Acids, Bases and Salts

All substances are acidic, neutral or basic (alkaline). How acidic or basic a substance is shown by its pH. There are several other ways by which we could find out whether a substance is acidic, neutral or basic.

pH Scale:

This is a scale that runs from 0 to 14. Substances with a pH below 7 are acidic. Substances with pH above 7 are basic. And those with pH 7 are neutral.


Indicators are substances that identify acidity or alkalinity of substances. They cannot be used in solid form.

Universal Indicator:

This is a substance that changes color when added to another substance depending on its pH. The indicator and the substance should be in aqueous form.

Litmus Paper or Solution:

This indicator is present in two colors: red and blue. We use blue litmus if we want test a substance for acidity. We use red litmus if we want to test a substance for alkalinity. Its results are:

Note: use damp litmus paper if testing gases.


This is an indicator that is used to test for alkalinity because it is colorless if used with an acidic or neutral substance and it is pink if it is used with a basic substance.

Methyl Orange:

This indicator gives fire colors: Red with acids, yellow with neutrals and orange with bases.



Acids are substances made of a hydrogen ion and non-metal ions.  They have the following properties:

All acids must be in aqueous form to be called an acid. For example Hydrochloric acid is hydrogen chloride gas dissolved in water. The most common acids are:

Strength of Acids:

One of the most important properties of acids is that it gives hydrogen ion H+ when dissolved in water. This is why the amount of H+ ions the acid can give when dissolved in water is what determines its strength. This is called ionization or dissociation. The more ionized the acid is the stronger it is, the lower its pH. The more H+ ions given when the acid is dissolved in water the more ionized the acid is.

Strong Acids:

  • Have pH’s: 0,1,2,3
  • They are fully ionized
  • When dissolved in water, they give large amounts of H+ ions
  • Examples:
  • Hydrochloric Acid
  • Sulfuric Acid
  • Nitric Acid
  • Weak Acids:

  • Have pH’s: 4,5,6
  • They are partially ionized
  • When dissolved in water, they give small amounts of H+ ions
  • Examples:
  • Ethanoic acid (CH3COOH)
  • Citric Acid
  • Carbonic Acid
  • Hydrochloric acid is a strong acid. When it is dissolved in water all HCl molecules are ionized into H+ and Cl-  ions. It is fully ionized.

    Ethanoic acid has the formula CH3COOH. It is a weak acid. When it is dissolved in water, only some of the CH3COOH molecules are ionized into CH3COO-  and H+ ions. It is partially ionized.

    Note: Acids with pH 3 or 4 can be considered moderate in strength.

    Solutions of strong acids are better conductors of electricity than solutions of weak acids. This is because they contain much more free mobile ions to carry the charge.

    Concentrated acids are not necessarily strong. The concentration of an acid only means the amount of molecules of the acid dissolved in water. Concentrated acids have a large amount of acid molecules dissolved in water. Dilute acids have a small amount of acid molecules dissolved in water. Concentration is not related to strength of the acids. Strong acids are still strong even if they are diluted. And weak acids are still weak even if they are concentrated.



    Bases are substances made of hydroxide OH- ions and a metal. Bases can be made of:

    Properties of bases:

    Some bases are water soluble and some bases are water insoluble. Water soluble bases are also called alkalis.

    Like acids, alkalis' strength is determined by its ability to be ionized into metal and hydroxide OH-  ions. Completely ionized alkalis are the strongest and partially ionized alkalis are the weakest. Ammonium hydroxide is one of the strongest alkalis while weak alkalis include the hydroxides of sodium, potassium and magnesium.


