Chapter 11: Acids and Bases
11.1 Brønsted-Lowry Acids and Bases
11.2 Lewis Acids and Bases
11.3 Acidic, Basic, and Amphoteric Oxides
11.4 Proton Exchange Between Water Molecules
11.5 The pH Scale
11.6 The pOH of Solutions
11.7 Acidity and Basicity Constants
11.8 The Conjugate Seesaw
11.9 Molecular Structure and Acid Strength
11.10 The Strengths of Oxoacids and Carboxylic Acids
11.11 Solutions of Weak Acids
11.12 Solutions of Weak Bases
11.13 The pH of Salt Solutions
11.14 The pH of a Polyprotic Acid Solution
11.15 Solutions of Salts of Polyprotic Acids
11.16 The Concentrations of Solute Species
11.7 Composition and pH
11.18 Very Dilute Solutions of Strong Acids and Bases
11.19 Very Dilute Solutions of Weak Acids
Terms : Hide Images
| 1585620090 | Brønsted Acid | A proton donor. | 0 | |
| 1585620091 | Brønsted Base | A proton acceptor. | 1 | |
| 1585620092 | Acid | A species containing an acidic hydrogen atom. | 2 | |
| 1585620093 | Acidic Hydrogen Atom | A hydrogen that can be transferred as a proton from one species to another. | 3 | |
| 1585620094 | Brønsted-Lowry Theory | That an acid is a proton donor and a base is a proton acceptor. | 4 | |
| 1585620095 | Proton Transfer Reaction | A reaction in which a proton is transferred from one species to another. | 5 | |
| 1585620096 | Deprotonation | When a molecule loses a proton (hydrogen atom). | 6 | |
| 1585620097 | Hydronium Ion | H₃O⁺ | 7 | |
| 1585620098 | Strong Acid | An acid that is completely deprotonated in solution. | 8 | |
| 1585620099 | Weak Acid | An acid that is incompletely deprotonated in solution. | 9 | |
| 1585620100 | Protonation | When a molecule accepts a proton (hydrogen atom). | 10 | |
| 1585620101 | Strong Base | A base that is completely protonated in solution. | 11 | |
| 1585620102 | Weak Base | A base that is incompletely protonated in solution. | 12 | |
| 1585620103 | Conjugate Base | The species associated with an acid that is left when the acid donates a proton. | 13 | |
| 1585620104 | Conjugate Acid | The species associated with a base that is left when the base accepts a proton. | 14 | |
| 1585620105 | Lewis Acid | An electron pair acceptor. | 15 | |
| 1585620106 | Lewis Base | An electron pair donor. | 16 | |
| 1585620107 | 11.2 Lewis Acids and Bases Summary | A Lewis acid is an electron pair acceptor; a Lewis base is an electron pair donor. A proton is a Lewis acid that attaches to a lone pair provided by a Lewis base. | 17 | |
| 1585620108 | Acidic Oxide | An oxide that reacts with water to form a solution of a Brønsted acid. Acidic oxides are molecular compounds. | 18 | |
| 1585620109 | Basic Oxide | An oxide that reacts with water to form a solution of hydroxide ions. Basic oxides are ionic compounds. | 19 | |
| 1585620110 | 11.3 Acidic, Basic, and Amphoteric Oxides Summary | Metals form basic oxides, nonmetals form acidic oxides; the elements on a diagonal line from beryllium to polonium and several d-block metals form amphoteric oxides. | 20 | |
| 1585620111 | Amphiprotic | Able to act as both a proton donor and proton acceptor. Water is amphiprotic. | 21 | |
| 1585620112 | Autoprotolysis | When a molecule transfers a proton to another molecule of the same kind. | 22 | |
| 1585620113 | Equation for the autoprotolysis of water | K = a₋H₃O⁺ a₋OH⁻ / (a₋H₂O)² | 23 | |
| 1585620114 | Autoprotolysis Constant of Water | K₋w = a₋H₃O⁺ a₋OH⁻ K₋w = [H₃O⁺][OH⁻] | 24 | |
| 1585620115 | 11.4 Proton Exchange Between Water Molecules Summary | In aqueous solutions, the concentration of H₃O⁺ and OH⁻ ions are related by the autoprotolysis equilibrium; if one concentration is increased, then the other must decrease to maintain the value of K₋w. | 25 | |
| 1585620116 | pH is the negative logarithm of the hydronium ion activity | pH = -log a₋H₃O⁺ | 26 | |
| 1585620117 | Simplifying -log a₋H₃O⁺ | a₋H₃O⁺ = [H₃O⁺]/cº strike out units pH = -log[H₃O⁺] | 27 | |
| 1585620118 | The pH of a basic solution is... | greater than 7 | 28 | |
| 1585620119 | The pH of pure water is... | 7 | 29 | |
| 1585620120 | The pH of an acidic solution is... | less than 7 | 30 | |
| 1585620121 | 11.5 The pH Scale Summary | The pH scale is used to report H₃O⁺ concentration: pH = -log[H₃O⁺]; pH > 7 denotes a basic solution, pH < 7 an acidic solution; a neutral solution has pH = 7. | 31 | |
