FIRST GROUP OF CATIONS (ACID-BASIC CLASSIFICATION OF CATIONS). THEORY OF STRONG ELECTROLYTES.

 

Analytical chemistry occupies one of major places in trade education of pharmacist. As science she consists of qualitative and quantitative analysis. A qualitative analysis can be executed iumber of different ways. Normatively-analytical document on substances and medications contain an obligatory test “Identification”. He foresees the conduction of qualitative reactions on cations, anions or functional groups which are in the molecule of organic matter. A study of qualitative reactions, study of technique of their conduction is a primary stage of education of future pharmacist.

The cations of first analytical group – K⁺, Na⁺, NH₄⁺ enter in the composition of medicinal preparations in the different prepared forms. Pharmacist–analyst must be able to execute an analysis for confirmation of qualitative composition of preparation.

 

1.                The regulations of the work and prevention of accidents in an analytical laboratory.

2.                Basic concepts of chemical analysis: specific and selective reactions, chemical reagent and demands to him, sensitivity of chemical reaction (determined minimum, minimum concentration, maximum dilution).

3.                Methods of implementation of analytical reactions are after the amount of the explored substance.

4.                Technique of execution of major operations of semimicroanalysis: heating, evaporation, precipitation, centrifugation, filtration, washing of precipitate.

5.                Qualitative analysis: fractional and systematic.

6.                Types of classifications of cations.

7.                Acid-basic classification. Qualitative reactions of cations of first analytical group.

8.                Basic concepts of theory of strong electrolytes: ionic force of solution, activity, concentration, coefficient of activity of ion and middle coefficient of activity.

10. A law of preservation the masses, his application in analytical chemistry.

11. Constants of equilibrium: thermodynamics, concentrations, conditional.

 

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

 

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

        

The natural sciences begin with observation, and this usually involves numerical measurements of quantities such as length, volume, density, and temperature. Most of these quantities have units of some kind associated with them, and these units must be retained when you use them in calculations. Measuring units can be defined in terms of a very small number of fundamental ones that, through “dimensional analysis”, provide insight into their derivation and meaning, and must be understood when converting between different unit systems.

The SI base units

         In principle, any physical quantity can be expressed in terms of only seven base units. Each base unit is defined by a standard which is described in the NIST Web site.

         A few special points about some of these units are worth noting:

·                     The base unit of mass is unique in that a decimal prefix (see below) is built-in to it; that is, it is not the gram, as you might expect.

·                     The base unit of time is the only one that is not metric. Numerous attempts to make it so have never garnered any success; we are still stuck with the 24:60:60 system that we inherited from ancient times. (The ancient Egyptians of around 1500 BC invented the 12-hour day, and the 60:60 part is a remnant of the base-60 system that the Sumerians used for their astronomical calculations around 100 BCE.)

·                     Of special interest to Chemistry is the mole, the base unit for expressing the quantity of matter. Although the number is not explicitly mentioned in the official definition, chemists define the mole as Avogadro’s number (approximately 6.0210²³) of anything.

The SI decimal prefixes

         Owing to the wide range of values that quantities can have, it has long been the practice to employ prefixes such as milli and mega to indicate decimal fractions and multiples of metric units. As part of the SI standard, this system has been extended and formalized.

Units outside the SI

         There is a category of units that are “honorary” members of the SI in the sense that it is acceptable to use them along with the base units defined above.

         These include such mundane units as the hour, minute, and degree (of angle), etc., but the three shown here are of particular interest to chemistry, and you will need to know them.

 

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.

Formation of a Precipitate

         The formation of a precipitate may be one of the most common signs of a chemical reaction taking place. A precipitate is defined to be a solid that forms inside of a solution or another solid. Precipitates should not be confused with suspensions, which are solutions that are homogeneous fluids with particles floating about in them. For instance, when a soluble carbonate reacts with Barium, a Barium Carbonate precipitate can be observed.

Test Tube

For further information, please refer to Classification of Matter.

Formation of Bubbles

         The formation of bubbles, or rather a gas, is another indicator of a chemical reaction taking place. When bubbles form, a temperature change could also be taking place. Temperature change and formation of bubbles often occur together. For example, in the following image, one can see a gas spewing. This is the formation of a gas.

Gas Formation

         However, most reactions are much more subtle. For instance, if the following reaction occurs, one may notice Carbon Dioxide bubbles forming. If there is enough Hydrochloric Acid, bubbles are visible. If there isn’t, one can’t readily notice the change.

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.

 

To the major operations of semi microanalysis belongs: heating, evaporation, sedimentation, centrifugation, filtration, washing of sediment.

Heating. As known, a lot of reactions take place at heating. To heat of solution which is in little test tubes, directly on the opened flame it is impossible as a liquid practically will always be thrown out a pair. They are heated on a water-bath. The term ‘water bath’ means a bath of boiling water, unless water at some other temperature is indicated in the text. An alternative form of heating may be employed providing that the required temperature is approximately maintained but not exceeded.

Evaporation. The evaporation of solutions for the concentration or complete delete of water conducts on the water-bath in the cups of porcelains.

Sedimentation. In the qualitative analysis the sedimentation often conduct in centrifugation’s test tubes. In a test tube place the necessary amount of drops of the explored solution and, if it is needed heat on a water-bath, whereupon add the necessary amount of drops of reagent a pipette. In connection with that in little test tubes which apply in semi microanalysis, solutions are mixed badly, it is necessary immediately after their uniting to mix maintenance of test tube a glass stick. It is yet better to mix after addition of every drop a reagent. It is instrumental in more slow growth of crystals of sediment, which are then formed large, better move away from solution at centrifugation, and having a less surface, less adsorb related matters from solution.

If sediment which fell out can form colloid solution, for the best coagulation usefully after addition a reagent a few minutes to heat maintenance of test tube on a water-bath.

 

FIRST ANALYTICAL GROUP OF CATIONS

(K⁺, Na⁺, NH)

Almost all salts to Potassium, Sodium and ammonium well water-solubles. Therefore cations of K⁺, Na⁺, NH₄⁺ not have a general reagent which would be capable synchronously to obtain of precipitate with them as insoluble in water compounds. The first analytical group of cations differs from cations of other groups.

