© 2015 Pearson Education, Inc. Chapter 17 Lecture presentation Acids and Bases Catherine E. MacGowan Armstrong Atlantic State University

© 2015 Pearson Education, Inc. General Properties of Acids Sour taste Ability to dissolve many metals Ability to neutralize bases Change blue litmus paper to red Acetic acid: vinegar

© 2015 Pearson Education, Inc. Table of Common Acids

© 2015 Pearson Education, Inc. Structure of Organic Acids: Carboxylic Acids Carboxylic acids have a COOH group. –Acetic acid: HC 2 H 3 O 2 –Citric acid: H 3 C 6 H 5 O 7 Only the H in the COOH (carboxylic acid) group is acidic. –The H is on the COOH. –H—O—C— || O

© 2015 Pearson Education, Inc. General Properties of Bases Taste bitter –Alkaloid = plant product that is alkaline Often poisonous Feel slippery to touch Ability to neutralize acids Change red litmus paper to blue

© 2015 Pearson Education, Inc. Table of Common Bases

© 2015 Pearson Education, Inc. Hydronium Ion: H 3 O + The H + ions (protons) produced by the acid are so reactive they cannot exist in water. Instead, they react with water molecules to produce complex ions, mainly hydronium ion, H 3 O +. H + + H 2 O → H 3 O +

© 2015 Pearson Education, Inc. Definitions of Acids and Bases Arrhenius definition (simplest and most restrictive) –Acids are substances that when dissolved in water produce a hydronium, H 3 O + (hydrogen ion H + ). –Bases are substances that when dissolved in water produce a hydroxide ion, OH –. Brønsted–Lowry definition (based on reactions in water) –Acids are substances that when dissolved in water donate hydronium, H 3 O + (hydrogen ion H + ). –Bases are substances that accept a hydronium, H 3 O + (hydrogen ion H + ). Lewis definition (most expansive and based on electron donors and acceptors) –Acids are substances that accept or need an electron pair. –Bases are substances that donate an electron pair to another substance.

© 2015 Pearson Education, Inc. Arrhenius Theory Acids produce H + ions in aqueous solution. HCl(aq) → H + (aq) + Cl − (aq)

© 2015 Pearson Education, Inc. Arrhenius Theory Bases produce OH − ions in aqueous solution. NaOH(aq) → Na + (aq) + OH(aq)

© 2015 Pearson Education, Inc. Arrhenius Acid–Base Reactions The H + from the acid combines with the OH − from the base to make a molecule of H 2 O. –Think of H 2 O as H—OH. The cation from the base combines with the anion from the acid to make a salt. acid + base → salt + water Example: HCl(aq) + NaOH(aq) → NaCl(aq) + H 2 O(l) (acid) (base) (salt) (water)

© 2015 Pearson Education, Inc. Problems with Arrhenius Theory: Too Narrow of a Definition It does not explain why molecular substances, such as ammonia, NH 3, dissolve in water to form basic solutions, even though they do not contain OH – ions. It does not explain how some ionic compounds, such as sodium carbonate (washing soda), Na 2 CO 3, or sodium oxide, Na 2 O, dissolve in water to form basic solutions, even though they do not contain OH – ions. It does not explain why molecular substances, such as CO 2, dissolve in water to form acidic solutions, even though they do not contain H + ions. It does not explain acid–base reactions that take place outside aqueous solution.

© 2015 Pearson Education, Inc. Brønsted–Lowry Acid–Base Theory It defines acids and bases based on what happens in a reaction. Any reaction involving an H + (proton) that transfers from one molecule to another is an acid–base reaction, regardless of whether it occurs in aqueous solution or if there is OH − present. All reactions that fit the Arrhenius definition also fit the Brønsted–Lowry definition.

