Acid and bases

Acids and Base reactions

Acids are substances that taste sour. eg. citric acid. Bases are substances that taste bitter. eg. Bi-carb of soda. There are many different types of acids and bases and most are too corrosive to taste. The word corrosive means it will eat away. There must be another way to determine whether a chemical substance is acidic or basic.

Indicators are chemicals that change color in the presence of an acid or a base. Pink litmus paper will change to blue in the presence of a base. Blue litmus paper will change to pink in the presence of an acid.

Acids and bases react with each other to produce water and a salt. This type of reaction is called a neutralization reaction. Acid + Base => Salt + Water

An alkali is a strong soluble base. All of the Group I metals or alkali metals produce alkali solutions on reaction with water. eg. lithium hydroixde, sodium hydroxide and potassium hydroxide.

A common neutralization reaction carried out in high school is the neutralization of hydrochloric acid with sodium hydroxide.

Hydrochloric acid + sodium hydroxide => Sodium chloride + Water

2HCl + NaOH => NaCl + H₂O

Bromothymol blue is an excellent indicator to add as it is yellow in acid, blue in base and green when neutral. The solution can be evaporated in a watch glass which is placed on top of a water bath. Salt crystals appear as the water evaporates.

All neutralization reactions are also exothermic meaning heat is given off. Hydrogen ions produced from the acid react with hydroxide ions produced from the base. Energy is released as new chemical bonds are formed as water is produced.

Hydrogen ion + hydroxide ion => water

H⁺(aq)+ OH^-(aq) => H₂O(l)

 

Acids and carbonates

Acid on carbonate

Carbonates are compounds that contain the carbonate or CO₃ group. Eg. Bi carb of soda, NaHCO₃

All carbonates react with acids producing carbon dioxide water and salt.

 

Acids and metals

Acid on metal

Acids react with active metals like zinc, magnesium and iron to produce hydrogen gas and salt.

 

Chemical change

Chemical change

A chemical change is a sign that indicates that a chemical reaction has occurred.

Chemical changes include

• The production of a gas. Eg. Fizz from sherbet
• A permanent color change. Eg. A red solution changing to a colorless solution
• The production of light. Eg. Glow in the dark light sticks which is an example of chemiluminesce
• The production of a flame. Eg. A Bunsen burner
• The dissolving of an insoluble solid. Eg. Magnesium dissolving in acid.
• The production of an insoluble solid. Eg. The production of a precipitate

These changes indicate that something new has been produced.

Physical changes on the other hand are easy to reverse and don't produce anything new. They only involve a change in the physical states (solid. liquid and gas). eg. the melting of an ice cube and steam being produced from a kettle.

 

Combustion

Combustion is the burning of a fuel with the production of energy or heat

Examples of combustion reactions

1. Cellular respiration

Cellular respiration is the process whereby cells in our bodies obtain energy. Glucose reacts with oxygen producing carbon dioxide and water with the release of energy.

Glucose + oxygen => carbon dioxide + water + energy

C₆H₁₂O₆ + 6O₂ => 6CO₂ + 6H₂O + energy

 

2. The Combustion of fossil fuels

Fossil fuels are made up of hydrocarbons. Hydrocarbons are compounds that contain the elements hydrogen and carbon. Examples are methane, CH₄, propane, C₃H₈, butane, C₄H₁₀ and octane, C₈H₁₈

Hydrocarbons readily burn or undergo combustion reactions.Combustion may be complete or incomplete.

Complete combustion of fossil fuels results the production carbon dioxide and water. The incomplete combustion of fossil fuels results in the formation of carbon monoxide and soot (carbon).

1. Complete combustion of methane: Burns with a blue flame

Methane + oxygen => carbon dioxide and water

CH4 + 2O2 => CO₂ + 2H₂O

2. Incomplete combustion of methane producing carbon monoxide

Methane + oxygen => carbon monoxide and water

CH₄ + 1.5O₂ => CO + 2H₂O

3. Incomplete combustion of methane producing soot: Burns with a yellow flame

Methane + oxygen => carbon and water

CH₄ + O₂ => C + 2H₂O

As the amount of oxygen is decreased from 2 to 1.5 to 1 in the above three equations the combustion products respectively move from carbon dioxide to carbon monoxide to carbon or soot. As the soot is heated in the flame it produces a yellow flame. In real life there is a mixture of carbon monoxide and soot from the incomplete combustion of hydrocarbons as well as unburnt hydrocarbons.

