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9.3 The acidic environment 4. Definitions of acid and base

Syllabus reference (October 2002 version)
4. Because of the prevalence and importance of acids, they have been used and studied for hundreds of years. Over time, the definitions of acid and base have been refined	Students learn to: outline the historical development of ideas about acids including those of: Lavoisier Davy Arrhenius outline the Brönsted-Lowry theory of acids and bases describe the relationship between an acid and its conjugate base and a base and its conjugate acid identify a range of salts which form acidic, basic or neutral solutions and explain their acidic, neutral or basic nature identify conjugate acid/base pairs identify amphiprotic substances and construct equations to describe their behaviour in acidic and basic solutions identify neutralisation as a proton transfer reaction which is exothermic describe the correct technique for conducting titrations and preparation of standard solutions qualitatively describe the effect of buffers with reference to a specific example in a natural system	Students: gather and process information from secondary sources to trace developments in understanding and describing acid/base reactions choose equipment and perform a first-hand investigation to identify the pH of a range of salt solutions perform a first-hand investigation and solve problems using titrations and including the preparation of standard solutions, and use available evidence to quantitatively and qualitatively describe the reaction between selected acids and bases perform a first-hand investigation to determine the concentration of a domestic acidic substance using computer based technologies analyse information from secondary sources to assess the use of neutralisation reactions as a safety measure or to minimise damage in accidents or chemical spills

Extract from Chemistry Stage 6 Syllabus (Amended October 2002). © Board of Studies, NSW.
[Edit: 27 Jun 08]

gather and process information from secondary sources to trace developments in understanding and describing acid/base reactions

• This is an opportunity for you to gather relevant information to summarise the developments in understanding and describing acid-base reactions by Arrhenius and Bronsted-Lowry. The information required is readily available in HSC chemistry level texts.
  
  
• Process the information by accessing a number of sources. When you have identified a source of information, evaluate its validity by checking the reputation of the source and by looking to see how the information compares to information from other sources. You might use a table like that following to record and compare information.

Scientist(s)	Acid definition	Base definition	Notes
Arrhenius	?	?	Water solutions only
Bronsted-Lowry	?	?	Acid must contain hydrogen

outline the historical development of ideas about acids including those of:

• Lavoisier
• Davy
• Arrhenius

• In the 1780s, the French chemist, Antoine Lavoisier, found that non-metal oxides reacted with water forming acidic solutions. He concluded that an acid must contain oxygen.
  
  
• In 1815, the English chemist, Humphry Davy, observed that all known acids contained hydrogen that could be replaced by reaction with a metal. He also noted that compounds of metal with oxygen were bases.
  
  
• Lavoisier and Davy's definitions were based on observable properties. In 1884, the Swedish chemist, Svante Arrhenius, put forward definitions based on concepts about particles too small to be directly observed. Arrhenius proposed that:

  An acid produced hydrogen ions H⁺ when dissolved in water. A base produced hydroxide ions OH^- when dissolved in water.

• The Arrhenius definitions are the ones generally used in junior high school (stages 4 and 5 Science).

outline the Brönsted-Lowry theory of acids and bases

• A theory, based on proton transfer, was independently outlined in 1923 by the Danish chemist, Johannes Bronsted, and the British chemist, Thomas Lowry. An acid is a proton donor and a base is a proton acceptor.
  
  
• An acid-base reaction involves proton transfer from acid to base.
  
  
• The Bronsted-Lowry theory is used to explain acids and bases in Stage 6 science courses.

describe the relationship between an acid and its conjugate base and a base and its conjugate acid

• When an acid donates a proton, it forms its conjugate base.
       HCl     +     H₂O                Cl^-           +     H₃O⁺
       acid                           conjugate base
  
  
• When a base accepts a proton, it forms its conjugate acid.
       HCl     +     H₂O         Cl^-     +     H₃O⁺
                       base                   conjugate acid

identify conjugate acid/base pairs

• Whenever an acid and a base react, they form their conjugates:
       HCl     +     H₂O             Cl^-          +         H₃O⁺
      acid₁          base₂       conjugate base₁    conjugate acid₂
  
  
• Hydrochloric acid and chloride ion are a conjugate acid-base pair.
  
  
• Water and hydronium ion are another conjugate acid-base pair.

identify neutralisation as a proton transfer reaction which is exothermic

• In junior high school, neutralisation is studied as the reaction between an acid and a base to form a salt and water. The solutions reacted to demonstrate neutralisation are usually of a strong acid, such as hydrochloric acid, and a strong base, such as sodium hydroxide.
  
  acid      +          base     salt      +     water
  
  HCl        +        NaOH      NaCl     +     H₂O
  
  H⁺ + Cl^- +  Na⁺ + OH^-    Na⁺ + Cl^- +     H₂O
  
  
• The net ionic equation for reaction is:
  
   H⁺       +         OH^-         H₂O
  
  
• The net ionic equation shows that neutralisation is a proton transfer reaction. A proton from the acid transfers to the hydoxide ion of the base.
  