© 2015 Pearson Education, Inc. Brønsted–Lowry Theory: Acid The acid is an H + donor. The base is an H + acceptor. –Base structure must contain an atom with an unshared pair of electrons. In a Brønsted–Lowry acid–base reaction, the acid molecule donates an H + to the base molecule. H—A + :B :A – + H—B + (acid) (base) (conjugate) (conjugate) base acid

© 2015 Pearson Education, Inc. Brønsted–Lowry Acids Brønsted–Lowry acids are H + donors. –Any material that has H can potentially be a Brønsted–Lowry acid. –Because of the molecular structure, often one H in the molecule is easier to transfer than others. When HCl dissolves in water, the HCl is the acid because HCl transfers an H + to H 2 O, forming H 3 O + ions. –Water acts as base, accepting H +. HCl(aq) + H 2 O(l) → Cl – (aq) + H 3 O + (aq) (acid) (base) (conjugate) (conjugate) base acid

© 2015 Pearson Education, Inc. Brønsted–Lowry Bases Brønsted–Lowry bases are H + acceptors. –Any material that has atoms with lone pairs can potentially be a Brønsted–Lowry base. –Because of the molecular structure, often one atom in the molecule is more willing to accept H + transfer than others. When NH 3 dissolves in water, the NH 3 (aq) is the base because NH 3 accepts an H + from H 2 O, forming OH – (aq). –Water acts as acid, donating H +. NH 3 (aq) + H 2 O(l) NH 4 + (aq) + OH – (aq) (base) (acid) (conjugate) (conjugate) acid base

© 2015 Pearson Education, Inc. Brønsted–Lowry Acid–Base Reactions One of the advantages of Brønsted–Lowry theory is that it illustrates equilibrium reactions to be as follows: H—A + :B :A – + H—B + (acid) (base) (conjugate) (conjugate) base acid –The original base has an extra H + after the reaction, so it will act as an acid in the reverse process. –And the original acid has a lone pair of electrons after the reaction, so it will act as a base in the reverse process: :A – + H—B + H—A + :B (conjugate) (conjugate) (acid) (base) base acid

© 2015 Pearson Education, Inc. Conjugate Acid–Base Pairs In a Brønsted–Lowry acid–base reaction, –the original base becomes a conjugate acid in the reverse reaction; –the original acid becomes a conjugate base in the reverse process. Each reactant and the product it becomes is called a conjugate pair.

© 2015 Pearson Education, Inc. Conjugate Pairs A base accepts a proton and becomes a conjugate acid. An acid donates a proton and becomes a conjugate base.

© 2015 Pearson Education, Inc. Practice Problem: Identifying Brønsted–Lowry Acids and Bases and Their Conjugates

© 2015 Pearson Education, Inc. Acid Strength and Molecular Structure of Acids Binary acids (H—Y) have acidic hydrogens attached to a nonmetal atom. –Example: HCl and HF The more polarized the bond ( δ+ H—X δ− ), the more acidic the bond. The stronger the H—X bond, the weaker the acid. Binary acid strength increases to the right across a period. –Acidity: H—C < H—N < H—O < H—F Binary acid strength increases down the column. –Acidity: H—F < H—Cl < H—Br < H—I

© 2015 Pearson Education, Inc. Relationship between Bond Strength and Acidity for Binary Acids

© 2015 Pearson Education, Inc. Strengths of Oxyacids, H—O—Y The more electronegative the Y atom, the stronger the oxyacid. –Example: HClO > HIO –Acidity of oxyacids decreases down a group. Same trend as binary acids –Helps weaken the H—O bond The larger the oxidation number of the central atom, the stronger the oxyacid. –Example: H 2 CO 3 > H 3 BO 3 –Acidity of oxyacids increases to the right across a period. Opposite trend of binary acids The more oxygens attached to Y, the stronger the oxyacid. –Further weakens and polarizes the H—O bond –HClO 3 > HClO 2

© 2015 Pearson Education, Inc. Structure of Oxyacids Oxyacids have acid hydrogens attached to an oxygen atom.

© 2015 Pearson Education, Inc. Relationship between Electronegativity and Acidity of Oxyacids The more electronegative Y is, the weaker and more polarized is the H—O bond in the acid. The more electronegative Y is, the more acidic is the acid.

