Edward Wen, PhD

Acids & Bases

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Learning Outcomes• Properties of acids and bases and definitions• pH scale and calculation of pH• Completing and balancing Neutralization

reactions• Titration calculations for neutralization

reactions• Defining Weak vs. Strong electrolytes (using

the concept of equilibrium)• Buffers – recognition of a buffer system, how a

buffer works

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Types of Ionic Compounds• Acids = form H+ ions in water solution

• Bases = combine with H+ ions in water solution increases the OH- concentration

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

• Salts = Ionic compounds formed from Acid and Base. all strong electrolytes Cation: except H+ Anion: except OH-

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

i.e. Al, Zn, Fe, but NOT w/ Ag, Au

Zn + 2 HCl ZnCl2 + H2

• react with Carbonates, producing CO2

marble, baking soda, 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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Most food contains acids

• Citric acid (HO2CCH(CO2H)COHCO2H):

citrus fruits, tomato

• Malic acid (HO2CCH2CHOHCO2H):

green apple, tomato, grape

• Ascorbic acid (aka Vitamin C)

• Folic acid

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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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Binary acids

• (HmX): acid hydrogens attached to a nonmetal atomHCl, HF, HBr, HIH2S, H2Se

Hydrofluoric acid

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Oxyacids

• acid hydrogens (H+) attached to an oxygen atomH2SO4, HNO3, H3PO4

HClO4

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Carboxylic acids • Many exist in food like vinegar,

tomato, citrus fruit• -COOH group

HC2H3O2, 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• 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

Common 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

Aluminum hydroxide

Al(OH)3 Antacid Weak

magnesium hydroxide

Mg(OH)2 milk of

magnesia antacid Weak

ammonium hydroxide

NH4OH, {NH3(aq)}

ammonia water

detergent, Windex fertilizer,

explosives, fibers Weak

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Structure of Bases

• most ionic bases contain OH- ionsDrano clog-remover: NaOH, Ca(OH)2

• some contain CO32- ion: it produces OH- with water

Baking soda: CaCO3

Alka-Seltzer: NaHCO3

• molecular bases that react with H+

Windex: Ammonia (NH3)

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Acid-Base Reactions (Neutralization, Double Displacement Reaction)

• H+ (from the acid) + OH- (from the base) H2O

it is often helpful to think of H2O as H-OH

• Cation (from base) + Anion (from acid) Salt

acid + base → salt + water

HCl(aq) + NaOH(aq) → NaCl(aq) + H2O(l)

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Acid Reactions. I.Reaction with Metals

• Reaction with many metals: Al, Zn, Fe, Mgbut not all!! Not for Cu, Au, Ag, etc.

• Producing a Salt and hydrogen gas H2

3 H2SO4(aq) + 2 Al(s) → Al2(SO4)3(aq) + 3 H2(g)

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Acid Reactions. IIReaction 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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Acid Reactions. IIIGas-evolving Reaction with Salts• when acids react with metal carbonate,

bicarbonate, sulfide, sulfite, and bisulfite, gas will be produced along with other products

2 HNO3 + FeCO3 → Fe(NO3)2 + CO2 + H2O

HCl + NaHCO3 → NaCl + CO2 + H2O

ZnS + 2 HBr → ZnBr2 + H2S

CaSO3 + 2 HI → CaI2 + SO2 + H2O

H2SO4 + 2 NH4HSO3 → (NH4)2SO4 + 2SO2 + 2H2O

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Base Reactions• Neutralization of acids

• Reaction with Nonmetal oxides, CO2

2 NaOH + CO2 → Na2CO3 + H2O

• Strong bases will react with Al metal to form sodium aluminate and hydrogen gas

Example: Dissolving recycled aluminum can with NaOH solution

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

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Relative Strength of Acids Acid Conjugate Base

Strong Acid Hydroiodic acid HI I-

Hydrobromic acid HBr Br-

Hydrochloric Acid HCl Cl-

Sulfuric Acid H2SO4 HSO4-

Nitric Acid HNO3 NO3-

Hydronium ion H3O+ H2O

Phosphoric Acid H3PO4 H2PO4-

Hydrofluoric Acid HF F-

Acetic Acid HC2H3O2 C2H3O2-

Carbonic Acid H2CO3 HCO3-

Ammonium ion NH4+ NH3

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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+

100% ionized in water

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

Stomach acid

HCl H+ + Cl-

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Strong Acids

Examples:• Binary Acid: HCl, HBr, HI• Oxyacid: HNO3, H2SO4, HClO4, HClO3

• Example:

HNO3 = H+ + NO3-

H2SO4 = 2H+ + SO42-

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Weak Acids

• Weak acids donate a small fraction of their H+

most of the weak acid molecules do not donate H+ to water

• [H3O+] << [weak acid]