All salts of cations of the first group have not a colour, and their compounds can have colour, only when they will form compounds together of colour anions (for example, CrO₄^2–, Cr₂O₇^2–, MnO₄^–, and others). The salts of Potassium and of Sodium are persistent during of heating. The ammonium salts are not persistent during of heating, especially these salts which have obtained together with volatile acids (HCl, HNO₃, H₂CO₃ and other).

 

Characteristic reactions of ions of K⁺

Sodium hexanitrocobaltate (III) of Na₃[Co(NO₂)₆] (pharmacopeia’s reaction). This complex compound with the ions of K⁺ ieutral or acetic-acid medium forms yellow crystalline precipitate of double salt of K₂Na[Co(NO₂)₆]:

2K⁺ + Na⁺ + [Co(NO₂)₆]^3- = K₂Na[Co(NO₂)₆]¯.

Alkali metals hydroxides interfere of this reaction, because lay out a reagent, as a result darkly-brown precipitate of Co(OH)₃ is selected:

[Co(NO₂)₆]^3– + 3OH^– = Co(OH)₃¯ + 6NO₂^–.

If there are strong acids, complex anion is also laid out:

[Co(NO₂)₆]^3– + 6H⁺ = Co³⁺ + 6HNO₂.

It should be remembered that Na₃[Co(NO₂)₆] is not stability, his colour can change on rose (color of ions of Co²⁺), and such reagent caot apply for the detection of ions of K⁺.

 

The ions of NH₄⁺ interfere with the exposure of ions of K⁺, as from Na₃[Co(NO₂)₆] form sediment, similar to in color sediment which appears at presence of to Potassium.

Implementation of reaction. To the drop of the explored solution (pH 5-7) add 2-3 drops solution of Na₃[Co(NO₂)₆]. If a reaction of solution is acidic, that it is necessary to add Sodium acetate for linkage of ions of H⁺. If there are ions of K⁺, yellow precipitate appears.

Sodium hydrotartratic of NaHC₄H₄O₆ (pharmacopeia’s reaction). This compound with ions Potassium hydrotartratic forms in a neutral medium white crystalline precipitate Potassium hydrotartratic:

                                           K⁺ + HC₄H₄O₆^– ® KHC₄H₄O₆¯.       

         Precipitate is soluble in mineral acids and alkalis. Solubility of precipitate increases at heating. Precipitate KHC₄H₄O₆ forms on the rubbing of wall-side of a test tube a glass stick and cooling (refregeration).

This reaction conduct with tartratic acid in a presence Sodium acetate (Pharmacopeia of Europe).

Implementation of reaction. To 3-4 drops of the explored solution add 3-4 drops solution of NaHC₄H₄O₆ and rub the wall-side of test tube a glass stick.  If a reaction of solution is acidic, that it is necessary to add Sodium acetate for linkage of ions of H⁺. If there are ions of K⁺, white precipitate forms.

Microcristaloscopic Reaction. Ions of K⁺ forms the black cubic crystals with the reagent of Na₂PbCu(NO₂)₆ :

2K⁺ + PbCu(NO₂)₆^2– ® K₂PbCu(NO₂)_6.

A reaction must run in a neutral medium; the ions of NH₄⁺ interfere with its conducting.

Implementation of reaction. Drop of the explored solution put on the glass and evaporate on water-bath. After that the dry remain is cooling, adds a drop of reagent. If there are ions of K⁺, that black or brown cubic crystals appear (fig. 1).

Flame test Potassium (pharmacopeia’s reaction). Salts of Potassium paint flame in a fleeting pale-violet color.

Implementation of reaction. A platinum wire moistens with solution HCl and heat in a flame. A clean platinum wire moistens with solution of salt K⁺and heat in a flame. If there are ions to Potassium, flame paints in a fleeting pale-violet color.

Characteristic reactions of ions of Na⁺

Potassium hexahydroxostibatе (V) K[Sb(OH)₆] (pharmacopeia’s reaction). This complex compound with the concentrated solutions of salts of Sodium in a neutral or weak basic medium white precipitate forms:

Na⁺ + [Sb(OH)₆]^–  ® Na[Sb(OH)₆] ¯.

         Precipitate Na[Sb(OH)₆] forms on the rubbing of wall-side of a test tube a glass stick and cooling (refregeration). Precipitate does not appear in a strongly basic medium. This reagent is not stability in a acidic medium:

[Sb(OH)₆]^– + H+ ® H[Sb(OH)₆]

H[Sb(OH)₆] ® HSbO₃¯ + 3H₂O.

Precipitate does not form in dilute solutions.

Precipitation is slowed in presence of nitrate ions. The ions of NH₄⁺, Mg²⁺, Li⁺ interfere.

Implementation of reaction. To 3-4 drops of the explored solution add 3-4 drops solution of reagent, the internal wall-side of test tube rubs a glass stick. If a reaction of solution is acidic, that it is necessary to add Potassium acetate for linkage of ions of H⁺. If there are ions of Na+, white precipitate appears.

Microcristaloscopic Reaction. Ions of Na⁺ forms the octahedral pale-yellow crystals with the reagent Zn(UO₂)₃(CH₃COO)₈:

Na⁺ + Zn(UO₂)₃×(CH₃COO)₈ + CH₃COO^– + 9H₂O ® NaZn(UO₂)₃×(CH₃COO)₉×9H₂O¯.

A reaction must run in a neutral or acetic medium; the ions of Ag+, Hg₂²⁺, Sb^(III), PO₄^3–, AsO₄^3– interfere with its conducting.

Implementation of reaction. Drop of the explored solution put on the glass and evaporate on water-bath. After that the dry remain is cooling, adds a drop of reagent. If there are ions of Na⁺, that octahedral pale-yellow crystals appear (fig. 2).

 

 

Flame test Sodium (pharmacopeia’s reaction). Salts of Sodium paint flame in a persistent yellow color.

Implementation of reaction. A platinum wire moistens with solution HCl and heat in a flame. A clean platinum wire moistens with solution of salt Sodium and heat in a flame. If there are ions to Sodium, flame paints in a persistent yellow color.

 

 

Metoxyphenilacetic acid (pharmacopeia’s reaction). This compound with ions Sodium forms white precipitate in presence tetramethylammonium hydroxide:

Implementation of reaction. To 5-6 drops of the explored solution add 10-12 drops of reagent Metoxyphenilacetic acid in the solution tetramethylammonium hydroxide and cool in icy water.

