CHEM, 2nd edition Cengage Learning Chapter 9 Acid-base reactions Acids and bases are chemical compounds that occur regularly in 'everyday life'. These two types of substances have opposite properties. They often occur in the foods we eat.
Shawn McDonald Linn-Benton Community College Types of Electrolytes
• Salts are water-soluble ionic compounds. All are strong electrolytes. Example: NaCl • Acids form H+1 ions in water solution. • Bases combine with H+1 ions in water solution. Bases increase the OH-1 concentration of the solution. -1 +1 May either directly release OH or pull H off H2O molecule. If the latter, this forms OH- ion and a different positively charged ion.
2 Properties of Acids • Sour taste. Like biting into a lemon... • React with “active” metals. I.e., Al, Zn, Fe, but not Cu, Ag or Au.
2 Al + 6 HCl AlCl3 + 3 H2 Corrosive. To Al (mid-right) and skin!
• React with carbonates, producing CO2. Marble, baking soda, chalk, limestone.
CaCO3 + 2 HCl CaCl2 + CO2(g) + H2O • Change color of vegetable dyes. Blue litmus turns red. Picture at right. • React with bases to form ionic salts and water. Called a neutralization. 3 Common Acids Chemical name Formula Uses Strength
Nitric acid HNO3 Explosive, fertilizer, dye, glue Strong Explosive, fertilizer, dye, glue, Sulfuric acid H SO Strong 2 4 batteries Metal cleaning, food prep, ore Hydrochloric acid HCl Strong refining, stomach acid Fertilizer, plastics and rubber, Phosphoric acid H PO Moderate 3 4 food preservation Plastics and rubber, food Acetic acid HC H O Weak 2 3 2 preservation, vinegar Hydrofluoric acid HF Metal cleaning, glass etching Weak
Carbonic acid H2CO3 Soda water Weak
Boric acid H3BO3 Eye wash Weak
4 Structures of Acids • Binary acids have acid hydrogens attached to a nonmetal atom. 2 types of elements only. HCl, HF Write the H atom first, then the nonmetal atom. Dissociate in water to form H+ ions and Hydrofluoric acid nonmetal anions (such as Cl- or F-) 5 Structure of Acids • Oxyacids have acid hydrogens attached to an oxygen atom.
H2SO4, HNO3 Also write the H atom(s) first, then the polyatomic ion group. Will dissociate when put into water, to give H+ ions and a polyatomic anion - 2- (like NO3 or SO4 )
6 Structure of Acids
• Carboxylic acids have COOH group.
HC2H3O2, H3C6H5O3 Component of vinegar • Only the first H in the formula is acidic. The H is on the COOH.
Lemons and limes Apples and wine
7 Properties of Bases • Also known as alkalis. • Taste bitter. Alkaloids = Plant product that is alkaline. Often poisonous. Potato and tomato shoots. • Solutions feel slippery. • Change color of vegetable dyes. Occurs in hemlock bark, Different color than acid. repels beetles. Red litmus turns blue. The dye in the paper reacts with OH- from base. • React with acids to form ionic salts. Neutralization. Negates taste and metal dissolving power of acids.
8 Common Bases Chemical Common Formula Uses Strength name name Sodium Lye, Soap, plastic, NaOH Strong hydroxide caustic soda petrol refining Potassium Caustic Soap, cotton, KOH Strong hydroxide potash electroplating Calcium Ca(OH) Slaked lime Cement Strong hydroxide 2 Sodium NaHCO Baking soda Cooking, antacid Weak bicarbonate 3 Magnesium Milk of Mg(OH) Antacid Weak hydroxide 2 magnesia Detergent, Ammonium NH OH, Ammonia 4 fertilizer, Weak hydroxide {NH (aq)} water 3 explosives, fibers
9 Structure of Bases • Most ionic bases contain OH ions.
NaOH, Ca(OH)2 2- • Some contain CO3 ions.
