Acids & Bases - Brown

7/29/2019 Acids & Bases - Brown

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Chapter 10

Acids and Bases

John D. Bookstaver

St. Charles Community College

St. Peters, MO

2006, Prentice Hall, Inc.

Chemistry, The Central Science, 10th edition

Theodore L. Brown; H. Eugene LeMay, Jr.; and Bruce E. Bursten


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Some Definitions

Arrhenius

Acid: Substance that, when dissolved in water, increases the

concentration of hydrogen ions.

Base: Substance that, when dissolved in water, increases the

concentration of hydroxide ions.


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Polyprotic Acids


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Arrhenius Acids and Bases

In 1884, Svante Arrhenius proposed these definitions

acid: a substance that produces H3O+ ions aqueous solution

base: a substance that produces OH- ions in aqueous solution

this definition of an acid is a slight modification of the original

Arrhenius definition, which was that an acid produces H+ in

aqueous solution today we know that H+ reacts immediately with a water molecule

to give a hydronium ion

H+(aq) + H2O( l) H3 O

+(aq)

Hydroniu m ion


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when HCl, for example, dissolves in water, its reacts with water to

give hydronium ion and chloride ion

HCl(aq)+H2 O(l) H3 O+(aq) + Cl

-(aq)


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With bases, the situation is slightly different

many bases are metal hydroxides such as KOH, NaOH, Mg(OH)2,

and Ca(OH)2

these compounds are ionic solids and when they dissolve in water,

their ions merely separate

other bases are not hydroxides; these bases produce OH- byreacting with water molecules

NaOH(s)H2O Na

+(aq) + OH

-(aq)

NH3

(aq) + H2

O(l) NH4

+(aq) + OH

-(aq)


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Acid and Base Strength Strong acid: one that reacts completely or almost completely with

water to form H3O+ ions

Strong base: one that reacts completely or almost completely with

water to form OH- ions

here are the six most common strong acids and the four mostcommon strong bases

HCl

HBrHI

HNO3H2 SO4HClO

4

LiOH

NaOHKOH

Ba( OH)2

Hydrochloric acid

Hydrobromic acidHydroiodic acid

Nitric acid

Sulfuric acid

Perchloric acid

Lithium hydroxide

Sodium hydroxidePotassium hydroxide

Barium hydroxide

Formula Name Formula Name


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Acids produce H+ in aqueous solutions waterHCl H+(aq) + Cl- (aq)

Bases produce OH-

in aqueous solutionswater

NaOH Na+(aq) + OH- (aq)


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Acids

Produce H+ (as H3O+) ions in water

Produce a negative ion (-) too

Taste sour

Corrode metals

React with bases to form salts and water


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Bases

Produce OH- ions in water

Taste bitter, chalky

Are electrolytes

Feel soapy, slippery

React with acids to form salts and water


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Learning Check AB1

Describe the solution in each of the following as: 1) acid 2) base

or 3)neutral.

A. ___soda

B. ___soapC. ___coffee

D. ___ wine

E. ___ water

F. ___ grapefruit


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Solution AB1

Describe each solution as:

1) acid 2) base or 3) neutral.

A. _1_ soda

B. _2_ soapC. _1_ coffee

D. _1_ wine

E. _3_ water

F. _1_ grapefruit


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Learning Check AB2

Identify each as characteristic of an A) acid or B) base

____ 1. Sour taste

____ 2. Produces OH- in aqueous solutions

____ 3. Chalky taste____ 4. Is an electrolyte

____ 5. Produces H+ in aqueous solutions


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Solution AB2

Identify each as a characteristic of an A) acid or B) base

_A_ 1. Sour taste

_B_ 2. Produces OH- in aqueous solutions

_B_ 3. Chalky taste

A, B 4. Is an electrolyte

_A_ 5. Produces H+ in aqueous solutions


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Acid and Base Strength

Strong acids are completelydissociated in water.

Weak acids only dissociatepartially in water.


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Substances with negligible aciditydo not dissociate in water.


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Weak acid: a substance that dissociates only partially in water

to produce H3O+ ions

acetic acid, for example, is a weak acid; in water, only 4 out every

1000 molecules are converted to acetate ions

Weak base: a substance that dissociates only partially in

water to produce OH- ions

ammonia, for example, is a weak base

CH3 COOH(aq) + H2O( l) CH3 COO-(aq) + H3O

+( aq)

Acetic acid Acetate ion

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

-( aq)


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Some Definitions

BrnstedLowry

Acid: Proton donor

Base: Proton acceptor


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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 asions in aqueous solution.

For the monoprotic strong acids,

[H3O+] = [acid].


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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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Strengths of Acids and Bases

Strong acids completely ionize (100%) in aqueous solutions

HCl + H2O H3O+ + Cl- (100 % ions)

Strong bases completely (100%) dissociate into ions in aqueous

solutions.

