Comprehensive Notes on Acids and Bases

Acids and Bases

Brønsted-Lowry Theory

• Brønsted-Lowry Acid: Any substance that can donate a proton, $H^+$.

  • An acid must contain a hydrogen that can be donated as $H^+$ .

  • In aqueous solution, $H^+$ ions immediately react with water molecules to form hydronium ions, $H_3O^+$ .

    • HNO3(g)+H2O(l)→H3O+(aq)+NO3−(aq)

    • HNO3(g)+H2O→H+(aq)+NO3−​(aq)

• Brønsted-Lowry Base: Any substance that can accept a proton, $H^+$.

  • A base must have at least one lone pair of electrons to form a new bond and accept a proton.

    • NH3(g)+H2O(l)⇌NH4+​(aq)+OH−(aq)

  Example with ammonia gas ($NH<em>3$) and hydrogen chloride gas ($HCl$) forming ammonium chloride ($NH</em>4Cl$):

  • $NH<em>3(g) + HCl(g) \rightarrow NH</em>4Cl(s)$

  • $HCl$ donates its proton (acid), and $NH_3$ accepts it (base).

• Summary:

  any substance that can donate a proton is considered an acid, while a substance that can accept a proton is classified as a base, according to the Brønsted-Lowry theory.

  • thus bases must contain a lone pair of electrons that can form a bond with the proton being donated.

  • while an acid mus contain a bonding electron pair to facilitate the donation of its proton.

Glossary

• Ionization: When a non-ionic (covalent) compound breaks apart in water to form ions.

  • All acids ionize in water.

  • HNO3(g)+H2O→H+(aq)+NO3−​(aq)

  • HCl2(g)+H2O(l)→H3​O+(aq)+2Cl−(aq)

• Dissociation: When an ionic substance breaks apart into its ions.

  • Most common bases dissociate in water.

  • $NaOH(s) + H_2O \rightarrow Na^+(aq) + OH^-(aq)$

Practice Examples

• $H<em>2CO</em>3(g) + H<em>2O(l) \rightarrow 2H</em>3O^+(aq) + CO_3^{2-}(aq)$

• $H<em>2CO</em>3(g) + H<em>2O \rightarrow 2H^+(aq) + CO</em>3^{2-}(aq)$

• $KOH(s) + H_2O \rightarrow K^+(aq) + OH^-(aq)$

Strength of Acids

• Strong Acid: Ionizes completely in aqueous solution (one-way arrow).

  • Example: $HCl(aq) + H_2O \rightarrow H^+(aq) + Cl^-(aq)$

• Weak Acid: Does not ionize completely (reversible arrow).

  • Example: $H<em>2CO</em>3(g) + H<em>2O \rightleftharpoons 2H^+(aq) + CO</em>3^{2-}(aq)$

• The number of hydrogen atoms in the acid does not indicate the strength of the acid.

• Strong acids fully ionize into ions in water.

  • $HCl \rightarrow H^+ + Cl^-$

• Weak acids partially ionize.

  • CH3COOH+H2O⇌H++CH3​COO−

Identifying Strong vs. Weak Acids

• Only a handful of strong acids exist; most others are weak.

• Common strong acids:

  • Perchloric acid: $HClO_4$

  • Hydroiodic acid: $HI$

  • Hydrobromic acid: $HBr$

  • Hydrochloric acid: $HCl$

  • Sulfuric acid: $H<em>2SO</em>4$

  • Nitric acid: $HNO_3$

• Weak acids only partially ionize, leaving some $H^+$ and the complete acid in solution.

Acid Nomenclature

• Acids without oxygen: hydro-(stem)-ic acid

• Acids with oxygen:

  • -ate: (stem)-ic acid

  • -ite: (stem)-ous acid

• Examples:

  • $HI(aq)$: hydroiodic acid (iodide)

  • $HCl(aq)$: hydrochloric acid (chloride)

  • $HBr(aq)$: hydrobromic acid (bromide)

  • $H<em>2CO</em>3$: carbonic acid (carbonate)

  • $H<em>2SO</em>4$: sulfuric acid (sulfate)

  • $H<em>2SO</em>3$: sulfurous acid (sulfite)

  • $HNO_3$: nitric acid (nitrate)

  • $HNO_2$: nitrous acid (nitrite)

  • $H_2S$: hydrosulfic acid (sulfide)

  • $HF$: hydrofluoric acid (fluoride)

  • $H<em>3PO</em>4$: phosphoric acid (phosphate)

Strength of Bases

• Strong Bases: Fully dissociate into ions in water (one-way arrow).

  • $NaOH \rightarrow Na^+ + OH^-$

• Weak Bases: Partially dissociate into ions (reversible arrows).

  • $NH<em>3 \rightleftharpoons NH</em>4^+ + OH^-$

• Strong bases are Group 1 and Group 2 metal hydroxides (e.g., $NaOH$, $KOH$, $Ca(OH)_2$).

