Chapter 14Acids and

Bases

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Types of Electrolytes

• salts = water soluble ionic compoundsall strong electrolytes

• acids = form H+1 ions in water solution

• bases = combine with H+1 ions in water solutionincreases the OH-1 concentration

may either directly release OH-1 or pull H+1 off H2O

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Properties of Acids• Sour taste• react with “active” metals

i.e. Al, Zn, Fe, but not Cu, Ag or Au

2 Al + 6 HCl AlCl3 + 3 H2

corrosive

• react with carbonates, producing CO2

marble, baking soda, chalk, limestone

CaCO3 + 2 HCl CaCl2 + CO2 + H2O

• change color of vegetable dyesblue litmus turns red

• react with bases to form ionic salts

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Common AcidsChemical Name Formula Uses Strength

Nitric Acid HNO3 explosive, fertilizer, dye, glue Strong

Sulfuric Acid H2SO4 explosive, fertilizer, dye, glue,

batteries Strong

Hydrochloric Acid HCl metal cleaning, food prep, ore

refining, stomach acid Strong

Phosphoric Acid H3PO4 fertilizer, plastics & rubber,

food preservation Moderate

Acetic Acid HC2H3O2 plastics & rubber, food preservation, Vinegar

Weak

Hydrofluoric Acid HF metal cleaning, glass etching Weak

Carbonic Acid H2CO3 soda water Weak

Boric Acid H3BO3 eye wash Weak

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Structures of Acids• binary acids have acid

hydrogens attached to a nonmetal atomHCl, HF

Hydrofluoric acid

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Structure of Acids• oxy acids have acid

hydrogens attached to an oxygen atomH2SO4, HNO3

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Structure of Acids• carboxylic acids have

COOH groupHC2H3O2, H3C6H5O3

• only the first H in the formula is acidic the H is on the COOH

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Properties of Bases• also known as alkalis• taste bitter

alkaloids = plant product that is alkalineoften poisonous

• solutions feel slippery• change color of vegetable dyes

different color than acid red litmus turns blue

• react with acids to form ionic saltsneutralization

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Common BasesChemical

Name Formula

Common Name

Uses Strength

sodium hydroxide

NaOH lye,

caustic soda soap, plastic,

petrol refining Strong

potassium hydroxide

KOH caustic potash soap, cotton, electroplating

Strong

calcium hydroxide

Ca(OH)2 slaked lime cement Strong

sodium bicarbonate

NaHCO3 baking soda cooking, antacid Weak

magnesium hydroxide

Mg(OH)2 milk of

magnesia antacid Weak

ammonium hydroxide

NH4OH, {NH3(aq)}

ammonia water

detergent, fertilizer,

explosives, fibers Weak

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Structure of Bases

• most ionic bases contain OH ionsNaOH, Ca(OH)2

• some contain CO32- ions

CaCO3 NaHCO3

• molecular bases contain structures that react with H+

mostly amine groups

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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

anionsbecause 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)

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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)

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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+ ionsH+ = a proton!H3O+ = hydronium ions

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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+

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Amphoteric Substances• amphoteric substances can act as either an

acid or a basehave both transferable H and atom with lone pair

• HCl(aq) is acidic because HCl transfers an H+ to H2O, forming H3O+ ionswater 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)

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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

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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

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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

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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

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Practice – Identify the Brønsted-Lowery Acids and Bases and their Conjugates in

each Reaction

H2SO4 + H2O HSO4– + H3O+

HCO3– + H2O H2CO3 + HO–

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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

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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)

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Acid ReactionsAcids React with Metals

• acids react with many metalsbut 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)

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Acid ReactionsAcids React with Metal Oxides

• when acids react with metal oxides, they produce a salt and water

3 H2SO4 + Al2O3 → Al2(SO4)3 + 3 H2O

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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

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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

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Titration

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TitrationThe base solution is thetitrant in the buret.

As the base is added tothe acid, the H+ reacts withthe OH– to form water. But there is still excess acid present so the colordoes not change.

At the titration’s endpoint,just enough base has been added to neutralize all theacid. At this point the indicator changes color.

