Chemistry Kenya
Salts
Salts
Salts
Methods of preparing salts
Preparation of soluble salts by reaction of acids with; metals, metal hydroxides, metal oxides, metal carbonates and metal hydrogen carbonates
Several methods are available for the preparation of salts. The solubilities of the prepared salts determine their methods of preparation. Hence, in the choice of a method of preparation of a particular salt, one has to be acquainted with its solubility properties.
Soluble salts are usually prepared by methods which involve crystallization. In this method, as the name suggests, resultant salts are in the form of crystals.
Insoluble salts are usually prepared by methods which involve precipitation. These methods are sometimes referred to as double decomposition. To precipitate an insoluble salt, you must mix a solution that contains its positive ions with the one that contains its negative ions.
Salts may also be prepared by direct combination (or synthesis). For example, magnesium chloride may be prepared in the laboratory by heating magnesium in a stream of chlorine.
Mg(s) + Cl2(g) → MgCl2(s)
Preparation of soluble salts
Soluble salts may be prepared by any of the following methods:
- Reaction between an acid and an alkali: In this method, a dilute acid is added to an alkali in the appropriate volume ratio. The reaction between an acid and an alkali is termed as neutralization. For example, sodium chloride may be prepared by the following neutralization reaction:NaOH(aq) + HCl(aq)→ NaCl(aq) + H2O(l). Both reactants are soluble, and no gas is given off during the reaction. So, it is difficult to know when the reaction is over. In this case, you have to use an indicator. A universal indicator or litmus could be used, but even better is phenolphthalein. This is pink in alkaline solution, but colourless in neutral or acidic solutions.
- Reaction of a metal with an acid: This is another general method for preparing salts. For example, zinc sulphate can be made by reacting dilute sulphuric acid with zinc:Zn(s) + H2SO4(aq) → ZnSO4(aq) + H2(g). However, this method is not suitable for all metals or all acids. It is good for preparing salts of fairly reactive metals such as magnesium, aluminium, zinc and iron. However, the reactions of highly reactive metals like sodium, potassium and calcium with acids are very violent and dangerous. The reaction with lead is too slow. Copper, silver and gold do not react at all.
- Reaction of a metal oxide with an acid. Metal oxides, as you studied early, react with dilute acids to produce salts. Copper oxide is an insoluble base. Although copper will not react with dilute sulphuric acid, copper (II) oxide will. The salt that forms is copper (II) sulphate.CuO(s) + H2SO4(aq) → CuSO4(aq) + H2O(l)
- Reaction of a metal carbonate with an acid. The reaction between metal carbonates and dilute acids are accompanied with evolution of carbon dioxide gas. The evolution of a gas can be used to indicate when the reaction is over. An example of such reactions is the reaction between calcium carbonate and dilute hydrochloric acid. CaCO3(s) + 2HCl(aq) → CaCl2(aq) + H2O(l) + CO2(g)
General methods of preparing soluble salts
The above methods for preparing soluble salts are specific for each method mentioned. Generally, soluble salts may be prepared by two broad methods.
Method 1: This route is essentially the same whether starting with a solid metal, a solid base (oxide) or a solid carbonate. The route can be divided into four stages:
- Stage 1: An excess (more than enough) of the solid is added to the acid and allowed to react. Using an excess of the solid makes sure that all the acid used up. If it is not used up at this stage, the acid would become more concentrated when the water is evaporated later (stage 3).
- Stage 2: The excess solid is filtered out after the reaction is completed.
- Stage 3: The filtrate is gently evaporated to concentrate the solution. This can be done on a heated water bath. Do not heat so strongly or “spitting” might take place.
- Stage 4: The concentrated solution is cooled down to let the crystals form. Filter off the crystals. Wash them with a little distilled water. Dry the crystals carefully between the filter papers.
Method 2: This method (titration method) involves the neutralization of an acid with an alkali (for example sodium hydroxide) or a soluble carbonate (for example sodium carbonate). Since both the reactants and the products are colourless, an indicator is used to find the neutralization point or end point (when all the acid has just been neutralized). Once the end point is reached, the resulting salt solution is evaporated and cooled to form crystals as described in method 1.
General methods for preparing insoluble salts
Some salts are insoluble in water (for example silver chloride and barium sulphate – see table 3.6). Such salts are generally prepared by ionic precipitation.Precipitation is the sudden formation of a solid either: when two solutions are mixed; or when a gas is bubbled into a solution.
