Electrolytes, Acids, Bases, and Ionic Equations

Electrolytes

• Definition: Substances that produce ions when dissolved in water, allowing for electrical conductivity.

• Strong Electrolytes:

  • Dissolve and ionize almost $100\%$ into products.

  • Represented by a forward arrow ($\rightarrow$) in chemical equations.

  • Lead to a relatively strong light in a conductivity test.

• Weak Electrolytes:

  • Dissolve and ionize less than $1\%$ into products.

  • Represented by an equilibrium or reversible arrow ($\rightleftharpoons$), indicating products can rejoin to form reactants.

  • Lead to a very dim light in a conductivity test.

• Non-Electrolytes:

  • Dissolve in water but do not form ions.

  • Do not conduct electricity (e.g., sugar in water, alcohol).

  • Pure water is considered a non-electrolyte for practical purposes, though technically a very, very weak electrolyte.

  • In a conductivity test, the light bulb does not turn on.

Strong Electrolyte Categories

Strong electrolytes fall into three main categories:

1. Acids: Defined as substances that produce $H^+$ or $H<em>3O^+$ (hydronium ion) in solution. These terms are interchangeable because an $H^+$ proton in water attaches to a water molecule to form $H</em>3O^+$ ($H<em>2O + H^+ \rightarrow H</em>3O^+$). Strong acids reduce the concentration of hydroxide ($OH^-$).

2. Bases: Do the opposite of acids; they decrease $H^+$ concentration and increase $OH^-$ concentration.

3. Salts: Soluble ionic compounds. Requires knowledge of solubility rules.

Acids: Definitions and Properties

• Proton Donors: Acids give an $H^+$ to something else, thereby increasing the $H^+$ concentration in water.

• Corrosive: They can burn on contact, especially with metals. Examples include hydrochloric acid.

• Taste: Tend to taste sour (e.g., Vitamin C is an acidic substance).

• pH: Have a low pH, specifically less than $7$ (pH < 7).

• Litmus Paper Test: Turn blue litmus paper red.

• Strong Acids vs. Weak Acids:

  • Strong Acids: Ionize $100\%$ with a forward arrow. The lack of a reversible reaction makes them strong.

    • Seven Strong Acids to Know:

      1. Hydrochloric Acid ($HCl$)

      2. Hydrobromic Acid ($HBr$)

      3. Hydroiodic Acid ($HI$)

      4. Chloric Acid ($HClO_3$)

      5. Perchloric Acid ($HClO_4$)

      6. Nitric Acid ($HNO_3$)

      7. Sulfuric Acid ($H<em>2SO</em>4$)

  • Weak Acids: Ionize less than $5\%$ with a reversible arrow ($\rightleftharpoons$). The reversible reaction makes them weak (e.g., acetic acid or vinegar).

• General Formula Characteristics:

  • Typically contain Group $6$ or Group $7$ anions (like chlorine, sulfur, bromine) with hydrogen, or hydrogen attached to polyatomic ions (like nitrate, sulfate, phosphate).

  • Often start their formula with hydrogen (e.g., $HCN$) or contain a $COOH$ group (e.g., $CH_3COOH$ for acetic acid).

  • Molecules like $NH<em>3$ (ammonia) or $CH</em>4$ (methane) are generally not acids, even though they contain hydrogen, as hydrogen loss from C-H, B-H, or N-H bonds is uncommon unless nitrogen is positively charged.

Bases: Definitions and Properties

• Hydroxide Donors or Proton Acceptors: Bases increase hydroxide ($OH^-$) and decrease hydrogen ($H^+$) concentrations.

• Feel: Feel slippery because they react with oils in the skin to create soap-like substances.

• Corrosive: Can also be corrosive.

• Taste: Tend to taste bitter. This dislike is a likely defense mechanism against naturally occurring poisons (organic alkaloids/bases like morphine, nicotine, caffeine).

• pH: Have a high pH, specifically above $7$ (pH > 7).

• Litmus Paper Test: Turn red litmus paper blue.