| 1585620122 | The quantity pX is a generalization of pH; pOH would then be... | pOH = -log[OH⁻] | 32 | |
| 1585620123 | pH + pOH = pKw | Kw = [H₃O⁺][OH⁻] logKw = log[H₃O⁺][OH⁻] logKw = log[H₃O⁺] + log[OH⁻] -logKw = -log[H₃O⁺] - log[OH⁻] pH + pOH = pKw | 33 | |
| 1585620124 | 11.6 The pOH of Solutions Summary | The pH and pOH of a solution are related by the experssion pH + pOH = pKw. | 34 | |
| 1585620125 | Acid Ionization Constant | Acid Dissociation Constant | 35 | |
| 1585620126 | Base Ionization Constant | AcBased Dissociation Constant | 36 | |
| 1585620127 | Acidity Constant Equation | Ka = [H₃O⁺][A⁻]/[HA] | 37 | |
| 1585620128 | Basicity Constant Equation | Kb = [OH⁻][HB⁺]/[B] | 38 | |
| 1585620129 | 11.7 Acidity and Basicity Constants Summary | The proton-donating strength of an acid is measured by its acidity constant; the proton-accepting strength of a base is measured by its basicity constant. The smaller the constants, the weaker the respective strengths. The larger the value of pK, the weaker the acid or base. | 39 | |
| 1585620130 | Ka and Kb are related by... | Kw = Ka × Kb *Kw = 14 | 40 | |
| 1585620131 | pKa and pKb are related by... | pKw = pKa + pKb | 41 | |
| 1585620132 | Leveled | All strong acids in water behave as though they were solutions of H₃O⁺, in water they are leveled to the strength of the acid H₃O⁺. | 42 | |
| 1585620133 | Sulfuric acid as a special case | The loss of its first acidic hydrogen leaves a conjugate base that is itself a weak acid, the HSO₄⁻. | 43 | |
| 1585620134 | 11.8 The Conjugate Seesaw Summary | The stronger the acid, the weaker its conjugate base; the stronger the base, the weaker the conjugate acid. | 44 | |
| 1585620135 | 11.9 Molecular Structure and Acid Strength Summary | Acid strengths of binary acids across a period correlate with electron affinities; acid strengths down a group correlate with bond strength. | 45 | |
| 1585620136 | Binary Acids | Acids composed of a hydrogen bonded to a single other atom of a nonmetallic element. | 46 | |
| 1585620137 | Acid strength and Polarity | The more polar the H-A bond, the stronger the acid. This effect is dominant for acids of the same period. | 47 | |
| 1585620138 | Acid strength and Bond strength | The weaker the H-A bond, the stronger the acid. This effect is dominant for acids of the same group. | 48 | |
| 1585620139 | Oxoacids | Acids containing at least one oxygen atom. | 49 | |
| 1585620140 | Hypohalous Acids | Acids of oxygen and a halogen; the more electronegative the halogen, the stronger the oxoacid. | 50 | |
| 1585620141 | Acid strength and Oxygen count | The greater the number of oxygen atoms attached to the central atom, the stronger the acid. | 51 | |
| 1585620142 | Acid strength and Oxidation number | The greater the oxidation number of the central atom, the stronger the acid. | 52 | |
| 1585620143 | Carboxylic Acids | Acids containing an -OOH group. | 53 | |
| 1585620144 | Acid strength and R group charges | The greater the electronegativities of the groups attached to the carboxyl group of a carboxylic acid, the stronger the acid. | 54 | |
| 1585620145 | 11.10 The Strengths of Oxoacids and Carboxylic Acids Summary | The greater the number of oxygen atoms and the more electronegative the atoms present in the molecules of an acid, the stronger the acid. | 55 | |
| 1585620146 | Percentage Deprotonation | The percentage of HA molecules that are deprotonated in the solution; = [A⁻]/[HA]ini | 56 | |
| 1585620147 | Initial Concentration, AKA... | Analytical Concentration Formal Concentration (F) | 57 | |
| 1585620148 | 11.11 Solutions of Weak Acids Summary | The calculate the pH and percentage deprotonation of a solution of a weak acid, set up an equilibrium table and determine the H₃O⁺ concentration by using the acidity constant. | 58 | |
| 1585620149 | Percent Protonated | The percentage of base molecules that have been protonated; = [HB⁺]/[B]ini | 59 | |
| 1585620150 | 11.12 Solutions of Weak Bases Summary | To calculate the pH of a solution of a weak base, set up an equilibrium table to calculate pOH from the value of Kb and convert that pOH into pH by using pH + pOH = 14.00. | 60 | |
