Solutions

General Definitions
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 * Solution - homogeneous mixtures – made up of a solute and a solvent.
 * Solvent – ‘does the dissolving’ – usually a liquid.
 * Solute – ‘is dissolved’ – usually the solid. For example: When sugar dissolves in water, solvent is the water and the solute is the sugar.
 * Saturated solution – one in which no more solute will dissolve under existing conditions of temperature and pressure.
 * Unsaturated solution – a solution which contains less than the quantity of solute needed to saturate it under existing conditions of temperature and pressure.
 * Supersaturated solution – solution which contains more solute than a saturated solution could normally hold under the existing conditions.
 * Solubility – the degree to which a solute dissolves in a solvent.
 * Dilute: contains very little solute in relation to the amount of solvent.
 * Concentrated: contains a large amount of solute in relation to the amount of solvent.
 * Concentration: refers to the amount of solute dissolved in a certain amount of solvent.

The Solution Process And Factors Affecting Solubility
For Sugar to dissolve in water the following must take place:
 * Intermolecular forces between the sugar molecules must be overcome as the crystals dissolve.
 * The intermolecular forces between some water molecules must be overcome to make room for the sugar molecules.
 * Intermolecular forces must form between sugar molecules and surrounding water molecules.
 * Intermolecular forces are forces of electrostatic attraction between molecules.

Nature of the solute and solvent:
In general, solute will dissolve in a solvent only if:


 * Intermolecular forces within the solute and those within the solvent are similar to those occurring between solute and solvent molecules.
 * Polar solvents tend to dissolve polar solutes and non-polar solvents tend to dissolve non-polar solutes.

Oil does not dissolve in water because suitable intermolecular forces do not occur between solute and solvent molecules. Oil will dissolve in non-polar solvents such as kerosene because suitable intermolecular forces exist between the oil and kerosene molecules.

Some ionic compounds dissolve in water and others do not because some ionic compounds are very stable – that is the force holding their ions in their crystal lattice are very strong.

CH3OH(l) HCl (g) H2S (g) ||= CH4 (g) C8H8 (l) ||
 * = **Readily soluble in water (polar)** ||= **Readily soluble in CCl4 (non-polar)** ||
 * = NH3 (g)

Temperature

 * An increase in temperature results in an increase in solubility.
 * The solubility of gases when dissolved in a liquid decreases with increasing temperature.

Gas Pressure
Gas pressure has little or no effect on the solubility of solids in liquids. The solubility of gases in liquids is directly proportional to the pressure of the gas above the liquid.

Aqueous Solutions And Electrolytes

 * Aqueous solutions – solutions in which the solute dissolves in water – denoted by symbol (aq).
 * Electrolytes are those substances that dissolve in water producing ions.
 * They are classified according to the extent to which they produce ions when dissolved in water.

Strong Electrolytes

 * Entirely or mostly present as ions when dissolved NaCl (s) → Na+ (as) + Cl-(aq)
 * All strong electrolytes dissolve to form ions only.
 * All ionic compounds are strong electrolytes.

Weak Electrolytes

 * Partly present as ions when dissolved – covalent molecular substances.
 * Dissolve mainly forming molecules with a much lower concentration of ions.

Non-electrolyte

 * Produce no ions when dissolved – covalent molecular substances.
 * Produce only molecules when dissolved in water.

An electrolyte will increase a solution’s electrical conductivity. This is due to the mobile ions releases from the electrolyte. A solutions electrical conductivity increases with increasing ion concentration. This in turns increases the electrolyte strength and concentration.

Dissociation
NaCl(s) → Na+(aq) + Cl-(aq) This is known as a dissociation reaction – the ions in the ionic solid have been separated in the solution/dissolving process.
 * Ionic compounds dissociated into their ions when they dissolve in water:

A solvent (e.g. distilled waer) which does not contain ions does not conduct electritity. An aqueous solution of NaCl is a good conductor of electritity – the Na+ and Cl- ions move through the solution when an electric current flows through it.
 * Evidence for the dissociation of NaCl is provided by conductivity measurements:
 * CATIONS (positive ions) – attracted to the negative terminal of power pack (ANODE)
 * ANIONS (negative ions) attracted to the positive terminal of power pack (CATHODE)

Ionisation
Covalent molecular substances such as strong acids undergo ionisation reactons when they dissolve in water. HCl(g) → H+(aq) + Cl-(aq) Weak covalent molecular substances show both ions and molecular substance present in solution: HF(aq) → H+(aq) + F-(aq) The difference between ionisation and dissociation:
 * Dissociation – ions that are already present in the solid are separated and move into solution.
 * Ionisation – covalent molecular substances react to form ions which are separated in solution.

Common acid reactions
Acid + base → salt + water e.g. hydrochloric acid + sodium hydroxide → sodium chloride + water HCl (aq) + NaOH(aq) → NaCl(aq) + H2O(l)

Acid + reactive metal → salt + hydrogen gas e.g. sulfuric acid + zinc metal → zinc sulfate + hydrogen gas H2SO4 (aq) + Zn (s) → ZnSO4 (aq) + H2 (g)

Acid + carbonate → salt + water + carbon dioxide gas e.g. nitric acid + sodium carbonate → sodium nitrate + carbon dioxide gas + water 2HNO3 (aq) + Na2CO3 (aq) → 2NaNO3 (aq) + CO2 (g) + H2O (l)

Acid + hydrogen carbonate → salt + water + carbon dioxide gas e.g. hydrochloric acid + sodium hydrogen carbonate → sodium chloride + carbon dioxide gas + water HCl (aq) + NaHCO3 (aq) → NaCl(aq) + CO2 (g) + H2O(l)

Solubility Rules And Precipitation Reactions

 * A precipitation reaction occurs when an insoluble solid forms within a previously clear solution.