CaCO3 NaHCO3 • Molecular bases contain structures that react with H+. Mostly amine groups (N atoms). Caffeine has three amine type
groups with CH2 group attached. 10 9-1b What is an acid or a base? • An acid–base reaction is any reaction in which an H+ is transferred. Does not have to take place in aqueous solution. Broader definition than Arrhenius. • An acid is a H+ donor; A base is a H+ acceptor. Either can be a molecule or an ion. Since H+ is a proton, acid is a proton donor and base is a proton acceptor. Base structure must contain an atom with an unshared pair of electrons to bond to H+. • In the reaction, the acid molecule gives an H+ to the base molecule. H–A + :B :A– + H–B+
11 Amphoteric Substances • Amphoteric substances can act as either an acid or a base. They have both a transferable H atom and an atom with a lone pair. • HCl(aq) is acidic because HCl transfers an H+ to + H2O, forming H3O ions. Water acts as base, accepting H+. – + HCl(aq) + H2O(l) → Cl (aq) + H3O (aq) + • NH3(aq) is basic because NH3 accepts an H from – H2O, forming OH (aq). Water acts as acid, donating H+. + – NH3(aq) + H2O(l) NH4 (aq) + OH (aq) Thus water is amphoteric, it can act as a base with an acid, or as an acid with a base. Its nature is the opposite of the compound with which it is interacting. 12 An example acid-base reaction – +: In the reaction H2O + NH3 HO + NH4 water ammonia hydroxide ion ammonium ion – H2O and HO constitute an acid/conjugate–base pair. If hydroxide ion accepts a proton it will revert to a water molecule.
+ NH3 and NH4 constitute a base/conjugate–acid pair. If the ammonium ion donates a proton to a base, it will revert to the ammonia molecule. 13 Example—Identify the Brønsted–Lowry acids and bases and their conjugates in this reaction.
+ - C. HNO3(aq) + H2O(l) H3O (aq) + NO3 (aq)
14 Neutralization Reactions + - • H + OH H2O net ionic eqn. • Acid + base salt + water • Double-displacement reactions. Salt = cation from base + anion from acid. Sometimes the salt in insoluble. Cation and anion charges stay constant.
H2SO4 + Ca(OH)2 → CaSO4 + 2 H2O • Some neutralization reactions are gas evolving, where H2CO3 (carbonic acid) decomposes into CO2 and H2O.
H2SO4 + 2 NaHCO3 → Na2SO4 + 2 H2O + 2 CO2 sulfuric acid sodium bicarbonate sodium sulfate water carbon dioxide gas
15 Example — Write the equation for the reaction of aqueous perchloric acid with strontium hydroxide.
1. Write the formulas of the reactants.
HClO4(aq) + Sr(OH)2(aq) 2. Determine the ions present when each reactant dissociates. + − 2+ − (H + ClO4 ) + (Sr + OH ) 3. Exchange the ions. H+1 combines with OH-1 to
make H2O(l). Like other double displacements. + − 2+ − 2+ − (H + ClO4 ) + (Sr + OH ) (Sr + ClO4 ) + H2O(l)
16 Write the equation for reaction of aqueous perchloric acid with strontium hydroxide, Continued. 4. Write the formulas of the products. Cross charges and reduce subscripts if possible.
HClO4(aq) + Sr(OH)2(aq) Sr(ClO4)2 + H2O(l) 5. Balance the equation. Each atom and group on left vs. right side of the equation. May be quickly balanced by matching the numbers of
H and OH to make H2O. Coefficient of the salt is always 1.
2 HClO4(aq) + Sr(OH)2(aq) Sr(ClO4)2 + 2 H2O(l) 17 Write the Equation for reaction of aqueous perchloric acid with strontium hydroxide.Continued 6. Determine the solubility of the salt.
Sr(ClO4)2 is soluble (look up in solubility table). 7. Write an (s) after the insoluble products and an (aq) after the soluble products.
2 HClO4(aq) + Sr(OH)2(aq) Sr(ClO4)2(aq) + 2 H2O(l) The reaction occurs since one of the products formed is water, and water molecules mainly stay as molecules and don't ionize very much.
18 9-2 Strong or Weak • A strong acid is a strong electrolyte. Practically all the acid molecules ionize, →. completely • A strong base is a strong electrolyte. Practically all the base molecules form OH– ions, either through dissociation or reaction with water, →. completely • A weak acid is a weak electrolyte. Only a small percentage of the molecules ionize, . • A weak base is a weak electrolyte. Only a small percentage of the base molecules form OH– ions, either through dissociation or reaction with water, .