    Types of Oxides:

    Basic Oxides

  • They are metal oxides
  • They react with acids forming a salt and water
  • They are solids
  • They are insoluble in water except group 1 metal oxides.
  • They react with an acid forming salt and water
  • Examples: Na2O, CaO and CuO
  • Amphoteric Oxides

  • These are oxides of Aluminum, Zinc & Lead
  • They act as an acid when reacting with an alkali & vice versa
  • Their element’s hydroxides are amphoteric too
  • They produce salt and water when reacting with an acid or an alkali.
  • Acidic Oxides

  • They are all non-metal oxides except non-metal monoxides
  • They are gases
  • They react with an alkali to form salt and water
  • Note: metal monoxides are neutral oxides
  • Examples: CO2, NO2, SO2 (acidic oxides) & CO, NO,
    H2O (neutral oxides)


    A salt is a neutral ionic compound. Salts are one of the products of a reaction between an acid and a base. Salts are formed in reactions I n which the H+ ion from the acid is replaced by any other metal ion. Some salts are soluble in water and some are insoluble.

    Soluble Salts:

  • All Nitrates
  • All halides EXCEPT AgCl and PbCl2
  • All sulfates EXCEPT CaSO4, BASO4, PbSO4
  • All group 1 metals salts
  • All ammonium salts
  • Insoluble Salts:

  • Silver and lead chlorides (AgCl & PbCl2)
  • Calcium, barium and lead sulphates (CaSO4, BASO4, PbSO4)
  • All carbonates EXCEPT group 1 metals and ammonium carbonates

    Preparing Soluble Salts:

    Displacement Method (Excess Metal Method):

    Metal + Acid → Salt + Hydrogen

    Note: this type of method is suitable to for making salts of moderately reactive metals because highly reactive metals like K, Na and Ca will cause an explosion. This method is used with the MAZIT (Magnesium, Aluminum, Zinc, Iron and Tin) metals only.

    Example: set up an experiment to obtain magnesium chloride salt.

    Mg + 2HCl → MgCl2 + H2

    1. Add 100 cm3 of dilute hydrochloric acid to a beaker
    2. Add excess mass of powdered magnesium
    3. When the reaction is done, filter the mixture to get rid of excess magnesium (residue)
    4. The filtrate is magnesium chloride solution
    5. To obtain magnesium chloride powder, evaporate the solution till dryness
    6. To obtain magnesium chloride crystals, heat the solution while continuously dipping a glass rod in the solution
    7. When you observe crystals starting to form on the glass rod, turn heat off and leave the mixture to cool down slowly
    8. When the crystals are obtained, dry them between two filter papers

    Observations of this type of reactions:

    You know the reaction is over when:

    Proton Donor and Acceptor Theory:

    When an acid and a base react, water is formed. The acid gives away an H+ ion and the base accepts it to form water by bonding it with the OH- ion. A hydrogen ion is also called a proton this is why an acid can be called Proton Donor and a base can be called Proton Acceptor.


    Neutralization Method:

    Acis + Base → Salt + Water

    Note: This method is used to make salts of metals below hydrogen in the reactivity series. If the base is a metal oxide or metal hydroxide, the products will be salt and water only. If the base is a metal carbonate, the products will be salt, water and carbon dioxide.

    Type 1:

    Acid + Metal Oxide → Salt + Water

    To obtain copper sulfate salt given copper oxide and sulfuric acid:

    CuO + H2SO4 → CuSO4 + H2O

    Observations of this reaction:

    Type 2:

    Acid + Metal Hydroxide → Salt + Water

    to obtain sodium chloride crystals given sodium hydroxide and hydrochloric acid:

    HCl + NaOH → NaCl + H2O


    You know the reaction is over when:

    Type 3:

    Acid + Metal Carbonate → Salt + Water + Carbon Dioxide

    To obtain copper sulfate salt given copper carbonate and sulfuric acid:

    CuCO3 + H2SO4 → CuSO4 + H2O + CO2


    You know the reaction is finished when:

    Titration Method:

    This is a method to make a neutralization reaction between a base and an acid producing a salt without any excess. In this method, the experiment is preformed twice, the first time is to find the amounts of reactants to use, and the second experiment is the actual one.