• Weak bases are usually nitrogen-containing compounds (e.g., ammonia).

Bases Recap

• Strong bases dissociate completely in water.

  • Group I metal hydroxides: $NaOH$, $LiOH$, etc.

  • Some Group II metal hydroxides: $Ca(OH)<em>2$, $Ba(OH)</em>2$, $Sr(OH)_2$

• Weak bases dissociate only a few ions (e.g., $NH_3$).

Polyprotic Acids

• Polyprotic acids have multiple protons they can donate.

  • poly = multiple

  • protic = protons/$H^+$

  • acids = can donate protons/$H^+$

• Monoprotic: Has one proton it can donate.

• Diprotic: Has two protons it can donate.

• Triprotic: Has three protons it can donate.

• Polyprotic acids lose protons one at a time in stepwise equations.

  • Sulfuric acid (diprotic) loses protons when added to water.

  • Phosphoric acid (triprotic) loses protons when added to water.

Polyprotic Acid Examples

• Monoprotic: donate a single proton ($H^+$) in an aqueous solution (e.g., $HBr$, $HF$, $HNO<em>3$, $CH</em>3COOH$).

• Diprotic: donate two protons in an aqueous solution (e.g., $H<em>2CO</em>3$, $H<em>2CrO</em>4$).

• Triprotic: donate three protons in an aqueous solution (e.g., $H<em>3AsO</em>4$).

Neutralization Reaction

• Occurs when an acid and a base react.

• Hydronium ions from the acid and hydroxide ions from the base form water molecules.

• Another product of these reactions is a salt.

• Example: $NaOH + HCl \rightarrow water + salt$

Reactions of Acids

• Acid + hydrogen carbonate → salt + water + carbon dioxide

• Example: ethanoic acid + sodium hydrogen carbonate → sodium ethanoate + water + carbon dioxide

  • $CH<em>3COOH + NaHCO</em>3 \rightarrow CH<em>3COONa + H</em>2O + CO_2$

Acid Reactions: Carbon Dioxide and Water

• Carbon dioxide dissolves in water to form carbonic acid, mostly in oceans and the atmosphere.

Acid Rain

• Pure water has a pH of 7 (neutral).

• Normal precipitation is slightly acidic (pH = 5.6) due to carbon dioxide forming carbonic acid.

Formation of Acid Rain

• Factories burning fossil fuels release sulfur dioxide ($SO<em>2$) and nitrous oxides ($NO</em>x$).

• Cars burning fossil fuels release nitrous oxides ($NO_x$).

• Sulfur and nitrous oxides combine with water in the atmosphere to form sulfuric acid and nitric acid & nitrous acid.

  • $H<em>2O + SO</em>3 \rightarrow H<em>2SO</em>4$

  • $H<em>2O + 2NO</em>2 \rightarrow HNO<em>3 + HNO</em>2$

• These strong acids lower the pH of rain.

• Acid rain follows the winds, starting in areas with coal-fired power plants and factories.

Effects of Acid Rain on Plants

• Acidic soil causes nutrients to be washed away.

• Sulfur dioxide clogs openings on plant surfaces, reducing photosynthesis.

• Acid rain releases toxic metals into the soil, causing root damage.

Acid Shock

• Sudden runoff of large amounts of highly acidic water into lakes, streams, and rivers when snow melts or after heavy rains.

• Causes large numbers of fish to die.

• Affects reproduction of fish and amphibians; fewer eggs are produced, and those often do not hatch.

• Surviving offspring often have birth defects and cannot reproduce.

Acid Rain and Humans

• Toxic metals (aluminum and mercury) released into the environment.

• These metals enter crops, water, and fish, poisoning humans who consume them.

• Fishermen are affected by decreased fish populations.

• Forestry is affected because trees are damaged by acid rain.

• Increases in respiratory problems in children.

• Acid rain dissolves calcium carbonate in building materials like concrete and marble.

Effects of Acid Rain

• Causes acidification of soils and water.

• Reduces photosynthesis, leading to plant death.

• Large volumes lead to fish death.

• Affects the reproduction of fish and amphibians.

• Releases toxic metals, damaging roots.

• Toxic metals are taken in by crops, water, and aquatic organisms, poisoning people.

• Increases respiratory problems (e.g., asthma).

• Dissolves calcium carbonate in building materials, damaging historic buildings/statues.

Preventing Acid Rain

• Global emissions of sulfur and nitrogen oxides have been declining.

• Higher global air quality monitoring.

• Implementation of environmental laws and international regulations.

• Development of new technologies to minimize vehicle and power station emissions.

• Vehicles are now equipped with catalytic converters that convert nitrogen oxides into harmless nitrogen gas.

• Power stations have flue gas desulfurization (FGS) to reduce emissions of sulfur oxides.

• Sulfur oxides react with calcium carbonate or calcium hydroxide to be converted into a solid sludge for disposal.