Comprehensive Notes on Acids and Bases

Acids and Bases

Chapter Overview

• Online Resources:
  • HMDScience.com for chemistry resources.
  • Online labs include: Is It an Acid or a Base?, Effects of Acid Rain on Plants.
• Chapter 14 Sections:
  • Section 1: Properties of Acids and Bases.
  • Section 2: Acid-Base Theories.
  • Section 3: Acid-Base Reactions.

Section 1: Properties of Acids and Bases

• Main Ideas:
  • Acids are identified by their properties.
  • Some acids are useful in industry.
  • The properties of bases differ from those of acids.
  • Arrhenius acids and bases produce ions in solution.
• Key Terms:
  • Binary acid
  • Arrhenius acid
  • Strong acid
  • Oxyacid
  • Arrhenius base
  • Weak acid

Common Acids and Bases

• Acids:
  • Citric acid (in citrus fruits): \text{H}3\text{C}6\text{H}5\text{O}7
  • Ascorbic acid (Vitamin C): \text{H}2\text{C}6\text{H}6\text{O}6
  • Benzoic acid: \text{HC}7\text{H}5\text{O}_2
  • Sorbic acid: \text{HC}6\text{H}7\text{O}_2
  • Phosphoric acid: \text{H}3\text{PO}4
  • Carbonic acid: \text{H}2\text{CO}3
  • Lactic acid (in sour milk)
  • Acetic acid (in vinegar)
• Bases:
  • Household ammonia (\text{NH}_3\text{(aq)}):
    • Useful for general cleaning.
  • Sodium hydroxide (NaOH, lye):
    • Present in some commercial cleaners.
  • Milk of magnesia (\text{Mg(OH)}_2):
    • Suspension in water, not very water-soluble.
    • Used as an antacid to relieve excess hydrochloric acid in the stomach.
  • Aluminum hydroxide (\text{Al(OH)}3) and sodium hydrogen carbonate (\text{NaHCO}3):
    • Also commonly found in antacids.

Properties of Acids

1. Sour Taste:
   • Aqueous solutions of acids have a sour taste.
   • Caution: Taste should NEVER be used as a test to evaluate any chemical substance due to the risk of corrosion and poisoning.
2. Acid-Base Indicators:
   • Acids change the color of acid-base indicators.
   • pH paper turns certain colors in acidic solution.
3. Reaction with Active Metals:
   • Acids react with active metals to release hydrogen gas (\text{H}_2).
   • Metals above hydrogen in the activity series undergo single-displacement reactions with certain acids.
   • Example: Barium with sulfuric acid:
     • \text{Ba(s)} + \text{H}2\text{SO}4\text{(aq)} \longrightarrow \text{BaSO}4\text{(s)} + \text{H}2\text{(g)}
4. Neutralization:
   • Acids react with bases to produce salts and water.
   • Properties of acids disappear when neutralized by bases.
   • Products are water and an ionic compound called a salt.
5. Electrical Conductivity:
   • Acids conduct electric current.
   • Strong acids completely separate into ions in water and are strong electrolytes.
   • Weak acids are weak electrolytes.

Acid Nomenclature

• Binary Acid:
  • An acid that contains only two different elements: hydrogen and one of the more electronegative elements.
  • Examples: Hydrogen halides (HF, HCl, HBr, HI).
  • Naming Rules:
    1. The name of a binary acid begins with the prefix hydro-.
    2. The root of the name of the second element follows this prefix.
    3. The name then ends with the suffix -ic.
• Oxyacid:
  • An acid that is a compound of hydrogen, oxygen, and a third element, usually a nonmetal.
  • Example: Nitric acid (\text{HNO}_3).
  • Considered one class of ternary acids (acids containing three different elements).
  • In oxyacids, H atoms are bonded to O atoms.
  • Naming based on the names of the anions.
    • Examples:
      • Acetic acid (\text{CH}3\text{COOH}) and acetate (\text{CH}3\text{COO}^-
      • Carbonic acid (\text{H}2\text{CO}3) and carbonate (\text{CO}_3^{2-}
      • Hypochlorous acid (\text{HClO}) and hypochlorite (\text{ClO}^-
      • Chlorous acid (\text{HClO}2) and chlorite (\text{ClO}2^-
      • Chloric acid (\text{HClO}3) and chlorate (\text{ClO}3^-
      • Perchloric acid (\text{HClO}4) and perchlorate (\text{ClO}4^-
      • Iodic acid (\text{HIO}3) and iodate (\text{IO}3^-
      • Nitrous acid (\text{HNO}2) and nitrite (\text{NO}2^-
      • Nitric acid (\text{HNO}3) and nitrate (\text{NO}3^-
      • Phosphorous acid (\text{H}3\text{PO}3) and phosphite (\text{PO}_3^{3-}
      • Phosphoric acid (\text{H}3\text{PO}4) and phosphate (\text{PO}_4^{3-}
      • Sulfurous acid (\text{H}2\text{SO}3) and sulfite (\text{SO}_3^{2-}
      • Sulfuric acid (\text{H}2\text{SO}4) and sulfate (\text{SO}_4^{2-}