© 2015 Pearson Education, Inc. Relationship between Number of Oxygens on the Central Atom and Acidity

© 2015 Pearson Education, Inc. Strong Acid/Base versus Weak Acid/Base A strong acid is a strong electrolyte. –Complete or near complete ionization of acid molecule in water A weak acid is a weak electrolyte. –Only a small percentage or partial ionization of the acid molecule in water –An equilibrium ( ) is established. A strong base is a strong electrolyte. –Complete or near complete ionization of base molecule in water –Produces OH – ions, either through dissociation or reaction with water A weak base is a weak electrolyte. –Only a small percentage or partial ionization of the base molecule in water –Produces OH – ions, either through dissociation or reaction with water –An equilibrium ( ) is established.

© 2015 Pearson Education, Inc. General Trends in Acidity The stronger an acid is at donating H, the weaker the conjugate base is at accepting H. Higher oxidation number = stronger oxyacid –H 2 SO 4 > H 2 SO 3 ; HNO 3 > HNO 2 Cation stronger acid than neutral molecule; neutral stronger acid than anion –H 3 O + > H 2 O > OH – ; NH 4 + > NH 3 > NH 2 − –Trend in base strength opposite

© 2015 Pearson Education, Inc. Strong Acids: K a > 1 Strong acids donate practically all their hydrogen atoms. –Near 100% ionization of acid molecule occurs in water –Strong electrolyte [H 3 O + ] = [strong acid] –[HA] = [H + ] –The brackets designate molarity There are six strong acids: –HCl, HBr, HI, HNO 3, HClO 4, and H 2 SO 4

© 2015 Pearson Education, Inc. Weak Acids: K a < 1 Weak acids donate a small fraction (partial ionization) of their hydrogen atoms. –Weak acid molecules do not donate much of their hydrogens to water. –Much less than 1% ionized in water [H 3 O + ] << [weak acid] –So the [weak acid] does NOT = [H + ] in solution

© 2015 Pearson Education, Inc. Examples of Weak Acids

© 2015 Pearson Education, Inc. Strengths of Acids (K a ) and Bases (K b ) Commonly, acid or base strength is measured by determining the equilibrium constant of a substance’s reaction with water. HAcid + H 2 O Acid − + H 3 O + Base: + H 2 O HBase + + OH − The farther the equilibrium position lies toward the products, the stronger the acid or base. The position of equilibrium depends on the strength of attraction between the base form and the H +. –Stronger attraction means stronger base or weaker acid.

© 2015 Pearson Education, Inc. Acid Ionization Constant, K a Acid strength is measured by the size of the equilibrium constant when it reacts with H 2 O. The equilibrium constant for this reaction is called the acid ionization constant, K a. –The larger the K a value, the stronger the acid.

© 2015 Pearson Education, Inc. Table of Acid Ionization Constants, K a

© 2015 Pearson Education, Inc. Autoionization of Water Water is amphoteric; it can act as either an acid or a base. –Therefore, there must be a few ions present. About two out of every one billion water molecules form ions through a process called autoionization. H 2 O H + + OH – H 2 O + H 2 O H 3 O + + OH – All aqueous solutions contain both H 3 O + and OH –. –The concentrations of H 3 O + and OH – are equal in water. –[H 3 O + ] = [OH – ] = 10 −7 M at 25 °C

© 2015 Pearson Education, Inc. Ion Product of Water, K w The product of the H 3 O + and OH – concentrations is always the same number at room temperature, 1 × 10 –14. This value is called the ion product of water and has the symbol K w. –It is also know as the dissociation constant of water. [H 3 O + ] × [OH – ] = K w = 1.00 × 10 −14 at 25 °C –If you measure one of the concentrations, you can calculate the other. As [H 3 O + ] increases, the [OH – ] must decrease so that the product stays constant. –Inversely proportional

© 2015 Pearson Education, Inc. Acidic and Basic Solutions All aqueous solutions contain both H 3 O + and OH – ions. Neutral solutions have equal [H 3 O + ] and [OH – ]. –[H 3 O + ] = [OH – ] = 1.00 × 10 −7 Acidic solutions have a larger [H 3 O + ] than [OH – ]. –[H 3 O + ] > 1.00 × 10 −7 ; [OH – ] < 1.00 × 10 −7 Basic solutions have a larger [OH – ] than [H 3 O + ]. –[H 3 O + ] 1.00 × 10 −7