Example 14.4

Acid-Base TitrationThe 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?

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Strong or Weak• a strong acid is a strong electrolyte

practically all the acid molecules ionize, →

• a strong base is a strong electrolytepractically all the base molecules form OH– ions, either

through dissociation or reaction with water, →

• a weak acid is a weak electrolyteonly a small percentage of the molecules ionize,

• a weak base is a weak electrolyteonly a small percentage of the base molecules form OH–

ions, either through dissociation or reaction with water,

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Strong Acids• The stronger the acid, the

more willing it is to donate Huse water as the standard base

• strong acids donate practically all their H’s100% ionized in waterstrong electrolyte

• [H3O+] = [strong acid] [ ] = molarity

HCl H+ + Cl-

HCl + H2O H3O+ + Cl-

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Strong Acids

Pure Water HCl solution

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Weak Acids• weak acids donate a small

fraction of their H’smost of the weak acid

molecules do not donate H to water

much less than 1% ionized in water

• [H3O+] << [weak acid]

HF H+ + F-

HF + H2O H3O+ + F-

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Weak Acids

Pure Water HF solution

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Strong Bases• The stronger the base, the more

willing it is to accept Huse water as the standard acid

• strong bases, practically all molecules are dissociated into OH– or accept H’sstrong electrolytemulti-OH bases completely

dissociated

• [HO–] = [strong base] x (# OH)

NaOH Na+ + OH-

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Weak Bases• in weak bases, only a small

fraction of molecules accept H’sweak electrolytemost of the weak base molecules

do not take H from watermuch less than 1% ionization in

water

• [HO–] << [strong base]

NH3 + H2O NH4+ + OH-

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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

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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 H+ + OH–

H2O + H2O 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

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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 constantinversely proportional

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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

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Example - Determine the [H+1] for a 0.00020 M Ba(OH)2 and determine whether the solution is

acidic, basic or neutralBa(OH)2 = Ba2+ + 2 OH– therefore [OH–] = 2 x 0.00020 = 0.00040 = 4.0 x 10-4 M

4

14

1004101

OHH

OHH

.w

w

K

K

[H+] = 2.5 x 10-11 M

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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

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pH

• the acidity/basicity of a solution is often expressed as pH

• pH = -log[H+], [H+] = 10-pH

exponent on 10 with a positive signpHwater = -log[10-7] = 7need to know the [H+] concentration to find pH

• pH < 7 is acidic; pH > 7 is basic, pH = 7 is neutral

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pH• the lower the pH, the more acidic the solution; the

higher the pH, the more basic the solution1 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 MpH can be negative (very acidic) or larger than 14

(very alkaline)

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pH of Common SubstancesSubstance 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 (sat’d Mg(OH)2) 10.5

household ammonia 10.5 to 11.5

1.0 M NaOH 14

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Example - Calculate the pH of a 0.0010 M Ba(OH)2 solution & determine if is acidic,

basic or neutral

[H+] = 1 x 10-14

2.0 x 10-3 = 5.0 x 10-12M

pH > 7 therefore basic

Ba(OH)2 = Ba2+ + 2 OH- therefore [OH-] = 2 x 0.0010 = 0.0020 = 2.0 x 10-3 M

pH = -log [H+] = -log (5.0 x 10-12)pH = 11.3

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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

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Sample - Calculate the concentration of [H+] for a solution with pH 3.7

[H+] = 10-pH

[H+] = 10-3.7

means 0.0001 < [H+1] < 0.001

[H+] = 2 x 10-4 M = 0.0002 M

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Practice - Determine the [H+] for each of the following

• pH = 2.7

• pH = 12

• pH = 0.60

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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 baseor weak base and it conjugate acid

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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

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What is Acid Rain?

• natural rain water has a pH of 5.6naturally 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

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What Causes Acid Rain?• many natural and pollutant gases dissolved in the air are

nonmetal oxidesCO2, 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 rainnatural – volcanoes and some bacterial actionman-made – combustion of fuel

• weather patterns may cause rain to be acidic in regions other than where the nonmetal oxide is produced

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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

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Damage from Acid Rain

circa 1935 circa 1995