Chapter 14 Acids and Bases Types of Electrolytes
Chapter 14 Acids and Bases
Types of Electrolytes • salts = water soluble ionic compounds üall strong electrolytes • acids = form H+1 ions in water solution • bases = combine with H+1 ions in water solution üincreases the OH-1 concentration Ømay either directly release OH-1 or pull H+1 off H 2 O 2
Properties of Acids • Sour taste • react with “active” metals ü i. e. Al, Zn, Fe, but not Cu, Ag or Au 2 Al + 6 HCl 2 Al. Cl 3 + 3 H 2 ü corrosive • react with carbonates, producing CO 2 ü marble, baking soda, chalk, limestone Ca. CO 3 + 2 HCl Ca. Cl 2 + CO 2 + H 2 O • change color of vegetable dyes ü blue litmus turns red • react with bases to form ionic salts 3
Common Acids 4
Structures of Acids • binary acids have acid hydrogens attached to a nonmetal atom üHCl, HF Hydrofluoric acid 5
Structure of Acids • oxy acids have acid hydrogens attached to an oxygen atom üH 2 SO 4, HNO 3 6
Structure of Acids • carboxylic acids have COOH group ü HC 2 H 3 O 2, H 3 C 6 H 5 O 3 • only the first H in the formula is acidic ü the H is on the COOH 7
Properties of Bases • also known as alkalis • taste bitter ü alkaloids = plant product that is alkaline Ø often poisonous • solutions feel slippery • change color of vegetable dyes ü different color than acid ü red litmus turns blue • react with acids to form ionic salts ü neutralization 8
Common Bases 9
Structure of Bases • most ionic bases contain OH ions üNa. OH, Ca(OH)2 • some contain CO 32 - ions üCa. CO 3 Na. HCO 3 • molecular bases contain structures that react with H+ ümostly amine groups 10
Arrhenius Theory • bases dissociate in water to produce OHions and cations ü ionic substances dissociate in water Na. OH(aq) → Na+(aq) + OH–(aq) • acids ionize in water to produce H+ ions and anions ü because molecular acids are not made of ions, they cannot dissociate ü they must be pulled apart, or ionized, by the water HCl(aq) → H+(aq) + Cl–(aq) ü in formula, ionizable H written in front HC 2 H 3 O 2(aq) → H+(aq) + C 2 H 3 O 2–(aq) 11
Arrhenius Acid-Base Reactions • the H+ from the acid combines with the OHfrom the base to make a molecule of H 2 O üit is often helpful to 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 HCl(aq) + Na. OH(aq) → Na. Cl(aq) + H 2 O(l) 12
Problems with Arrhenius Theory • does not explain why molecular substances, like NH 3, dissolve in water to form basic solutions – even though they do not contain OH– ions • does not explain acid-base reactions that do not take place in aqueous solution • H+ ions do not exist in water. Acid solutions contain H 3 O+ ions ü H+ = a proton! ü H 3 O+ = hydronium ions 13
Brønsted-Lowery Theory • in a Brønsted-Lowery Acid-Base reaction, an H+ is transferred ü does not have to take place in aqueous solution ü broader definition than Arrhenius • acid is H donor, base is H acceptor ü base structure must contain an atom with an unshared pair of electrons • in the reaction, the acid molecule gives an H+ to the base molecule H–A + : B : A– + H–B+ 14
Amphoteric Substances • amphoteric substances can act as either an acid or a base ühave both transferable H and atom with lone pair • HCl(aq) is acidic because HCl transfers an H+ to H 2 O, forming H 3 O+ ions üwater acts as base, accepting H+ HCl(aq) + H 2 O(l) → Cl–(aq) + H 3 O+(aq) • NH 3(aq) is basic because NH 3 accepts an H+ from H 2 O, forming OH–(aq) üwater acts as acid, donating H+ NH 3(aq) + H 2 O(l) NH 4+(aq) + OH–(aq) 15
