Reactions in Aqueous Solution
Introduction
Many chemical and biological reactions occur in water; such reactions take place in aqueous solution. Water's abundance and unique properties make it an excellent medium for dissolving substances and enabling reactions. Almost all reactions in aqueous solutions involve ions. The main types of reactions studied here are precipitation reactions, acid-base reactions and redox reactions. Before discussing these reaction types, we examine how ions behave in water and how aqueous solutions conduct electricity.
Ions in Aqueous Solution
Dissociation and the polar nature of water
The water molecule is polar: one end (near the oxygen) has higher electron density and is slightly negative (δ-), while the hydrogen ends are slightly positive (δ+). Because of this polarity, water molecules interact strongly with charged species and can surround and stabilise ions.
How ionic compounds dissolve
When an ionic solid such as sodium chloride (NaCl) is placed in water, the polar water molecules orient themselves about the ions. The negatively charged side of water surrounds cations, and the positively charged side surrounds anions. The water molecules pull ions out of the crystal lattice into solution. This process is called dissociation. Dissolving is a physical, generally reversible change: the solute can be recovered by evaporating the solvent.
Key definitions
Dissociation - A process in which ionic compounds separate into their component ions, usually in a reversible manner.
Dissolution - The process where ionic crystals break up into ions and become incorporated into water (they dissolve).
Hydration - The process by which ions become surrounded by water molecules.
Representative equations
NaCl(s) → Na+(aq) + Cl-(aq)
K2SO4(s) → 2 K+(aq) + SO42-(aq)
Glucose (molecular solute that does not ionise): C6H12O6(s) → C6H12O6(aq)
Hydrogen chloride (ionises in water): HCl(g) + H2O(l) → H3O+(aq) + Cl-(aq)
Exercise 18 -1
For each of the following, say whether the substance is ionic or molecular.
- potassium nitrate (KNO3)
- ethanol (C2H5OH)
- sucrose (a type of sugar) (C12H22O11)
- sodium bromide (NaBr)
Write a balanced equation to show how each of the following ionic compounds dissociate in water.
- sodium sulphate (Na2SO4)
- potassium bromide (KBr)
- potassium permanganate (KMnO4)
- sodium phosphate (Na3PO4)
Draw a diagram to show how KCl dissolves in water.
Electrolytes, Ionisation and Conductivity
Electrolytes and conductivity
An electrolyte is a substance that contains free ions and behaves as an electrically conductive medium. Solutions that contain many ions are called strong electrolytes and conduct electricity well. Solutions with few ions are weak electrolytes and conduct poorly. Non-electrolytes (for example many covalent molecular solutes such as sugar) do not produce ions and do not conduct electricity.
Conductivity is a measure of a solution's ability to carry electric current; it increases with the concentration of mobile ions and depends on the nature of the dissolved substance and the temperature.
Factors affecting conductivity
- Concentration of ions - higher ion concentration increases conductivity.
- Type of solute - whether it is a strong electrolyte (e.g. KNO3), weak electrolyte (e.g. CH3COOH) or non-electrolyte (e.g. sugar) affects the number of ions produced.
- Temperature - higher temperature generally increases solubility and ion mobility, so conductivity increases.
General experiment: Electrical conductivity (outline)
Aim: To investigate the electrical conductivities of different substances and solutions.
Apparatus: Solid salt crystals (e.g. NaCl); liquids such as distilled water, tap water, seawater, sugar solutions, oil, alcohol; aqueous solutions of salts (e.g. NaCl, KBr, CaCl2, NH4Cl); acid solution (e.g. HCl) and base solution (e.g. NaOH); torch cells or battery, ammeter, conducting wire, crocodile clips, two carbon rods.
Method (outline): Connect a simple circuit with an ammeter and place the test substance as the conducting element (solid or solution). For solids attach clips to opposite ends; for liquids use carbon rods immersed in the solution. Allow current to flow and observe the ammeter reading.