 If there are ions of Na⁺, white precipitate forms.

 

Characteristic reactions of ions of NH₄⁺

Solutions of alkalis (NaOH or KOH) (pharmacopeia’s reaction) during heating with solutions of salts an ammonia is selected an ammonia:

NH₄⁺ + OH^– = NH₃­ +H₂O.

Implementation of reaction. 3-5 drops of the explored solution place in test tube and add some drops of NaOH or KOH. Phenolphtalein’s paper moisten a water and keep its above test tube. If there are ammonia ions, that  Phenolphtalein’s paper will be rose-colour.

Potassium tetraiodomercurate (II) K₂[HgI4] in KOH (reagent of Nessler) This compound with ions NH₄⁺ forms red-brown precipitate in presence Potassium hydroxide:

NH₄⁺ + 2[HgI₄]^2– + 4OH^– = [NH₂Hg₂O]I¯ + 7I^– + 3H₂O.

The ions of Fe3+, Bi3+, Cu2+, Cd2+, Ag+, Pb2+, As (V) interfere with its conducting.

This reaction is sensible and it application for determination of “tracks” quantity of the ammonium or ammonia. Then solution is painted in yellow.

Implementation of reaction. To 1-2 drops of the explored solution add some drops of Nessler’s reagent. If there are ions of NH₄⁺, red-brown precipitate forms or solution appears yellow colour.

Sodium hexanitrocobaltate (III) of Na₃[Co(NO₂)₆] (pharmacopeia’s reaction). This complex compound with the ions of NH₄⁺ ieutral or acetic-acid medium forms yellow crystalline precipitate of double salt of (NH₄)₂Na[Co(NO₂)₆]:

2 NH4+  + Na+ + [Co(NO2)₆]^3- = (NH4)_2Na[Co(NO2)₆]¯

Implementation of reaction. To the drop of the explored solution (pH 5-7) add 2-3 drops solution of Na₃[Co(NO₂)₆]. If a reaction of solution is acidic, that it is necessary to add Sodium acetate for linkage of ions of H⁺. If there are ions of NH₄⁺, yellow precipitate appears.

 

Systematic analysis of cation’s mixtures of the first analytical group

1.     Determination ions of NH₄⁺ with NaOH or Nessler’s reagent.

2.     If there aren’t NH₄⁺-ions, that the ions of K⁺ and Na⁺ determinate in two separate portions of the explored solution.

3.     If there are NH₄⁺-ions, that them extract (heat with NaOH to dry remain a few minutes). Then it cool and dissolve in some drops of distilled water (verification is with the reagent of Nessler). If there aren’t NH₄⁺-ions, that the ions of K⁺ determinate in this solution.

4.     Use 3 for determination of Na⁺, but KOH necessary to add for extraction NH₄⁺-ions.

 Analytical chemistry is the study of the separation, identification, and quantification of the chemical components of natural and artificial materials. Qualitative analysis gives an indication of the identity of the chemical species in the sample and quantitative analysis determines the amount of one or more of these components. The separation of components is often performed prior to analysis.

Analytical methods can be separated into classical and instrumental.^[2] Classical methods (also known as wet chemistry methods) use separations such as precipitation, extraction, and distillation and qualitative analysis by color, odor, or melting point. Quantitative analysis is achieved by measurement of weight or volume. Instrumental methods use an apparatus to measure physical quantities of the analyte such as light absorption, fluorescence, or conductivity. The separation of materials is accomplished using chromatography, electrophoresis or Field Flow Fractionation methods.

Analytical chemistry is also focused on improvements in experimental design, chemometrics, and the creation of new measurement tools to provide better chemical information. Analytical chemistry has applications in forensics, bioanalysis, clinical analysis, environmental analysis, and materials analysis.

History

Analytical chemistry has been important since the early days of chemistry, providing methods for determining which elements and chemicals are present in the object in question. During this period significant analytical contributions to chemistry include the development of systematic elemental analysis by Justus von Liebig and systematized organic analysis based on the specific reactions of functional groups.

The first instrumental analysis was flame emissive spectrometry developed by Robert Bunsen and Gustav Kirchhoff who discovered rubidium (Rb) and caesium (Cs) in 1860.

Gustav Kirchhoff (left) and Robert Bunsen (right)

Most of the major developments in analytical chemistry take place after 1900. During this period instrumental analysis becomes progressively dominant in the field. In particular many of the basic spectroscopic and spectrometric techniques were discovered in the early 20th century and refined in the late 20th century.

The separation sciences follow a similar time line of development and also become increasingly transformed into high performance instruments. In the 1970s many of these techniques began to be used together to achieve a complete characterization of samples.

Starting in approximately the 1970s into the present day analytical chemistry has progressively become more inclusive of biological questions (bioanalytical chemistry), whereas it had previously been largely focused on inorganic or small organic molecules. Lasers have been increasingly used in chemistry as probes and even to start and influence a wide variety of reactions. The late 20th century also saw an expansion of the application of analytical chemistry from somewhat academic chemical questions to forensic, environmental, industrial and medical questions, such as in histology.

Modern analytical chemistry is dominated by instrumental analysis. Many analytical chemists focus on a single type of instrument. Academics tend to either focus oew applications and discoveries or oew methods of analysis. The discovery of a chemical present in blood that increases the risk of cancer would be a discovery that an analytical chemist might be involved in. An effort to develop a new method might involve the use of a tunable laser to increase the specificity and sensitivity of a spectrometric method. Many methods, once developed, are kept purposely static so that data can be compared over long periods of time. This is particularly true in industrial quality assurance (QA), forensic and environmental applications. Analytical chemistry plays an increasingly important role in the pharmaceutical industry where, aside from QA, it is used in discovery of new drug candidates and in clinical applications where understanding the interactions between the drug and the patient are critical.

Classical methods

 

Although modern analytical chemistry is dominated by sophisticated instrumentation, the roots of analytical chemistry and some of the principles used in modern instruments are from traditional techniques many of which are still used today. These techniques also tend to form the backbone of most undergraduate analytical chemistry educational labs.

Qualitative analysis

A qualitative analysis determines the presence or absence of a particular compound, but not the mass or concentration. That is, if it is not related to quantity.