Brønsted-Lowery Acid-Base Reactions • one of the advantages of Brønsted-Lowery theory is that it allows reactions to be reversible H–A + : B → : A– + H–B+ • the original base has an extra H+ after the reaction – so it could act as an acid in the reverse process • and the original acid has a lone pair of electrons after the reaction – so it could act as a base in the reverse process : A– + H–B+ → H–A + : B • a double arrow, , is usually used to indicate a process that is reversible 16
Conjugate Pairs • In a Brønsted-Lowery Acid-Base reaction, the original base becomes an acid in the reverse reaction, and the original acid becomes a base in the reverse process • each reactant and the product it becomes is called a conjugate pair • the original base becomes the conjugate acid; and the original acid becomes the conjugate base 17
Brønsted-Lowery Acid-Base Reactions : B base HCHO 2 + H 2 O acid base H 2 O + acid H–A acid + NH 3 base : A– conjugate base CHO 2– conjugate base + H–B+ conjugate acid + H 3 O+ conjugate acid HO– conjugate base + NH 4+ conjugate acid 18
Conjugate Pairs In the reaction H 2 O + NH 3 HO– + NH 4+ H 2 O and HO– constitute an Acid/Conjugate Base pair NH 3 and NH 4+ constitute a Base/Conjugate Acid pair 19
Practice – Identify the Brønsted-Lowery Acids and Bases and their Conjugates in each Reaction H 2 SO 4 + H 2 O HSO 4– + H 3 O+ HCO 3– + H 2 O H 2 CO 3 + HO– 20
Neutralization Reactions • H+ + OH- H 2 O • acid + base salt + water • double displacement reactions ü salt = cation from base + anion from acid ü cation and anion charges stay constant H 2 SO 4 + Ca(OH)2 → Ca. SO 4 + 2 H 2 O • some neutralization reactions are gas evolving where H 2 CO 3 decomposes into CO 2 and H 2 O H 2 SO 4 + 2 Na. HCO 3 → Na 2 SO 4 + 2 H 2 O + 2 CO 2 21
Nonmetal Oxides are Acidic • nonmetal oxides react with water to form acids • causes acid rain CO 2 (g) + H 2 O(l) → H 2 CO 3(aq) 2 SO 2(g) + 2 H 2 O(l) → 2 H 2 SO 4(aq) 4 NO 2(g) + 2 H 2 O(l) → 4 HNO 3(aq) 22
Acid Reactions Acids React with Metals • acids react with many metals übut not all!! • when acids react with metals, they produce a salt and hydrogen gas 3 H 2 SO 4(aq) + 2 Al(s) → Al 2(SO 4)3(aq) + 3 H 2(g) 23
Acid Reactions Acids React with Metal Oxides • when acids react with metal oxides, they produce a salt and water 3 H 2 SO 4 + Al 2 O 3 → Al 2(SO 4)3 + 3 H 2 O 24
Base Reactions • the reaction all bases have is common is neutralization of acids • strong bases will react with Al metal to form sodium aluminate and hydrogen gas 2 Na. OH + 2 Al + 6 H 2 O → 2 Na. Al(OH)4 + 3 H 2 25
Titration • using reaction stoichiometry to determine the concentration of an unknown solution • Titrant (unknown solution) added from a buret • indicators are chemicals added to help determine when a reaction is complete • the endpoint of the titration occurs when the reaction is complete 26
Titration 27
Titration The base solution is the titrant in the buret. As the base is added to the acid, the H+ reacts with the OH– to form water. But there is still excess acid present so the color does not change. At the titration’s endpoint, just enough base has been added to neutralize all the acid. At this point the indicator changes color. 28
Example 14. 4 Acid-Base Titration The titration of 10. 00 m. L of HCl solution of unknown concentration requires 12. 54 m. L of 0. 100 M Na. OH solution to reach the end point. What is the concentration of the unknown HCl solution?
