolytes

A substance, that dissolves in water to give an electrically conducting solution is called an electrolyte. A substance, that dissolves in water to give nonconducting or very poorly conducting solutions is called a nonelectrolyte.

When electrolytes dissolve in water they produce ions, but they do so to varying extents. A strong electrolyte is an electrolyte that exists in solution almost entirely ions. A weak electrolyte is an electrolyte that dissolves in water to give equilibrium between a molecular substance and a small concentration of ions.

According to Svante Arrhenius concept:

Acid is any substance that, when dissolved in water, increase the concentration of hydrogen ion H⁺.

Base is any substance that, when dissolved in water, increase the concentration of hydroxide ion OH^–.

NaOH ® Na⁺ + OH^–

HCl ® H⁺ + Cl^–

The most short comings of Arrhenius concept:

1. Arrhenius concept (theory) does not explain the cause of dissociation of electrolytes on ions.

2. Arrhenius concept (theory) does not explain an acid or base property of organic substances, which not produced ions in water solution.

3. Arrhenius concept (theory) does not take account of interaction between solvent and dissolved substance. 

According to Johannes N. Brønsted and Thomas M. Lowry concept:

Acid is the species (molecule or ion) that donates a proton to another species in a proton-transfer reaction.

Base is the species (molecule or ion) that accepts a proton in a proton-transfer reaction.

HCl + NH₃ ® NH₄Cl

    acid          base

NH₃ + H₂O ® NH₄⁺ + OH^–

  base        acid    acid     base

A conjugate acid-base pair consists of two species in an acid-base equilibrium, one acid and one base, which differ by the gain or loss of a proton. The acid in such a pair is called the conjugate acid of the base, whereas the base is the conjugate base is the conjugate base of the acid.

The Brønsted-Lowry concept of acids and bases has greater scope than the Arrhenius concept:

1. A base is a species that accept protons; the OH^– ions is only one example of a base.

2. Acids and bases can be ions as well as molecular substances.

3. Acid-base reactions are not restricted to aqueous solutions.

4. Some species can act as either acids or bases, depending on what the other reactant is.

Such species, which can act either as an acid or a base (it can lose or gain a proton), called an amphiprotic species:

HCO₃^– + HF ® H₂CO₃ + F^–

base acid   acid       base

HCO₃^– + OH^– ® CO₃^2– + H₂O

      acid      base  base   acid

 

According to G. N. Lewis concept:

Lewis acid is a species that can form a covalent bond by accepting an electron pair from another species.

Lewis base is a species that can form a covalent bond by donating an electron pair to another species.

H⁺ + :NH₃ –® NH₄⁺

electron-pair                 electron-pair

acceptor              donor

Lewis acid           Lewis base

 

The Lewis and the Brønsted-Lowry concepts are simply different ways of looking at certain chemical reactions. The Lewis concept could be generalised to include many other reactions, as well as proton-transfer reactions.

Acids and bases are classified as strong or weak.

Strong acids are acids that ionise completely in water (that is, they react completely to give ions).

Weak acids are acids that are only partly ionised as the result of equilibrium reaction with water.

Strong bases are bases that are present in aqueous solution entirely as ions, one of which is OH^–.

Weak bases are bases that are only partly ionised as the result of equilibrium reactions with water.

The strongest acids have the weakest conjugate bases, and the strongest bases have the weakest conjugate acids. The terms strong and weak are used only in a comparative sense. The strengths of acids and bases are relative. In acid base interaction the water (or another solvent) exhibits a levelling effect on the strength of the strong acids.

Acid and base with water produce hydrogen ion or hydroxide ion (relatively) and its conjugated ions. The process is called electrolyte ionisation or electrolyte dissociation.

For the strong electrolyte (acid or base), which completely ionise in solution, the concentration of ions are determined by the stoichiometry of the reaction from the initial concentration of electrolyte:

[H⁺] » [HA]

[OH^–] » [BOH]

The weak electrolyte (acid and base) ionises or dissociates to a small extent in water (about 1 % or less, depending on concentration of electrolyte). For the weak electrolyte (acid or base) the concentration of ions in solution are determined from the acid ionisation (or dissociation) constant (K_a) or the base ionisation (or dissociation) constant (K_b), which is the equilibrium constant from the ionisation of a weak electrolyte.