Chemical tests

For more details on this topic, see Chemical test.

There are numerous qualitative chemical tests, for example, the acid test for gold and the Kastle-Meyer test for the presence of blood.

 

Types of Chemical Reactions

Precipitation, or double-replacement reaction

         A reaction that occurs when aqueous solutions of anions (negatively charged ions) and cations (positively charged ions) combine to form a compound that is insoluble is known as precipitation. The insoluble solid is called the precipitate, and the remaining liquid is called the supernate. See Figure2.1

         Example. An example of a precipitation reaction is the reaction between silver nitrate and sodium iodide. This reaction is represented by the chemical equation :

AgNO₃ (aq)+ NaI (aq) → AgI (s) + NaNO₃ (aq)

         Since all of the above species are in aqueous solutions, they are written as ions, in the form:

Ag⁺ +NO₃^– (aq)+ Na⁺ (aq) + I^– (aq) → AgI (s) + Na⁺ (aq) + NO₃^– (aq)

         Ions that appear on both sides of the equation are called spectator ions. These ions do not affect the reaction and are removed from both sides of the equation to reveal the net ionic equation, as written below:

Ag⁺ (aq) + I^– (aq) → AgI (s)

In this reaction, the solid, AgI, is known as the precipitate. The formation of a precipitate is one of the many indicators that a chemical reaction has taken place.

         Real life example: The white precipitate formed by acid rain on a marble statue. Chemical reaction: CaCO₃(aq)+H₂SO₄(aq)=CaSO₄(s)+H₂O(l)+CO₂(g)

 

Acid-base, or neutralization reaction

         A neutralization reaction occurs when an acid and base are mixed together. An acid is a substance that produces H⁺ ions in solution, whereas a base is a substance that that produces OH^– ions in solution. A typical acid-base reaction will produce an ionic compound called a salt and water. A typical acid-base reaction is the reaction between hydrochloric acid and sodium hydroxide. This reaction is represented by the equation:

HCl (aq) + NaOH (aq) → NaCl (aq)+ H₂O (l)

         In this reaction, HCl is the acid, NaOH is the base, and NaCl is the salt.

         Real life example: Baking soda reacts with vinegar is a neutralization reaction.

 

Oxidation-reduction, or redox reaction

         A redox reaction occurs when the oxidatioumber of atoms involved in the reaction are changed. Oxidation is the process by which an atom’s oxidation number is increased, and reduction is the process by which an atom’s oxidation number is decreased. If the oxidation states of any elements in a reaction change, the reaction is an oxidation-reduction reaction. An atom that undergoes oxidation is called the reducing agent, and the atom that undergoes reduction is called the oxidizing agent. An example of a redox reaction is the reaction between hydrogen gas and fluorine gas:

H₂ (g) + F₂ (g)→2HF (g)

         In this reaction, hydrogen is oxidized from an oxidation state of 0 to +1, and is thus the reducing agent. Fluorine is reduced from 0 to -1, and is thus the oxidizing agent.

 

 

In this redox reaction, a H₂ molecule donate electrons to F₂ resulting in two HF molecules

Real life example: The cut surface of an apple turns brownish after exposed to the air for a while.

 

Combustion reaction

         A combustion reaction is a type of redox reaction during which a fuel reacts with an oxidizing agent, resulting in the release of energy as heat. Such reactions are exothermic, meaning that energy is given off during the reaction. An endothermic reaction is one which aborbs heat. A typical combustion reaction has a hydrocarbon as the fuel source, and oxygen gas as the oxidizing agent. The products in such a reaction would be CO₂ and H₂O.

C_xH_yO_z+O₂→ CO₂+H₂O (unbalanced)

Such a reaction would be the combustion of glucose in the following (unbalanced) equation:

C₆H₁₂O₆ (s) + O₂ (g) → CO₂ (g) +H₂O (g)

Real life example: explosion; buring

 

Synthesis reaction

A synthesis reaction occurs when one or more compounds combines to form a complex compound. The simplest eqaution of synthesis reaction is illustrated below.

         An example of such a reaction is the reaction of silver with oxygen gas to form silver oxide:

2Ag (s) +O₂(g)→ 2AgO (s)

         Real life example: Hydrogen gas is burned in air, reacts with oxygen gas to form water.          Chemical equation: 2H₂(g)+O₂(g)→ 2H₂O(l)

 

Decomposition reaction

         A decomposition reaction is the opposite of a synthesis reaction. During a decomposition reaction, a more complex compound breaks down into multiple simpler compounds. A classic example of this type of reaction is the decomposition of hydrogen peroxide into oxygen and hydrogen gas:

2H₂O (l) → 2H₂ (g) + O₂ (g)

The molecule AB is decomposing into A and B

         Real life examples: Electrolysis of water; Carbonic acid in soda

 

Single replacement reaction

         A type of oxidation-reduction reaction in which an element in a compound is replaced another element.

         An example of such a reaction is:

Cu (s) + AgNO₃ (aq) → Ag(s) + Cu(NO₃)₂ (aq)

 

Flame test

For more details on this topic, see Flame test.

Flame Test Introduction

The flame test is used to visually determine the identity of an unknown metal or metalloid ion based on the characteristic color the salt turns the flame of a bunsen burner. The heat of the flame excites the metals ions, causing them to emit visible light. The characteristic emission spectra can be used to differentiate between some elements.

Classic Wire Loop Method

         First, you need a clean wire loop. Platinum or nickel-chromium loops are most common. They may be cleaned by dipping in hydrochloric or nitric acid, followed by rinsing with distilled or deionized water. Test the cleanliness of the loop by inserting it into a gas flame. If a burst of color is produced, the loop is not sufficiently clean. The loop must be cleaned between tests.

         The clean loop is dipped in either a powder or solution of an ionic (metal) salt. The loop with sample is placed in the clear or blue part of the flame and the resulting color is observed.

 

The presence of copper in this qualitative analysis is indicated by the bluish-green color of the flame.

 

Wooden Splint or Cotton Swab Method

Wooden splints or cotton swabs offer an inexpensive alternative to wire loops. To use wooden splints, soak them overnight in distilled water. Pour out the water and rinse the splints with clean water, being careful to avoid contaminating the water with sodium (as from sweat on your hands). Take a damp splint or cotton swab that has been moistened in water, dip it in the sample to be tested, and wave the splint or swab through the flame. Do not hold the sample in the flame as this would cause the splint or swab to ignite. Use a new splint or swab for each test.