| 1585620151 | All cations that are the conjugate acids of weak bases produce acidic solutions | Conjugate acids of weak bases, such as NH⁴⁺, act as proton donors, and so we can expect them to form acidic solutions. | 61 | |
| 1585620152 | Small, highly charged metal cations that can act as Lewis acids in water produce acidic solutions, even though the cations themselves have no hydrogen ions to donate | Protons that come from the water molecules which then hydrate these metal cations in solution; the water molecules act as Lewis bases and share electrons with the metal cations. The partial loss of electrons weakens the O-H bond and allows one or more hydrogen ions to be lost from the water molecules. Small, highly charged cations exert the greatest pull on the electrons and so form the most acidic solutions. | 62 | |
| 1585620153 | Cations of Group 1 and 2 metals, as well as those of charge +1 from other groups, are such weak Lewis acids that the hydrated ions do not act as acids | These metal cations are too large or have too low a charge to have an appreciable polarizing effect on the hydrating water molecules that surround them, and so the water molecules do not readily release their protons. | 63 | |
| 1585620154 | Very few anions that contain hydrogen produce acidic solutions | It is difficult for a positively charged proton to leave a negatively charged anion. The few anions that do act as acids include H₂PO₄⁻ and HSO₄⁻. | 64 | |
| 1585620155 | All anions that are the conjugate bases of weak acids produce basic solutions | ex. formic acid, HCOOH, is a weak acid and so the formate ion acts as a base in water: H₂O(l) + HCO₂⁻(aq) ↔ HCOOH(aq) + OH⁻(aq) | 65 | |
| 1585620156 | The anions of strong acids are such weak bases that they have no significant effect on the pH of a solution | They are considered "neutral" in water. Includes: Cl⁻, Br⁻, I⁻, NO₃⁻, ClO₄⁻ | 66 | |
| 1585620157 | Determining the pH of a salt solution | First, determine if it will be acidic, basic, or neutral using the cation and anion of the dissociated salt. Then, use known information (pKa, etc.) to calculate the pH. | 67 | |
| 1585620158 | 11.13 The pH of Salt Solutions Summary | Salts that contain the conjugate acids of weak bases produce acidic aqueous solutions; so do salts that contain small, highly charged metal cations. Salts that contain the conjugate bases of weak acids produce basic aqueous solutions. | 68 | |
| 1585620159 | Polyprotic Acid | A compound that can donate more than one proton. | 69 | |
| 1585620160 | Polyprotic Base | A compound that can accept more than one proton. | 70 | |
| 1585620161 | Trends in acidity constant for polyprotic acids | The acidity constant decreases significantly with each donated proton, usually by a factor of 10³ or more. Ka₁ >>> Ka₂ >>> Ka₃ ... | 71 | |
| 1585620162 | Sulfuric Acid | Strong acid; deprotonates twice easily. H₂SO₄ → HSO₄⁻ → SO₄²⁻ | 72 | |
| 1585620163 | 11.14 The pH of a Polyprotic Acid Solution Summary | Estimating the pH of a polyprotic acid for which all deprotonations are weak by using only the first deprotonation equilibrium and assuming that further deprotonation is insignificant. An exception is sulfuric acid, the only common polyprotic acid that is a strong acid in its first deprotonation. | 73 | |
| 1585620164 | Amphiprotic | Acting as either an acid or base. | 74 | |
| 1585620165 | 11.15 Solutions of Salts of Polyprotic Acids Summary | The pH of the aqueous solution of an amphiprotic salt can be estimated from the average of the pKas of the salt and its conjugate acid. The pH of a solution of a salt of the final conjugate base of a polyprotic acid is found from the reaction of the anion with water. | 75 | |
| 1585620166 | How to calculate the concentrations of all species in a polyprotic acid solution; diprotic acid example | - From the deprotonation equilibrium of the acid (H₂A), determine the concentrations of conjugate base (HA⁻) and H₃O⁺. - Find the concentration of A²⁻ from the second deprotonation equilibrium (that of HA⁻) by substituting the concentrations of H₃O⁺ and HA⁻ into the expression for Ka₂. - Find the concentration of OH⁻ by dividing Kw by the concentration of H₃O⁺. | 76 | |