K_a =            K_b =

Value of ionisation constants depends on:

1) nature of solvent,

2) nature of electrolyte,

3) temperature.

And not depends from electrolyte concentration.

[H⁺] =      (pH = ½pK_a – ½lgC_a)

[OH^–] =    (pOH = ½pK_b – ½lgC_b)

The degree of ionisation (a) of a weak electrolyte is the fraction of molecules that react with water to give ions. This also may be expressed as a percentage, giving the percent ionisation:

a_a =             [H⁺] = [A^–] = a×[HA] = a×C_a

a_b =                    [OH^–] = [B⁺] = a×[BOH] = a×C_b

For very small concentration of electrolyte a have very small value, and percent of ionisation can be shown approximately on Ostwald’s dilution rule:

 

K_c =             a =

The aqueous solutions of strong electrolytes and concentrated solutions of weak electrolytes not submit to classic law of mass action in full. Peter Debye and Erich Hückel were able to show that the properties of electrolyte solutions could be explained by assuming. The electrolyte is completely ionised in solution but that the activities, or effective concentrations, of the ions are less than their actual concentrations as a result of the electrical interaction of the ions in solution. The Debye-Hückel theory allows us to calculate these activities. When this is done, excellent agreement is obtained for dilute solutions:

a = C×g                 lgg = – A

a – active concentration of ions;

C – relative concentration of ions;

g – activity index;

A – value, calculate theoretically, depends from temperature, ion-dipole force etc.; for water solutions at t = 25 °C A = 0,509;

I – ionic strength;

z – charge of ion.

 

In solutions ion is a charged particle, surrounded ionic atmosphere with solvent ions. Ionic atmosphere parameters are definite by ionic strength:

 

I = ½     C_i – ions concentration (M)

 

Thus, we have seen that equilibrium-constant of electrolyte solutions change value accordance to activities of ions and depend from ionic strength of solution.

 

Protolytic balance in electrolytes solutions

Equilibrium in Solutions with Ions of the Same Kind

The effect of adding another solute to a solution of a weak acid or base called the common-ion effect. It is significant effect acid or base ionisation – that is, strong acids or bases and salts that contain an ion in common with the weak acid or base.

The common-ion effect is the shift in an ionic equilibrium caused by the addition of a solute that provides an ion that takes part in the equilibrium.

Thus strong acid provide an ion H⁺ common to an acid ionisation equilibrium. For example, ionisation of acetic acid:

CH₃COOH « CH₃COO^– + H⁺

In this solution is added a solution of HCl:

HCl « H⁺ + Cl^–

Because HCl is a strong acid, it provide H⁺ ion, which is present on the right side of equation for acetic acid ionisation. According to LeChateilier’s principle, the equilibrium composition should shift to the left. Thus the degree of ionisation of acetic acid is decreased by the addition of a strong acid. The repression of ionisation of acetic acid by HCl is an example of the common-ion effect.

 

Equilibrium in Solutions with Ions of Various Kinds

Addition a strong electrolyte to solution of weak acid or base increased the common concentration of ions in solution. Because the strong electrolyte ionises in solution completely, amount of all ions increases and, as a cause, change the activity of ions. This increasing of ionic strength of solution called salting effect.

According to LeChateilier’s principle, when shift in general concentration of reactants, should shift the equilibrium to the right. Thus the concentration of all ions and, respectively, the ionic strength of solution is increased. Increasing the concentration of all ions caused the increasing the H⁺ ion concentration. Thus, the value pH is changes, becomes more.

 

Hydrolysis

One of the successes of Brønsted-lowry concept is its explanation of the acid-basic properties of salt solutions. The reaction of ions with water called hydrolysis. The hydrolysis reaction produces either hydrogen-ion or hydroxide-ion. It is a typical acid-base reaction, which change protolytic (protons) equilibrium.

Such ions may produce H⁺ or OH^– ions, so they may give acidic or basic solutions.

 

Hydrolysis of various types salts.