 * Precipitation will happen if any pair of oppositely charged ions present in a solution combine and form an insoluble compound. This will happen when two solutions containing different ions are combined and mixed.

Example: Solutions of Pb(NO3)2 (aq) and NaI(aq) are combined and mixed. Will a precipitate form? If so, give its formula. All solution concentrations are 1.0 mol L-1.

Identify and list all ions present in the combined mixture of the two solutions.

Pb2+(aq) NO3-(aq) Na+(aq) I-(aq)

Write the formula for possible precipitates by combing oppositely charged ions.

PbI2 or NaNO3

Refer to the solubility rules and determine the solubility of these substances. One is given on the next page but you must learn to use the one on your data table as this is the one in your tests. Any low or slightly solubility compound can be expected to form a precipitate.

NaNO3 is soluble (all nitrates are) so the only solid formed is PbI2 (s)

Write a balanced equation to summarise the change. Only include the ions that form the precipitate.

Pb2+(aq) + 2I(aq) → PbI2 (s)

Solubility rules for ionic solids in water

Colours of aqueous ions of selected transition elements

Ionic Equations When two solutions, such as silver nitrate solution and sodium chloride solution are mixed, a precipitate of silver chloride will form.

Ag+(aq) + NO3(aq) + Na+(aq) + Cl-(aq)  AgCl (s) + Na+(aq) + NO3(aq) This reaction occurs due to the silver ion and chloride ions. The nitrate ion and sodium ions are not involved in this reaction as shown by the fact that they are the same on both sides of the reaction. These ions are called “spectator ions” and are not included in the final equation. An equation that shows only the species that react is called an ionic equation. An ionic equation for this reaction is Ag+(aq) + Cl-(aq)  AgCl (s) The spectator ions are not shown in the reaction. An ionic equation is the most accurate way of representing what occurs when the reaction takes place. Ionic equations can be used for reactions other than precipitation reactions. However, they are mainly used for reactions which take place in solution. When writing ionic equations the following rules should be observed: Strong electrolytes are written in ionic form. Weak electrolytes are written in molecular form. Non-electrolytes are written in molecular form. Insoluble substances are written as the formulas of the substances. Gases are written in molecular form. Ionic equations should not include spectator ions. Equations must be balanced in atoms and electrical charge. Example 1: Write an ionic equation which shows the preparation of carbon dioxide gas by the reaction of sodium carbonate solution with hydrochloric acid.

The full equation is: Na2CO3 (aq) + 2HCl(aq) → 2NaCl (aq) + CO2 (g) + H2O (l) However an ionic equation gives shows us the actual reactions taking place and should be used instead. Checking the solubility tables we can determine the solid produced and the ions in solution:

Na+(aq) + CO32-(aq) + H+(aq) + Cl(aq) → Na+(aq) + Cl(aq) + CO2 (g) + H2O (l) It can be seen that the spectator ions are Na+(aq) and Cl(aq) and these take no part in the reaction so are removed and the ionic equation is:

CO32- (aq) + 2H+(aq)  CO2 (g) + H2O(l)

NOTE: Depending on whether the acid is added to a solid carbonate or a solution containing carbonate ions, the ionic equation can be written as follows: (This is the case for all substances.)

Acid + solid carbonate: Na2CO3 (s) + 2H+(aq)  2Na+(aq) + CO2 (g) + H2O(l)

Acid + carbonate in solution: CO32- (aq) + 2H+(aq)  CO2 (g) + H2O(l)

Example 2: Write an ionic equation for the preparation of hydrogen gas by the reaction of zinc metal with an acid.

Note: All acids have hydrogen ions so it doesn’t matter which acid is used as the reaction is only concerned with the hydrogen ion as shown below.

Zn(s) + 2H+(aq)  Zn2+(aq) + H2 (g)

CONCENTATIONS OF SOLUTIONS 8.1 Grams per Litre (g L-1) Solution concentration can be expressed as the number of grams dissolved in a litre of solution. For example, the sucrose concentration in cola drinks is usually 0.106 g mL-1. Hence a 375 mL can of this drink contains 39.8 g of sugar.

8.2 Percentage Composition by Mass The percentage composition by mass is the mass of solute in grams dissolved in 100 g of solution Concentration (% composition) = mass of solute (g) x 100 Mass of solution (g) For example concentrated hydrochloric acid is labelled as 32% w/w. This means that 1.00 kg of this acid solution would contain 320 g of hydrochloric.

8.3 Parts per Million (ppm) This method is used for very dilute solutions. The concentration in parts per million is the mass of solute, in milligrams, in one kilogram of solution. This is because 1 mg is one-millionth the mass of 1 kg.

Concentration (ppm) = mass of solute (mg) Mass of solution (kg) This unit is often used for very dilute solutions such as the concentration of salt in tap water or impurities/additives in a produce. Natural mineral water typical analysis (ppm) hydrogencarbonates 380 calcium 145 magnesium 25 chloride 20 sodium 10

8.4 Moles per Litre (mol L-1) The concentration of solution in moles per litre indicates the number of moles of solute in a litre of solution.

Concentration (mol L-1) = number of moles of solute (mol) Volume of solution (L)