Salts

Salts

 
 Salts are crystalline Compounds composed of the negative ions of an acid and the positive ions of the base.
The reaction of a base and an acid to produce a salt and water is neutralization.
NaOH + HCl → NaCl + H2O

While the reaction above usually comes to mind when salt production is mentioned, there are many reactions that produce salts. It is even possible to have salts that do not produce neutral solutions.
Salts can also be formed by the reaction of an acidic or basic anhydride with a corresponding base, acid, or anhydride.

• acidic anhydride + base → salt

SO₃ + 2NaOH → Na₂SO₄ + H₂O

• basic anhydride + acid → salt

Na₂O + H₂SO₄ → Na₂SO₄ + H₂O

• basic anhydride + acidic anhydride → salt

Na₂O + SO₃ → Na₂SO₄

Certain acids and bases react to produce only a partial neutralization. These reactions produce either acidic salts or basic salts.

• This reaction produces an acidic salt:

H₂SO_{4 (aq)} + NaOH _(aq) → NaHSO_{4 (aq)} + H₂O _(l)

Sodium hydrogen sulfate is an acidic salt because it still contains an ionizable hydrogen atom.

 

Naming salts: the name of a salt is related to the name of the acid that forms it.

• Binary acids produce salts ending with -ide.
   
• Ternary acids ending in -ic produce salts ending with -ate.
   
• Ternary acids ending in -ous produce salts ending with -ite.
   
• Any prefixes in the ternary acid remain in the salt name.
   
• In naming acidic and basic salts, each ion in the salt is named separately.

 

• Hydrogen is named immediately before the names of any negative ions.

A prefix is used to indicate more than one hydrogen.
 

• Hydroxide is named immediately after the names of any positive ions.

The hydroxide is commonly placed in parenthesis.

Examples of Salt Names:

• CaCl₂ - calcium chloride
• K₂SO₄ - potassium sulfate
• NaHC₂O₄ - sodium hydrogen oxalate
• NaHS - sodium hydrogen sulfide
• NaH₂PO₄ - sodium dihydrogen phosphate
• Sn(OH)NO₃ - tin (II) hydroxide nitrate

 
Naming Salts

 
For about 95% of all compounds, solubility in water increases with increasing temperature. Many compounds can have their solubility in water increased or decreased by the presence of another solute.
Solubilities can be broken into four general classes:

• soluble - all of the material dissolves and does so fairly quickly.
   
• slightly soluble - some of the material visibly dissolves over a period of time.
   
• sparingly soluble - the materials has a very low solubility, such as 0.5 g per liter.
   
• Insoluble - none of the material dissolves.

 

Use these general solubility rules to predict the solubility of salts.

• Salts of group 1 and ammonia are soluble.
   
• Acetates and nitrates are soluble.
   
• Binary compounds of group 17, except F, are soluble with metals, except Ag, Hg⁺, and Pb.
   
• All sulfates are soluble, except those of Ba, Sr, Pb, Ca, Ag, and Hg⁺.
   
• Except for those in rule 1, carbonates, hydroxides, oxides, sulfides, and phosphates are insoluble.

 
Salt Solubility

 
Solubility Product Constant, K_sp

• Given this equilibrium equation:

AgBr_(cr) Ag⁺_(aq) + Br⁻_(aq)

 

• The equilibrium constant expression for the equation is:

 

• K_eq = [products] / [reactants]
   
• K_eq = [Ag⁺] [Br⁻] / [AgBr]
   
• Remember that the brackets, [ ], indicate "concentration".

 

• Since AgBr is a solid substance, its concentration is constant. The equilibrium constant expression can therefore be manipulated to read:

K_eq[AgBr] = [Ag⁺] [Br⁻]

 

• K_eq[AgBr] is the new constant, the solubility product constant, K_sp

If solid silver bromide is placed in water and allowed to stand, it dissolves until an equilibrium exists between the undissolved solid and the ions in solution.