NaOH Na+ (aq) + OH-(aq)

(100 % ions)


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Strong and WeakAcids and Bases

Strong acids

HCl, HNO3 ,H2SO4

Most other acids are weak.

Strong bases

NaOH, KOH, and Ca(OH)2

Most other bases are weak.


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What Happens When an AcidDissolves in Water?

Water acts as a Brnsted

Lowry base and abstracts a

proton (H+) from the acid.

As a result, the conjugate base

of the acid and a hydronium

ion are formed.


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Conjugate Acids and Bases:

From the Latin word conjugare, meaning to join together.

Reactions between acids and bases always yield their conjugate bases andacids.


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NH3, A Bronsted-Lowry Base

When NH3 reacts with water, most of the reactants remain dissolved as

molecules, but a few NH3 reacts with water to form NH4+ and hydroxide

ion.

NH3 + H2O NH4+

(aq) + OH-

(aq)acceptor donor

+ +


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Bronsted-Lowry Acids

Acids are hydrogen ion (H+) donors

Bases are hydrogen ion (H+) acceptors

HCl + H2O H3O+

+ Cl-

donor acceptor

+ -

+ +


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Brnsted-Lowry Acids & Bases

Acid: a proton donor

Base: a proton acceptor

Acid-base reaction: a proton transfer reaction

Conjugate acid-base pair: any pair of molecules or ions that can beinterconverted by transfer of a proton

HCl(aq) + H2 O(l) H3O+( aq)+Cl

-(aq)

WaterHydrogenchloride

Hydroniumion

Chlorideion

(base)(acid) (conjugate

acid of water)

(conjugate

base of HCl)

conjugate acid -base pair

conjugate acid -base pair


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Brnsted-Lowry definitions do not require water as a reactant

NH4+

CH3 COOH CH3 COO-

NH3

(base) (conjugate baseacetic acid)

(conjugate acidof ammonia)

conjugate acid-base pair

+ +Acetic acid Ammonia

(acid)

conjugate acid-base pair

Acetateion

Ammoniumion


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Note the following about the conjugate acid-base pairs in the

table1. an acid can be positively charged, neutral, or negatively

charged; examples of each type are H3O+, H2CO3, and

H2PO4-

2. a base can be negatively charged or neutral; examples areOH-, Cl-, and NH3

3. acids are classified a monoprotic, diprotic, or triprotic

depending on the number of protons each may give up;

examples are HCl, H2CO3, and H3PO4


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carbonic acid, for example can give up one proton to become

bicarbonate ion, and then the second proton to become carbonate

ion

4. several molecules and ions appear in both the acid and

conjugate base columns; that is, each can function as either

an acid or a base

Carbonicacid

Bicarbonateion

Bicarbonateion

Carbonateion

H2 CO3 H2 O

HCO3-

H2 O

HCO3-

CO32 -

H3 O+

H3 O+

+ +

+ +


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If it can be either

...it is amphiprotic.

HCO3

HSO4

H2O


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pH

pH is defined as the negative base-10 logarithm of the

hydronium ion concentration.

pH = log [H3O+]


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pH

Therefore, in pure water,pH = log (1.0 107) = 7.00

An acid has a higher [H3O+] than pure water, so its pH is 7.


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pH Range

0 1 2 3 4 5 6 7 8 9 10 11 12 13 14

Neutral

[H+]>[OH-] [H+] = [OH-] [OH-]>[H+]

Acidic Basic


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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 logKw


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pH of Some Common Acids

gastric juice 1.0

lemon juice 2.3

vinegar 2.8

orange juice 3.5

coffee 5.0

milk 6.6


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pH of Some Common Bases

blood 7.4

tears 7.4

seawater 8.4

milk of magnesia 10.6

household ammonia 11.0


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pH

These are the pH

values for severalcommon

substances.


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How Do We Measure pH?

For less accuratemeasurements, one can use

Litmus paper

Red paper turns

blue above ~pH = 8

Blue paper turns redbelow ~pH = 5

An indicator


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How Do We Measure pH?

For more accurate measurements,

one uses a pH meter, whichmeasures the voltage in the

solution.


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Learning Check SW1

Identify each of the following as a

1) strong acid or base 2) weak acid

3) weak base

A. ___ HCl (aq)

B. ___ NH3(aq)

C. ___ NaOH (aq)

D. ___ H2CO

3(aq)


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the HCO3- ion, for example, can give up a proton to become CO3

2-

, or it can accept a proton to become H2CO3 a substance that can act as either an acid or a base is said to be

amphiprotic

the most important amphiprotic substance in Table 8.2 is H2O; it

can accept a proton to become H3O+, or lose a proton to become

OH-

5. a substance cannot be a Brnsted-Lowry acid unless it

contains a hydrogen atom, but not all hydrogen atoms in most

compounds can be given up

acetic acid, for example, gives up only one proton


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6. there is an inverse relationship between the strength of an

acid and the strength of its conjugate base the stronger the acid, the weaker its conjugate base

