Introduction to Acids and Bases 15.2-15.4 by Mia Tsang

15.2: The Nature of Acids and Bases Name Occurrence/Uses Hydrochloric Acid (HCl) Metal cleaning; food preparation; ore refining; primary component of stomach acid Sulfuric Acid (H2SO4) Fertilizer/explosives manufacturing; dye/glue production; automobile batteries; electroplating of copper Nitric Acid (HNO3) Fertilizer/explosives manufacturing; dye/glue production Acetic Acid (HC2H3O2) Plastic/rubber manufacturing; food preservative; active component of vinegar Citric Acid (H3C6H5O7) Present in citrus fruits like lemons/limes; used to adjust pH in foods and beverages Carbonic Acid (H2CO3) Found in carbonated beverages as a product of the reaction between carbon dioxide and water Hydrofluoric Acid (HF) Metal cleaning; glass frosting and etching Phosphoric Acid (H3PO4) Fertilizer manufacture; biological buffering; preservative in beverages Properties of Acids: sour taste ability to dissolve metals ability to turn blue litmus paper red ability to neutralize bases Table at the right shows some common acids

Table at the right shows some common bases Properties of Bases bitter taste slippery to the touch this is caused by reactions between the base and the oils on your skin, forming soap-like compounds ability to turn red litmus paper blue ability to neutralize acids Table at the right shows some common bases Name Occurrence/Uses Sodium Hydroxide (NaOH) Petroleum processing; soap and plastic manufacturing Potassium Hydroxide (KOH) Cotton processing; electroplating; soap production; batteries Sodium Bicarbonate (NaHCO3) Antacid; ingredient of baking soda; source of CO2 Sodium Carbonate (Na2CO3) Manufacture of glass and soap; general cleanser; water softener Ammonia (NH3) Detergent; fertilizer/explosives manufacturing; synthetic fiber production

15.3: Definitions of Acids and Bases There are 3 different definitions of acids and bases Arrhenius definition Brønsted-Lowry definition Lewis definition In this section we just discuss the first two.

Arrhenius Definition Proposed by Swedish chemist Svante Arrhenius in the 1880’s States that: An acid is a substance that produces H+ ions in an aqueous solution A base is a substance that produces OH— ions in an aqueous solution By this definition, HCl is an acid because it ionizes to produce H+ ions in solution, as seen in the reaction below: HCl (aq) → H+ (aq) + Cl- (aq) Even though HCl is a covalent compound, it can still ionize; the term just means that it separates into individual ions, which covalent compounds can still do. The H+ ions produced in the above reaction are highly reactive and will bond with water molecules in aqueous solution to form H3O+, otherwise known as the hydronium ion. H+ ions in solution will always associate with H2O molecules to form either hydronium ions or other molecules with the general formula of H(H2O)n+ where n is the number of H2O molecules If an H+ ion associates with 2 water molecules, you get H(H2O)2+ If an H+ ion associates with 3 water molecules, you get H(H2O)3+

By the Arrhenius definition, NaOH is a base because it produces OH- ions in solution, as seen in the reaction below: NaOH (aq) → Na+ (aq) + OH- (aq) Since NaOH is an ionic compound, it dissociates into its constituent ions This definition also states that acids and bases combine to form water and neutralize each other in the process, as seen in the reaction below: H+ (aq) + OH- (aq) → H2O (l)

Brønsted-Lowry Definition (the important one) More widely applicable definition of acids and bases introduced in 1923 Focuses on the transfer of H+ ions (protons) An acid is a proton donor A base is a proton acceptor By this definition HCl is an acid because it donates a proton to water in this reaction: HCl (aq) + H2O (l) → H3O+ (aq) + Cl- (aq) This definition is preferred because it’s easier to see exactly what’s happening: The H+ ion clearly associates with the water molecule to form a hydronium ion

The Brønsted-Lowry definition is also preferred because it includes bases that don’t inherently contain OH- ions but still produce them in solution, such as NH3 NH3 is a base under this definition because it accepts a proton from water, as seen in this reaction: NH3 (aq) + H2O (l) → NH4+ (aq) + OH- (aq) Notice that OH- ions are still produced in the reaction, despite the fact that the base doesn’t contain any All Arrhenius acids and bases are also Brønsted-Lowry acids and bases, but not all Brønsted-Lowry acids and bases are Arrhenius acids and bases.

Under the Brønsted-Lowry definition, acids (proton donors) and bases (proton acceptors) always occur together in an acid-base reaction. HCl (aq) + H2O (l) → H3O+ (aq) + Cl- (aq) In the above reaction, HCl is the proton donor, so it is the acid. H2O is the proton acceptor, so it is the base. NH3 (aq) + H2O (l) ⇌ NH4+ (aq) + OH- (aq) In the above reaction, H2O is the proton donor, so it is the acid. NH3 is the proton acceptor, so it is the base. Since water can act as either an acid or a base, it is amphoteric.