Strong or Weak • a strong acid is a strong electrolyte ü practically all the acid molecules ionize, → • a strong base is a strong electrolyte ü practically all the base molecules form OH– ions, either through dissociation or reaction with water, → • a weak acid is a weak electrolyte ü only a small percentage of the molecules ionize, • a weak base is a weak electrolyte ü only a small percentage of the base molecules form OH– ions, either through dissociation or reaction with water, 30
Strong Acids • The stronger the acid, the more willing it is to donate H ü use water as the standard base HCl ® H+ + Cl. HCl + H 2 O® H 3 O+ + Cl- • strong acids donate practically all their H’s ü 100% ionized in water ü strong electrolyte • [H 3 O+] = [strong acid] ü [ ] = molarity 31
Strong Acids Pure Water HCl solution 32
Weak Acids • weak acids donate a small fraction of their H’s ümost of the weak acid molecules do not donate H to water ümuch less than 1% ionized in water HF Û H+ + FHF + H 2 O Û H 3 O+ + F- • [H 3 O+] << [weak acid] 33
Weak Acids Pure Water HF solution 34
Strong Bases • The stronger the base, the more willing it is to accept H Na. OH ® Na+ + OH- ü use water as the standard acid • strong bases, practically all molecules are dissociated into OH– or accept H’s ü strong electrolyte ü multi-OH bases completely dissociated • [HO–] = [strong base] x (# OH) 35
Weak Bases • in weak bases, only a small fraction of molecules accept H’s NH 3 + H 2 O Û NH 4+ + OH- ü weak electrolyte ü most of the weak base molecules do not take H from water ü much less than 1% ionization in water • [HO–] << [strong base] 36
Relationship between Strengths of Acids and their Conjugate Bases • the stronger an acid is, the weaker the attraction of the ionizable H for the rest of the molecule is • the better the acid is at donating H, the worse its conjugate base will be at accepting a H strong acid HCl + H 2 O → Cl– + H 3 O+ weak conj. base weak acid HF + H 2 O F– + H 3 O+ strong conj. base 37
Autoionization of Water • Water is actually an extremely weak electrolyte ütherefore there must be a few ions present • about 1 out of every 10 million 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+ and OH– üthe concentration of H+ and OH– are equal in water ü[H+] = [OH–] = 10 -7 M @ 25°C 38
Ion Product of Water • the product of the H+ and OH– concentrations is always the same number • the number is called the ion product of water and has the symbol Kw • [H+] x [OH–] = 1 x 10 -14 = Kw • as [H+] increases the [OH–] must decrease so the product stays constant üinversely proportional 39
Acidic and Basic Solutions • neutral solutions have equal [H+] and [OH–] ü[H+] = [OH–] = 1 x 10 -7 • acidic solutions have a larger [H+] than [OH–] ü[H+] > 1 x 10 -7; [OH–] < 1 x 10 -7 • basic solutions have a larger [OH–] than [H+] ü[H+] < 1 x 10 -7; [OH–] > 1 x 10 -7 40
Example - Determine the [H+1] for a 0. 00020 M Ba(OH)2 and determine whether the solution is acidic, basic or neutral Ba(OH)2 = Ba 2+ + 2 OH– therefore [OH–] = 2 x 0. 00020 = 0. 00040 = 4. 0 x 10 -4 M [H+] = 2. 5 x 10 -11 M 41
Practice - Determine the [H+1] concentration and whether the solution is acidic, basic or neutral for the following • [OH–] = 0. 000250 M • [OH–] = 3. 50 x 10 -8 M • Ca(OH)2 = 0. 20 M 42
Complete the Table [H+] vs. [OH-] [H+] 100 10 -1 + H OH- 10 -3 10 -5 + H OH- 10 -7 + H 10 -9 10 -11 H+ 10 -13 10 -14 H+ OH OH OH [OH-] 43
Complete the Table + [H ] vs. [OH ] [H+] 100 10 -1 + H OH- Acid 10 -3 10 -5 + H OH- [OH-]10 -14 10 -13 10 -11 10 -9 10 -7 10 -9 Base 10 -11 H+ + H OH OH 10 -7 10 -5 10 -3 10 -14 H+ OH 10 -1 100 even though it may look like it, neither H+ of OH- will ever be 0 the sizes of the H+ and OH- are not to scale because the divisions are powers of 10 rather than units 44