Industrial Acids

• Common Industrial Acids:
  • Sulfuric acid (\text{H}2\text{SO}4)
  • Nitric acid (\text{HNO}_3)
  • Phosphoric acid (\text{H}3\text{PO}4)
  • Hydrochloric acid (HCl)
  • Acetic acid (\text{CH}_3\text{COOH})
• Sulfuric Acid (\text{H}2\text{SO}4):
  • Most commonly produced industrial chemical worldwide (over 37 million metric tons per year in the United States).
  • Used in:
    • Petroleum refining
    • Metallurgy
    • Manufacture of fertilizer
    • Production of metals, paper, paint, dyes, detergents, and chemical raw materials
    • Automobile batteries
  • Concentrated sulfuric acid is an effective dehydrating (water-removing) agent.
    • Can cause serious burns on the skin.
• Nitric Acid (\text{HNO}_3):
  • Pure nitric acid is a volatile, unstable liquid.
  • Solutions of nitric acid are widely used in industry.
  • Stains proteins yellow.
  • Has a suffocating odor, stains skin, and can cause serious burns.
  • Used in making:
    • Explosives
    • Rubber
    • Plastics
    • Dyes
    • Pharmaceuticals
  • Nitric acid solutions gradually become yellow due to slight decomposition to brown nitrogen dioxide gas.
• Phosphoric Acid (\text{H}3\text{PO}4):
  • Essential element for plants and animals (along with nitrogen and potassium).
  • Used directly for manufacturing fertilizers and animal feed.
  • Dilute phosphoric acid:
    • Has a pleasant but sour taste and is not toxic.
    • Used as a flavoring agent in beverages and as a cleaning agent for dairy equipment.
  • Important in the manufacture of detergents and ceramics.
• Hydrochloric Acid (HCl):
  • The stomach produces HCl to aid in digestion.
  • Industrially, hydrochloric acid is important for “pickling” iron and steel (immersion of metals in acid solutions to remove surface impurities).
  • Used in:
    • Industry as a general cleaning agent
    • Food processing
    • Activation of oil wells
    • Recovery of magnesium from seawater
    • Production of other chemicals
  • Concentrated solutions (muriatic acid) are used to maintain acidity in swimming pools and clean masonry.
• Acetic Acid (\text{CH}_3\text{COOH}):
  • Pure acetic acid is a clear, colorless, and pungent-smelling liquid known as glacial acetic acid (freezing point of 17°C).
  • The fermentation of certain plants produces vinegars containing acetic acid (white vinegar contains 4% to 8% acetic acid).
  • Important industrially in synthesizing chemicals used in the manufacture of plastics.
  • A raw material in the production of food supplements (e.g., lysine, an essential amino acid).
  • Also used as a fungicide.

Properties of Bases

1. Bitter Taste:
   • Aqueous solutions of bases taste bitter.
   • Caution: Taste should NEVER be used to test if a substance is a base, as many are caustic and cause severe burns.
2. Acid-Base Indicators:
   • Bases change the color of acid-base indicators.
   • The color will differ in a basic solution compared to an acidic solution.
3. Slippery Feel:
   • Dilute aqueous solutions of bases feel slippery (e.g., soap).
4. Neutralization:
   • Bases react with acids to produce salts and water.
   • The properties of a base disappear with the addition of an equivalent amount of an acid.
5. Electrical Conductivity:
   • Bases form ions in aqueous solutions and are electrolytes.

