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Acidsand

Bases

Chapter 16Acids and Bases

John D. Bookstaver

St. Charles Community College

St. Peters, MO

2006, Prentice Hall, Inc.

Chemistry, The Central Science, 10th editionTheodore L. Brown; H. Eugene LeMay, Jr.;

and Bruce E. Bursten

• sour taste

• react with “active” metals recall the activity series (any metals above H are “active”)

2 Al (s) + 6 HCl (aq) 2 AlCl3 (aq) + 3 H2 (g)

• corrosive has the ability to eat away/destroy other substances

• react with carbonates (containing CO32-)

produce CO2 gas

marble, baking soda, chalk, limestone

CaCO3 (A) + 2 HCl (aq) CaCl2 (aq) + CO2 (g) + H2O (l)

• change color of vegetable dyesblue litmus paper turns red

• react with bases to form ionic salts (ionic compounds)

Properties of Acids

Properties of Bases• also known as alkalis

• taste bitter

• solutions feel slipperyMany soaps are manufactured with the use of

bases

• change color of vegetable dyesdifferent color than acid

red litmus paper turns blue

• react with acids to form ionic salts (ionic compounds)This is known as “neutralization”

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Acidsand

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What are Acids and Bases?

• There are several different ways to define the terms acid and base

• Each definition comes from a different acid/base theory:

Arrhenius Theory

Brønsted-Lowry Theory

Lewis Theory

• Acid: when dissolved in water, the hydrogen ion concentration

increases (basically, an Arrhenius acid produces H+ ions)

Ex) HCl (aq) → H+ (aq) + Cl– (aq)

HCl is an acid b/c it is ionized (molecular acid that is pulled apart by water) to form H+

• Base: when dissolved in water, the hydroxide ion concentration

increases (basically, an Arrhenius base produces OH–

ions)

Ex) NaOH (aq) → Na+ (aq) + OH– (aq)

NaOH is a base b/c it dissociates (ionic compound breaking up into its ions) to form OH–

Arrhenius Theory

Problems with Arrhenius Theory…

• Does not explain why molecular substances, like NH3, and some ionic compounds, like Na2CO3 or Na2O, are bases since they do not contain OH– ions, they can’t

be considered Arrhenius bases

• Does not explain why molecular substances, like CO2, dissolve in water to form acidic solutions since it doesn’t even contain H+ ions it can’t be

considered an Arrhenius acid

• Does not explain acid-base reactions that take place outside of aqueous solutions

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Hydronium Ion: H3O+

• the H+ ions produced by the acid are so reactive they cannot exist in waterH+ ions are just protons w/ no valence electrons!!!

• To be more stable, H+ ions react with water molecule(s) to produce hydronium ions, H3O+

• Note: for our purposes, H+ and H3O+ will be used interchangeably

• The Brønsted-Lowry theory looks at the entire acid-base chemical reaction

• In an acid-base reaction, an H+ is transferred from the acid to the base

Broader, more general, definition than Arrhenius

does not have to take place in aqueous solution

Brønsted-Lowry Theory

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Brønsted-Lowry Acids & Bases• Brønsted-Lowry acids are H+ donors (aka:

proton donors)Must have a removable (acidic) hydrogenbecause of the molecular structure, often one H

in the molecule is easier to transfer than others (this is known as the “ionizable hydrogen”)

• Brønsted-Lowry bases are H+ acceptors (aka: proton acceptors)Must have a lone pair/ nonbonding pair of

electronsbecause of the molecular structure, often one

atom in the molecule is more willing to accept H+ transfer than others

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• HCl transfers an H+ to H2O, forming H3O+ ions water acts as base, accepting H+

• NH3 (aq) is a Brønsted-Lowry base because NH3accepts an H+ from H2O (it is also an Arrhenius base because it forms OH–)water acts as acid, donating H+

base acid

acid base

Brønsted-Lowry Acid/Base Rxns

Ex) NH3 (aq) + H2O (l) NH4+ (aq) + OH– (aq)

Ex) HCl (aq) + H2O (l) → Cl– (aq) + H3O+ (aq)