NH4+ (aq) + OH- (aq) ⇌ NH3 (aq) + H2O (l) What happens when this, NH3 (aq) + H2O (l) → NH4+ (aq) + OH- (aq) , is reversed? NH4+ (aq) + OH- (aq) ⇌ NH3 (aq) + H2O (l) NH4+ donates a proton to OH-, making it an acid under the Brønsted-Lowry definition OH- accepts a proton from NH4+, making it a base under the Brønsted-Lowry definition Let’s go back to the original equation: NH3 (aq) + H2O (l) ⇌ NH4+ (aq) + OH- (aq) The substance that started off as the base becomes an acid, and vice versa The base and the acid are together known as a conjugate acid-base pair, two substances related to each other by a transfer of a proton A conjugate acid is any base to which a proton has been added A conjugate base is any acid from which a proton has been removed In this reaction: NH3 is a base and NH4+ is its conjugate acid H2O is an acid and OH- is its conjugate base

To summarize the Brønsted-Lowry definition: A base accepts a proton and becomes a conjugate acid. An acid donates a proton and becomes a conjugate base.

Practice problem! In each reaction, identify the Brønsted-Lowry acid, the Brønsted-Lowry base, the conjugate acid, and the conjugate base. Acid Base Conjugate Base Conjugate Acid H2SO4 (aq) + H2O (l) → HSO4- (aq) + H3O+ (aq) Base Acid Conjugate Acid Conjugate Base HCO3- (aq) + H2O (l) → H2CO3 (aq) + OH- (aq)

15.4: Acid Strength and the Acid Ionization Constant (Ka) Just as the strength of an electrolyte depends on how much it dissociates into its constituent ions in solution, the strength of an acid depends on how completely the acid ionizes in solution. A strong acid completely ionizes in solution A weak acid only partially ionizes in solution The strength of an acid depends on the equilibrium of the following general equation, where HA is a generic formula for an acid: HA (aq) + H2O (l) ⇌ H3O+ (aq) + A- (aq) If the equilibrium lies far to the right, the products are favored and the acid completely ionizes, meaning it is a strong acid If the equilibrium lies to the left, the reactants are favored and a smaller percentage of the acid molecules ionize, meaning it is a weak acid

Strong Acids Strong Acids Hydrochloric acid is an example of a strong acid: HCl (aq) + H2O (l) → H3O+ (aq) + Cl- (aq) Notice the single arrow; this means that HCl ionizes completely. A 1.0 M HCl solution actually contains an H3O+ concentration of 1.0 M, because the HCl will dissociate entirely into H3O+ and Cl-, and there will be no HCl left. So the concentration has to be measured in terms of H3O+. Because of this, the concentration of a 1.0 M HCl solution can be abbreviated to [H3O+] = 1.0 M Strong Acids Hydrochloric Acid (HCl) Nitric Acid (HNO3) Hydrobromic Acid (HBr) Perchloric Acid (HClO4) Hydriodic Acid (HI) Sulfuric Acid (H2SO4) You might want to memorize these!

Weak Acids Unlike HCl, HF is a weak acid: HF (aq) + H2O (l) ⇌ H3O+ (aq) + F- (aq) Notice the equilibrium arrow; this means that HF only partially ionizes. A 1.0 M HF solution has an H3O+ concentration much lower than 1.0 M, because much less HF actually ionizes, so less is formed. Weak Acids Hydrofluoric Acid (HF) Sulfurous Acid (H2SO3) Acetic Acid (HC2H3O2) Carbonic Acid (H2CO3) Formic Acid (HCHO2) Phosphoric Acid (H3PO4)

Degree of Strength The degree to which an acid is strong or weak depends on the attraction between the anion (conjugate base) and the hydrogen ion, relative to their attractions to water HA (aq) + H2O (l) ⇌ H3O+ (aq) + A- (aq) If the attraction between H+ and A- is weak, the forward reaction is favored and the acid will ionize completely, because it’s easier to separate the ions → the acid will be strong If the attraction between H+ and A- is strong, the reverse reaction is favored and the acid will only partially ionize, because it’s harder to separate the ions → the acid will be weak When the forward reaction (that of the acid) has a high tendency to occur, the reverse reaction (that of the conjugate base) has a low tendency to occur. In general, the stronger the acid, the weaker the conjugate base, and vice versa.

The Acid Ionization Constant (Ka) The relative strength of a weak acid can be quantified with the acid ionization constant (Ka), which is basically the equilibrium constant for the ionization reaction of a weak acid Given the two reactions below: HA (aq) + H2O (l) ⇌ H3O+ (aq) + A- (aq) HA (aq) ⇌ H+ (aq) + A- (aq) The equilibrium constant is : Ka= [H3O+][A-]/[HA] or [H+][A-]/[HA] Since [H3O+] is equivalent to [H+], both forms of the expression are equal The ionization constants for all weak acids are relatively small, but they vary in magnitude The smaller the constant, the less the acid ionizes, and the weaker the acid is

Practice problem! Consider these two acids and their Ka values: HF Ka= 3.5 x 10-4 HClO Ka= 2.9 x 10-8 Which acid is stronger? HF is stronger because it has a larger Ka value, meaning it dissociates in water more than HClO does.