Arrhenius Acids and Bases

• Arrhenius Definition:
  • Arrhenius acid: A chemical compound that increases the concentration of hydrogen ions (\text{H}^+) in aqueous solution.
  • Arrhenius base: A substance that increases the concentration of hydroxide ions (\text{OH}^-) in aqueous solution.
• Aqueous Solutions of Acids:
  • Acids are molecular compounds with ionizable hydrogen atoms.
  • Their water solutions are known as aqueous acids, which are all electrolytes.
  • Hydrogen ion in aqueous solution is best represented as \text{H}_3\text{O}^+, the hydronium ion.
    • \text{HNO}3\text{(l)} + \text{H}2\text{O(l)} \longrightarrow \text{H}3\text{O}^+\text{(aq)} + \text{NO}3^-\text{(aq)}
• Strength of Acids:
  • Strong acid: Ionizes completely in aqueous solution and is a strong electrolyte (e.g., perchloric acid, hydrochloric acid, and nitric acid).
  • Weak acid: Releases few hydrogen ions in aqueous solution, with hydronium ions, anions, and dissolved acid molecules (e.g., hydrocyanic acid).
    • \text{HCN(aq)} + \text{H}2\text{O(l)} \rightleftharpoons \text{H}3\text{O}^+\text{(aq)} + \text{CN}^-\text{(aq)}
  • Acid strength increases with increasing polarity and decreasing bond energy.
  • Organic acids containing the carboxyl group (COOH) are generally weak acids (e.g., acetic acid).
    • \text{CH}3\text{COOH(aq)} + \text{H}2\text{O(l)} \rightleftharpoons \text{H}3\text{O}^+\text{(aq)} + \text{CH}3\text{COO}^-\text{(aq)}
• Aqueous Solutions of Bases:
  • Most bases are ionic compounds containing metal cations and the hydroxide anion, \text{OH}^-.
  • They dissociate when dissolved in water.
  • Solutions where a base completely dissociates in water to yield aqueous \text{OH}^- ions are considered strongly basic (e.g., sodium hydroxide, NaOH).
    • \text{NaOH(s)} \longrightarrow \text{Na}^+\text{(aq)} + \text{OH}^-\text{(aq)}
  • Group 1 elements (alkali metals) like Li, Na, K, Rb, and Cs form alkaline (basic) solutions.
  • Ammonia (\text{NH}_3) is a molecular base that produces hydroxide ions when it reacts with water.
    • \text{NH}3\text{(aq)} + \text{H}2\text{O(l)} \rightleftharpoons \text{NH}_4^+\text{(aq)} + \text{OH}^-\text{(aq)}
• Strength of Bases:
  • The strength of a base depends on the extent to which the base dissociates or adds hydroxide ions to the solution.
  • Strong bases are strong electrolytes (e.g., potassium hydroxide, KOH).
    • \text{KOH(s)} \longrightarrow \text{K}^+\text{(aq)} + \text{OH}^-\text{(aq)}
  • Bases that are not very soluble do not produce a large number of hydroxide ions (e.g., \text{Cu(OH)}_2).
  • Alkalinity depends on the concentration of \text{OH}^- ions in solution.
  • Many organic compounds that contain nitrogen atoms are also weak bases (e.g., codeine).

Section 2: Acid-Base Theories

• Key Terms:
  • Brønsted-Lowry acid
  • Brønsted-Lowry base
  • Brønsted-Lowry acid-base reaction
  • Monoprotic acid
  • Polyprotic acid
  • Diprotic acid
  • Triprotic acid
  • Lewis acid
  • Lewis base
  • Lewis acid-base reaction

Brønsted-Lowry Acids and Bases

• A Brønsted-Lowry acid is a molecule or ion that is a proton donor.
• A Brønsted-Lowry base is a molecule or ion that is a proton acceptor.
• In a Brønsted-Lowry acid-base reaction, protons are transferred from one reactant (the acid) to another (the base).
  • Example: \text{HCl} + \text{NH}3 \longrightarrow \text{NH}4^+ + \text{Cl}^-
• Water can act as a Brønsted-Lowry acid by donating a proton or as a Brønsted-Lowry base by accepting a proton.