I.   Salts, formed by strong acid and strong bases not hydrolyse. They completely ionise (dissociate):

KCl ® K⁺ + Cl^–

II. Salts, formed by strong acid and weak base (NH₄Cl):

Cat⁺ + H₂O « CatOH + H⁺

Solutions of these salts are acid – pH < 7

III. Salts, formed by weak acid and strong base (CH₃COONa):

An^– + H₂O « HAn + OH^–

Solutions of these salts are base – pH > 7

IV. Salts, formed by weak acid and weak base (NH₄CN):

Cat⁺ + An^– + H₂O « CatOH + HAn

[H⁺] » [OH^–]

In general, a solution of this salt is acidic, base or neutral by comparing the hydrolysis constants of the two ions from the salt. The hydrolysis constant of the cation will be its K_a, and that for the anion it’s K_b.

The solution will be acidic if K_a (cation) > K_b (anion).

The solution will be neutral if K_a (cation) = K_b (anion).

The solution will be basic if K_a (cation) < K_b (anion).

The concentration of ions in solution is determined from the constant of hydrolysis of salt. From this constant we may determine concentration of hydrogen and hydroxide ions and, accordingly, the pH of solution. And we may calculate the degree of hydrolysis.

The degree of hydrolysis (a_h) of a salt is the fraction of molecules that react with water to give ions.

 

Using hydrolysis in analysis

1.   Detecting some ions. The salts of this ions during hydrolysis forms insoluble compounds. This phenomenon is thypical for salts of metalloids or salts of very weak bases or acids:

SbCl₃ + H₂O ® SbOCl¯ + HCl

Bi(NO₃)₃ + H₂O ® BiOH(NO₃)₂¯ + HNO₃

Some salts, formed by weak bases and weak acids, hydrolyse completely with producing another chemical compounds:

2CrCl₃ + 3(NH₄)₂S ® Cr₂S₃ + 6NH₄Cl

Cr₂S₃ + 6H₂O ® 2Cr(OH)₃¯ + 3H₂S­

2. Separation of ions. For example, Al⁺³ and Cr⁺³:

CrCl₃ + 4KOH ® KCrO₂ + 3KCl + 2H₂O

AlCl₃ + 4KOH ® KAlO₂ + 3KCl + 2H₂O

    t°

KCrO₂ + 2H₂O ® Cr(OH)₃¯ + KOH

KAlO₂ not hydrolyses

3. Changing the concentration of hydrogen or hydroxide ions:

         2baCl₂ + K₂Cr₂O₇ + H₂O « 2BaCrO₄¯ + 2KCl + 2HCl

Formed strong acid HCl may dissolves precipitate BaCrO4. In presence of CH3COONa:

2CH₃COONa + 2HCl ® 2CH₃COOH + 2NaCl

Formed weak acid CH₃COOH, concentration of H⁺ ions decreases, and precipitate not dissolves.

2)      K₃AlO₃ + 3H₂O « Al(OH)₃¯ + 3KOH

Formed strong base KOH may dissolves precipitate Al(OH)₃. In presence of NH₄Cl:

 

3NH₄Cl + 3KOH ® 3NH₄OH + 3KCl

Formed weak base NH₄OH, concentration of OH^– ions decreases, and precipitate not dissolves.

 

Repressing and intensification of hydrolysis

Sometime hydrolysis prevents to run an analytical reaction. In this case we may to repress or to intensify the hydrolysis. As any chemical equilibrium process, the hydrolysis submits to LeChateilier’s principle. Accordance to our purposes we may do next:

1.   Add to solution the salt of another hydrolysed electrolyte (salt, acid or base).

2.   Change the salt concentration.

3.   Heat or freeze the solution.

 

Equations for Hydrolysis Parameters Calculation

 

 

 

Buffers

A buffer is a solution characterised by the ability to resist changes in pH when limited amounts of acid or base are added to it. Buffer contains either a weak acid and its conjugate base or a weak base and its conjugate acid.

Suppose a buffer contains approximately equal molar amount of weak acid HA and its conjugate base A^–. When a strong acid is added to the buffer, it supplies hydrogen ions that react with the base A^–:

H⁺ + A^– ® HA

On the other hand, when a strong base is added to the buffer, it supplies hydroxide ions. Then ions react with the acid HA:

OH^– + HA ® H₂O + A^–

Thus a buffer solution resists changes in pH through its ability to combine with both H⁺ and OH^– ions.

There are three types of buffers which distinguish its components:

 

I.    Buffer contains weak acid and its salt (pH of buffer < 7):

HCOOH + HCOONa;            CH₃COOH + CH3COONa.