    1st Experiment:

    • Add 50 cm3 of sodium hydroxide using a pipette to be accurate to flask
    • Add 5 drops of phenolphthalein indicator to the sodium hydroxide. The solution turns pink indicating presence of a base
    • Fill a burette to zero mark with hydrochloric acid
    • Add drops of the acid to conical flask
    • The pink color of the solution becomes lighter
    • When the solution turns colorless, stop adding the acid (End point: is the point at which every base molecule is neutralized by an acid molecule)
    • Record the amount of hydrochloric acid used and repeat the experiment without using the indicator
    • After the 2nd experiment, you will have a sodium chloride solution. Evaporate it till dryness to obtain powdered sodium chloride or crystalize it to obtain sodium chloride crystals


    Preparing Insoluble Salts:

    Precipitation Method:

    A precipitation reaction is a reaction between two soluble salts. The products of a precipitation reaction are two other salts, one of them is soluble and one is insoluble (precipitate).

    Example: To obtain barium sulfate salt given barium chloride and sodium sulfate:

    BaCl2 + Na2SO4 → BaSO4 + 2NaCl
    Ionic Equation: Ba2+ + SO42- → BaSO4


    You know the reaction is over when:


    Controlling Soil pH:

    If the pH of the soil goes below or above 7, it has to be neutralized using an acid or a base. If the pH of the soil goes below 7, calcium carbonate (lime stone) is used to neutralize it. The pH of the soil can be measured by taking a sample from the soil, crushing it, dissolving in water then measuring the pH of the solution.

    Colors of Salts:

    Salt Formula Solid In Solution
    Hydrated copper sulfate CuSO4.5H2O Blue crystals Blue
    Anhydrous copper sulfate CuSO4 White powder Blue
    Copper nitrate Cu(NO3)2 Blue crystals Blue
    Copper chloride CuCl2 Green Green
    Copper carbonate CuCO3 Green Insoluble
    Copper oxide CuO Black Insoluble
    Iron(II) salts E.g.: FeSO4, Fe(NO3)2 Pale green crystals Pale green
    Iron(III) salts E.g.: Fe(NO3)3 Reddish brown Reddish brown


    Tests for Gases:

    Gas Formula Tests
    Ammonia NH3 Turns damp red litmus paper blue
    Carbon dioxide CO2 Turns limewater milky
    Oxygen O2 Relights a glowing splint
    Hydrogen H2 ‘Pops’ with a lighted splint
    Chlorine Cl2 Bleaches damp litmus paper
    Nitrogen dioxide NO2 Turns damp blue litmus paper red
    Sulfur dioxide SO2 Turns acidified aqueous potassium dichromate(VI) from orange to green


    Tests for Anions:

    Anion Test Result
    Carbonate (CO32-) Add dilute acid Effervescence,
    carbon dioxide produced

    Chloride (Cl-)
    (in solution)

    Acidify with dilute nitric acid, then add
    aqueous silver nitrate
    White ppt.
    Iodide (I-)
    (in solution)
    Acidify with dilute nitric acid, then add
    aqueous silver nitrate
    Yellow ppt.
    Nitrate (NO3-)
    (in solution)
    Add aqueous sodium hydroxide, then
    aluminium foil; warm carefully
    Ammonia produced
    Sulfate (SO42-) Acidify, then add aqueous barium nitrate White ppt.


    Tests for aqueous cations:

    Cation Effect of aqueous sodium hydroxide Effect of aqueous ammonia
    Aluminium (Al3+) White ppt., soluble in excess giving a
    colourless solution
    White ppt., insoluble in excess
    Ammonium (NH4+) Ammonia produced on warming
    Calcium (Ca2+) White ppt., insoluble in excess No ppt. or very slight white ppt.
    Copper (Cu2+) Light blue ppt., insoluble in excess Light blue ppt., soluble in excess,
    giving a dark blue solution
    Iron(II) (Fe2+) Green ppt., insoluble in excess Green ppt., insoluble in excess
    Iron(III) (Fe3+) Red-brown ppt., insoluble in excess Red-brown ppt., insoluble in excess
    Zinc (Zn2+) White ppt., soluble in excess,
    giving a colourless solution
    White ppt., soluble in excess,
    giving a colourless solution