Results and explanation: Solutions that contain free-moving ions show an ammeter reading; pure distilled water, oil and alcohol (non-electrolytes) show little or no current. Solid ionic crystals do not conduct electricity because ions are fixed in the lattice. Dissolved salts dissociate into ions which can move and carry charge.
Representative dissociation/ionisation equations
KBr(s) → K+(aq) + Br-(aq)
NaCl(s) → Na+(aq) + Cl-(aq)
HCl(l) + H2O(l) → H3O+(aq) + Cl-(aq)
NaOH(s) → Na+(aq) + OH-(aq)
Precipitation Reactions
Definition
Precipitate - a solid that forms in a solution as the result of a chemical reaction. A precipitation reaction is any reaction in which an insoluble solid is formed from ions in solution.
General experiment: Reaction of ions in solution (outline)
Aim: To investigate reactions between aqueous ions and observe precipitate formation.
Apparatus: Test tubes; copper(II) chloride solution (dilute); sodium carbonate solution; sodium sulphate solution.
Method (outline): Mix sodium carbonate solution with copper(II) chloride solution in one test tube and observe. Mix sodium sulphate solution with copper(II) chloride in another test tube and observe.
Observations: A light-blue precipitate forms when sodium carbonate is added to copper(II) chloride. No precipitate forms with sodium sulphate; the solution remains light blue.
Solubility rules (general)
| Salt | Solubility |
|---|---|
| Nitrates | All are soluble |
| Potassium, sodium and ammonium salts | All are soluble |
| Chlorides, bromides and iodides | All are soluble except silver, lead(II) and mercury(II) salts |
| Sulphates | All are soluble except lead(II) sulphate, barium sulphate and calcium sulphate |
| Carbonates | All are insoluble except those of potassium, sodium and ammonium |
| Compounds with fluorine | Almost all are soluble except those of magnesium, calcium, strontium, barium and lead(II) |
| Perchlorates and acetates | All are soluble |
| Chlorates | All are soluble except potassium chlorate |
| Metal hydroxides and oxides | Most are insoluble |
Illustration using the experiment
In the reaction of CuCl2 with Na2CO3, the ions in solution are Cu2+, Cl-, Na+ and CO32-. Carbonates (except those of K+, Na+ and NH4+) are generally insoluble, so CuCO3 precipitates. The ionic equation can be represented as:
2 Na+(aq) + CO32-(aq) + Cu2+(aq) + 2 Cl-(aq) → CuCO3(s) + 2 Na+(aq) + 2 Cl-(aq)
In contrast, sulphates of many metals are soluble, so no precipitate forms when Na2SO4 is mixed with CuCl2; the ions remain in solution.
Tests for Anions
Qualitative tests rely on characteristic precipitates and secondary reactions to distinguish ions.
Test for chloride
Add silver nitrate solution to the aqueous sample. Formation of a white precipitate indicates either chloride or carbonate (AgCl or Ag2CO3).
Ag+(aq) + Cl-(aq) → AgCl(s) (white)
2 Ag+(aq) + CO32-(aq) → Ag2CO3(s) (white)
Treat the precipitate with dilute nitric acid: AgCl does not react (precipitate remains), while Ag2CO3 dissolves with evolution of CO2, indicating carbonate.
Test for bromide and iodide
Silver bromide and silver iodide also give precipitates when AgNO3 is added; these are typically pale yellow. To distinguish bromide from iodide, oxidising halogen (chlorine water) can liberate bromine or iodine from their salts; the liberated halogen colours carbon tetrachloride:
2 Br-(aq) + Cl2(aq) → 2 Cl-(aq) + Br2(g) (reddish-brown in CCl4)
2 I-(aq) + Cl2(aq) → 2 Cl-(aq) + I2(g) (purple in CCl4)
Test for sulphate
Add barium chloride solution to the sample. A white precipitate of barium sulphate indicates sulphate; carbonate would also give a white precipitate with Ba2+, so treat with dilute nitric acid to distinguish:
Ba2+(aq) + SO42-(aq) → BaSO4(s) (white)
BaCO3(s) + 2 HNO3(l) → Ba2+(aq) + 2 NO3-(aq) + H2O(l) + CO2(g) (carbonate dissolves with CO2 evolution)
Test for carbonate
Treat a sample of the dry salt with a small amount of acid; evolution of carbon dioxide indicates a carbonate:
2 HCl(aq) + K2CO3(aq) → CO2(g) + 2 KCl(aq) + H2O(l)
Passing the gas through limewater (Ca(OH)2(aq)) produces an insoluble calcium carbonate precipitate (CaCO3) making the limewater milky:
CO2(g) + Ca(OH)2(aq) → CaCO3(s) + H2O(l)
Exercise 18 -2
Silver nitrate (AgNO3) reacts with potassium chloride (KCl) and a white precipitate is formed.