© 2015 Pearson Education, Inc. Practice Problem: K a Calculations

© 2015 Pearson Education, Inc. Measuring Acidity: pH The acidity or basicity of a solution is often expressed as pH. –pH = −log[H 3 O + ] Exponent on 10 with a positive sign pH water = −log[10 −7 ] = 7 We need to know the [H 3 O + ] concentration to find the pH. pH 7 is basic. pH = 7 is neutral. – If we know the pH, we can determine [H 3 O + ]. [H 3 O + ] = 10 −pH

© 2015 Pearson Education, Inc. What Does the pH Number Imply? The lower the pH, the more acidic the solution; the higher the pH, the more basic the solution. –One pH unit corresponds to a factor of 10 difference in acidity. Normal range of pH is 0 to 14. –pH 0 is [H 3 O + ] = 1 M; pH 14 is [OH – ] = 1 M. –pH can be negative (very acidic) or larger than 14 (very alkaline).

© 2015 Pearson Education, Inc. Practice Problem: Calculating pH from [H 3 O + ] or [OH – ]

© 2015 Pearson Education, Inc. Practice Problem: Calculating [H 3 O + ] from pH

© 2015 Pearson Education, Inc. What is pOH? Another way of expressing the acidity/basicity of a solution is pOH. –pOH = −log[OH – ] If you know pOH, then you can determine [OH – ]. –[OH – ] = 10 −pOH –pOH water = −log[10 −7 ] = 7 You need to know the [OH – ] concentration to find pOH. pOH 7 is acidic; pOH = 7 is neutral.

© 2015 Pearson Education, Inc. Relationship between pH and pOH pH + pOH = at 25 °C. –If you know pOH, you can determine pH; if you know pH, you can determine pOH.

© 2015 Pearson Education, Inc. The pK’s: pK a and pK b A way of expressing the strength of an acid or base is through its pK. pK a = −log(K a ), K a = 10 −pKa –The stronger the acid, the smaller the pK a. Larger K a = smaller pK a –Because pK a is −log(K a ) pK b = −log(K b ), K b = 10 −pKb –The stronger the base, the smaller the pK b. Larger K b = smaller pK b

© 2015 Pearson Education, Inc. [H 3 O + ] and [OH − ] in a Strong Acid or Strong Base Solution There are two sources of H 3 O + in an aqueous solution of a strong acid—the acid and the water. There are two sources of OH − in an aqueous solution of a strong acid—the base and the water. For a strong acid or base, the contribution of the water to the total [H 3 O + ] or [OH − ] is negligible. –The [H 3 O + ] acid shifts the K w equilibrium so far that [H 3 O + ] water is too small to be significant. Except in very dilute solutions, generally < 1 × 10 −4 M

© 2015 Pearson Education, Inc. Finding pH of a Strong Acid There are six strong acids: Five are monoprotic and one is diprotic. –Monoprotic: HCl, HBr, HI, HClO 4, and HNO 3 –Diprotic: H 2 SO 4 For a monoprotic strong acid, the acid concentration equals the hydronium concentration. –[H 3 O + ] = [HAcid] –Example: 0.10 M HCl has [H 3 O + ] = 0.10 M and pH = 1.00 For H 2 SO 4, the second ionization can generally be ignored. –However, you must account for the TWO H+ ions when determining the pH. Example: 0.10 M H 2 SO 4 has [H 3 O + ] = 0.20 M and pH = 0.70

© 2015 Pearson Education, Inc. Finding the pH of a Weak Acid: Using ICE There are also two sources of H 3 O + in an aqueous solution of a weak acid—the acid and the water. However, finding the [H 3 O + ] is complicated by the fact that the acid undergoes only partial ionization. Calculating the [H 3 O + ] requires solving an equilibrium problem (using an ICE format) for the reaction that defines the acidity of the acid. HAcid + H 2 O Acid – + H 3 O +

