Acids, Bases, and Salts

• Acids, Bases, and Salts are fundamental concepts in chemistry.

Acid Definition

• An acid is a compound with an H+ cation dissolved in water.
• The phase label (aq) indicates a substance is dissolved in water, derived from "aqueous".
• HCl(aq) is an example of a substance dissolved in water.
• A compound lacking the (aq) label is treated as a molecular compound, not an acid.

Nomenclature of Acids with Monoatomic Anions

• Name acids with hydrogen and one other element using "hydro-" + stem of the other element + "-ic acid".
• HCl(aq) is hydrochloric acid.
• H2S(aq) is hydrosulfuric acid.
• Acids not dissolved in water are named as molecular compounds.
• Hydrogen chloride and hydrogen sulfide are examples, respectively.

Nomenclature of Acids with Polyatomic Anions

• Name acids with hydrogen ions and a polyatomic anion by deriving it from the polyatomic ion's stem name.
• Anion names ending in "-ate" become acids with the stem of the anion name + "-ic acid."
• Anion names ending in "-ite" become acids with the stem of the anion name + "-ous acid."

Practice Naming Acids

• Perchloric acid formula: HClO4
• Hydriodic acid formula: HI
• Hydrosulfuric acid formula: H2S
• Phosphorous acid formula: H3PO3
• HF(aq) is hydrofluoric acid.
• HNO3(aq) is nitric acid.
• H2SO4(aq) is sulfuric acid.
• H3PO4(aq) is phosphoric acid.
• HCl(aq) is hydrochloric acid.

Arrhenius Acids

• Svante Arrhenius provided the first chemical definitions of acids and bases.
• An Arrhenius acid increases the concentration of hydrogen ions (H+) in an aqueous solution.
• The hydronium ion (H3O+) is the chemical species that actually represents an H+ ion in an aqueous environment.
• H+ and H3O+ ions are often used interchangeably in chemical equations.

Arrhenius Acids and Monoprotic/Polyprotic Acids

• Acids like HNO3 and HCl give one proton upon dissociation and are called monoprotic acids.
• Acids like H2SO4 and H3PO4 have more than one hydrogen atom and give more than one H+ ion upon dissociation, they are called polyprotic acids.
• Whether an acid is polyprotic does not determine how strong it is.

Uses of Acids

• HCl is present in the stomach.
• H2SO4 is found in car batteries as a drying agent.
• HNO3 is used in fertilizer production.
• Ethanoic acid is utilized in the food industry.
• Fatty acids are used in soap making.
• Ascorbic acid is utilized in medicine.

Arrhenius Bases

• An Arrhenius base increases the OH- ion concentration in an aqueous solution or has at least one OH- ion in its formula.

Uses of Bases

• Sodium hydroxide (caustic soda) is used to manufacture soap, in petroleum refining, medicine, and rayon.
• Calcium hydroxide (slaked lime) is used to neutralize acid in water, in bleaching powder and fungicide manufacture, for acid burns, with sand for mortar, and to neutralize acid rain.
• Ammonium hydroxide removes ink spots and grease and is used in cosmetics.
• Alkalis are used in alkaline batteries.
• Potassium hydroxide is generally used in alkaline batteries.

Properties of Acids

• Acids turn litmus paper red.
• Acids react with some metals to produce H2 gas.
• Acids react with carbonate and hydrocarbonate salts to produce CO2 gas.
• Edible acids generally have a sour, sharp taste.

Properties of Bases

• Bases turn litmus paper blue.
• Bases are slippery to the touch.
• Bases have a bitter flavor if ingested.

Salts Definition

• Acids and bases react to form water and an ionic compound called a salt.
• A salt is any ionic compound formed by combining an acid and a base in chemistry.
• The hydrogen ion from the acid and hydroxide ion from the base form water.
• The cation from the base and the anion from the acid form a salt.
• Neutralization reaction: acid + base → water + salt

Salts Properties

• Most salts are crystalline solids, and they can be transparent or opaque.
• Most salts are water-soluble.
• Salt solutions and molten salts conduct electricity.
• Salts can be salty, sour, sweet, or bitter.
• Neutral salts are odorless.

Uses of Salts

• Ammonium chloride is used in torch batteries.
• Ammonium nitrate is used in fertilizers.
• Calcium chloride is used as a drying agent.
• Iron sulfate is used in iron tablets.
• Magnesium sulfate is used in medicine.
• Potassium nitrate is used in gunpowder.
• Silver bromide is used in photography.
• Sodium chloride is used to make NaOH.
• Sodium stearate is used to make soap.

pH Scale

• The pH scale expresses acidity.
• It is the negative logarithm of the hydronium ion concentration in an aqueous solution, ranging from 1 to 14.
• A pH of 1-6 is acidic.
• A pH of 7 is neutral.
• A pH of 8-14 is basic.