 

Corrosion

Corrosion reactions

Corrosion is the loss of metallic properties of a metal as the metal reacts with the atmosphere or water. e.g. strength, lustre or shine and electrical conductivity.

Rust is brownish red in color and is formed from the corrosion of iron. Other metals like copper and aluminium also corrode or weaken.

For corrosion to occur both water and oxygen need to be present.

Examining the corrosion of iron using nails:

Oxygenated water (shaken)	Humid air	Dry air  (anhydrous CaCl2)	Deoxygenated water (Boiled water/ oil layer)	Salt water
Rust	Rust	No rust	No rust	Extensive rust

If a nail is placed in deoxygenated water it will not rust. If a nail is placed in dry air it will not rust. For rust to occur both oxygen and water need to be present. The presence of salt accelerates or speeds up the rate of corrosion.  

The chemical formula of rust is Fe₂O₃.nH₂O

The process of rust formation is due to a series of chemical reactions. Points of stress acts in the nail allow the iron to be easily oxidised. Carbon impurities in the nail allow water to be reduced. An internal battery or galvanic cell is set up.

Oxidation of iron at points of stress in the crystal lattice:  2Fe(s) ==> 2Fe²⁺(aq) + 4e^-
Reduction of water at the site of carbon impurities: O₂(g) + H₂O(l) + 4e^- ==>4OH^-(aq)  

Overall equation: 2Fe(s) + O₂(g) + H₂O(l) ==>Fe(OH)₂

The iron(II) hydroxide is converted to rust through a serious of reactions.

The ion(II) hydroixde firstly oxides to iron(III) oxide.
1. Fe(OH)₂(s) =oxidation-=> Fe(OH)₃

The iron(III) oxide then changes to rust through a dehydration reaction.
2. Fe(OH)₃(s) =dehydration=> Fe₂O₃.nH₂O(s) or rust

Rust does adheres loosely to the surface of the metal. This exposes the metal to more and more water and oxygen allowing rust to continue to proceed.

 

 

Decomposition

Decomposition reactions

In a decomposition reaction a chemical substance is broken down into simpler substances.

The general equation for the reaction is A => B + C

Substance A decomposes or breaks apart producing substances B and C.

Decomposition can be achieved by 1. heat, 2. electricity and 3. a catalyst.

Note: Some extremely unstable chemicals can be set of by sound waves!

Examples of decomposition reactions

1. Thermal decomposition of copper carbonate.

If a test tube containing green copper carbonate is heated in a Bunsen burner the powder turns black and occasionally ‘jumps’ as carbon dioxide gas is given off.

Copper carbonate =heat=> Copper oxide + carbon dioxide

CuCO₃ =heat=> CuO + CO₂

2. The electrolytic decomposition of water

Electrolysis is the process using electrical energy to speed up or cause a chemical reaction.

Pure water is a poor electrical conductor. Sodium sulfate salt or sulfuric acid needs to be added to the water first. Note: DO NOT use sodium chloride salt or hydrochloric acid as toxic chlorine gas is produced!

Two inert electrodes (carbon or platinum) are placed into the water and are connected to an external D.C power pack or series of batteries. Hydrogen gas is produced at the negative electrode and oxygen gas at the positive electrode in a ratio of 2:1.

Water =electrolysis=> Hydrogen and Oxygen

2H₂O =electrolysis=> 2H₂ + O₂

3. The catalytic decomposition of hydrogen peroxide

A catalyst is a substance that speeds up a chemical reaction without being used up in the reaction.

Hydrogen peroxide is broken done quite rapidly into oxygen gas and water when a catalyst like potassium iodide or manganese dioxide is added. Potassium iodide salt is preferred if available as it makes less mess. The black manganese dioxide powder is difficult to clean up.

