15.2-15.4 The Nature of Acids and Bases - Acid Strength and the Acid Ionization Constant (Ka) Rachel Pietrow

The Nature of Acids and Bases General properties of acids Sour taste, ability to dissolve many metals, turn blue litmus paper red, neutralize bases Some common acids Hydrochloric acid, HCl Sulfuric Acid, H2SO4 Nitric Acid, HNO3 Acetic Acid, HC2H3O2 General properties of bases Bitter taste, slippery feel (react with oils on skin to form soap-like substances), turn red litmus paper blue, neutralize acids Alkaloids - organic bases found in plants, often poisonous Some common bases Sodium hydroxide, NaOH Potassium hydroxide, KOH Sodium bicarbonate, NaHCO3 Ammonia, NH3

Structure of Acids and Bases Binary acids have hydrogens attached to a nonmetal Ex: HCl, HF Oxy acids have hydrogens attached to an oxygen atom Ex: H2SO4, HNO3 Carboxylic acids have the -COOH group Ex: HC2H3O2 Most bases contain OH- ions Ex: NaOH, KOH Some bases contain CO32- ions Ex: NaHCO3, CaCO3

The Arrhenius Definition According to Arrhenius Acid: A substance that produces H+ ions in aqueous solution Base: A substance that produces OH- ions in aqueous solution HCl (aq) → H+ (aq) + Cl- (aq) According to this definition, HCl is an acid because it produces H+ ions in solution Since HCl is a covalent compound it does not contain ions. However, in water it ionizes completely to form H+ and Cl- ions These H+ ions are highly reactive and in aqueous solutions they bond to water to form H3O+ (hydronium ion) H+ (aq) and H3O+ (aq) are often used interchangeably to mean the same thing - an H+ has been dissolved in water NaOH (aq) → Na+ (aq) + OH- (aq) NaOH is a base according to this definition because it produces OH- ions in solution NaOH is ionic and when added to water dissociates into Na+ and OH- ions H+ (aq) + OH- (aq) →H2O (l) Under this definition, acids and bases combine to form water This is known as a neutralization reaction

The Brønsted-Lowry Definition This definition focuses on the transfer of H+ ions Acid: proton (H+ ion) donor Base: proton (H+ ion) acceptor HCl (aq) + H2O (l) →H3O+ (aq) + Cl- (aq) According to this definition, HCl is an acid because it donates a proton to water The H+ ion associates with a water molecule to form a hydronium ion (H3O+) NH3 (aq) + H2O (l) ⇌ NH4+ (aq) + OH- (aq) Applies nicely to bases such as NH3 that do not inherently contain OH- but still produce them in solutions. According to this definition, NH3 is a base because it accepts a proton from water Acids (proton donors) and bases (proton acceptors) always occur together in an acid-base reaction Ex: HCl (aq) + H2O (l) →H3O+ (aq) + Cl- (aq) HCl is the proton donor (acid) and H2O is the proton acceptor (base) Ex: NH3 (aq) + H2O (l) ⇌ NH4+ (aq) + OH- (aq) H2O is the proton donor (acid) and NH3 is the proton acceptor (base) An H+ ion is referred to as a proton because it is a hydrogen atom without its one electron.

Brønsted-Lowry Definition Continued NH3 (aq) + H2O (l) ⇌ NH4+ (aq) + OH- (aq) HCl (aq) + H2O (l) →H3O+ (aq) + Cl- (aq) Water is acting as an acid in the first equation and as a base in the second According to the Brønsted-Lowry definition, substances that can act as acids or bases (such as water) are amphoteric Let’s reverse the first equation: NH4+ (aq) + OH- (aq) ⇌ NH3 (aq) + H2O (l) Here, NH4+ is the proton donor (acid) and OH- is the proton acceptor (base). The substance that was the base (NH3) has become the acid (NH4+ ) and vice versa NH4+ and NH3 are often referred to as a conjugate acid-base pair - two substances related to each other by the 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 Summary: A base accepts a proton and becomes a conjugate acid An acid donates a proton and becomes a conjugate base Base Acid Conjugate Conjugate acid base

Practice Identify the Brønsted-Lowry acid, the Brønsted-Lowry base, the conjugate acid, and the conjugate base for this reaction. H2SO4 (aq) + H2O (l) → HSO4- (aq) + H3O+ (aq) H2SO4 (aq) + H2O (l) → HSO4- (aq) + H3O+ (aq) Since H2SO4 donates a proton to H2O it is the acid (proton donor). After it donates this proton, it becomes HSO4-, the conjugate base. Since H2O accepts a proton, it is the base (proton acceptor). After it accepts the proton, it becomes H3O+, the conjugate acid. Acid Base Conjugate Conjugate base acid