For example, barium sulphate can be prepared by adding a solution of a soluble sulphate (for example sodium sulphate) to a solution of a soluble barium salt (for example barium chloride). The insoluble barium sulphate is formed immediately. This solid falls to the bottom of the container as a precipitate (figure 3.2). The precipitate can be filtered off. It is then washed with distilled water and dried in a warm oven. The equation for the reaction is:
BaCl2(aq) + Na2SO4(aq)→ BaSO4(s) + 2NaCl(aq)
This shows how important the state symbols can be - it is only through state symbols that we can tell this equation shows a precipitation.
Barium sulphate could also be made from barium nitrate and sodium sulphate, for example, since these salts are both soluble. As long as barium and sulphate ions are present, barium sulphate will be precipitated.
Ba2+(aq) + SO42-(aq) →BaSO4(s)
Precipitation reactions are often used in a qualitative analysis to identify salts such as chlorides, iodides and sulphates.
Preparation of salts by direct combination (synthesis)
Some soluble and insoluble salts can be made directly by reacting two elements together. This is called combination (or synthesis). This type of reaction is mainly possible for metal chlorides, bromides and iodides. For instance, if a piece of burning sodium is lowered into a gas jar of chlorine, the two react violently to produce a white powder of sodium chloride:
2Na(s) + Cl2(g) → 2NaCl(s)
Other chlorides can also be prepared by combination, for example, iron (III) chloride and aluminium chloride can be made by heating iron and aluminium metals in stream of chlorine:
2Fe(s) + 3Cl2(g) → 2FeCl3(s)
2Al(s) + 3Cl2(g) → 2AlCl3(s)
The reaction between ammonia gas and hydrogen chloride gas to produce ammonium chloride is also a synthesis reaction.
NH3(g) + HCl(g) → NH4Cl(s)
Direct combination reactions do not produce crystals of the salt, but only a powder.
Preparation of insoluble salts by precipitation (ionic equations required)
Direct combination reaction (e.g. sodium with chlorine, iron with sulphur)
Types of salts; normal, acid and double salts.
Solubility of salts
Some salts are more soluble in water than others are. However, other salts are insoluble in water. The knowledge of solubility of different salts in water is very important because it can help us prepare different salts in the laboratory by such methods as precipitation, direct combination (synthesis), crystallization and so forth.
As regards to solubilities, salts can be classified into two groups: salts which are soluble in water (soluble salts) and salts which do not dissolve in water (insoluble salts). Table 3.6 summarizes the solubility of different salts in water.
| Soluble salts | Insoluble salts |
| 1. All sodium, potassium and ammonium salts. | silver, mercury(I) and lead chlorides barium, lead (II) and calcium sulphates but other common carbonates are insoluble.but other common hydroxides are insoluble. |
| 2. All nitrates of metals | |
| 3. All chlorides except .……………........ | |
| 4. All sulphates except………………….. | |
| 5. Sodium, potassium, and ammonium carbonates…………………………... hydroxides………………………….…… | |
| 6. Sodium, potassium and ammonium |
Solubility of sulphates, chlorides, nitrates and carbonates in water
Relationship between method of preparation and solubility. NOTE: The solubility of hydroxides and oxides should be considered along with others.
Action of heat on salts
Effects of heat on the following salts; carbonates, nitrates, sulphates and hydrated salts (include ammonium salts)
When different salts are heated, they behave in different manners. The crystals of some salts contain water of crystallization. When these hydrated salts are heated, their water of crystallization is driven off as steam. The crystals then lose their shape and become a powder. The following are few examples of hydrated salts:
| Salt formula | Chemical name |
| CuSO4.5H2O | Copper (II) sulphate five water |
| Na2CO3.I0H2O | Sodium carbonate ten water |
| MgCl2.6H2O | Magnesium chloride six water |
| FeCl3.6H2O | Iron (III) chloride six water |
| FeSO4.7H2O | Iron (II) sulphate seven water |
| CoCl2.6H2O | Cobalt (II) chloride six water |
| MgSO4.7H2O | Magnesium sulphate seven water |
| CaSO4.2H2O | Calcium sulphate two water |
Sulphates
Sulphates of potassium, sodium, calcium, lithium and magnesium are stable to heat and do not decompose when heated. Other sulphates decompose to give the oxide and sulphur trioxide gas except iron (III) sulphate which decomposes to give sulphur dioxide and sulphur trioxide.