One to two spatulas of potassium iodide salt is added to a conical flask containing 50mL of 5% hydrogen peroxide solution. The solution is gently mixed and bubbles of gas are observed to form. If a glowing splint is placed into the conical flask it will re-ignite due to the production of oxygen gas.

Hydrogen peroxide =catalyst=> Oxygen and water

2H₂O₂ =catalyst=> O₂ + 2H₂O

 

Displacement

Displacement reactions

Metals differ in their reactivities. Magnesium is more reactive than zinc and zinc is more reactive than iron.

Metals can be placed in order of decreasing chemical activity. This is called the activity series of metals.

Activity series of metals

Metal ion	Metal	Symbol
Mg2+	Magnesium	Mg	Most active metal
Zn2+	Zinc	Na
Fe2+	Iron	Fe
Pb2+	Lead	Pb
Cu2+	Copper	Cu
Ag+	Silver	Ag	Least active metal

An active metal will displace the ion of a less active metal from solution.

A metal higher on the table will react with a metal ion solution lower on the table.

An example of a displacement reaction is placing a piece of steel wool into a solution of copper sulfate.

Note: Don't use the commercial brands of steel wool that are protected from rusting. They will not react.

The steel wool is turned into pink copper wool. Over time the solution also looses its blue color and turns olive green.

Iron (steel wool) + Copper sulfate => Iron sulfate + Copper

Fe +	CuSO4	=>	FeSO4 +	Cu
	Blue		Green	Pink

The reaction is also exothermic producing heat.

 

Precipitation

Precipitation reactions

In a precipitation reaction two solutions are mixed together resulting in the formation of an insoluble solid which is called the precipitate.

 

Redox

Reduction oxidation reactions

Reduction oxidation reactions are called redox reactions. They involve the transfer of electrons from one species to another and are represented by two half equations, the oxidation half equation and the reduction half equation.

Oxidation half equation: Eg. Mg ==> Mg²⁺ + 2e^-   OIL: Oxidation Is Loss of electrons by a chemical species. The Mg species looses electrons and is oxidised.

Reduction half equation: Eg. Cu²⁺ + 2e^- ==> Cu    RIG: Reduction Is Gain of electrons by a species.The Cu²⁺ ion gains electrons and is reduced.

Notice the electrons, symbol e^-, are on different sides of the two equations in equal numbers.

Oxidation reduction reactions always occur together. One chemical species looses electrons and the other species gains electrons.

The resulting net reaction is called a redox reaction. Eg. Mg + Cu²⁺ ==> Mg²⁺ + Cu

A galvanic cell is a device which separates these two half reactions allowing the electrons to flow through an external wire. We call this device a battery.

 

 

Synthesis

Synthesis reaction

The word synthesis means to put together. In a synthesis reaction two elements combine to produce a compound. A + B => C

The reaction of a metal with a non-metal to produce a compound is an example of a synthesis reaction.

Examples of synthesis reactions

1. The burning of magnesium metal

If a piece of magnesium ribbon is ignited a very bright light is given off and a white powder or ash is produced/ Note: the burning of magnesium produces UV light which may damage the eyes if looked upon directly. Wear safety goggles and do not directly look at the light.

Magnesium + oxygen => magnesium oxide

2Mg + O₂ => 2MgO

2. The reaction of sodium with chlorine

If a piece of sodium metal is melted in a deflagrating spoon and placed into a gas jar containing chlorine gas a violent reaction occurs. If the jar is observed crystals of sodium chloride salt are seen which have formed on the sides of the glass.

Sodium + chlorine => Sodium chloride

2Na + Cl₂ => 2NaCl

3. The reaction of alumnium with bromine

Aluminum is resistant towards reaction because of a tough impermeable oxide layer which strongly adheres to the metal. If however, a sheet of aluminum foil is scratched with steel wool the aluminum is exposed it will react violently with pure liquid bromine

Aluminum + bromine => aluminum bromide

2Al + 3Br₂ => 2AlBr₃

4. The burning of iron

A nail will not react when heated over a Bunsen burner. However, if the surface area of the iron is increased by using steel wool a reaction is seen to occur. Increasing the surface area increases the rate of reaction.

Iron + oxygen => Iron(II) oxide

2Fe + O₂ => 2FeO