Vinegar

HC2H3O2 H+ + C2H3O2-

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Weak Acids

Examples:• Binary Acid: HF, H2S, H2Se• Oxyacid: HNO2, H2SO3, H3PO4, HClO• Most carboxylic acids, such as acetic

acid

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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+

1 mol NaOH = 1 mol OH-

1 mol Ca(OH)2 = 2 mol OH-

• [OH–] = [strong base] x (# OH)

DranoTM

NaOH Na+ + OH-

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Weak Bases

• Definition: a small fraction of molecules accept H+

most of the weak base molecules do not take H+ from water

• [HO–] << [weak base]

WindexTM

NH3 + H2O NH4+ + OH-

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Autoionization of Water• Water: extremely Weak electrolyte

therefore there must be a few ions present

• about 2 out of every 1 billion water molecules form Ions: Autoionization

H2O + H2O H3O+ + OH–

H2O H+ + OH–

• ALL aqueous solutions contain both H+ and OH–

the concentration of H+ and OH– are equal in water@ 25°C: [H+] = [OH–] = 10-7M

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Ion Product of Water

• [H+] x [OH–] = constant: Ion Product of water, Kw

• At 25°C, [H+] x [OH–] = 1 x 10-14 = Kw

• as [H+] increases, [OH–] must decrease so the product stays constant

][OH

101][H

14

][H

101][OH

14

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Acidic and Basic Solutions

• Neutral solutions have equal [H+] and [OH–][H+] = [OH–] = 1 x 10-7 M

• Acidic solutions : [H+] > [OH–][H+] > 1 x 10-7 M [OH–] < 1 x 10-7 M

• Basic solutions: [OH–] > [H+][H+] < 1 x 10-7 M [OH–] > 1 x 10-7 M

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Practice - Determine the [H+] concentration and whether the solution is acidic, basic or

neutral for the following

All [H+] compared to 1 x 10-7 M

•[OH–] = 3.50 x 10-8 M

•[NaOH] = 0.000250 M

•[HCl] = 0.50 M

Practice - Determine the [H+] concentration and whether the solution is acidic, basic or

neutral for the following

• [OH–] = 3.50 x 10-8 M

• NaOH = 0.000250 M

• [HCl] = 0.50 M

[H+] = 1 x 10-14 3.50 x 10-8

= 2.86 x 10-7 M [H+] >[OH-], therefore acidic

[H+] = 1 x 10-14 0.000250

= 4.00 x 10-11 M [H+] < [OH-], therefore basic

[H+] > 1.0 x 10-7 M therefore acidic

[H+] = 0.50 M

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Acidic/Basic: [H+] vs. [OH-]

OH-H+ H+ H+ H+ H+

OH-OH-OH-OH-

[OH-]10-14 10-13 10-11 10-9 10-7 10-5 10-3 10-1 100

[H+] 100 10-1 10-3 10-5 10-7 10-9 10-11 10-13 10-14

even though it may look like it, neither H+ of OH- will ever be 0the sizes of the H+ and OH- are not to scale

because the divisions are powers of 10 rather than units

Acid Base

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pH• The measure of the acidity/basicity of a solution

• pH = -log[H+], [H+] = 10-pH

exponent on 10 with a positive signpHwater = -log[10-7] = 7

need to know the [H+] concentration to find pH

• pH < 7 : Acidic; pH > 7 : Basic• pH = 7 : Neutral

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pH scale• pH↓, Acidity↑ • pH↑, basicity↑

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

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pH measurement

pH can be measured by pH meter:

• The change in [H+] affects the voltage of a standard cell

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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 the following strong acid or base solutions

• 0.0020 M HCl

• 0.010 M NaOH

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Example - Calculate the pH of the following strong acid or base solutions

• 0.0020 M HClHCl as strong acid, so [H+] = 0.0020 M

pH = - log (2.0 x 10-3) = 2.7

pH = - log (1.0 x 10-12) = 12

[H+] = 1 x 10-14 1 x 10-2

= 1 x 10-12 M

NaOH as strong base, so [OH-] = 0.010 M

• 0.010 M NaOH

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pH in everyday life

OH-H+ H+ H+ H+ H+

OH-OH-OH-OH-

[OH-]10-14 10-13 10-11 10-9 10-7 10-5 10-3 10-1 100

[H+] 100 10-1 10-3 10-5 10-7 10-9 10-11 10-13 10-14

pH 0 1 3 5 7 9 11 13 14Acid Base

Stomach acid Vinegar Pure water Windex Drano

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

[H+] = 10-pH

[H+] = 10-3.7

= 2 x 10-4 M = 0.0002 M

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Find concentration of Acid or Base? Titration

• Purpose: using Reaction Stoichiometry to determine the Concentration of an unknown solution

• Titrant (solution of known concentration) added from a Buret

• Indicators: 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: Color change w/ Indicator

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TitrationStart: The base solution as titrant in the buret.

Titrating: As the Base is added to the Acid, H+ + OH– HOH. But still excess Acid present so the color does not change.

