Chapter 14 Acids and Bases

Types of Electrolytes salts = water soluble ionic compounds all strong electrolytes acids = form H+1 ions in water solution bases = combine with H+1 ions in water solution increases the OH-1 concentration may either directly release OH-1 or pull H+1 off H2O

Properties of Acids Sour taste react with “active” metals i.e. Al, Zn, Fe, but not Cu, Ag or Au 2 Al + 6 HCl ® 2 AlCl3 + 3 H2 corrosive react with carbonates, producing CO2 marble, baking soda, chalk, limestone CaCO3 + 2 HCl ® CaCl2 + CO2 + H2O change color of vegetable dyes blue litmus turns red react with bases to form ionic salts

Common Acids

Structures of Acids binary acids have acid hydrogens attached to a nonmetal atom HCl, HF Hydrofluoric acid

Structure of Acids oxy acids have acid hydrogens attached to an oxygen atom H2SO4, HNO3

Structure of Acids carboxylic acids have COOH group HC2H3O2, H3C6H5O3 only the first H in the formula is acidic the H is on the COOH

Properties of Bases also known as alkalis taste bitter alkaloids = plant product that is alkaline often poisonous solutions feel slippery change color of vegetable dyes different color than acid red litmus turns blue react with acids to form ionic salts neutralization

Common Bases

Structure of Bases most ionic bases contain OH ions NaOH, Ca(OH)2 some contain CO32- ions CaCO3 NaHCO3 molecular bases contain structures that react with H+ mostly amine groups

Arrhenius Theory bases dissociate in water to produce OH- ions and cations ionic substances dissociate in water NaOH(aq) → Na+(aq) + OH–(aq) acids ionize in water to produce H+ ions and anions because molecular acids are not made of ions, they cannot dissociate they must be pulled apart, or ionized, by the water HCl(aq) → H+(aq) + Cl–(aq) in formula, ionizable H written in front HC2H3O2(aq) → H+(aq) + C2H3O2–(aq)

Arrhenius Acid-Base Reactions the H+ from the acid combines with the OH- from the base to make a molecule of H2O it is often helpful to think of H2O 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) + NaOH(aq) → NaCl(aq) + H2O(l)

Problems with Arrhenius Theory does not explain why molecular substances, like NH3, 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 H3O+ ions H+ = a proton! H3O+ = hydronium ions

Brønsted-Lowery Theory in a Brønsted-Lowery Acid-Base reaction, an H+ is transferred does not have to take place in aqueous solution broader definition than Arrhenius acid is H donor, base is H acceptor 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+

Amphoteric Substances amphoteric substances can act as either an acid or a base have both transferable H and atom with 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)

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

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

Brønsted-Lowery Acid-Base Reactions H–A + :B  :A– + H–B+ acid base conjugate conjugate base acid HCHO2 + H2O  CHO2– + H3O+ acid base conjugate conjugate base acid H2O + NH3  HO– + NH4+ acid base conjugate conjugate base acid

Conjugate Pairs In the reaction H2O + NH3  HO– + NH4+ H2O and HO– constitute an Acid/Conjugate Base pair NH3 and NH4+ constitute a Base/Conjugate Acid pair

Practice – Identify the Brønsted-Lowery Acids and Bases and their Conjugates in each Reaction H2SO4 + H2O  HSO4– + H3O+ HCO3– + H2O  H2CO3 + HO–

Neutralization Reactions H+ + OH- H2O acid + base salt + water double displacement reactions salt = cation from base + anion from acid cation and anion charges stay constant H2SO4 + Ca(OH)2 → CaSO4 + 2 H2O some neutralization reactions are gas evolving where H2CO3 decomposes into CO2 and H2O H2SO4 + 2 NaHCO3 → Na2SO4 + 2 H2O + 2 CO2

Nonmetal Oxides are Acidic nonmetal oxides react with water to form acids causes acid rain CO2 (g) + H2O(l) → H2CO3(aq) 2 SO2(g) + O2(g) + 2 H2O(l) → 2 H2SO4(aq) 4 NO2(g) + O2(g) + 2 H2O(l) → 4 HNO3(aq)

Acid Reactions Acids React with Metals acids react with many metals but not all!! when acids react with metals, they produce a salt and hydrogen gas 3 H2SO4(aq) + 2 Al(s) → Al2(SO4)3(aq) + 3 H2(g)

Acid Reactions Acids React with Metal Oxides when acids react with metal oxides, they produce a salt and water 3 H2SO4 + Al2O3 → Al2(SO4)3 + 3 H2O

2 NaOH + 2 Al + 6 H2O → 2 NaAl(OH)4 + 3 H2 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 NaOH + 2 Al + 6 H2O → 2 NaAl(OH)4 + 3 H2

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

Titration

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.