II.               Buffer contains weak base and its salt (pH of buffer > 7):

H₃BO₃ + Na₂B₄O₇;                  NH₄OH + NH₄Cl

III. Buffer contains salts of polyprotic acids (pH of buffer » 7):

Na₂HPO₄ + NaH₂PO₄             Na₂CO₃ + NaHCO₃

Two important characteristics of a buffer are the pH and the buffer capacity.

 

The buffer capacity – is the amount of acid or base the buffer can react with before giving a significant pH change.

Buffer capacity depends on the amount of acid and conjugated base in the solution. The ratio of amounts of acid and conjugated base is also important. Unless this ratio is approximately 1 (between 1:10 and 10:1), the buffer capacity will be too low to be useful.

 

 

The other important characteristic of a buffer is its pH. Buffer always must be prepared from a conjugated acid-base pair in which the acid ionisation constant is approximately equal to the desired H⁺ ion concentration.

The Henderson-Hasselbalch equation relates the pH of a buffer for different concentrations of conjugate acid and base:

pH = pK_a + lg [base]/[acid]

By substituting the value of pK_a for the conjugate acid and the ratio [base]/[acid], we obtain the pH of the buffer.

 

Equations for Calculation [H⁺] and pH of Buffers

 

It must be remembered, however, that pH is not entirely established by ratio of conjugate base to conjugate acid bat can be affected by concentration. For typical buffers (i.e. concentration less than 0.1 M or with K values of 10^-3 or less) the Henderson-Hasselbalch equation can be used.

 

Using law mass action due to equations in homogenous systems

Ampholytes

The term amphoteric refers to a substance that has both acidic and basic properties. For example, aluminium oxide dissolves in acids to produce the cation Al⁺³, as expected for a metal oxide:

Al₂O₃ + 6HCl ® 2AlCl₃ + 3H₂O

But the oxide also dissolves in strong base:

Al₂O₃ + 3H₂O + 2KOH ® 2K[Al(OH)₄]

In this case the aluminate anion, Al(OH)₄^–, is formed.

In more common sense, accordance to Brønsted-Lawry concept of electrolytes, amphoteric substances are concluded to class (type) of species called ampholytes.

Ampholytes are species that may to accept and to donate the protons. They are both neutral and charged particles (substances).

There are three types of ampholytes:

I.   Ampholyte that contain hydrogen ion (HCO₃^–, H₂PO₄^–, HSO₃^–):

Dissociation equations:          HX^– « H⁺ + X^–2

HX^– + H⁺ « H₂X

H₂O « H⁺ + OH^–

Matter balance:             [H⁺] = [X^­2] + [OH^–] – [H₂X]

II. Ampholyte that contain hydroxide ion [Al(OH)₆⁺³, Ni(OH)⁺]:

Dissociation equations:          MOH⁺ « M⁺² + OH^–

                                               MOH⁺ + OH^– « M(OH)₂

                                               H₂O « H⁺ + OH^–

Matter balance:             [OH^–] = [M⁺²] + [H⁺] – [M(OH)₂]

III. Ampholyte – salt, containing both protons donor and acceptor (CH₃COONH₄, NH₄CN):

Dissociation equations:          MH⁺ « H⁺ + M

                                               X^– + H⁺ « HX

                                               H₂O « H⁺ + OH^–

Matter balance:             [H⁺] = [M] + [OH^–] – [HX]

 

Equations for calculation ampholytes solutions parameters

 

 

Using ampholytes in analysis:

1. Dissolving insoluble hydroxides:

Al(OH)₃ + 3KOH ® K₃AlO₃ + 3H₂O

2. Changing degree ionisation of cation-ampholyte:

2Cr(OH)₃ + 3H₂O₂ + 4KOH ® 2K₂CrO₄ + 8H₂O

K_HCrO2 = 9×10^–17                      K_H2CrO4 = 1,8×10^–1

3. Separation of cation (in insoluble hydroxides):

To mixture of sediments Fe(OH)₃, Al(OH)₃, Mn(OH)₂ add mix NH₄OH + NH₄Cl – to solution pass MnCl₂. To sediment Fe(OH)₃, Al(OH)₃ add KOH – to solution pass K₃AlO₃.