by iyell4

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Define acids and bases  according to the Bronsted–Lowry and Lewis theories.    

      Acid:  proton donor
      Base: proton acceptor


      Acid:  electron pair acceptor
      Base: electon pair donor
Deduce whether or not a species could act as a Brønsted–Lowry and/or a Lewis acid or base.
Bronsted-Lowry acids and bases must be able to accept or donate an acidic hydrogen (not one bonded directly to carbon).

Lewis acids and bases involve the complete donation of a pair of electrons by one species (the ligand) to form a coordinate covalent bond.
Deduce the formula of the conjugate acid (or base) of any Brønsted–Lowry base (or acid).
Students should make clear the location of the proton transferred, for example, CH3COOH/CH3COOrather than C2H4O2/C2H3O2.  
Outline the characteristic properties of acids and bases in aqueous solution.
Bases that are not hydroxides, such as ammonia, soluble carbonates and hydrogencarbonates, should be included. Alkalis are bases that dissolve in water. Students should consider the effects on indicators and the reactions of acids with bases, metals and carbonates.
Distinguish between strong and weak acids and bases in terms of the extent of dissociation, reaction with water and electrical conductivity.
Strong acids and bases completely dissociate and maximize the electrical conductivity of the solution.

Weak acids and bases only partially dissociate and are less conductive than the strong acids and bases.
State whether a given acid or base is strong or weak.
Students should consider hydrochloric acid, nitric acid and sulfuric acid as examples of strong acids, and carboxylic acids and carbonic acid (aqueous carbon dioxide) as weak acids. Students should consider all group 1 hydroxides and barium hydroxide as strong bases, and ammonia and amines as weak bases.
Distinguish between strong and weak acids and bases, and determine the relative strengths of acids and bases, using experimental data.
Besides comparing pH measurements to concentration values, titration curves can be examined to determine the strength of the acid or base.
Distinguish between aqueous solutions that are acidic, neutral, or alkaline using the pH scale.
 A pH below 7 is acidic, at 7 is neutral, and above 7 is basic. 
Identify which of two or more aqueous solutions is more acidic or alkaline using pH values.
Students should be familiar with the use of a pH meter and universal indicator.
State that each change of one pH unit represents a 10-fold change in the hydrogen ion concentration [H+(aq)].  
Relate integral values of pH to [H+(aq)] expressed as powers of 10. Calculation of pH from [H+(aq)] is not required.  
Deduce changes in [H+ (aq)] when the pH of a solution changes by more than one pH unit.    
This is a matter of using the pH-related equations to calculate a new concentration of hydrogen ions.
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