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Acids, Bases and Alkalis Cheat Sheet by


Acids are compounds which ionise­/di­sso­ciate in water to produce hydrogen ions (H+).
Bases are compounds that are metal oxides or hydroxides that react with an acid to give a salt and water only.
Alkalis are bases that ionise­/di­sso­ciate in water to produce hydroxide ions (OH-).

Examples of Acids & Bases

Chemical Formula
Chemical Formula
Hydroc­hloric Acid
Magnesium Oxide
Sulfuric Acid
Copper (II) Oxide
Nitric Acid
Sodium Hydroxide
Citric Acid
Potassium Hydroxide
Ethanoic Acid
Calcium Hydroxide
Lactic Acid
Aqueous Ammonia
Acids 1 to 3 are known as mineral / inorganic acids while Acids 4 to 6 are known as organic acids.

Bases 1 & 2 are insoluble bases while Bases 3 to 6 are soluble bases / alkalis.

Metal Reactivity Series

Types of Reactions

Metal + Acid Salt + Hydrogen Gas
Metal Carbonate + Acid Salt + Water + Carbon Dioxide
Metal Oxide + Acid Salt + Water
Metal Hydroxide + Acid Salt + Water
Base + Acid Salt + Water (Neutr­ali­sation)
Alkali + Acid Salt + Water (Neutr­ali­sation)
Alkali + Ammonium Salt Salt + Water + Ammonia Gas
Alkali + Salt Metal Hydroxide + Salt
Tests for Gases:
Hydrogen Gas - Exting­uishes a lighted splinter with a 'pop' sound.
Carbon Dioxide Gas - Released as efferv­esc­ence. Reacts with limewater to form a white precip­itate.
Ammonia Gas - Pungent odour. Turns red litmus paper blue.

Base / Alkali + Acid is an exothermic reaction.
Pb (s) + H2SO4 / HCl PbSO4 / PbCl2 + H2
Lead reacts slowly then stops. Salt forms on the surface of the lead. The salt formed is insoluble.

pH Scale

Acidic solutions have pH values < 7.
They contain more H+ ions and fewer OH- ions.
Neutral solutions have pH values = 7.
They contain equal amounts of H+ ions and OH- ions.
Alkaline solutions have pH values > 7.
They contain more OH- ions and fewer H+ ions.

Ionic Equations

1. Write a balanced chemical equation with state symbols.
2. Check which reactants and products can form ions in water. (Aqueous)
3. Split up these reactants and products into their respective ions.
4. Check for ions that appear in both LHS & RHS of the equation, these are spectator ions that can be removed from the equation.
5. For those reactants and products which are unable to form ions, do not split the compounds.
6. What is left will be the net ionic equation. The coeffi­cients must be in the lowest ratio.

Polyatomic Ions

Chemical Formula
Silver ion: Ag+
Zinc ion: Zn2+

Properties of Acids

1. Acids are corrosive.
2. Acids have a sour taste.
3. Acidic solutions conduct electr­icity. (Elect­rol­ytes)
4. Acids change the colour of indica­tors.

Litmus Paper: Blue to Red
Methyl Orange Solution: Orange to Red
Universal Indicator Paper: Orange to Red
Universal Indicator Solution: Green to Red

Properties of Alkalis

1. Alkalis have a soapy feeling and a bitter taste.
2. Alkaline solutions conduct electr­icity. (Elect­rol­ytes)
3. Alkalis change the colour of indica­tors.

Litmus Paper: Red to Blue
Methyl Orange Solution: Orange to Yellow
Universal Indicator Paper: Orange to Violet
Universal Indicator Solution: Green to Violet

Balancing Chemical Equations

Step 1: Write down the chemical equation.
Step 2: List down the atoms (or polyatomic ions) involved in both sides.
Step 3: Count the number of atoms on both sides.
Step 4: Compare both sides and change the coeffi­cients (not subscr­ipts) so that the atoms on the left side are equal to the atoms on the right side.
(Tip: Balance the Metals first, then the Non-Metals, and then the Oxygen atoms and Hydrogen atoms.)
Step 5: Double check both sides to make sure the atoms on both sides are equal.

Soluble Salts

All nitrates
Most sulfates
Lead sulfate, barium sulfate and calcium sulfate
Most chlorides, bromides and iodides
Silver chloride, silver bromide, silver iodide, lead chloride, lead bromide, lead iodide
Sodium carbonate, potassium carbonate, ammonium carbonate
Most other carbonates
Sodium hydroxide, potassium hydroxide, ammonium hydroxide
Most other hydroxides

Uses of Acids

Citric Acid
Used as a sour flavouring agent in food
Hydroc­hloric Acid
Used as a rust remover
Sulfuric Acid
Used in car batteries
Nitric Acid
Used in fertil­isers
Ethanoic Acid
Used as a food preser­vative
Carbonic Acid
Used in making soft drinks

Uses of Alkalis

Sodium Hydroxide
Used in making soap
Calcium Hydroxide
Used in making toothpaste and to reduce acidity in soil
Aqueous Ammonia
Used in making fertil­isers and as a bleaching agent
Aqueous Ammonia
Used in making fertil­isers and as a bleaching agent
Potassium Hydroxide
Used in electr­opl­ating and in making cement and plaster
Magnesium Hydroxide
Used as a detergent

Strength of Acids

Strong Acids
Weak Acids
Hydroc­hloric Acid
Citric Acid
Sulfuric Acid
Tartaric Acid
Nitric Acid
Ethanoic Acid
Strong Acids:
React very fast & vigorously
Ionise completely to produce large amounts of H+ ions

Weak Acids:
React slowly & less vigorously
Ionise partially to produce small amounts of H+ ions

Do not confuse the strength of an acid with the concen­tration of an acid. The strength tells you how many H+ ions are produced while the concen­tration tells you how much of an acid is dissolved in water.

Strength of Alkalis

Strong Alkalis
Weak Alkalis
Sodium Hydroxide
Aqueous Ammonia
Potassium Hydroxide
Calcium Hydroxide
Strong Alkalis ionise completely to produce large amounts of OH- ions.
Weak Alkalis ionise partially to produce small amounts of OH- ions.

How to Carry Out Titration

1. For solid samples, weigh the solid and dissolve in a known volume of solution (usually 100cm3).
2. Use a pipette to measure a known volume of the solution (e.g 10cm3) and empty into an Erlenmeyer flask.
3. Add a few drops of indicator into the solution.
4. Put the second chemical into a burette. This other solution will react with the synthe­sised chemical sample in the flask. Often the solution in the burette is an acid or alkali, and it must be of a precise, known concen­tra­tion.
5. Drop by drop, mix the chemical in the burette into the Erlenmeyer flask until the end point is reached. A colour change indicates the correct amount has been added to react completely with the chemical in the sample.
6. Take note of the volume of the solution added from the burette.


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