How to Interpret the Results

The sample is identified by comparing the observed flame color against known values from a table or chart.

Red
Carmine to Magenta: Lithium compounds. Masked by barium or sodium.

Scarlet or Crimson: Strontium compounds. Masked by barium.

Red: Rubidium (unfiltered flame)

Yellow-Red: Calcium compounds. Masked by barium.

Yellow
Gold: Iron

Intense Yellow: Sodium compounds, even in trace amounts. A yellow flame is not indicative of sodium unless it persists and is not intensified by addition of 1% NaCl to the dry compound.

White
Bright White: Magnesium
White-Green: Zinc

Green
Emerald: Copper compounds, other than halides. Thallium.
Bright Green: Boron
Blue-Green: Phosphates, when moistened with H₂SO₄ or B₂O₃.
Faint Green: Antimony and NH₄ compounds.
Yellow-Green: Barium, manganese(II), molybdenum.

Blue
Azure: Lead, selenium, bismuth, cesium, copper(I), CuCl₂ and other copper compounds moistened with hydrochloric acid, indium, lead.
Light Blue: Arsenic and come of its compounds.
Greenish Blue: CuBr₂, antimony

Purple
Violet: Potassium compounds other than borates, phosphates, and silicates. Masked by sodium or lithium.
Lilac to Purple-Red: Potassium, rubidium, and/or cesium in the presence of sodium when viewed through a blue glass.

Limitations of the Flame Test

·                     The test cannot detect low concentrations of most ions.

·                     The brightness of the signal varies from one sample to another. For example, the yellow emission from sodium is much brighter than the red emission from the same amount of lithium.

·                     Impurities or contaminants affect the test results. Sodium, in particular, is present in most compounds and will color the flame. Sometimes a blue glass is used to filter out the yellow of sodium.

·                     The test cannot differentiate between all elements. Several metals produce the same flame color. Some compounds do not change the color of the flame at all.

Primary Reference: Lange’s Handbook of Chemistry, 8th Edition, Handbook Publishers Inc., 1952.

Flame Test Colors

Symbol	Element	Color
As	Arsenic	Blue
B	Boron	Bright green
Ba	Barium	Pale/Yellowish Green
Ca	Calcium	Orange to red
Cs	Cesium	Blue
Cu(I	Copper(I)	Blue
Cu(II)	Copper(II) non-halide	Green
Cu(II)	Copper(II) halide	Blue-green
Fe	Iron	Gold
In	Indium	Blue
K	Potassium	Lilac to red
Li	Lithium	Magenta to carmine
Mg	Magnesium	Bright white
Mn(II)	Manganese(II)	Yellowish green
Mo	Molybdenum	Yellowish green
Na	Sodium	Intense yellow
P	Phosphorus	Pale bluish green
Pb	Lead	Blue
Rb	Rubidium	Red to purple-red
Sb	Antimony	Pale green
Se	Selenium	Azure blue
Sr	Strontium	Crimson
Te	Tellurium	Pale green
Tl	Thallium	Pure green
Zn	Zinc	Bluish green to whitish green

 

Inorganic qualitative analysis generally refers to a systematic scheme to confirm the presence of certain, usually aqueous, ions or elements by performing a series of reactions that eliminate ranges of possibilities and then confirms suspected ions with a confirming test. Sometimes small carbon containing ions are included in such schemes. With modern instrumentation these tests are rarely used but can be useful for educational purposes and in field work or other situations where access to state-of-the-art instruments are not available or expedient.

Gravimetric analysis

For more details on this topic, see Gravimetric analysis.

 

Gravimetric analysis involves determining the amount of material present by weighing the sample before and/or after some transformation. A common example used in undergraduate education is the determination of the amount of water in a hydrate by heating the sample to remove the water such that the difference in weight is due to the loss of water.

 

Volumetric analysis

For more details on this topic, see Titration.

 

A Winkler titration to determine the concentration of dissolved oxygen in a water sample

 

Titration involves the addition of a reactant to a solution being analyzed until some equivalence point is reached. Often the amount of material in the solution being analyzed may be determined. Most familiar to those who have taken chemistry during secondary education is the acid-base titration involving a color changing indicator. There are many other types of titrations, for example potentiometric titrations. These titrations may use different types of indicators to reach some equivalence point.

 

Instrumental methods

Instrumental analysis

Block diagram of an analytical instrument showing the stimulus and measurement of response

 

Spectroscopy

For more details on this topic, see Spectroscopy.

Spectroscopy measures the interaction of the molecules with electromagnetic radiation. Spectroscopy consists of many different applications such as atomic absorption spectroscopy, atomic emission spectroscopy, ultraviolet-visible spectroscopy, x-ray fluorescence spectroscopy, infrared spectroscopy, Raman spectroscopy, dual polarisation interferometry, nuclear magnetic resonance spectroscopy, photoemission spectroscopy, Mössbauer spectroscopy and so on.

 

Mass spectrometry

For more details on this topic, see Mass spectrometry.

An accelerator mass spectrometer used for radiocarbon dating and other analysis.

 

Mass spectrometry measures mass-to-charge ratio of molecules using electric and magnetic fields. There are several ionization methods: electron impact, chemical ionization, electrospray, fast atom bombardment, matrix assisted laser desorption ionization, and others. Also, mass spectrometry is categorized by approaches of mass analyzers: magnetic-sector, quadrupole mass analyzer, quadrupole ion trap, time-of-flight, Fourier transform ion cyclotron resonance, and so on.

 

Electrochemical analysis

For more details on this topic, see Electroanalytical method.

Electroanalytical methods measure the potential (volts) and/or current (amps) in an electrochemical cell containing the analyte.^[7][8] These methods can be categorized according to which aspects of the cell are controlled and which are measured. The three main categories are potentiometry (the difference in electrode potentials is measured), coulometry (the cell’s current is measured over time), and voltammetry (the cell’s current is measured while actively altering the cell’s potential).