19 Strong Acids • The stronger the acid, the more willing it is to donate HCl H+ + Cl- + - H. HCl + H2O H3O + Cl Use water as the standard base. • Strong acids donate practically all their H’s. 100% ionized in water. Strong electrolyte. + • [H3O ] = [strong acid]. [ ] = molarity. No HCl is left after you dissolve the compound in water. 20 Strong Acids, Continued
Hydrochloric acid HCl Hydrobromic acid HBr Hydroiodic acid HI
Nitric acid HNO3
Perchloric acid HClO4
Sulfuric acid H2SO4
21 Strong Acids, Continued
Pure water HCl solution HCl is a strong electrolyte. Its Water is not an electrolyte and solution will conduct current. will not conduct current. 22 Weak Acids • Weak acids donate a small HF H+ + F- fraction of their Hs. + - HF + H2O H3O + F Most of the weak acid molecules do not donate H to water. Often less than 1% ionized in water. + • [H3O ] << [weak acid]. Most HF stays in this molecular form, does not ionize in water to + + form H or H3O 23 Weak Acids, Continued
Hydrofluoric acid HF
Acetic acid HC2H3O2
Formic acid HCHO2
Sulfurous acid H2SO3
Carbonic acid H2CO3
Phosphoric acid H3PO4 Stronger attraction between H and F than between H+ and water 24 Weak Acids, Continued
HF in water is a weak electrolyte and only conducts electricity poorly – note dimness of light bulb. Pure water HF solution
25 Degree of Ionization • The extent to which an acid ionizes in water depends in part on the strength of the bond between the acid H+ and anion compared to the strength of the bond between the acid H+ and the O of water. − + HA(aq) + H2O(l) A (aq) + H3O (aq) In other words, which is stronger? the H-A bond or the H-O bond in the hydronium ion. If the former, then the acid will only ionize slightly. Fluorine for example forms a stronger bond with H than does Cl, thus HF is a weak acid and HCl is a strong acid. 26 Strong Bases • The stronger the base, the more willing it is to accept H. NaOH Na+ + OH- Use water as the standard acid. • Strong bases, practically all molecules are dissociated into OH– or accept Hs. Strong electrolyte. Multi-OH bases completely dissociated. • [HO–] = [strong base] x (# OH). All the dissolved NaOH has • Molarity will be discussed dissociated into sodium and shortly.... hydroxide ions. 27 Strong Bases, Continued
Lithium hydroxide LiOH Sodium hydroxide NaOH Potassium hydroxide KOH
Calcium hydroxide Ca(OH)2
Strontium hydroxide Sr(OH)2
Barium hydroxide Ba(OH)2
28 Weak Bases
• In weak bases, only a small NH + H O NH + + OH- fraction of molecules accept 3 2 4 Hs. Weak electrolyte. Most of the weak base molecules do not take H from water. Much less than 1% ionization in water. – • [OH ] << [weak base]. Most of the ammonium hydroxide molecules do not accept H from water molecules.
29 Weak Bases, Continued
+ − Ammonia NH3(aq) + H2O(l) NH4 (aq) + OH (aq)
+ − Pyridine C5H5N(aq) + H2O(l) C5H5NH (aq) + OH (aq)
+ − Methyl amine CH3NH2(aq) + H2O(l) CH3NH3 (aq) + OH (aq)
+ − Ethyl amine C2H5NH2(aq) + H2O(l) C2H5NH3 (aq) + OH (aq)
− − Bicarbonate HCO3 (aq) + H2O(l) H2CO3 (aq) + OH (aq) Most of the base molecules or species stay in the unprotonated form.
30 Autoionization of Water • Water is actually an extremely weak electrolyte. Therefore, there must be a few ions present. • About 1 out of every 10 million water molecules form ions through a process called autoionization. + – H2O + H2O H3O + OH + – • All aqueous solutions contain both H3O and OH . + – The concentration of H3O and OH are equal in DI water. + – -7 [H3O ] = [OH ] = 1 x 10 M at 25 °C in pure water. These are important concentrations to remember (related to pH that will we study shortly).
31 Ion Product of Water + – • The product of the H3O and OH concentrations is always the same number for solutions at 25 Celsius. • The number is called the ion product of water and has the symbol Kw. + – -14 • [H3O ] x [OH ] = 1 x 10 = Kw. + – • As [H3O ] increases, the [OH ] must decrease so the product stays constant. Inversely proportional.
