Acids and Bases

Acids

Bases

  • Tastes sour
  • Corrosive
  • Turns blue litmus paper red
  • Tastes bitter
  • Corrosive
  • Turns red limus paper blue
  • Feels slippery to touch
  • Be able to identify some examples of laboratory acids and bases

Acids:

  1. Hydrochloric Acid (HCl)
  2. Sulphuric Acid (H2SO4)
  3. Nitric Acid (HNO3)

Bases:

  1. Sodium Hydroxide (NaOH)
  2. Potassium Hydroxide (KOH)
  3. Ammonium Hydroxide (NH4OH)
  • Know what indicators are and how we can use them

Indicators are substances that change colour when in an acidic or basic solution.

Universal indicator:

  • Universal indicators are mixtures of several different indicators that change color over a range of pH values.
  • When added to a solution, the universal indicator changes color depending on the pH of the solution. It typically changes through a spectrum of colors from red (indicating acidity) to purple (neutral) to blue or green (indicating alkalinity).
  • Universal indicators are often used for rough estimations of pH in a wide range of solutions, from strongly acidic to strongly basic.

Litmus paper:

  • Litmus paper is a simple pH indicator made from paper treated with litmus dye, which is extracted from lichens.
  • Litmus paper turns red in acidic solutions (pH less than 7) and blue in basic solutions (pH greater than 7). It remains purple in neutral solutions.
  • Litmus paper is commonly used for quick tests to determine whether a solution is acidic or basic. It's especially useful for educational purposes and field testing.

pH probe:

  • pH probes are electronic devices equipped with a pH-sensitive electrode that measures the hydrogen ion concentration in a solution.
  • The pH probe generates a voltage proportional to the pH of the solution. This voltage is converted into a digital readout or displayed on a pH meter.
  • pH probes provide accurate and precise measurements of pH in a solution. They are widely used in laboratory settings, industrial processes, and environmental monitoring where precise pH measurements are required.
  • Know the meaning of pH and the range of the pH scale
  • pH stands for "power of hydrogen" or "potential of hydrogen."
  • It is a measure of the acidity or basicity of a solution.
  • pH quantifies the concentration of hydrogen ions (𝐻+) in a solution.
  • The pH scale ranges from 0 to 14, with 7 considered neutral.
  • pH values below 7 are acidic, the lower the pH the higher the acidity.
  • pH values above 7 are basic/alkaline, the higher the pH the more basic it is.
  • Pure water has a pH of 7, which is considered neutral. This means it has an equal concentration of 𝐻+ and hydroxide ions (π‘‚π»βˆ’)
  • Predict the products when acids react with metals, metal carbonates, and bases (metal oxides or hydroxides)
  • Acid + Metal β†’ Salt + Hydrogen

E.g: Zinc + Sulphuric acid β†’ Zinc Sulphate + Hydrogen

  • Acid + Metal carbonate β†’ Salt + Water + Carbon dioxide

E.g: Copper carbonate + Sulphuric acid β†’ Copper sulphate + carbon dioxide + water

  • Acid + Metal oxide/metal hydroxide (base) β†’ Salt + Water

E.g: Copper oxide + Sulphuric acid β†’ Copper sulphate + water

  • Be able to name the salts produced in acid reactions, given the names of the reacting acid and metal, metal carbonate, or base
  • Acid-Metal Reactions:

Name the salt using the metal's name and the non-metal part of the acid with "-ide".

Example: Hydrochloric acid + Sodium β†’ Sodium chloride.

  • Acid-Metal Carbonate Reactions:

Name the salt using the metal's name and the non-metal part of the acid with "-ate".

Example: Sulfuric acid + Calcium carbonate β†’ Calcium sulfate.

  • Acid-Base Reactions:

Name the salt using the metal from the base and the non-metal part of the acid with "-ate".

Example: Nitric acid + Potassium hydroxide β†’ Potassium nitrate.

  • Give balanced equations for simple acid-base reactions
  • Acid + Base β†’ Salt + Water
  • Ensure that the equation is balanced
  • E.g: HCl + NaOH β†’ NaCl + H2O
  • Describe how to test for H2 and CO2 gas

Hydrogen Gas (H2)

  • Method:

Collect the gas in a test tube by displacement of water or another suitable method. Insert a lighted splint into the test tube.

  • Observation:

A squeaky pop sound or a sudden burst of flame is observed.

  • Explanation:

Hydrogen gas rapidly reacts with oxygen in the air, producing water vapor and heat. The sudden combustion of hydrogen results in a distinctive sound or flame.

Carbon dioxide Gas (CO2)

  • Method:

Bubble the gas through a solution of limewater (πΆπ‘Ž(𝑂𝐻)2)

  • Observation:

Limewater turns cloudy or milky.

  • Explanation:

Carbon dioxide reacts with the calcium hydroxide in limewater to form insoluble calcium carbonate (πΆπ‘ŽπΆπ‘‚3), which precipitates out of the solution, causing the limewater to become cloudy.

  • Describe and explain the steps involved in making salts in the lab- using precipitation reactions or reactions between metals and acids/ soluble metal oxides and acids

1. Precipitation Reactions:

  • Choose soluble salts containing desired ions.
  • Mix solutions to form insoluble salt (precipitate).
  • Filter, wash, dry, and collect the precipitate.

2. Reactions Between Metals and Acids/Soluble Metal Oxides and Acids:

  • Select metal/metal oxide and acid.
  • React them to form salt and hydrogen gas (if metal).
  • Filter to remove impurities, evaporate, and crystallize the salt.
  • Appreciate the use of different definitions for acids- the Arrhenius and Bronsted-Lowry's definitions of an acid

The Arrhenius definition of an acid:

  • A compound that increases the concentration of hydrogen ion (H +) in aqueous solution.

Bronsted-Lowry's definition of an acid:

  • A substance that gives up or donates hydrogen ions during a chemical reaction
  • Explain the difference between a strong and weak acid or base.

Strong Acids:

  • Strong acids completely dissociate into their ions in water, producing a high concentration of hydrogen ions (H+)
  • High conductivity in solution due to the abundance of ions.
  • Low pH values (typically close to 0).
  • Strongly corrosive properties.

Weak Acids:

  • Weak acids only partially dissociate into their ions in water, resulting in a relatively low concentration of hydrogen ions (H+)
  • Lower conductivity in solution compared to strong acids.
  • pH values closer to neutral (typically above 3).
  • Less corrosive than strong acids.

Strong Bases:

  • Strong bases completely dissociate into their ions in water, producing a high concentration of hydroxide ions (OH-)
  • High conductivity in solution due to the abundance of ions.
  • High pH values (typically close to 14).
  • Strongly caustic and corrosive properties.

Weak Bases:

  • Weak bases only partially dissociate into their ions in water, resulting in a relatively low concentration of hydroxide ions (OH-)
  • Lower conductivity in solution compared to strong bases.
  • pH values closer to neutral (typically above 7).
  • Less caustic than strong bases.
  • Know what is meant by the molarity or concentration of a solution (in mol/dm3)

Molarity/concentration:

  • The concentration is a measure of the amount of solute that is present in the solution.
  • Be able to calculate the number of moles of a substance in a solution and its

concentration using n= c x V

  • Calculate concentrations of solutions given the mass of solutes dissolved