- Write a balanced equation for the reaction that takes place. Include the state symbols.
- What is the name of the insoluble salt that forms?
- Which of the salts in this reaction are soluble?
Barium chloride reacts with sulphuric acid to produce barium sulphate and hydrochloric acid.
- Write a balanced equation for the reaction that takes place. Include the state symbols.
- Does a precipitate form during the reaction?
- Describe a test that could be used to test for the presence of barium sulphate in the products.
A test tube contains a clear, colourless salt solution. A few drops of silver nitrate solution are added to the solution and a pale yellow precipitate forms. Chlorine water and carbon tetrachloride were added, which resulted in a purple solution. Which one of the following salts was dissolved in the original solution? Write the balanced equation for the reaction that took place between the salt and silver nitrate.
- NaI
- KCl
- K2CO3
- Na2SO4
Other Types of Reactions in Aqueous Solution
Ion-exchange reactions
Ion-exchange reactions are those in which ions exchange partners. They can be represented generally as:
AB(aq) + CD(aq) → AD + CB
Either AD or CB may be an insoluble solid (precipitate) or a gas; when a precipitate forms the reaction is a precipitation reaction, and when a gas forms it is a gas-forming reaction. Acid-base reactions are a special class of ion-exchange reactions.
Definition: An ion-exchange reaction is one in which the positive ions exchange their respective negative ions due to a driving force (for example formation of a precipitate, gas, or weakly dissociated molecule).
Ion exchange processes are applied practically in ion-exchange chromatography and water softening.
Gas-forming reactions
These are ion-exchange reactions where a gaseous product is produced and escapes the solution, favouring the forward reaction. Example:
Na2CO3(s) + 2 HCl(aq) → CO2(g) + 2 NaCl(aq) + H2O(l)
Acid-base reactions (neutralisation)
Acid-base reactions generally produce water and an ionic salt. They are ion-exchange reactions in which a cation from the base pairs with an anion from the acid.
Example:
NaOH(aq) + HCl(aq) → NaCl(aq) + H2O(l)
Redox Reactions
Definition and identification
Redox (oxidation-reduction) reactions involve transfer of electrons between species. One species becomes more positive (oxidised) by losing electrons, and another becomes more negative (reduced) by gaining electrons. A change in oxidation state indicates that a redox reaction has occurred.
Example: Oxidation of sodium metal
Reaction of sodium metal with oxygen produces sodium oxide (and under certain conditions some sodium peroxide):
4 Na + O2 → 2 Na2O
In the products Na has oxidation state +1 and O has -2; therefore sodium has lost electrons (oxidised) and oxygen has gained electrons (reduced).
Demonstration: Oxidation of sodium metal (safety note)
Warning: Sodium metal is very reactive, especially with water. Handle only under qualified supervision and never allow sodium to contact water.
When a small piece of sodium is heated in air it reacts to form a white powder containing Na2O and Na2O2.
Classroom experiments and classification
Typical experiments to observe different reaction types include dissolving salts, mixing solutions to produce precipitates or gases, neutralising acids with bases using indicators (e.g. bromothymol blue) and displacement reactions (placing a more reactive metal in a solution of a less reactive metal's salt to observe metal deposition). By observing precipitate formation, gas evolution, colour change with indicators or metal displacement, reactions can be classified as precipitation, gas-forming, acid-base or redox.