Monoprotic and Polyprotic Acids

• A monoprotic acid can donate only one proton (hydrogen ion) per molecule (e.g., perchloric acid (\text{HClO}4), hydrochloric acid (HCl), and nitric acid (\text{HNO}3)).
  • \text{HCl(g)} + \text{H}2\text{O(l)} \longrightarrow \text{H}3\text{O}^+\text{(aq)} + \text{Cl}^-\text{(aq)}
• A polyprotic acid is an acid that can donate more than one proton per molecule (e.g., sulfuric acid (\text{H}2\text{SO}4) and phosphoric acid (\text{H}3\text{PO}4)).
• The ionization of a polyprotic acid occurs in stages.
• Diprotic Acid: An acid that can donate two protons per molecule (e.g., sulfuric acid).
  • \text{H}2\text{SO}4\text{(l)} + \text{H}2\text{O(l)} \longrightarrow \text{H}3\text{O}^+\text{(aq)} + \text{HSO}_4^-\text{(aq)}
  • \text{HSO}4^-\text{(aq)} + \text{H}2\text{O(l)} \rightleftharpoons \text{H}3\text{O}^+\text{(aq)} + \text{SO}4^{2-}\text{(aq)}
• Triprotic Acid: An acid able to donate three protons per molecule (e.g., phosphoric acid).
  • \text{H}3\text{PO}4\text{(aq)} + \text{H}2\text{O(l)} \rightleftharpoons \text{H}3\text{O}^+\text{(aq)} + \text{H}2\text{PO}4^-\text{(aq)}
  • \text{H}2\text{PO}4^-\text{(aq)} + \text{H}2\text{O(l)} \rightleftharpoons \text{H}3\text{O}^+\text{(aq)} + \text{HPO}_4^{2-}\text{(aq)}
  • \text{HPO}4^{2-}\text{(aq)} + \text{H}2\text{O(l)} \rightleftharpoons \text{H}3\text{O}^+\text{(aq)} + \text{PO}4^{3-}\text{(aq)}

Lewis Acids and Bases

• A Lewis acid is an atom, ion, or molecule that accepts an electron pair to form a covalent bond.
• A Lewis base is an atom, ion, or molecule that donates an electron pair to form a covalent bond.
• A Lewis acid-base reaction is the formation of one or more covalent bonds between an electron-pair donor and an electron-pair acceptor.
• Examples:
  • \text{H}^+\text{(aq)} + :\text{NH}3\text{(aq)} \longrightarrow [\text{H--NH}3]^+\text{(aq)}
  • \text{Ag}^+\text{(aq)} + 2:\text{NH}3\text{(aq)} \longrightarrow [\text{H}3\text{N--Ag--NH}_3]^+\text{(aq)}
  • \text{BF}3\text{(aq)} + \text{F}^-\text{(aq)} \longrightarrow \text{BF}4^-\text{(aq)}

Section 3: Acid-Base Reactions

• Key Terms:
  • Conjugate base
  • Conjugate acid
  • Amphoteric
  • Neutralization
  • Salt

Brønsted-Lowry Reactions and Conjugate Acid-Base Pairs

• The species that remains after a Brønsted-Lowry acid has given up a proton is the conjugate base of that acid.
• The species that is formed when a Brønsted-Lowry base gains a proton is the conjugate acid of that base.
• In every conjugate acid-base pair, the acid has one more proton than its conjugate base.
• Example:
  • \text{HF(aq)} + \text{H}2\text{O(l)} \rightleftharpoons \text{F}^-\text{(aq)} + \text{H}3\text{O}^+\text{(aq)}
• The stronger an acid is, the weaker its conjugate base; the stronger a base is, the weaker its conjugate acid.
• Proton-transfer reactions favor the production of the weaker acid and the weaker base.

