SCH 3U - Strong and Weak Electrolytes

Weak & Strong Electrolytes

Electrolytes are substances that release ions into solution when dissolved in water (acids – H⁺; bases – OH^-).

Dissociation is the separation of a substance into its ions when dissolved in water.

NaCl(s)  →  Na⁺(aq)  +  Cl^-(aq)

Electrolyte solutions will conduct electricity due to the dissociated ions in solution.

Strong electrolytes (strong acid/base) are electrolytes that completely ionize in solution.

HCl(aq)  ⇒  H⁺(aq)  +  Cl^-(aq)     100% dissociation

• common strong acids: HCl(aq), HNO₃(aq), H₂SO₄(aq)
• common strong bases: group I & II metal hydroxides

Weak electrolytes (weak acid/base) are electrolytes that partially ionize in solution.

 HC₂H₃O₂(aq)  ↔  H⁺(aq)  +  C₂H₃O₂^-(aq)     1.3% dissociation (reversible rxn)

• common weak acids: HC₂H₃O₂(aq), H₂CO₃(aq)
• common weak bases: non-group I & II metal hydroxides

Concentration vs Strength

A concentrated electrolyte solution contains a large amount of solute in the solution.  A  dilute electrolyte solution contains a small amount of solute in the solution.

Monoprotic & Polyprotic Acids

HCl is a monoprotic acid because it has one ionizable hydrogen, H₂SO₄ is a diprotic acid and H₃PO₄ is a triprotic acid (any acid with more than one ionizable hydrogen is polyprotic).  

Polyprotic acids dissociate in a series of steps, losing one hydrogen ion per step, rather than losing all ionizable hydrogens in one fell swoop:

H₂SO₄(aq) → HSO₄^-(aq) + H⁺(aq)    strong acid

HSO₄^-(aq) ↔ SO₄^2-(aq) + H⁺(aq)      weak acid

The Autoionization of Water

Chemists realized that acidic solutions contain both H⁺ and OH^-, with [H⁺]＞[OH^-]. Conversely, for basic solutions [OH^-]＞[H⁺].

Chemists wondered if pure water contained OH^- and H⁺ ions as well.  Very sensitive equipment shows water is very slightly conductive, so the concentration of the ions must be very small.

 

This observation, combined with other evidence, led chemists to conclude that water can autoionize:

H₂O(l)  +  H₂O(l)  →  H₃O⁺(aq)  +  OH^-(aq)

or simply written, since a hydronium ion is just a hydrated hydrogen ion:

H₂O(l)  →  H⁺(aq)  +  OH^-(aq) 

Pure water is neutral because [H⁺] = [OH^-] = 1.0 × 10^-7 mol/L.

[H⁺] & [OH^-]

Recall that strong electrolytes ionize completely in solution.  This knowledge makes the following calculations quite easy.

ex. Calculate the concentration of hydrogen ions in a 0.25 mol/L solution of hydrochloric acid.

 

              1 HCl(aq)  ⇒  1 H⁺(aq)  +  1 Cl^-(aq)

Initial          0.25                 0                  0

Final         0                     0.25             0.25

∴ the concentration of hydrogen ions is 0.25 mol/L

ex. Calculate the concentration of all the ions in 1.0 M magnesium hydroxide solution.

            1 Mg(OH)₂(aq)  ⇒  1 Mg²⁺(aq)  +  2 OH^-(aq)

Initial          1.0                           0                    0

Final            0                           1.0                  2.0

∴ the concentration of the magnesium and hydroxide ions are 1.0 mol/L and 2.0 mol/L respectively.

In both of the above examples, initially the concentration of the dissociated ions is 0 mol/L because the ionization has not begun to occur. 

Finally, after the complete ionization has occurred, notice that there is no longer any undissociated electrolyte and only aqueous ions. 

Also notice that the amount of each dissociated ion is dependent on the stoichiometric relationships inherent to the balanced dissociation equation.

Homework:

Learning Check, p. 457 # 4

Learning Check, p. 462 # 12

Review Questions, p. 463 # 1, 5, 10, 16

 

 

 

Student Questions:

1.  Hey miss, can you do the questions on p.463?

Sure I can.  I'm really rather smart.