Thermal analysis

Further information: Calorimetry, thermal analysis

 

The world’s first ice-calorimeter, used in the winter of 1782-83, by Antoine Lavoisier and Pierre-Simon Laplace, to determine the heat involved in various chemical changes; calculations which were based on Joseph Black’s prior discovery of latent heat. These experiments mark the foundation of thermochemistry.

 

Indirect calorimetry metabolic cart measuring oxygen uptake and CO₂ production of a spontaneously breathing subject (dilution method with canopy hood).

Snellen direct calorimetry chamber, University of Ottawa.^[1]

 

Calorimetry and thermogravimetric analysis measure the interaction of a material and heat.

 

Separation

 

Separation of black ink on a thin layer chromatography plate.

Further information: Separation process, Chromatography, electrophoresis

Separation processes are used to decrease the complexity of material mixtures. Chromatography, electrophoresis and Field Flow Fractionation are representative of this field.

Hybrid techniques

Combinations of the above techniques produce a “hybrid” or “hyphenated” technique. Several examples are in popular use today and new hybrid techniques are under development. For example, gas chromatography-mass spectrometry, gas chromatography-infrared spectroscopy, liquid chromatography-mass spectrometry, liquid chromatography-NMR spectroscopy. liquid chromagraphy-infrared spectroscopy and capillary electrophoresis-mass spectrometry.

Hyphenated separation techniques refers to a combination of two (or more) techniques to detect and separate chemicals from solutions. Most often the other technique is some form of chromatography. Hyphenated techniques are widely used in chemistry and biochemistry. A slash is sometimes used instead of hyphen, especially if the name of one of the methods contains a hyphen itself.

 

Microscopy

Fluorescence microscope image of two mouse cell nuclei in prophase (scale bar is 5 µm).

 

For more details on this topic, see Microscopy.

 

The visualization of single molecules, single cells, biological tissues and nanomaterials is an important and attractive approach in analytical science. Also, hybridization with other traditional analytical tools is revolutionizing analytical science. Microscopy can be categorized into three different fields: optical microscopy, electron microscopy, and scanning probe microscopy. Recently, this field is rapidly progressing because of the rapid development of the computer and camera industries.

Lab-on-a-chip

A glass microreactor

 

Further information: microfluidics, lab-on-a-chip

Devices that integrate (multiple) laboratory functions on a single chip of only millimeters to a few square centimeters in size and that are capable of handling extremely small fluid volumes down to less than pico liters.

 

Standards

Analytical quality control

Standard curve

A calibration curve plot showing limit of detection (LOD), limit of quantification (LOQ), dynamic range, and limit of linearity (LOL).

 

A general method for analysis of concentration involves the creation of a calibration curve. This allows for determination of the amount of a chemical in a material by comparing the results of unknown sample to those of a series known standards. If the concentration of element or compound in a sample is too high for the detection range of the technique, it can simply be diluted in a pure solvent. If the amount in the sample is below an instrument’s range of measurement, the method of addition can be used. In this method a known quantity of the element or compound under study is added, and the difference between the concentration added, and the concentration observed is the amount actually in the sample.

 

Internal standards

Sometimes an internal standard is added at a known concentration directly to an analytical sample to aid in quantitation. The amount of analyte present is then determined relative to the internal standard as a calibrant. An ideal internal standard is isotopically-enriched analyte which gives rise to the method of isotope dilution.

Standard addition

 

The method of standard addition is used in instrumental analysis to determine concentration of a substance (analyte) in an unknown sample by comparison to a set of samples of known concentration, similar to using a calibration curve. Standard addition can be applied to most analytical techniques and is used instead of a calibration curve to solve the matrix effect problem.

 

Signals and noise

One of the most important components of analytical chemistry is maximizing the desired signal while minimizing the associated noise. The analytical figure of merit is known as the signal-to-noise ratio (S/N or SNR).

Noise can arise from environmental factors as well as from fundamental physical processes.

 

Thermal noise Johnson–Nyquist noise

Thermal noise results from the motion of charge carriers (usually electrons) in an electrical circuit generated by their thermal motion. Thermal noise is white noise meaning that the power spectral density is constant throughout the frequency spectrum.

The root mean square value of the thermal noise in a resistor is given by^[15]

where k_B is Boltzmann’s constant, T is the temperature, R is the resistance, and is the bandwidth of the frequency .

 

Shot noise   Shot noise

Shot noise is a type of electronic noise that occurs when the finite number of particles (such as electrons in an electronic circuit or photons in an optical device) is small enough to give rise to statistical fluctuations in a signal.

Shot noise is a Poisson process and the charge carriers that make up the current follow a Poisson distribution. The root mean square current fluctuation is given by^[15]

where e is the elementary charge and I is the average current. Shot noise is white noise.

 

Flicker noise flicker noise

Flicker noise is electronic noise with a 1/ƒ frequency spectrum; as f increases, the noise decreases. Flicker noise arises from a variety of sources, such as impurities in a conductive channel, generation and recombination noise in a transistor due to base current, and so on. This noise can be avoided by modulation of the signal at a higher frequency, for example through the use of a lock-in amplifier.

 

Environmental noise

Noise in a thermogravimetric analysis; lower noise in the middle of the plot results from less human activity (and environmental noise) at night.

 

Environmental noise arises from the surroundings of the analytical instrument. Sources of electromagnetic noise are power lines, radio and television stations, wireless devices, Compact fluorescent lamps^[16] and electric motors. Many of these noise sources are narrow bandwidth and therefore can be avoided. Temperature and vibration isolation may be required for some instruments.

 

Noise reduction

Noise reduction can be accomplished either in computer hardware or software. Examples of hardware noise reduction are the use of shielded cable, analog filtering, and signal modulation. Examples of software noise reduction are digital filtering, ensemble average, boxcar average, and correlation methods.^[15]

 

Applications

Analytical chemistry research is largely driven by performance (sensitivity, selectivity, robustness, linear range, accuracy, precision, and speed), and cost (purchase, operation, training, time, and space). Among the main branches of contemporary analytical atomic spectrometry, the most widespread and universal are optical and mass spectrometry. In the direct elemental analysis of solid samples, the new leaders are laser-induced breakdown and laser ablation mass spectrometry, and the related techniques with transfer of the laser ablation products into inductively coupled plasma. Advances in design of diode lasers and optical parametric oscillators promote developments in fluorescence and ionization spectrometry and also in absorption techniques where uses of optical cavities for increased effective absorption pathlength are expected to expand. The use of plasma- and laser-based methods is increasing. An interest towards absolute (standardless) analysis has revived, particularly in emission spectrometry.

great effort is put in shrinking the analysis techniques to chip size. Although there are few examples of such systems competitive with traditional analysis techniques, potential advantages include size/portability, speed, and cost. (micro Total Analysis System (µTAS) or Lab-on-a-chip). Microscale chemistry reduces the amounts of chemicals used.