Endpoint: just enough Base to neutralize all the acid. The indicator changes color.

Calculations in Titration• At the Endpoint of the titration, acid base

neutralization reaction is complete. The mole ratio between acid and base in the reaction mixture is the same as in the balanced equation.

• Given the concentration of titrant, the mole of titrant can be calculated as: mole = Molarity x Volume (L)

• Then the mole of the other reactant can be calculated from the mole of titrant and the mole ratio in the equation (review stoichiometry: mole-to-mole).

• Finally, the molarity of other reactant can be determined.

Example: Acid-Base Titration

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Example:• The titration of 10.00 mL of HCl solution of unknown

concentration requires 12.51 mL of 0.100 M Ba(OH)2 solution to reach the endpoint. What is the concentration of the unknown HCl solution?

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• First, write balanced equation:

2 HCl(aq) + Ba(OH)2(aq) → BaCl2 (aq) + 2H2O(l)

2 mole HCl = 1 mole Ba(OH)2

0.100 M Ba(OH)2 0.100 mol Ba(OH)2 1 L sol’n

Information

Given: 10.00 mL HCl

12.54 mL 0.100 M Ba(OH)2

Find: M HCl

Example:The titration of 10.00 mL of HCl solution of unknown concentration requires 12.54 mL of 0.100 M Ba(OH)2 solution to reach the end point. What is the concentration of the unknown HCl solution?

solution liters

solute molesM olarity

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• Write a Solution Map:

mLBa(OH)2

LBa(OH)2

molBa(OH)2

2

2

Ba O0.100 mol

1 L

H

Ba OHmL 1

L 0010.

molHCl

2

2 mol HCl

1 mol Ba(OH)

Information

Given: 10.00 mL HCl

12.51 mL Ba(OH)2

Find: M HCl

CF: 2 mol HCl = 1 mol Ba(OH)2

0.100 mol Ba(OH)2 = 1 L

M = mol/L

Example:The titration of 10.00 mL of HCl solution of unknown concentration requires 12.51 mL of 0.100 M Ba(OH)2 solution to reach the end point. What is the concentration of the unknown HCl solution?

mLHCl

LHCl

mL 1

L 0010. HC l liters

HC l molesM olarity

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= 2.50 x 10-3 mol HCl

InformationGiven: 10.00 mL HCl

12.51 mL Ba(OH)2

Find: M HCl

CF: 2 mol HCl = 1 mol Ba(OH)2

0.100 mol Ba(OH)2 = 1 L M = mol/L

SM: mL Ba(OH)2 → L Ba(OH)2 →

mol Ba(OH)2 → mol HCl; mL HCl → L HCl & mol M

Example:The titration of 10.00 mL of HCl solution of unknown concentration requires 12.51 mL of 0.100 M Ba(OH)2 solution to reach the end point. What is the concentration of the unknown HCl solution?

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InformationGiven: 10.00 mL HCl

12.51 mL Ba(OH)2

Find: M HCl

CF: 2 mol HCl = 1 mol Ba(OH)2

0.100 mol Ba(OH)2 = 1 L M = mol/L

SM: mL Ba(OH)2 → L Ba(OH)2 →

mol Ba(OH)2 → mol HCl; mL HCl → L HCl & mol M

Example:The titration of 10.00 mL of HCl solution of unknown concentration requires 12.51 mL of 0.100 M Ba(OH)2 solution to reach the end point. What is the concentration of the unknown HCl solution?

-32.50 x 10 moles HClMolarity 0.250 M

0.01000 L HCl

H Cl L 01000.0mL 1

L 0 .001H Cl mL 0 .001

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How does pH change?

Initial pH pH after adding 1 mL 1 M HCl

pH after adding 1 mL 1 M NaOH

1 L Pure water 7.00

4.00 10.00

1 L 0.14 M K2HPO4 + 0.10 M KH2PO4

7.00

6.99 7.01

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Buffers• Definition: solutions that resist changing pH when

small amounts of acid or base are added• The mixture of 0.14 M K2HPO4 + 0.10 M KH2PO4

solution has much smaller pH change when strong acid or base is added, thus is called Buffer.

• Ingredient: mixing together a weak acid and its conjugate baseor weak base and it conjugate acid

Online demo: https://www.youtube.com/watch?v=P-R-Cqvb5yo

• Human body fluid as buffer: H2CO3/HCO3-

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BufferComposition:• a weak acid + its salt;

example: HC2H3O2 / NaC2H3O2, HF/KFWhen acid is added:

C2H3O2- + H+ HC2H3O2

When base is added:

OH- + HC2H3O2 C2H3O2- + H2O

• OR, a weak base + its salt

example: NH3 / NH4Cl

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Acetic Acid/Acetate Buffer

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Treasure Hunt: Which two can combine into a

Buffer?HCl

NH4+

C2H3O2-

Cl-

HCO3-

CO32-

HC2H3O2

NH3

H2CO3