32 Acidic and Basic Solutions
+ – • Neutral solutions have equal [H3O ] and [OH ]. + – -7 [H3O ] = [OH ] = 1 x 10 + – • Acidic solutions have a larger [H3O ] than [OH ]. + -7 – -7 [H3O ] > 1 x 10 ; [OH ] < 1 x 10 – + • Basic solutions have a larger [OH ] than [H3O ]. + -7 – -7 [H3O ] < 1 x 10 ; [OH ] > 1 x 10
We can measure the concentration of hydronium ions in a solution by using a pH meter. These devices are used in the general chemistry courses to study acid-base behavior and reactions. pH paper also works to give us the rough pH of a given aqueous solution. 33 + Example—Determine the [H3O ] for a 0.00020 M Ba(OH)2 solution and Determine Whether the Solution Is Acidic, Basic, or Neutral. 2+ – Ba(OH)2 = Ba + 2 OH therefore: [OH–] = 2 x 0.00020 = 0.00040 = 4.0 x 10−4 M
Kw H3O OH 14 Kw 110 H3O OH 4.0104 + -11 [H3O ] = 2.5 x 10 M. + −7 Since [H3O ] < 1 x 10 , the solution is basic.
34 9-3a The pH scale
• The acidity/basicity of a solution is often expressed as pH. + + −pH • pH = ─log[H3O ], [H3O ] = 10 The exponent on 10, but with a positive sign. -7 pHwater = −log[10 ] = 7. + Need to know the [H3O ] concentration to find pH. • 3 cases: case 1: pH < 7 is acidic; case 2: pH > 7 is basic; case 3: pH = 7 is neutral.
35 pH, Continued • The lower the pH, the more acidic the solution; the higher the pH, the more basic the solution. 1 pH unit corresponds to a factor of 10 fold difference in acidity of that solution. • Normal range is 0 to 14. pH 0 is [H+] = 1 M, pH 14 is [OH–] = 1 M. pH can be negative (very acidic) or larger than 14 (very alkaline, at high concentration).
36 pH of Common Substances Substance pH 1.0 M HCl 0.0 0.1 M HCl 1.0 Stomach acid 1.0 to 3.0 Lemons 2.2 to 2.4 Soft drinks 2.0 to 4.0 Plums 2.8 to 3.0 Apples 2.9 to 3.3 Cherries 3.2 to 4.0 Unpolluted rainwater 5.6 Human blood 7.3 to 7.4 Egg whites 7.6 to 8.0
Milk of magnesia (saturated Mg(OH)2) 10.5 Household ammonia 10.5 to 11.5 1.0 M NaOH 14 37 Example—Calculate the pH of a 0.0010 M
Ba(OH)2 Solution and Determine if It Is Acidic, Basic, or Neutral. 2+ − Ba(OH)2 = Ba + 2 OH therefore, [OH-] = 2 x 0.0010 = 0.0020 = 2.0 x 10-3 M. 1 x 10-14 [H O+] = = 5.0 x 10-12M 3 2.0 x 10-3
+ -12 pH = −log [H3O ] = −log (5.0 x 10 ) pH = 11.3
pH > 7 therefore, basic.
38 Practice—Calculate the pH of the Following Strong Acid or Base Solutions. • 0.0020 M HCl
• 0.0050 M Ca(OH)2
• 0.25 M HNO3
39 Practice—Calculate the pH of the Following Strong Acid or Base Solutions, Continued. + • 0.0020 M HCl strong acid therefore, [H3O ] = 0.0020 M. pH = −log (2.0 x 10-3) = 2.70 acidic – • 0.0050 M Ca(OH)2 strong base, [OH ] = 0.010 M. −14 + 1 x 10 −12 [H3O ] = = 1.0 x 10 1 x 10−2 pH = −log (1.0 x 10−12) = 12.00 basic + • 0.25 M HNO3 a strong acid, therefore, [H3O ] = 0.25 M. pH = −log (2.5 x 10−1) = 0.60 acidic
40 9-4 Acid-base buffers • Buffers are solutions that resist changing pH when small amounts of acid or base are added. • They resist changing pH by neutralizing added acid or base. • Buffers are made by mixing together a weak acid and its conjugate base. Or weak base and its conjugate acid. For example, a mixture of ammonia and ammonium chloride is a buffer.
41 How Buffers Work • The weak acid present in the buffer mixture can neutralize added base. • The conjugate base present in the buffer mixture can neutralize added acid. • The net result is little to no change in the solution pH. As long as you don’t over come the capacity of the buffer. Then the pH will change drastically. Can change by several pH units after capacity is overcome.
42 Acetic Acid/Acetate Buffer
The conjugate base neutralizes any added The weak acid neutralizes acid…. makes weak acid the added base, forms H2O. (acetic acid) 43