Many development improve the analysis of biological systems. Examples of rapidly expanding fields in this area are:

·                     Genomics – DNA sequencing and its related research. Genetic fingerprinting and DNA microarray are important tools and research fields.

·                     Proteomics – the analysis of protein concentrations and modifications, especially in response to various stressors, at various developmental stages, or in various parts of the body.

·                     Metabolomics – similar to proteomics, but dealing with metabolites.

·                     Transcriptomics – mRNA and its associated field

·                     Lipidomics – lipids and its associated field

·                     Peptidomics – peptides and its associated field

·                     Metalomics – similar to proteomics and metabolomics, but dealing with metal concentrations and especially with their binding to proteins and other molecules.

Analytical chemistry has played critical roles in the understanding of basic science to a variety of practical applications, such as biomedical applications, environmental monitoring, quality control of industrial manufacturing, forensic science and so on.

The recent developments of computer automation and information technologies have extended analytical chemistry into a number of new biological fields. For example, automated DNA sequencing machines were the basis to complete human genome projects leading to the birth of genomics. Protein identification and peptide sequencing by mass spectrometry opened a new field of proteomics.

Analytical chemistry has been an indispensable area in the development of nanotechnology. Surface characterization instruments, electron microscopes and scanning probe microscopes enables scientists to visualize atomic structures with chemical characterizations.

Classification of Matter

         Matter can be identified by its characteristic inertial and gravitational mass and the space that it occupies. On earth matter is commonly found in three different states: solid, liquid, and gas.

Introduction

         A substance is a sample of matter whose physical and chemical properties are the same throughout the sample because the matter has a constant composition. It is common to see substances changing from one state of matter to another. To differentiate the states of matter at least at a particle level, we look at the behavior of the particles within the substance. When substances change state, it is because the spacing between the particles of the substances is changing due to a gain or loss of energy. For example, we all have probably observed that water can exist in three forms with different characteristic ways of behaving: the solid state (ice), liquid state (water), and gaseous state (water vapor and steam). Due to water’s prevalence, we use it to exemplify and describe the three different states of matter. As ice is heated and the particles of matter that make up water gain energy, eventually the ice melts in to water that eventually boils and turns into steam.

         Before we examine the states of matter, we will consider some ways samples of matter have been classified by those who have studied how matter behaves.

Classifying Matter

         Evidence suggests that substances are made up of smaller particles that are ordinarily moving around. Some of those particles of matter can be split into smaller units using fairly strong heat or electricity into smaller rather uniform bits of matter called atoms. Atoms are the building blocks of elements. Elements are all those substances that have not ever been decomposed or separated into any other substances through chemical reactions, by the application of heat, or by attempting to force an direct electric current through the sample.

         Atoms in turn have been found to be made up of yet smaller units of matter called electrons, protons, and neutrons.

         Elements can be arranged into what is called the periodic table of elements based on observed similarities in chemical and physical properties among the different elements. When atoms of two or more elements come together and bond, a compound is formed. The compound formed can later be broken down into the pure substances that originally reacted to form it.

         Compounds such as water are composed of smaller units of bonded atoms called molecules. Molecules of a compound are composed of the same proportion of elements as the compound as a whole since they are the smallest units of that compound. For example, every portion of a sample of water is composed of water molecules. Each water molecule contains two hydrogen atoms and one oxygen atom, and so water as a whole has, in a combined state, twice as many hydrogen atoms as oxygen atoms..

         Water can still consist of the same molecules, but its physical properties may change. For instance, water at a temperature below 0 degrees Celsius (32 degrees Fahrenheit) is ice, whereas water above the temperature of 100. degrees Celsius (212 degrees Fahrenheit) is a gas, water vapor. When matter changes from one state to another, temperature and pressure may be involved in the process and the density and other physical properties change. The temperature and pressure exerted on a sample of matter determines the resulting form of that the matter takes, whether solid, liquid, or gas.

         Since the properties of compounds and elements are uniform, they are classified as substances. When two or more substances are mixed together, the result is called a mixture.          Mixtures can be classified into two main categories: homogeneous and heterogeneous. A homogeneous mixture is one in which the composition of its constituents are uniformly mixed throughout. A homogeneous mixture in which on substance, the solute, dissolves completely in another substance, the solvent, may also be called a solution. Usually the solvent is a liquid, however the solute can be either a liquid, solid, or a gas. In a homogeneous solution, the particles of solute are spread evenly among the solvent particles and the extremely small particles of solute cannot be separated from the solvent by filtration through filter paper because the spaces between paper fibers are much greater than the size of the solute and solvent particles. Other examples of homogeneous mixtures include sugar water, which is the mixture of sucrose and water, and gasoline, which is a mixture of dozens of compounds.

         A heterogeneous mixture is a nonuniform mixture in which the components separate and the composition varies. Unlike the homogeneous mixture, heterogeneous mixtures can be separated through physical processes. An example of a physical process used is filtration, which can easilty separate the sand from the water in a sand-water mixture by using a filter paper. Some more examples of heterogeneous mixtures include salad dressing, rocks, and oil and water mixtures. Heterogeneous mixtures involving at least one fluid are also called suspension mixtures and separate if they are left standing long enough. Consider the idea of mixing oil and water together. Regardless of the amount of time spent shaking the two together, eventually oil and water mixtures will separate with the oil rising to the top of the mixture due to its lower density.