| 1585620167 | How to calculate the concentrations of all species in a polyprotic acid solution; triprotic acid example | - From the deprotonation equilibrium of the acid (H₃A), determine the concentrations of conjugate base (H₂A⁻) and H₃O⁺. - Find the concentration of HA²⁻ from the second deprotonation equilibrium (that of H₂A⁻) by substituting the concentrations of H₃O⁺ and H₂A⁻ into the expression for Ka₂. - Find the concentration of A³⁻ from the deprotonation equilibrium of HA²⁻ by substituting the concentrations of H₃O⁺ and HA²⁻ into the equation for Ka₃. The concentration of H₃O⁺ stays the same through all the calculations because only the first deprotonation makes a significant contribution to its value. - Find the concentration of OH⁻ by dividing Kw by the concentration of H₃O⁺. | 77 | |
| 1585620168 | 11.16 The Concentrations of Solute Species Summary | The concentrations of all species in a solution of a polyprotic acid can be calculated by assuming that species present in smaller amounts do not affect the concentrations of species present in larger amounts. | 78 | |
| 1585620169 | 11.7 Composition and pH Summary | The fraction of deprotonated species increases as the pH is increased, as summarized in Figs. 11.21 and 11.22. * see the book, pg. 461 | 79 | |
| 1585620170 | The contribution of autoprotolysis to pH is only taken into consideration in strong acid/base solutions when... | the concentration of strong acid or base is less than ~10⁻⁶ mol/L. | 80 | |
| 1585620171 | To calculate the pH when taking autoprotolysis into account... | all species in solution must be taken into consideration, ex. H₃O⁺, OH⁻, and Cl⁻ (from HCl). | 81 | |
| 1585620172 | There are three unknown concentrations... | [H₃O⁺], [OH⁻], and [Cl⁻]. Thus, three equations are needed. | 82 | |
| 1585620173 | The first equation and Charge Balance | The first equation takes into account the requirement that the solution must be electrically neutral overall; the concentration of cations must equal the concentration of anions. The only cation is [H₃O⁺], thus [H₃O⁺] = [OH⁻] + [Cl⁻] [OH⁻] = [H₃O⁺] - [Cl⁻] | 83 | |
| 1585620174 | The second equation and Material balance | The second equation takes into account the requirement that all the added solute must be accounted for. Since Cl⁻:HCl is 1:1, [Cl⁻] = [HCl]ini and [OH⁻] = [H₃O⁺] - [HCl]ini | 84 | |
| 1585620175 | The third equation and the Autoprotolysis Constant | Kw = [H₃O⁺][OH⁻] Kw = [H₃O⁺]([H₃O⁺] - [HCl]ini) [H₃O⁺]² - [HCl]ini[H₃O⁺] - Kw = 0 Solve the quadratic equation. | 85 | |
| 1585620176 | 11.18 Very Dilute Solutions of Strong Acids and Bases Summary | In very dilute solutions of strong acids and bases, the pH is significantly affected by the autoprotolysis of water. The pH is determined by solving three simultaneous equations: the charge-balance equation, the material-balance equation, and the expression for Kw. | 86 | |
| 1585620177 | The contribution of autoprotolysis to pH is only taken into consideration in weak acid solutions... | when the acid is so weak that autoprotolysis contributes a fair amount to the pH. | 87 | |
| 1585620178 | All the species considered in dilute solutions of weak acids | HA, A⁻, H₃O⁺, OH⁻. | 88 | |
| 1585620179 | Four unknown solutions, so four equations needed. The equations are: | The autoprotolysis constant of water. The acidity constant of the acid HA. The charge balance. The material balance. | 89 | |
| 1585620180 | Weak acid autoprotolysis constant of water eq. | Kw = [H₃O⁺][OH⁻] | 90 | |
| 1585620181 | Weak acid acidity constant of the acid HA eq. | Ka = [H₃O⁺][A⁻]/[HA] | 91 | |
| 1585620182 | Weak acid charge balance eq. | [H₃O⁺] = [OH⁻] + [A⁻] | 92 | |
| 1585620183 | Weak acid material balance eq. | [HA]ini = [HA] + [A⁻] | 93 | |
| 1585620184 | Solving for weak acid pHs when considering autoprotolysis of water can lead to a very large equation (after much rearrangement, etc.); however, to simplify... | If [H₃O⁺] > 10⁻⁶ (i.e. pH < 6), then Kw/[H₃O⁺] < 10⁻⁸, which is so small that it can be ignored. However, if [H₃O⁺] ≤ 10⁻⁶, the whole equation still needs to be solved. | 94 | |
| 1585620185 | 11.19 Very Dilute Solutions of Weak Acids Summary | In aqueous solutions of very weak acids, the autoprotolysis of water must be taken into account if the hydronium ion concentration is less than 10⁻⁶ mol/L. The expressions for Kw and Ka are combined with the equations for charge balance and material balance to find the pH. | 95 |