© 2015 Pearson Education, Inc. Practice Problem: Calculating [H 3 O + ] for a Weak Acid using K a

© 2015 Pearson Education, Inc. Practice Problem continued

© 2015 Pearson Education, Inc. Practice Problem: Calculating pH for a Weak Acid

© 2015 Pearson Education, Inc. Practice Problem continued

© 2015 Pearson Education, Inc. Practice Problem: Calculating pH for a Weak Acid when 5% Approximation Does Not Work

© 2015 Pearson Education, Inc. Practice Problem continued

© 2015 Pearson Education, Inc. Practice Problem continued

© 2015 Pearson Education, Inc. Practice Problem: Finding Equilibrium Constant K a when pH Is Known

© 2015 Pearson Education, Inc. Percent Ionization Another way to measure the strength of an acid is to determine the percentage of acid molecules that ionize when dissolved in water; this is called the percent ionization. –The higher the percent ionization, the stronger the acid. Because [ionized acid] equil = [H 3 O + ] equil × 100 = percent ionization molarity of ionized acid initial molarity of weak acid × 100 = percent ionization [H 3 O] + equil [HA] init

© 2015 Pearson Education, Inc. Relationship between [H 3 O + ] equil and [HA] init Increasing the initial concentration of acid results in increased [H 3 O + ] at equilibrium. Increasing the initial concentration of acid results in decreased percent ionization. This means that the increase in [H 3 O + ] concentration is slower than the increase in acid concentration. × 100 = percent ionization [H 3 O] + equil [HA] init

© 2015 Pearson Education, Inc. Practice Problem: Percent Ionization of a Weak Acid

© 2015 Pearson Education, Inc. Why Doesn’t the Increase in H 3 O + Keep Up with the Increase in HA? The reaction for ionization of a weak acid is as follows: HA(aq) + H 2 O(l) A − (aq) + H 3 O + (aq) According to Le Châtelier’s principle, if we reduce the concentrations of all the (aq) components, the equilibrium should shift to the right to increase the total number of dissolved particles. –The (aq) concentrations can be reduced by using a more dilute initial acid concentration. The result will be a larger [H 3 O + ] in the dilute solution compared to the initial acid concentration. –This will result in a larger percent ionization.

© 2015 Pearson Education, Inc. Finding the pH of Mixtures of Acids Generally, you can ignore the contribution of the weaker acid to the [H 3 O + ] equil. For a mixture of a strong acid with a weak acid, the complete ionization of the strong acid provides more than enough [H 3 O + ] to shift the weak acid equilibrium to the left so far that the weak acid’s added [H 3 O + ] is negligible. For mixtures of weak acids, you generally need to consider only the stronger for the same reasons, as long as one is significantly stronger than the other and their concentrations are similar.

© 2015 Pearson Education, Inc. Practice Problem: Mixtures of Weak Acids

© 2015 Pearson Education, Inc. Strong Bases: K b > 1 Hydroxide compounds are strong bases. The stronger the base, the more willing it is to accept H in water (where water is acting as an acid). For ionic bases, practically all units are dissociated into OH – or accept H’s. –Strong electrolyte –Multi-OH strong bases completely dissociated

© 2015 Pearson Education, Inc. Finding pOH and pH for a Strong Base Solution For a strong mono hydroxyl ionic base, the [BOH] = [OH − ]. –Example: 0.10 M KOH has [OH – ] = 0.10 M To find the pH of a mono hydroxyl ionic base, first find the pOH and then determine the pH. –Example: 0.10 M KOH has 0.10 M [OH – ] ions. pOH = –log [OH – ] or pOH = –log [0.10], so pOH is 1.0. pH + pOH = 14, so pH = 14; thus, the pH is 13. For strong poly hydroxyl ionic base compounds, the [OH – ] is equal to number of OH – ions in the base. –Example: 0.10 M Ca(OH) 2 has [OH − ] = 0.20 M and pH = pOH = –log [OH – ] or pOH = –log [0.20], so pOH is pH + pOH = 14, so pH = 14; thus, the pH is 13.