• Strong Bases vs. Weak Bases:

  • Strong Bases: Any soluble hydroxides. Solubility rules for hydroxide determine if a base is strong.

  • Weak Bases: Insoluble hydroxides.

• General Formula Characteristics:

  • Tend to have negative charges on an electronegative atom (like oxygen or sulfur) or contain a neutral or negatively charged nitrogen (e.g., $NH<em>3$ ammonia, or organic amines like $NH</em>2CH<em>2CH</em>3$).

  • Hydroxide is the strongest base observed in water.

Amphoteric and Amphiprotic Substances

• Definition: These substances can act as both acids and bases. They often possess both a hydrogen atom and a negative charge.

• Examples: Hydrogen carbonate ($HCO_3^-$) can gain an $H$ to form carbonic acid or lose an $H$ to form carbonate.

• Major Exception: The bisulfate ion ($HSO<em>4^-$) only acts as an acid. It cannot go back to sulfuric acid (a strong acid) by accepting an $H^+$; it can only further dissociate to sulfate ($SO</em>4^{2-}$) and a proton.

Importance of Identifying Strong Electrolytes

• Understanding which substances are strong electrolytes is crucial because they are the only category (among strong, weak, and non-electrolytes) that exists mostly as ions in solution.

• Weak electrolytes are mostly neutral molecules that split up only a little bit to make ions.

• Non-electrolytes are just molecules.

• This identification is the main focus for writing net ionic equations.

Types of Chemical Equations

There are three ways to write a chemical equation:

1. Formula Unit (Molecular) Equation: This is the standard, normal way chemical equations are typically written, showing compounds as entire neutral units.

   • Example: $2AgNO<em>3(aq) + Cu(s) \rightarrow Cu(NO</em>3)_2(aq) + 2Ag(s)$

2. Total (Complete) Ionic Equation: This equation separates all strong electrolytes into their individual ions, showing how compounds mostly exist in solution.

   • Rules for Splitting:

     1. Only Strong Electrolytes Split: Weak electrolytes, non-electrolytes, and insoluble compounds (solids, liquids, gases) are written as intact molecules.

     2. Conditions for Splitting: A substance must be aqueous ($(aq)$) AND a strong electrolyte (one of the 7 strong acids, a soluble hydroxide, or any soluble ionic compound).

     3. Coefficients: A coefficient in front of a compound distributes to all ions formed from that compound.

     4. Subscripts: A subscript within a compound that applies to an ion or polyatomic ion becomes a coefficient for that ion in the ionic equation.

        • Example: $Cu(NO<em>3)</em>2$ splits into $Cu^{2+}(aq) + 2NO_3^-(aq)$.

   • Example from above: $2Ag^+(aq) + 2NO<em>3^-(aq) + Cu(s) \rightarrow Cu^{2+}(aq) + 2NO</em>3^-(aq) + 2Ag(s)$

3. Net Ionic Equation: Derived from the total ionic equation by removing spectator ions.

Spectator Ions

• Definition: Ions that appear in identical form on both the reactant and product sides of a total ionic equation.

• Role: They do not actively participate in the chemical reaction; they merely serve as counter ions to maintain charge neutrality.

• Removal: Spectator ions are cancelled out to obtain the net ionic equation.

  • Example from above: $2Ag^+(aq) + Cu(s) \rightarrow Cu^{2+}(aq) + 2Ag(s)$

• Purpose of Net Ionic Equation: To show only the actual chemistry occurring in the reaction.

Types of Reactions

1. Acid-Base Reactions: Reaction between an acid and a base. Can be double replacement or combination reactions.

2. Precipitation Reactions:

   • Always double replacement reactions.

   • Identified by starting with two aqueous ionic compounds (both strong electrolytes).

   • Occur when the cations and anions from the initial compounds can combine to form an insoluble ionic compound, which solidifies as a precipitate.

   • Determined using solubility rules.

   • If both products formed are soluble, then no reaction occurs.

   • Example: Mixing potassium iodide and lead nitrate solutions. Lead iodide forms a yellow solid precipitate, while nitrate and potassium ions remain dissolved.

   • The term