Amphiprotic Molecules• Amphiprotic (or amphoteric) means have qualities

of both… substance can act as both an acid and a base

depending what it is reacting with

After Donated

H+

Amphoteric Molecule

After Accepted

H+

OH– H2O H3O+

CO32− HCO3

− H2CO3

SO42− HSO4

− H2SO4

• Brønsted-Lowry theory allows A-B rxns to be reversible (shown w/ double-sided arrows)

General A-B rxn: H–A + :B :A– + H–B+

• the original base has an extra H+ after the reaction, so it will act as an acid in the reverse process

• and the original acid has a lone pair of electrons after the reaction, so it will act as a base in the reverse process

:A– + H–B+ H–A + :B

Brønsted-Lowry Reversible Acid-Base Reactions

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Conjugate Pairs

• In a Brønsted-Lowry 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 are called a conjugate pair

• the original base becomes the conjugate acid; and the original acid becomes the conjugate base

Acidsand

Bases

Conjugate Pairs

Brønsted-Lowry Acid-Base Reactions

H–A + :B :A– + H–B+

acid base conjugate conjugatebase acid

HCHO2 + H2O CHO2– + H3O+

acid base conjugate conjugatebase acid

H2O + NH3 OH– + NH4+

acid base conjugate conjugatebase acid

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Acidsand

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Conjugate Pairs Practice

• Determine the acid, base, conjugate acid, and conjugate base in the following reaction:

H2O + H2SO4 HSO4– + H3O+

base acid conjugate conjugatebase acid

Arrow Conventions

• chemists commonly use two kinds of arrows in reactions to indicate the degree of completion of the reactions

• a single arrow indicates all the reactant molecules are converted to product molecules at the end

• a double arrow indicates the reaction is in equillibrium. The reaction stops when someof the reactant molecules have been converted into products. Some reactant molecules remain and some products molecules have formedLooks like in these notes

• a strong acid is a strong electrolytepractically 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 % of the molecules ionize ()

• a weak base is a weak electrolyteonly a small % of the base molecules form

OH– ions ()

Strong or Weak

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Autoionization of Water• As we have seen, water is amphoteric.In pure water, a few molecules act as bases and a

few act as acids.about 1 out of every 10 million water molecules

form ions through a process called autoionization

• all aqueous solutions contain both H3O+ and OH–

the concentration of H3O+ and OH– are equal in pure water (which is neutral)

H2O(l) + H2O(l) H3O+(aq) + OH−(aq)

• the product of the H3O+ and OH–

concentrations is a constant number• the number is called the ion product constant

of water (Kw)At 25ºC, [H3O+] [OH−] = [H+] [OH−] = Kw = 1 x 10-14

if you measure one of the concentrations (H+ or OH–) , you can calculate the other

• as [H3O+] increases the [OH–] must decrease so the product stays constantinversely proportional

Ion Product of Water

• all aqueous solutions contain both H+

and OH– ions• neutral solutions [H+] = [OH–] = 1 x 10-7

• acidic solutions [H+] is greater than [OH–][H+] > 1 x 10-7; [OH–] < 1 x 10-7

• basic solutions [OH–] is greater than [3O+][OH–] > 1 x 10-7; [H+] < 1 x 10-7

Acidic and Basic Solutions

Be careful! We are

dealing with NEGATIVE exponents. the larger

the negative exponent,

the SMALLER the number actually is!

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Acidsand

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Concentration Calculation Practice

• Calculate the [OH] at 25ºC when the [H+] = 1.5 x 10-9 M, and determine if the solution is acidic, basic, or neutral

1.0 x 10-14 M = [OH] [H+]

1.0 x 10-14 M = [OH] [1.5 x 10-9 M]

[OH] = 6.67 x10-6 M

The solution is basic… [OH] > [H+]

• pH is a measure of how acidic or basic a solution is

pH = –log [H+]; [H+] = 10–pH

Ex) pHwater = –log [1.0 x 10-7 M] = 7Similarly, [H+]water = 10–7 = 1.0 x 10-7 M

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

pH

pH

• An acid has a higher [H+] than pure water, so its pH is < 7

• A base has a lower [H3O+] than pure water, so its pH is >7

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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 14pH 0 is [H+] = 1 M, pH 14 is [OH–] = 1 MpH can be negative (very acidic) or larger than 14

(very alkaline)

Acidsand

Bases

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 (sat’d Mg(OH)2) 10.5

household ammonia 10.5 to 11.5

1.0 M NaOH 14

pH of Common Substances

Acidsand

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Other “p” Scales

• The “p” in pH tells us to take the negative log of the quantity (in this case, hydrogen ions).