General experiment: Reaction types (class activity)
Aim: To use experiments to determine what type of reaction occurs.
Apparatus: Soluble salts (e.g. potassium nitrate, ammonium chloride, sodium carbonate, silver nitrate, sodium bromide); hydrochloric acid (HCl); sodium hydroxide (NaOH); bromothymol blue; zinc metal; copper(II) sulphate; beakers; test tubes.
Method:
- For each salt, dissolve a small amount in water and observe what happens.
- Try dissolving pairs of salts (e.g. potassium nitrate and sodium carbonate) and observe any reaction.
- Dissolve some sodium carbonate in hydrochloric acid and observe the result.
- Carefully measure out 20 cm3 of sodium hydroxide into a beaker.
- Add a few drops of bromothymol blue to the sodium hydroxide.
- Carefully add a few drops of hydrochloric acid to the sodium hydroxide and swirl until you notice the colour change.
- Place zinc metal into copper sulphate solution and observe.
Questions to answer (class):
- What did you observe when you dissolved each salt in water?
- What did you observe when you dissolved pairs of salts together?
- What happened when sodium carbonate was dissolved in hydrochloric acid?
- Why was bromothymol blue used when mixing hydrochloric acid and sodium hydroxide?
- What happened when zinc metal was placed into copper sulphate solution?
- Classify each observed reaction as precipitation, gas-forming, acid-base or redox.
- What drives each reaction (for example, formation of a precipitate or a gas)?
- What criteria would you use to determine the kind of reaction that has occurred?
- Write balanced chemical equations for each reaction observed.
Conclusion: Reactions can be classified experimentally by characteristic observations: bubbles for gas formation, a solid forming for precipitation, indicator colour change for acid-base reactions, and metal deposition or disappearance for redox/displacement reactions.
Summary
- The polar nature of water causes ionic compounds to dissociate into component ions in aqueous solution.
- Dissociation is the reversible separation of ionic compounds into ions; dissolution describes ionic crystals breaking up into ions in water; hydration is the surrounding of ions by water molecules.
- Conductivity measures a solution's ability to carry electric current; solutions containing free ions (electrolytes) conduct electricity.
- Electrolytes are classified as strong or weak depending on the extent of ionisation; non-electrolytes do not ionise and do not conduct electricity.
- Conductivity depends on the type of substance, ion concentration and temperature.
- Main reaction types in aqueous solutions: precipitation reactions, acid-base reactions and redox reactions.
- Precipitation and acid-base reactions are examples of ion-exchange reactions; gas-forming reactions are also ion-exchange reactions.
- Solubility rules help predict when a precipitate will form; qualitative tests are used to identify common anions (chloride, bromide, iodide, carbonate, sulphate).
- Redox reactions involve transfer of electrons, indicated by a change in oxidation states of the species involved.
End of Chapter Exercises
1.
Give one word for each of the following descriptions:
- the change in phase of water from a gas to a liquid
- a charged atom
- a term used to describe the mineral content of water
- a gas that forms sulphuric acid when it reacts with water
2.
Match the information in column A with the information in column B by writing only the letter (A to I) next to the question number (1 to 7).
Column A
- A polar molecule
- Molecular solution
- Mineral that increases water hardness
- Substance that increases the hydrogen ion concentration
- A strong electrolyte
- A white precipitate
- A non-conductor of electricity
Column B
- H2SO4
- CaCO3
- NaOH
- salt water
- calcium
- carbon dioxide
- potassium nitrate
- sugar water
- O2
3.
Explain the difference between a weak electrolyte and a strong electrolyte. Give a generalised equation for each.
4.
What factors affect the conductivity of water? How do each of these affect the conductivity?
5.
For each of the following substances state whether they are molecular or ionic. If they are ionic, give a balanced reaction for the dissociation in water.
- methane (CH4)
- potassium bromide