Limitations of Arrhenius Theory

• Arrhenius definitions are useful for common acids and bases.
• Arrhenius definitions do not apply when the solvent is not water or when reactions occur between gases.
• Most chemists use the Brønsted-Lowry theory for its generalized approach.

Neutralization Reaction Practice

• H2SO4 + KOH --> K2SO4 + H20
• HCl + Mg(OH)2 --> MgCl2 + H2O
• Sr(OH)2+ H3PO4 --> Sr3(PO4)2 + H2O
• CaCl2 is the product of hydrochloric acid and calcium hydroxide
• Zn(NO3)2 is the product of nitric acid and zinc
• Na3PO4 is the product of phosphoric acid and sodium hydroxide

Brønsted-Lowry Theory

• A major limitation of the Arrhenius theory is that it mainly describes acid-base behavior in water, while the Brønsted-Lowry theory is applicable in a variety of chemical reaction.
• A Brønsted-Lowry acid donates a proton.
• A Brønsted-Lowry base accepts a proton.
• A Brønsted-Lowry acid contains a hydrogen that dissociates as H+.
• A Bronsted-Lowry base has at least one lone pair of electrons to accept a proton.

Brønsted-Lowry Theory Examples

• Acid proton donor: HCl → H+(aq) + Cl-(aq)
• Base proton acceptor: NH3(aq) + H+(aq) → NH4+(aq)
• Acid-base reactions are processes where a proton transfers from an acid to a base.
• Reactions of ammonia gas, NH3(g), with hydrogen chloride gas, Cl(g), form solid ammonium chloride, NH4Cl(s).

Relationship of Bronsted-Lowry and Arrhenius

• A Bronsted-Lowry reaction is not considered an acid-base reaction by Arrhenius theory because the species do not form H+ or OH- in water.
• However, the chemistry (proton transfer from HCl to NH3 to form NH4Cl) is very similar to those that would occur in an aqueous phase.

Water as Both Acid and Base

• In the presence of nitric acid and water, nitric acid, HNO3, donates a proton to water, making it a Brønsted-Lowry acid.
• Since water accepts the proton from nitric acid to form H3O+, water acts as a Brønsted-Lowry base.
• The reaction of nitric acid and water highly favors product formation.
• In the presence of ammonia and water, ammonia, NH3, reacts with water to produce NH4+(aq) + OH-(aq).
• In this reaction, water donates a proton to ammonia, leading it to a Brønsted-Lowry acid.
• In the presence of a proton, ammonia acts as a Brønsted-Lowry base and accepts it from water.
• Water is amphoteric or amphiprotic because it can both accept and donate protons.
• Water can act as either a Brønsted-Lowry acid or base.

Strong and Weak Acids

• Strong acids will dissociate completely into constituent ions in an aqueous solution.
• Hydrochloric acid (HCl), nitric acid (HNO3), and sulfuric acid (H2SO4) are all strong acids.
• Sulfuric acid is a diprotic strong acid.
• Hydrochloric and nitric acids are monoprotic strong acids.
• Weak acids do not dissociate completely into constituent ions.
• Ethanoic acid is a weak monotropic acid written as: CH3COOH(aq) ⇌ CH3COO-(aq) + H+(aq)
• Once the common strong acids are known, identifying weak and strong acids becomes easier.

Strong and Weak Bases

• A strong base ionizes completely in aqueous solution.
• Group 1 and Group 2 hydroxides (NaOH) are strong bases.
• Ammonia is a weak base that becomes partially ionized in water
• Neutral nitrogen-containing compounds like ammonia, trimethylamine, and pyridine are common weak bases.

Conjugate Acid-Base Pairs

• Species formed after a base accepts a proton is a conjugate acid.
• Species remaining after an acid donates a proton is a conjugate base.
• Conjugate acid-base pairs have the same molecular formula except the acid has an extra H+ compound to the conjugate base.
• HCl(aq) + H2O(l) --> H3O+(aq) + Cl-(aq), Conjugate pair 1 = HCl and Cl-`, Conjugate pair 2 = H2O and H3O+
• NH3(aq) + H2O(l) = NH+(aq) + OH-(aq), Conjugate pair 1 = NH3 and NH4, Conjugate pair 2 = H2O and OH-

Bronsted-Lowry Summary

• A Brønsted-Lowry acid donates a proton.
• A Brønsted-Lowry base accepts a proton and has a lone pair of electrons to bond with the proton.
• Water is amphoteric and acts as both a Brønsted-Lowry acid and base.
• Strong acids and bases fully ionize in aqueous solutions, while weak acids and bases only partially ionize.
• A conjugate base results from the species formed after an acid donates its proton.
• A conjugate acid results from the species formed after a base accepts a proton.
• Two species in a conjugate acid-base pair share the same molecular formula except the acid has an extra H+.