Example 14.4 Acid-Base Titration The titration of 10.00 mL of HCl solution of unknown concentration requires 12.54 mL of 0.100 M NaOH 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 practically all the acid molecules ionize, → a strong base is a strong electrolyte practically all the base molecules form OH– ions, either through dissociation or reaction with water, → 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, 

Strong Acids The stronger the acid, the more willing it is to donate H use water as the standard base strong acids donate practically all their H’s 100% ionized in water strong electrolyte [H3O+] = [strong acid] [ ] = molarity HCl ® H+ + Cl- HCl + H2O® H3O+ + Cl-

Strong Acids Pure Water HCl solution

Weak Acids weak acids donate a small fraction of their H’s most of the weak acid molecules do not donate H to water much less than 1% ionized in water [H3O+] << [weak acid] HF Û H+ + F- HF + H2O Û H3O+ + F-

Weak Acids Pure Water HF solution

Strong Bases The stronger the base, the more willing it is to accept H use water as the standard acid strong bases, practically all molecules are dissociated into OH– or accept H’s strong electrolyte multi-OH bases completely dissociated [HO–] = [strong base] x (# OH) NaOH ® Na+ + OH-

Weak Bases 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 [HO–] << [strong base] NH3 + H2O Û NH4+ + OH-

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 + H2O → Cl– + H3O+ weak conj. base weak acid HF + H2O  F– + H3O+ strong conj. base

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 Û H+ + OH– H2O + H2O Û H3O+ + OH– all aqueous solutions contain both H+ and OH– the concentration of H+ and OH– are equal in water [H+] = [OH–] = 10-7M @ 25°C

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 inversely proportional

Acidic and Basic Solutions neutral solutions have equal [H+] and [OH–] [H+] = [OH–] = 1 x 10-7 acidic solutions have a larger [H+] than [OH–] [H+] > 1 x 10-7; [OH–] < 1 x 10-7 basic solutions have a larger [OH–] than [H+] [H+] < 1 x 10-7; [OH–] > 1 x 10-7

Ba(OH)2 = Ba2+ + 2 OH– therefore 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 = Ba2+ + 2 OH– therefore [OH–] = 2 x 0.00020 = 0.00040 = 4.0 x 10-4 M [H+] = 2.5 x 10-11 M

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

pH the acidity/basicity of a solution is often expressed as pH pH = -log[H+], [H+] = 10-pH exponent on 10 with a positive sign pHwater = -log[10-7] = 7 need to know the [H+] concentration to find pH pH < 7 is acidic; pH > 7 is basic, pH = 7 is neutral

pH 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 difference in acidity normal range 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)

pH of Common Substances 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 NaOH 14

Example - Calculate the pH of a 0 Example - Calculate the pH of a 0.0010 M Ba(OH)2 solution & determine if is acidic, basic or neutral Ba(OH)2 = Ba2+ + 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 10-12M pH = -log [H+] = -log (5.0 x 10-12) pH = 11.3 pH > 7 therefore basic

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

Sample - Calculate the concentration of [H+] for a solution with pH 3 means 0.0001 < [H+1] < 0.001 [H+] = 2 x 10-4 M = 0.0002 M

Practice - Determine the [H+] for each of the following pH = 2.7 pH = 12 pH = 0.60

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 it conjugate acid

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

What is Acid Rain? natural rain water has a pH of 5.6 naturally slightly acidic due mainly to CO2 rain water with a pH lower than 5.6 is called acid rain acid rain is linked to damage in ecosystems and structures

What Causes Acid Rain? many natural and pollutant gases dissolved in the air are nonmetal oxides CO2, SO2, NO2 nonmetal oxides are acidic CO2 + H2O  H2CO3 2 SO2 + O2 + 2 H2O  2 H2SO4 processes that produce nonmetal oxide gases as waste increase the acidity of the rain natural – volcanoes and some bacterial action man-made – combustion of fuel weather patterns may cause rain to be acidic in regions other than where the nonmetal oxide is produced

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

Damage from Acid Rain circa 1935 circa 1995