Practice Identify the Brønsted-Lowry acid, the Brønsted-Lowry base, the conjugate acid, and the conjugate base for this reaction. HCO3- (aq) + H2O (l) ⇌ H2CO3 (aq) + OH- (aq) HCO3- (aq) + H2O (l) ⇌ H2CO3 (aq) + OH- (aq) Since H2O donates a proton to HCO3- it is the acid (proton donor). After it donates this proton, it becomes OH-, the conjugate base. Since HCO3- accepts a proton, it is the base (proton acceptor). After it accepts the proton, it becomes H2CO3, the conjugate acid. Base Acid Conjugate Conjugate acid base

Acid Strength and the Acid Ionization Constant (Ka) A strong acid completely ionizes in solution A weak acid only partially ionizes The strength of an acid depends on the equilibrium HA (aq) + H2O (l) ⇌ H3O+ (aq) + A- (aq) HA is a generic formula for an acid If the equilibrium lies far to the right, the acid completely ionizes and is strong If the equilibrium lies to the left, only a small percentage of the molecules ionizes and the acid is weak

Strong Acids HCl is an example of a strong acid, one that completely ionizes in solution HCl (aq) + H2O (l) → H3O+ (aq) + Cl- (aq) So since HCl is a strong acid, an HCl solution contains basically no intact HCl; it has essentially completely ionized to form H3O+ (aq) and Cl- (aq) A 1.0 M HCl solution has [H3O+] = 1.0 M Examples of strong acids (know these!) Hydrochloric acid, HCl Hydrobromic acid, HBr Hydriodic acid, HI Nitric acid, HNO3 Perchloric acid, HClO4 Sulfuric acid, H2SO4 (diprotic) Except for sulfuric acid, all of these examples are monoprotic acids; they only contain one ionizable proton Sulfuric acid is diprotic, meaning it contains two ionizable protons

Weak Acids HF is a weak acid, one that only partially ionizes in solution HF (aq) + H2O (l) ⇌ H3O+ (aq) + F- (aq) Since HF is a weak acid, an HF solution contains a large number of intact HF molecules and only some H3O+ (aq) and F- (aq) molecules A 1.0 M HF solution has a [H3O+] that is much less than one, since only some of the HF molecules ionized to form H3O+ Examples of weak acids Hydrofluoric acid, HF Acetic acid, HC2H3O2 Formic acid, HCHO2 Sulfurous acid, H2SO3 (diprotic) Carbonic acid, H2CO3 (diprotic) Phosphoric acid, H3PO4 (triprotic) Triprotic - three ionizable protons

Why? The strength or weakness of an acid depends on the attraction between the anion (conjugate base) and the H+ ion HA (aq) + H2O (l) ⇌ H3O+ (aq) + A- (aq) The degree to which this generic reaction proceeds in the forward direction (how completely the molecule ionizes) depends on the strength of the attraction between H+ and A- If the attraction is weak, the reaction favors the forward direction and the acid is strong If the attraction is strong, the reaction favors the reverse reaction and the acid is weak In general, the stronger the acid, the weaker the conjugate base and vice versa. In HCl (a strong acid), the conjugate base (Cl-) has a fairly weak attraction to H+ In HF (a weak acid), the conjugate base (F-) has a greater attraction to H+ If the forward reaction (that of the acid) is likely to occur, then the reverse reaction (that of the conjugate base) has a lower tendency to occur. In HCl, the reverse reaction has a lower tendency to occur In HF, the reverse reaction occurs to a significant degree

The Acid Ionization Constant (Ka) The acid ionization constant (Ka), the equilibrium constant for the ionization reaction of a weak acid, is used to quantify the relative strength of an acid HA (aq) + H2O (l) ⇌ H3O+ (aq) + A- (aq) HA (aq) ⇌ H+ (aq) + A- (aq) The equilibrium constant for these two equivalent reactions is: Ka= [H3O+][A-]/[HA] = [H+][A-]/[HA] Since [H3O+] is equal to [H+], these expressions are equal The ionization constants for all weak acids are relatively small, but there is some variance in magnitude The smaller Ka, the less the acid ionizes and the weaker the acid The larger Ka, the more the acid ionizes and the stronger the acid