• At room temperature, the K_sp of silver bromide is 5.01 X 10⁻¹³
• Using the equilibrium constant expression, [Ag⁺] [Br⁻] = 5.01 X 10⁻¹³
• The concentration of both ions are the same, [Ag⁺]² = 5.01 X 10⁻¹³
• The square root of 5.01X10⁻¹³ gives:

[Ag⁺] = 7.08 X 10⁻⁷M, which is also the [Br⁻]

 

Common Ion Effect:

The addition of a substance containing an ion already at equilibrium in a saturated solution will shift the equilibrium toward the un dissolved substance. Another way to say this is that the addition of a common ion decreases the solubility of a substance in solution.
 

 Net Ionic Equations

 
Net ionic equations attempt to show only the particles involved in a chemical reaction.

A balanced equation in "molecular" form:

2 AgNO_3(aq) + ZnCl_2(aq) → 2 AgCl_(cr) + Zn(NO₃)_2(aq)

The "net ionic" form of the same equation:

Ag ⁺_(aq) + Cl⁻_(aq) → AgCl_(cr)

To write net ionic equation, you must first decide - are substances written as molecules or ions?
Molecules or Ions? Slovene rules to help decide:
1. Binary acids:

• Strong acids are written in ionic form.

Examples: HCl, HBr, HI

• Weak acids are written in molecular form.

Examples: All other binary acids.

2. Ternary acids:

• Strong ternary acids are written as ions: the number of oxygen atoms exceeds the number of hydrogen atoms by two or more.

Examples: H₂SO₄, HNO₃

• Weak ternary acids are written as molecules.

Examples: H₃PO₄, HNO₂

3. Polyprotic acids:

• Those acids have more than one ionizable hydrogen.
   
• The second and all other ionizations are always weak.

Examples: H₂SO₄ is written in ionic form according to Rule #2.
One H is removed leaving HSO₄⁻.
Rule #3 assures us that this particle will not ionize farther.

4. Bases:

• Hydroxides of groups 1 and 2, except Be, are strong bases and are written in ionic form.
   
• All others are weak and written in molecular form.

5. Salts:

• Salts are written in molecular form if insoluble.
   
• Salts are written in ionic form if soluble.

Salt Solubility Rules:

1. Salts of group 1 and ammonium (NH₄⁺) are soluble.
2. Acetates and nitrates are soluble.
3. Binary compounds of group 17, except F, are soluble with metals, except Ag, Hg⁺, and Pb.
4. All sulfates are soluble, except those of Ba, Sr, Pb, Ca, Ag, and Hg⁺.
5. Except for those in Rule 1, carbonates, hydroxides, oxides, sulfides, and phosphates are insoluble.

6. Oxides:

• Oxides are always written in molecular form.

7. Gases:

• Gases are always written in molecular form.

   

Physical state symbols in an equation can help some with these rules.

For the molecular equation:

H₂PO_{4 (aq)} + MnCO_{3 (s)} → Mg₃(PO₄)_{2 (cr)} + CO_{2 (g)} + H₂O _(l)

• (aq) shows H₂PO₄ is dissolved in water.
• (s) shows MnCO₃ is a solid - insoluble in aqueous solution.
• (cr) shows Mg₃(PO₄)₂ is in crystal form - insoluble in aqueous solution.
• (g) shows CO₂ is a gas.
• (l) shows H₂O is a liquid.

The net ionic equation is:

H ⁺_(aq) + HPO₄⁻_(aq) + MnCO_{3 (s)} → Mg₃(PO₄)_{2 (cr)} + CO_{2 (g)} + H₂O _(l)

 

 
Spectator ions are ions that appear on both sides of the equation.

• They are assumed to take no part in the reaction, and are canceled in the process of writing net ionic equations.
   
• A proper net ionic equation has NO spectator ions in it.

 Step used to write net ionic equation:

1. Using the 7 rules, write each substance as molecules or ions.
2. Cancel ions that are exactly the same on both sides of the equation.

    These ions would cancel: Cu⁺ → Cu⁺
    These ions would not cancel: Cu⁺ → Cu⁺²

3. You must cancel the same number on both sides of the equation.

    Before canceling the spectator ions: 4 Fe⁺² → 6 Fe⁺²
    after canceling the spectator ions: no iron ions → 2 Fe⁺²

4. Cancel any other particles found on both sides of the equation.

5. Rewrite the equations in net ionic form.
6. If all coefficients are divisible by a common number, reduced them.