Copper (II) sulphate five water crystals are blue in colour, but when heated, they are dehydrated to form a white powder:
CuSO4.5H2O(s)hydrated (blue)→ CuSO4(s)anhydrous (white) + 5H2O(g)
Crystals that have lost their water of crystallization are called anhydrous. If water is added back to the anhydrous copper (II) sulphate powder, the powder turns into blue crystals again and heat is evolved. This can be used as a qualitative test for water.If the white, anhydrous powder is further heated strongly, it decomposes to black copper (II) oxide:
CuSO4(s)white→ CuO(s)black+ SO3(g)
Hydrated iron (II) sulphate is green in colour. When heated, it loses all its water of crystallization and changes colour from green to white:
FeSO4.7H2O(s)green→ FeSO4(s)white + 7H2O(g)
When heated even more strongly, the white powder decomposes to form a black oxide:
2FeSO4(s)white→ Fe2O3(s)black+ SO2(g) + SO3(g)
Iron (III) sulphate decomposes on heating to form slightly different products:
Fe2(SO4)3(s) → Fe2O3 + 3SO3(g)
Chlorides
The chlorides of most metals are hydrated except those of potassium, lead, mercury and silver. Hydrated chlorides do not usually give the anhydrous salt when heated. Instead, a chemical change termed as hydrolysis normally occurs. The reaction is accompanied by the evolution of steam and hydrogen chloride gas, and the formation of the basic chloride or oxide. When, for example, hydrated magnesium chloride is heated, its basic chloride is formed:
MgCl2.6H2O(s) → Mg (OH)Cl(s) + HCl(g) + 5H2O(g)
The same case applies when hydrated calcium chloride is heated. However, when hydrated aluminum chloride is heated, it does not produce the anhydrous salt. Instead, the oxide is formed thus:
2AlCl3.6H2O(s) → Al2O3(s) + 6HCl(g) + 3H2O(g)
Ammonium chloride sublimes when heated. The reaction is reversible and the products may recombine on cooling to form the salt back.
NH4Cl=⇔NH3(g) + HCl(g)
Carbonates and hydrogencarbonates
The carbonates of potassium and sodium are very stable to heat. They do not decompose even when heated to very high temperatures. All other carbonates decompose when heated to give the oxide and carbon dioxide:
CaCO3(s) ⇔CaO(s) + CO2(g)
CuCO3(s) → CuO(s) + CO2(g)
However, there are very few and exceptional carbonates that do not behave like this. Ammonium carbonate, for example, decomposes readily when heated to give ammonia gas, water vapour and carbon dioxide gas:
(NH4)2CO3(s) → 2NH3(g) + H2O(g) + CO2(g)
All hydrogencarbonates decompose on heating to give the carbonates, water vapour and carbon dioxide:
2NaHCO3(s) → Na2CO3(s) + H2O(g) CO2(g)
Nitrates
When heated, potassium and sodium nitrates decompose to give the nitrite and oxygen:
2KNO3(s) → 2KNO2(s) + O2(g)
2NaNO3(s) → 2NaNO2(s) + O2(g)
The nitrates of common heavy metals (such as Pb, Al, Ca, Mg, Zn and Cu) decompose on heating to give the oxide, nitrogen dioxide and oxygen:
2Pb(NO3)2(s) → 2PbO(s) + 4NO2 (g) + O2 (g)
2Ca(NO3)2(s) → 2CaO(s) + 4NO2 (g) + O2 (g)
The nitrates of silver and mercury are completely decomposed to the metal, nitrogen dioxide and oxygen:
2AgNO3(s) → 2Ag (g) + 2NO2 (g) + O2 (g)
Hg(NO3)2(s) → Hg(l) + 2NO2(g) + O2(g)
Ammonium nitrate is decomposed by heat into dinitrogen oxide and water:
NH4NO3(s) → N2O (g) + 2H2O (l)
Hydroxides
Potassium and sodium hydroxides are very stable to heat. They do not decompose even when heated strongly. All other hydroxides decompose to give the oxide and water vapour, e.g.:
Ca(OH)2(s) → CaO(s) + H2O(g)
DELIQUESCENCE, EFFLORESCENCE AND HYGROSCOPY
Deliquescence
Deliquescence is the absorbing of moisture from the atmosphere by a solid to form a solution.If calcium chloride (CaCl2) is exposed to air, it absorbs water vapour from the atmosphere and eventually dissolves. Its tendency to absorb water vapour explains why it is used as a drying agent for gases (not ammonia, because it combines with the gas).
Solid sodium hydroxide is also deliquescent. On exposure to air, pellets of sodium hydroxide quickly become shiny and then sticky as they absorb water vapour from the atmosphere. Eventually the sodium hydroxide pellets absorb moisture from the atmosphere so much that they dissolve to form a solution of sodium hydroxide.
Applications:- use of lime to change pH, use of salts as anti-acids, use of salts as inorganic fertilizers
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