© 2015 Pearson Education, Inc. Practice Problem: Finding pOH and pH of a Strong Base

© 2015 Pearson Education, Inc. Weak Bases: K b < 1 In weak bases, only a small fraction of molecules accept H’s. –Weak electrolyte –Most of the weak base molecules do not take H from water. –Much less than 1% ionization in water [OH – ] << [weak base] Finding the pH of a weak base solution is similar to finding the pH of a weak acid using ICE.

© 2015 Pearson Education, Inc. Base Ionization Constant, K b Base strength is measured by the size of the equilibrium constant when it reacts with H 2 O. :Base + H 2 O (l) OH − + H:Base + The equilibrium constant is called the base ionization constant, K b. –The larger the K b, the stronger the base.

© 2015 Pearson Education, Inc. Table of Common Weak Bases

© 2015 Pearson Education, Inc. Structure of Amines: Organic Base Compounds

© 2015 Pearson Education, Inc. Practice Problem: Finding [OH – ] and pH of a Weak Base

© 2015 Pearson Education, Inc. Practice Problem continued

© 2015 Pearson Education, Inc. Acid–Base Properties of Ions and Salts Salts are water-soluble ionic compounds. Salts that contain the cation of a strong base and an anion that is the conjugate base of a weak acid are basic. –NaHCO 3 solutions are basic. Na + is the cation of the strong base NaOH. HCO 3 − is the conjugate base of the weak acid H 2 CO 3. Salts that contain cations that are the conjugate acid of a weak base and an anion of a strong acid are acidic. –NH 4 Cl solutions are acidic. NH 4 + is the conjugate acid of the weak base NH 3. Cl − is the anion of the strong acid HCl.

© 2015 Pearson Education, Inc. Table of Conjugate Acid–Base Pair Strength

© 2015 Pearson Education, Inc. Anions as Weak Bases Every anion can be thought of as the conjugate base of an acid. –Therefore, every anion can potentially be a base. A − (aq) + H 2 O(l) HA(aq) + OH − (aq) The stronger the acid, the weaker the conjugate base. An anion that is the conjugate base of a strong acid is pH neutral. Cl − (aq) + H 2 O(l) ← HCl(aq) + OH − (aq) An anion that is the conjugate base of a weak acid is basic. F − (aq) + H 2 O(l) HF(aq) + OH − (aq)

© 2015 Pearson Education, Inc. Practice Problem: Determining whether an Anion Is Basic or Neutral in Water

© 2015 Pearson Education, Inc. Relationship between K a of an Acid and K b of Its Conjugate Base Many reference books give tables of only K a values because K b values can be found from them by using the relationship K a × K b = K w When you add equations, you multiply the K’s.

© 2015 Pearson Education, Inc. Practice Problem: Finding pH of an Anion Acting as a Base

© 2015 Pearson Education, Inc. Practice Problem continued

© 2015 Pearson Education, Inc. Cations as Weak Acids Some cations can be thought of as the conjugate acid of a weak base. –Others are the counterions of a strong base. Therefore, some cations can potentially be acidic. MH + (aq) + H 2 O(l) MOH(aq) + H 3 O + (aq) The stronger the base, the weaker the conjugate acid. –A cation that is the counterion of a strong base is pH neutral. –A cation that is the conjugate acid of a weak base is acidic. NH 4 + (aq) + H 2 O(l) NH 3 (aq) + H 3 O + (aq) Because NH 3 is a weak base, the position of this equilibrium favors the right.