         Mixtures that fall between a solution and a heterogeneous mixture are called colloidal suspensions (or just colloids). A mixture is considered colloidal if it typically does not spontaneously separate or settle out as time passes and cannot be completely separated by filtering through a typical filter paper. It turns out that a mixture is colloidal in its behavior if one or more of its dimensions of length, width, or thickness is in the range of 1-1000 nm. A colloidal mixture can also be recognized by shining a beam of light through the mixture. If the mixture is colloidal, the beam of light will be partially scattered by the suspended nanometer sized particles and can be observed by the viewer. This is known as the Tyndall effect. In the case of the Tyndall effect, some of the light is scattered since the wavelengths of light in the visible range, about 400nm to 700 nm, are encountering suspended colloidal sized particles of about the same size. In contrast, if the beam of light were passed through a solution, the observer standing at right angles to the direction of the beam would see no light being reflected from either the solute or solvent formula units that make up the solution because the particles of solute and solvent are so much smaller than the wavelength of the visible light being directed through the solution.

• Solutions: molecules ~0.1-2 nm in size

• Colloids: molecules ~ 2-1000 nm in size

• Suspensions: molecules greater than ~ 1000 nm in size

Examples of Homogeneous Mixtures, also known as Solutions

Filtered seawater is solution of the compounds of water, salt (sodium chloride), and other compounds.

Examples of Heterogeneous Mixtures

separation of sand and water separation of salad dressing various mixtures within a rock

Example of a Colloidal Mixture Whose Components Tend Not to Settle Out

milk is a colloid of liquid butterfat globules suspended in water

 

Separation of Mixtures

         Most substances are naturally found as mixtures, therefore it is up to the chemist to separate them into their natural components. One way to remove a substance is through the physical property of magnetism. For example, separating a mixture of iron and sulfur could be achieved because pieces of iron would be attracted to a magnet placed into the mixture, removing the iron from the remaining sulfur.

         Filtration is another way to separate mixtures. Through this process, a solid is separated from a liquid by passing through a fine pored barrier such as filter paper. Sand and water can be separated through this process, in which the sand would be trapped behind the filter paper and the water would strain through. Another example of filtration would be separating coffee grounds from the liquid coffee through filter paper.

         Distillation is another technique to separate mixtures. By boiling a solution of a non-volatile solid disolved in a liquid in a flask, vapor from the lower boiling point solvent can be driven off from the solution by heat, be condensed back into the liquid phase as it comes in contact with cooler surfaces, and be collected in another container. Thus a solution such as this may be separated into its original components, with the solvent collected in a separate flask and the solute left behind in the original distillation flask. An example of a solution being separated through distillation would be the distillation of a solution of copper(II) sulfate in water, in which the water would be boiled away and collected and the copper(II) sulfate would remain behind in the disllation flask.

         The picture above depicts the equipment needed for a distillation process. The homogeneous mixture starts out in the left flask and is boiled. The vapor then travels down chilled tube on the right and condenses back into a liquid and drips into the flask.

States of Matter

         Everything that is familiar to us in our daily lives – from the land we walk on, to the water we drink and the air we breathe – is based upon the states of matter called gases, liquids, and solids.

Solid

         When the temperature of a liquid is lowered to the freezing point of the substance (for water the freezing point is 0^oC), the movement of the particles slows with the spacing between the particles changing until the attractions between the particles lock the particles into a solid form. At the freezing point, the particles are closely packed together and tend to block the motions of each other. The attractions between the particles hold the particles tightly together so that the entire ensemble of particles takes on a fixed shape. The volume of the solid is constant and the shape of a solid is constant unless deformed by a sufficiently strong external force. (Solids are thus unlike liquids whose particles are slightly less attracted to one another because the particles of a liquid are a bit further apart than those in the corresponding solid form of the same substance.) In a solid the particles remain in a relatively fixed positions but continue to vibrate. The vibrating particles in a solid do not completely stop moving and can slowly move into any voids that exist within the solid.

The diagram on the left represents a solid whose constituent particles are arranged in an orderly array, a crystal lattice. The image on the right is a ice cube. It has changed from liquid into a solid as a result of absorbing energy from its warmer environment.

Liquid

         When the temperature of a sample increases above the melting point of a solid, that sample can be found in the liquid state of matter. The particles in the liquid state are much closer together than those in the gaseous state, and still have a quite an attraction for each other as is apparent when droplets of liquid form. In this state, the weak attractive forces within the liquid are unable to hold the particles into a mass with a definite shape. Thus a liquid’s shape takes on the shape of any particular container that holds it. A liquid has a definite volume but not a definite shape. Compared to to the gaseous state there is less freedom of particle movement in the liquid state since the moving particles frequently are colliding with one another, and slip and slide over one another as a result of the attractive forces that still exist between the particles, and hold the particles of the liquid loosely together. At a given temperature the volume of the liquid is constant and its volume typically only varies slightly with changes in temperature.

The diagram on the left represents a container patially filled with a liquid. The image on the right is of water being poured out of a glass. This shows that liquid water has no particular shape of its own.

Gas

         In the gas phase, matter does not have a fixed volume or shape. This occurs because the molecules are widely separated with the spaces between the particles typically around ten times further apart in all three spatial directions, making the gas around 1000 times less dense than the corresponding liquid phase at the same temperature. (A phase is a uniform portion of mater.) As the temperature of a gas is increased, the particles to separate further from each other and move at faster speeds. The particles in a gas move in a rather random and independent fashion, bouncing off each other and the walls of the container. Being so far apart from one another, the particles of a real gas only weakly attract each other such that the gas has no ability to have a shape of its own. The extremely weak forces acting between the particles in a gas and the greater amount of space for the particles to move in results in almost independent motion of the moving, colliding particles. The particles freely range within any container in which they are put, filling its entire volume with the net result that the sides of the container determine the shape and volume of gas. If the container has an opening, the particles heading in the direction of the opening will escape with the result that the gas as a whole slowly flows out of the container.

The image on the left represents an enclosed container filled with gas. The images are meant to suggest that the gas particles in the container are moving freely and randomly in myriad directions.The image on the right shows condensing water forming from the water vapor that escaped from the container.

Other States of Matter

         Besides of the three classical states of matter, there are many other states of matter that share characteristics of one more of the classical states of matter. Most of these states of matter can be put into three categories according to the degrees in varying temperature. At room temperature, the states of matters include liquid crystal, amorphous solid, and magnetically ordered states. At low temperatures the states of matter include superconductors, superfluids, and Bose-Einstein condensate state of matter. At high temperatures the states of matter include, plasma and Quark-gluon plasma. These other states of matter are not typically studied in general chemistry.