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 H3O^+ (hydronium ion) in solution. These terms are interchangeable because an H^+ proton in water attaches to a water molecule to form H3O^+ (H2O + H^+ \rightarrow H3O^+). 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 (H2SO4)

    • 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 NH3 (ammonia) or CH4 (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., NH3 ammonia, or organic amines like NH2CH2CH3).

    • 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 (HSO4^-) 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 (SO4^{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: 2AgNO3(aq) + Cu(s) \rightarrow Cu(NO3)_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(NO3)2 splits into Cu^{2+}(aq) + 2NO_3^-(aq).

    • Example from above: 2Ag^+(aq) + 2NO3^-(aq) + Cu(s) \rightarrow Cu^{2+}(aq) + 2NO3^-(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