p. H • the acidity/basicity of a solution is often expressed as p. H • p. H = -log[H+], [H+] = 10 -p. H üexponent on 10 with a positive sign üp. Hwater = -log[10 -7] = 7 üneed to know the [H+] concentration to find p. H • p. H < 7 is acidic; p. H > 7 is basic, p. H = 7 is neutral 45
p. H • the lower the p. H, the more acidic the solution; the higher the p. H, the more basic the solution ü 1 p. H unit corresponds to a factor of 10 difference in acidity • normal range 0 to 14 ü p. H 0 is [H+] = 1 M, p. H 14 is [OH–] = 1 M ü p. H can be negative (very acidic) or larger than 14 (very alkaline) 46
p. H of Common Substances Substance p. H 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 (sat’d Mg(OH)2) 10. 5 household ammonia 10. 5 to 11. 5 1. 0 M Na. OH 14 47
Example - Calculate the p. H of a 0. 0010 M Ba(OH)2 solution & determine if is acidic, basic or neutral Ba(OH)2 = Ba 2+ + 2 OH- therefore [OH-] = 2 x 0. 0010 = 0. 0020 = 2. 0 x 10 -3 M [H+] = 1 x 10 -14 2. 0 x 10 -3 = 5. 0 x 10 -12 M p. H = -log [H+] = -log (5. 0 x 10 -12) p. H = 11. 3 p. H > 7 therefore basic 48
Practice - Calculate the p. H of the following strong acid or base solutions • 0. 0020 M HCl • 0. 0050 M Ca(OH)2 • 0. 25 M HNO 3 49
Complete the Table p. H [H+] 100 10 -1 + H OH- 10 -3 10 -5 + H OH- [OH-]10 -14 10 -13 10 -11 10 -9 10 -7 10 -9 10 -11 H+ H+ + H OH OH OH 10 -7 10 -13 10 -14 10 -5 10 -3 10 -1 100 50
Complete the Table p. H 0 1 [H+] 100 10 -1 + H OH- Acid 3 10 -3 5 7 9 10 -5 10 -7 10 -9 + H OH- [OH-]10 -14 10 -13 10 -11 10 -9 + H 13 10 -11 14 10 -13 10 -14 H+ H+ OH OH OH 10 -7 Base 11 10 -5 10 -3 10 -1 100 51
Sample - Calculate the concentration of [H+] for a solution with p. H 3. 7 [H+] = 10 -p. H [H+] = 10 -3. 7 means 0. 0001 < [H+1] < 0. 001 [H+] = 2 x 10 -4 M = 0. 0002 M 52
Practice - Determine the [H+] for each of the following • p. H = 2. 7 • p. H = 12 • p. H = 0. 60 53
Buffers • buffers are solutions that resist changing p. H when small amounts of acid or base are added • they resist changing p. H by neutralizing added acid or base • buffers are made by mixing together a weak acid and its conjugate base üor weak base and it conjugate acid 54
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 p. H 55
A Buffer made from Acetic acid and Sodium Acetate • a buffer solution with a p. H of 4. 75 can be made by mixing equal volumes of 1 M HC 2 H 3 O 2 and 1 M Na. C 2 H 3 O 2 • adding 10 m. L of 0. 1 M HCl to 1 L of this solution will give a solution with a p. H of 4. 75 ü adding 10 m. L of 0. 1 M HCl to 1 L of distilled water will give a solution with p. H of 3. 0 • adding 10 m. L of 0. 1 M Na. OH to 1 L of this solution will give a solution with a p. H of 4. 75 ü adding 10 m. L of 0. 1 M Na. OH to 1 L of distilled water will give a solution with p. H of 11. 0 56
Acetic Acid/Acetate Buffer 57
What is Acid Rain? • natural rain water has a p. H of 5. 6 ünaturally slightly acidic due mainly to CO 2 • rain water with a p. H lower than 5. 6 is called acid rain • acid rain is linked to damage in ecosystems and structures 58
What Causes Acid Rain? • many natural and pollutant gases dissolved in the air are nonmetal oxides ü CO 2, SO 2, NO 2 • nonmetal oxides are acidic CO 2 + H 2 O H 2 CO 3 2 SO 2 + 2 H 2 O 2 H 2 SO 4 • processes that produce nonmetal oxide gases as waste increase the acidity of the rain ü natural – volcanoes and some bacterial action ü man-made – combustion of fuel • weather patterns may cause rain to be acidic in regions other than where the nonmetal oxide is produced 59
p. H of Rain in Different Regions 60
Sources of SO 2 from Utilities 61
Damage from Acid Rain • acids react with metals, and materials that contain carbonates • acid rain damages bridges, cars and other metallic structures • acid rain damages buildings and other structures made of limestone or cement 62
Damage from Acid Rain circa 1935 circa 1995 63
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