© 2015 Pearson Education, Inc. Metal Cations as Weak Acids Cations of small, highly charged metals are weakly acidic. –Alkali metal cations and alkali earth metal cations are pH neutral. –Cations are hydrated. Al(H 2 O) 6 3+ (aq) + H 2 O(l) Al(H 2 O) 5 (OH) 2+ (aq) + H 3 O + (aq)

© 2015 Pearson Education, Inc. Practice Problem: Determining whether a Cation Is Acidic or Neutral in Water

© 2015 Pearson Education, Inc. Classifying Salt Solutions as Acidic, Basic, or Neutral If the salt cation is the counterion of a strong base and the anion is the conjugate base of a strong acid, a neutral solution will be formed. –NaCl Ca(NO 3 ) 2 KBr If the salt cation is the counterion of a strong base and the anion is the conjugate base of a weak acid, a basic solution will be formed. –NaF Ca(C 2 H 3 O 2 ) 2 KNO 2 If the salt cation is the conjugate acid of a weak base and the anion is the conjugate base of a strong acid, an acidic solution will be formed. –NH 4 Cl If the salt cation is a highly charged metal ion and the anion is the conjugate base of a strong acid, an acidic solution will be formed. –Al(NO 3 ) 3

© 2015 Pearson Education, Inc. Classifying Salt Solutions as Acidic, Basic, or Neutral If the salt cation is the conjugate acid of a weak base and the anion is the conjugate base of a weak acid, the pH of the solution depends on the relative strengths of the acid and base. –In NH 4 F, because HF is a stronger acid than NH 4 +, K a of NH 4 + is larger than K b of F − ; therefore, the solution will be acidic.

© 2015 Pearson Education, Inc. Table of pH of Salt Solutions

© 2015 Pearson Education, Inc. Practice Problem: Is the Salt Solution Acidic, Basic, or Neutral?

© 2015 Pearson Education, Inc. Practice Problem continued

© 2015 Pearson Education, Inc. Ionization in Polyprotic Acids Because polyprotic acids ionize in steps, each H has a separate K a value. –K a1 > K a2 > K a3 Generally, the difference in K a values is great enough so that the second ionization does not happen to a large enough extent to affect the pH. –Most pH problems just do first ionization. –[A 2− ] = K a2 as long as the second ionization is negligible.

© 2015 Pearson Education, Inc. Table of Polyprotic Acids and Their K a Values

© 2015 Pearson Education, Inc. Practice Problem: Determining the pH of Polyprotic Acids

© 2015 Pearson Education, Inc. Practice Problem: Determining the pH of Polyprotic Acids

© 2015 Pearson Education, Inc. Practice Problem continued

© 2015 Pearson Education, Inc. Practice Problem: Finding the Concentration of Anions in a Weak Diprotic Acid

© 2015 Pearson Education, Inc. Lewis Acid–Base Theory Lewis acid–base theory focuses on transferring an electron pair. –Lone pair → bond –Bond → lone pair It does NOT require H atoms to be classified as an acid. The electron donor is called the Lewis base. –It is electron rich; therefore, it is a nucleophile. The electron acceptor is called the Lewis acid. –It is electron deficient; therefore, it is an electrophile.

© 2015 Pearson Education, Inc. Lewis Acids: Electron Pair Acceptors They are electron deficient, either from being attached to electronegative atom(s) or as a result of not having an octet. They must have an empty orbital willing to accept the electron pair. –H + has an empty 1s orbital. –B in BF 3 has an empty 2p orbital and an incomplete octet. –Many small, highly charged metal cations have empty orbitals they can use to accept electrons. Atoms that are attached to highly electronegative atoms and have multiple bonds can be Lewis acids.

© 2015 Pearson Education, Inc. Lewis Bases: Electron Pair Donors The Lewis base has electrons it is willing to give away to or share with another atom. The Lewis base must have a lone pair of electrons on it that it can donate. Anions are better Lewis bases than neutral atoms or molecules. –N: < N: − Generally, the more electronegative an atom, the less willing it is to be a Lewis base. –O: < S:

© 2015 Pearson Education, Inc. Lewis Acid–Base Reactions The base donates a pair of electrons to the acid. It generally results in a covalent bond forming. H 3 N: + BF 3 → H 3 N—BF 3 The product that forms is called an adduct. Arrhenius and Brønsted–Lowry acid–base reactions are also Lewis acid–base reactions.

© 2015 Pearson Education, Inc. Examples of Lewis Acid–Base Reactions