  
• All neutralisations are exothermic. When the heat of neutralisation is measured for a range of strong acids and strong bases, the amount of heat released is always about 57 kJ per mole of water formed. This is the heat change for the following reaction:
  
  H⁺ + OH^- H₂O     DH = - 57 kJ mol^-1

perform a first-hand investigation and solve problems using titrations and including the preparation of standard solutions, and use available evidence to quantitatively and qualitatively describe the reaction between selected acids and bases

• There are many successful procedures readily available to describe how to perform titrations, including how to prepare of standard solutions. When performing the investigation, be careful to minimise hazards and wastage of resources. The notes that follow the next syllabus point provide you with a description of the issues that must be addressed.
  
  
• Use identified problem-solving strategies to develop quantitative descriptions of the reactions you study. Some of these are described in the notes for the next syllabus point. You should be able to match your quantitative descriptions to the qualitative observations you have made.
  
  
• When using evidence from your titrations, propose ideas that demonstrate coherence and logical progression and include correct use of scientific principles and ideas.
  • A titration can be carried out between a solid water soluble acid or base and a standard solution. For example an aspirin (acetylsalicylic acid) tablet can be titrated against a solution of NaOH using phenolphthalein indicator. Commercially bought tablets have labels indicating the amount of monoprotic acetylsalicylic acid in each tablet. Before titrating, the aspirin tablet needs to be crushed in about 10 mL of ethanol or methylated spirits.
    
    
  • The first step in the calculations for a neutralisation reaction is to write out a balanced equation for the reaction in this format:

            aAcid + bBase salt + water
    e.g. H₂SO₄ + 2NaOH  Na₂SO₄ + 2H₂O

    Here, a = 1 and b = 2

  • For the two solutions that have been used, the unknown acid concentration, c_a , or unknown base concentration, c_b , can be calculated using the relationship:

    

    where c = molar (moles per litre) concentration and v = volume (v_a and v_b must be in the same units).

    e.g. if 25.0 mL of 0.0124 M NaOH solution reacts with 15.6 mL of H₂SO₄ solution, the calculation is:

    c_a =     = 0.00994

    If one of the reactants is a solid, then convert the mass of solid reactant in grams to moles. Replace c_av_a or c_bv_b with the number of moles. In this case, the volume of the other reactant must be in L because cv is in moles. When c is in moles per litre, v is in litres.

    e.g. if 37.9 mL of sulfuric acid solution is required to neutralise 1.56 g of CaCO₃ , the calculation is:

    CaCO₃ + H₂SO₄ CaSO₄ + H₂O + CO₂
    1.56 g CaCO₃ = = 0.0156 mole

    = 0.0156

    c_a = = 0.412 M 

    The amount of product can also be calculated using the mole concept.

• For instance, as we can see from the previous example, the amount of CO₂ product is 0.0156 mole because each mole of CaCO₃ produces one mole of CO₂.
  0.0156 mole of CO₂ = 0.0156 mol x 44.0 g mol^-1 = 0.686 g
  0.0156 mole of CO₂ = 0.0156 x 24.8 L of gas at 100 kPa and 298 K = 0.387 L

describe the correct technique for conducting titrations and preparation of standard solutions

• A solution of accurately known concentration is called a standard solution.
  
  
• For a chemical to be suitable to prepare as a standard solution, it must:
  1. be a water soluble solid
  2. have high purity - usually Analytical Reagent (A.R.) grade
  3. have an accurately known formula
  4. be stable in air, i.e. it does not lose or gain water or react with oxygen or carbon dioxide in air.
     
     
• The solution is prepared by:
  1. accurately weighing a calculated amount of solid
  2. dissolving it in water
  3. transferring all of the dissolved solid to a volumetric flask
  4. adding water to the flask to prepare a fixed volume of solution.

  The concentration is usually calculated in mol L^-1.

• A standard solution can be reacted with a solution of unknown concentration using titration technique. One reactant in solution is slowly added to another reactant in solution until an end point is reached.
  
  
• The end point of the titration is usually indicated by a change in colour of a small amount of indicator solution added to the mixture of reactants. For an acid-base titration an indicator is selected that changes colour at the pH of the salt solution formed at the point of neutralisation. This is known as the equivalence point.
  
  
• At senior high school level equipment such as burettes, pipettes and volumetric flasks give readings to three significant figures. Calculations are carried out to three significant figures.

choose equipment and perform a first-hand investigation to identify the pH of a range of salt solutions

• Choose the most appropriate equipment available to you. If possible, a pH meter or data logger with probe would be most suitable, otherwise select indicator solution or indicator paper to measure the pH of the salt solutions.
  