Amphoteric Substances

• Any species that can react as either an acid or a base is described as amphoteric (e.g., water).
  • \text{H}2\text{SO}4\text{(aq)} + \text{H}2\text{O(l)} \longrightarrow \text{H}3\text{O}^+\text{(aq)} + \text{HSO}_4^-\text{(aq)}
  • \text{NH}3\text{(g)} + \text{H}2\text{O(l)} \rightleftharpoons \text{NH}_4^+\text{(aq)} + \text{OH}^-\text{(aq)}
• A substance acts as either an acid or a base depending on the strength of the acid or base with which it is reacting.

Neutralization Reactions

• Neutralization reactions produce water and a salt (an ionic compound composed of a cation from a base and an anion from an acid).
• Strong Acid-Strong Base Neutralization:
  • In aqueous solutions, neutralization is the reaction of hydronium ions and hydroxide ions to form water molecules.
  • \text{HCl(aq)} + \text{NaOH(aq)} \longrightarrow \text{NaCl(aq)} + \text{H}_2\text{O(l)}
  • \text{H}3\text{O}^+\text{(aq)} + \text{Cl}^-\text{(aq)} + \text{Na}^+\text{(aq)} + \text{OH}^-\text{(aq)} \longrightarrow \text{Na}^+\text{(aq)} + \text{Cl}^-\text{(aq)} + 2\text{H}2\text{O(l)}
  • \text{H}3\text{O}^+\text{(aq)} + \text{OH}^-\text{(aq)} \longrightarrow 2\text{H}2\text{O(l)}

Acid Rain

• Many industrial processes produce gases such as \text{NO}, \text{NO}2, \text{CO}2, \text{SO}2, and \text{SO}3.
• These compounds can dissolve in atmospheric water to produce acidic solutions that fall to the ground in the form of rain or snow (acid rain).
  • \text{SO}3\text{(g)} + \text{H}2\text{O(l)} \longrightarrow \text{H}2\text{SO}4\text{(aq)}
• Acid rain can erode statues and affect ecosystems.
• Amendments to the Clean Air Act in 1990 have decreased but not eliminated acid rain in the United States.

Math Tutor: Writing Equations for Ionic Reactions

• All dissolved substances in ionic reactions are dissociated into ions.
• Soluble ionic compounds and strong acids/bases are shown as separated ions in the full ionic equation.
• Ions that do not take part in the reaction are called spectator ions and are labeled as (aq) on both sides of the equation.
• Eliminating spectator ions produces the net ionic equation.
• Example:
• \text{(NH}4\text{)}2\text{SO}4\text{(aq)} + \text{Ba(NO}3\text{)}2\text{(aq)} \longrightarrow 2\text{NH}4\text{NO}3\text{(aq)} + \text{BaSO}4\text{(s)} Full Formula
• 2\text{NH}4\text{}^+\text{(aq)} + \text{SO}4\text{}^{2-}\text{(aq)} + \text{Ba}^{2+}\text{(aq)} + 2\text{NO}3^-\text{(aq)} \longrightarrow 2\text{NH}4\text{}^+\text{(aq)} + 2\text{NO}3^-\text{(aq)} + \text{BaSO}4\text{(s)} Full Ionic equation.
• \text{SO}4\text{}^{2-}\text{(aq)} + \text{Ba}^{2+}\text{(aq)} \longrightarrow \text{BaSO}4\text{(s)} Net Ionic equation.

Cross-Disciplinary Connection: Acid Water—A Hidden Menace

• Acidic water can cause the amount of lead in the water to rise.
• Older homes may have lead pipes or lead-containing solder.
• Acidic water can leach out lead from solder joints and copper from pipes.
• Lead poisoning is of particular concern in young children.
• Lead poisoning and other effects of acidic water can be prevented by:
  1. Monitoring the pH of your water on a regular basis.
  2. Letting your water tap run for about half a minute before use.
  3. Installing an alkali-injection pump.

Review Questions

1. What are five general properties of aqueous acids?
2. Name the following acids:
   • a. HBrO
   • b. HBrO3
3. What are five general properties of aqueous bases?
4. Why are strong acids also strong electrolytes? Is every strong electrolyte also a strong acid?
5. A classmate states, “All compounds containing H atoms are acids, and all compounds containing OH groups are bases.” Do you agree? Give examples.