• Some similar examples are

pOH = −log [OH−]

pKw = −log Kw

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• another way of expressing the acidity/basicity of a solution is pOH

pOH = -log [OH]; [OH] = 10-pOH

• pOH measures are the exact opposite of pH:• pOH < 7 is basic, pOH > 7 is acidic• pOH = 7 is neutral

pOH

Acidsand

Bases

Because

[H3O+] [OH−] = Kw = 1.0 10−14,

we know that

−log [H3O+] + −log [OH−] = −log Kw = 14.00

or, in other words,

pH + pOH = pKw = 14.00

Watch This!

TheAcid/ Base

Square

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Relationships: pH, pOH, [H+], and [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

pH: 0 1 3 5 7 9 11 13 14

pOH: 14 13 11 9 7 5 3 1 0

[H+][OH–] = 1 x 10-14 = kw

pH = -log[H+]

[H+] = 10-pH

pOH = -log[OH-]

[OH–] = 10-pOH

pH + pOH = 14

EQUATIONS TO REMEMBER!

Acidsand

Bases

How Do We Measure pH?

• For less accurate measurements, one can useLitmus paper

• “Red” paper turns blue above ~pH = 8

• “Blue” paper turns red below ~pH = 5

An indicator

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How Do We Measure pH?

For more accurate measurements, one uses a pH meter, which measures the voltage in the solution.

Acidsand

Bases

Strong Acids

• You will recall that the seven strong acids are HCl, HBr, HI, HNO3, H2SO4, HClO3, and HClO4.

• These are, by definition, strong electrolytes and exist totally as ions in aqueous solution.

• For the monoprotic strong acids,[H3O+] = [acid].

Acidsand

Bases

Strong Bases

• Strong bases are the soluble hydroxides, which are the alkali metal and heavier alkaline earth metal hydroxides (Ca2+, Sr2+, and Ba2+).

• Again, these substances dissociate completely in aqueous solution.

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Acidsand

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Dissociation Constants

• For a generalized acid dissociation,

the equilibrium expression would be

• This equilibrium constant is called the acid-dissociation constant, Ka.

[H3O+] [A−][HA]

Kc =

HA(aq) + H2O(l) A−(aq) + H3O+(aq)

Acidsand

Bases

Dissociation Constants

The greater the value of Ka, the stronger the acid.

Acidsand

Bases

Calculating Ka from the pH

• The pH of a 0.10 M solution of formic acid, HCOOH, at 25°C is 2.38. Calculate Ka for formic acid at this temperature.

• We know that

[H3O+] [COO−][HCOOH]

Ka =

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Acidsand

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Calculating Ka from the pH

• The pH of a 0.10 M solution of formic acid, HCOOH, at 25°C is 2.38. Calculate Ka for formic acid at this temperature.

• To calculate Ka, we need the equilibrium concentrations of all three things.

• We can find [H3O+], which is the same as [HCOO−], from the pH.

Acidsand

Bases

Calculating Ka from the pH

pH = −log [H3O+]

2.38 = −log [H3O+]

−2.38 = log [H3O+]

10−2.38 = 10log [H3O+] = [H3O+]

4.2 10−3 = [H3O+] = [HCOO−]

Acidsand

Bases

Calculating Ka from pH

Now we can set up a table…

[HCOOH], M [H3O+], M [HCOO−], M

Initially 0.10 0 0

Change −4.2 10-3 +4.2 10-3 +4.2 10−3

At Equilibrium

0.10 − 4.2 10−3

= 0.0958 = 0.10

4.2 10−3 4.2 10−3

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Acidsand

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Calculating Ka from pH

[4.2 10−3] [4.2 10−3][0.10]

Ka =

= 1.8 10−4