  
• To ensure you perform a valid investigation, prepare solutions of salts of equal concentration. Suitable salts are ammonium chloride (NH₄Cl), sodium chloride (NaCl), sodium carbonate (Na₂CO₃) and potassium acetate (KCH₃COO).

identify a range of salts which form acidic, basic or neutral solutions and explain their acidic, neutral or basic nature

• Salt ions formed from weak acids or weak bases can react with water to reform the acid or base. In undergoing these hydrolysis reactions, they release OH^- or H⁺, which can produce basic or acidic salt solutions.
  
  
• Ammonium salt solutions are acidic, because
  NH₄⁺ + H₂O NH₃ + H₃O⁺
  
  
• Sodium chloride solution is neutral, because Na⁺ and Cl^- (ions from the strong base NaOH and the strong acid HCl) do not undergo hydrolysis.
  
  
• Sodium carbonate solution is basic, because the carbonate ion from the weak acid carbonic acid can hydrolyse.
  CO₃^2- + H₂O HCO₃^- + OH^-
  
  
• Similarly, potassium acetate solution is basic.
  CH₃COO^- + H₂O CH₃COOH + OH^-
  
  
• If a salt is made up of two ions that hydrolyse to the same extent, the salt solution could be close to neutral, e.g. ammonium acetate NH₄CH₃COO.
  NH₄⁺ + H₂O NH₃ + H₃O⁺
  CH₃COO^- + H₂O CH₃COOH + OH^-
  The resulting reaction, H₃O⁺ + OH^- 2H₂O, results in a neutral solution.

identify amphiprotic substances and construct equations to describe their behaviour in acidic and basic solutions

• A molecule or ion that can behave as a proton donor or acceptor is called amphiprotic. Amphiprotic means protons on both sides. An amphiprotic molecule or ion can donate or accept a proton.
  
  
• Water is an amphiprotic molecule:

  Water as an acid: H₂O H⁺ + OH^- or more fully H₂O +H₂O H₃O⁺ + OH^-
  Water as a base: H⁺ + H₂O H₃O⁺ or more fully H₃O⁺+H₂O H₂O + H₃O⁺

• The hydrogen carbonate (bicarbonate) ion is an amphiprotic ion:

  As an acid HCO₃^- H⁺ + CO₃^2- or more fully HCO₃^-+H₂O H₃O⁺ + CO₃^2-
  As a base H⁺ + HCO₃^- H₂CO₃ or more fully H₃O⁺+ HCO₃^- H₂O +H₂CO₃

perform a first-hand investigation to determine the concentration of a domestic acidic substance using computer-based technologies

• Vinegar contains acetic acid which can be titrated against standardised NaOH(aq) using a pH probe attached to a data logger. The data recorded can be used to draw a graph. The endpoint is where the pH changes most rapidly.
  
  OR
  
  
• A standard solution of NaOH cannot be prepared by directly weighing out NaOH as the solid absorbs water and reacts with carbon dioxide in the air. A NaOH(aq) solution prepared in this way needs to be standardised by titration against a suitable acid, such as oxalic acid (COOH)₂.2H₂O.
  For the titration use a pH probe attached to a data logger to produce a graph that shows the endpoint where the pH changes most rapidly.

analyse information from secondary sources to assess the use of neutralisation reactions as a safety measure or to minimise damage in accidents or chemical spills

• A substance containing an amphiprotic ion, such as the hydrogen carbonate ion in NaHCO₃, is quite suitable for neutralising chemical spills.

  If the chemical spill contains an acid, H⁺ + HCO₃^- H₂O + CO₂
  If the spill contains a base, HCO₃^- + OH^- CO₃^2- + H₂O.

  Thus, NaHCO₃ is suitable for neutralising chemical spills of acids, bases and unknown acidity or basicity.

• Analyse information from secondary sources to suggest another amphiprotic substance suitable for neutralising a chemical spill, making sure you include the reason for your judgement.

qualitatively describe the effect of buffers with reference to a specific example in a natural system

• A buffer controls the level of acidity or basicity in a solution. If an acid or a base is added to a buffer solution, there is hardly any change in pH.
  
  
• A buffer solution is usually a mixture of a weak acid and its conjugate base, such as hydrogen carbonate ions, HCO₃^-, and carbonate ions, CO₃^2-.
  
  
• If an acid is added to the buffer, the hydrogen ions are removed by
  H⁺ + HCO₃^- H₂CO₃
  
  
• If a base is added to the buffer, hydroxide ions are removed by
  OH^- + HCO₃^- H₂O + CO₃^2-
  
  
• The net effect is that the pH of the solution containing buffer changes only slightly.
  
  
• Hydrogen carbonate ions are important in maintaining the pH of human blood at about 7.4.