HI, for example, is the strongest acid in Table 8.2, and its

conjugate base, I-, is the weakest base in the table

CH3COOH (acetic acid) is a stronger acid that H2CO3 (carbonicacid); conversely, CH3COO

- (acetate ion) is a weaker base that

HCO3- (bicarbonate ion)


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Properties of Acids & Bases

Neutralization

acids and bases react with each other in a process calledneutralization; these reactions are discussed in Section 8.10

Reaction with metals

strong acids react with certain metals (called active metals) toproduce a salt and hydrogen gas, H2

reaction of a strong acid with a metal is a redox reaction; the metalis oxidized to a metal ion and H+ is reduced to H2

Mg(s) + 2HCl(aq) MgCl2 (aq) + H2 (g)

Magnesium Hydrochloric

acid

Magnesium

chloride

Hydrogen


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Reaction with metal hydroxides reaction of an acid with a metal hydroxide gives a salt plus water

the reaction is more accurately written as

omitting spectator ions gives this net ionic equation

HCl(aq)

Hydrochloric

acid

+ KOH( aq)

Water

+KCl(aq)

Potassium

chloride

Potassium

hydroxide

H2 O

H3O+

Cl-

K+

OH-

2H2O Cl-

K++ + + + +

H3 O+

OH-

2H2 O+


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Reaction with metal oxides strong acids react with metal oxides to give water plus a salt

2H3 O+(aq) + CaO(s) 3H2O(l) + Ca

2 +( aq)

Calciumoxide


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Reaction with carbonates and bicarbonates

strong acids react with carbonates to give carbonic acid, which

rapidly decomposes to CO2 and H2O

strong acids react similarly with bicarbonates

2H3 O+(aq) + CO3

2 -(aq) H2 CO3 (aq) + 2H2 O(l)

H2 CO3 (aq) CO2 (g) + H2 O( l)

2H3 O+(aq) + CO3

2 -(aq) CO2 (g) + 3H2 O(l)

H3O+( aq) + HCO3

-( aq) H2 CO3 (aq) + H2 O( l)

H2CO3 (aq) CO2 (g) + H2O(l)

H3O+

( aq) + HCO3-

( aq) CO2 (g) + 2H2 O(l)


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Neutralization Reactions

When acid and bases with equal amounts of hydrogen ion H+ and

hydroxide ions OH- are mixed, the resulting solution is neutral.

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

base acid salt water

Ca(OH)2 + 2 HCl CaCl2 + 2H2O

base acid salt water


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Neutralization

H3O+ and OH- combine to produce water

H3O+ + OH- 2 H2O

from acid from base neutral

Net ionic equation:

H+ + OH- H2O


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Acid-Base Titrations

Titration: an analytical procedure in which a solute in a

solution of known concentration reacts with a knownstoichiometry with a substance whose concentration is to be

determined

in this chapter, we are concerned with titrations in which we use

an acid (or base) of known concentration to determine the

concentration of a base (or acid) in another solution


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Titration

A known concentration of base (or

acid) is slowly added to a solution

of acid (or base).


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TitrationA pH meter or indicators are used

to determine when the solution

has reached the equivalence point,

at which the stoichiometric

amount of acid equals that of

base.


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As an example, let us use 0.108 M H2SO4 to determine the

concentration of a NaOH solution

requirement 1: we know the balanced equation

requirement 2: the reaction between H3O+ and OH- is rapid and

complete

requirement 3: we can use either an acid-base indicator or a pH

meter to observe the sudden change in pH that occurs at the end

point of the titration

requirement 4: we use volumetric glassware

2NaOH(aq)+H2SO4 (aq) Na2SO4 (aq) + 2H2O( l)(concentration

known)(concentration

not known)


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experimental measurements

doing the calculations

Trial I

Volume of0.108 M H2SO4

Volumeof NaOH

25.0 mL 33.48 mL

Trial II 25.0 mL 33.46 mL

Trial III 25.0 mL 33.50 mL

average = 33.48 mL

2 mol NaOH

1 mol H2 SO4

=0.161 mol NaOH

L NaOH= 0.161 M

mol NaOHL NaOH

=0.108 mol H2 SO4

1 L H2 SO4x x

0.02 50 L H2SO4

0.03348 L NaOH


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Buffers:

Solutions of a weak conjugate

acid-base pair.

They are particularly resistant to

pH changes, even when strong

acid or base is added.


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Blood Buffers

The average pH of human blood is 7.4

any change larger than 0.10 pH unit in either direction can causeillness

To maintain this pH, the body uses three buffer systems

carbonate buffer: H2CO3 and its conjugate base, HCO3-

phosphate buffer: H2PO4-

and its conjugate base, HPO42-

proteins: discussed in Chapter 21

Acids & Bases - Brown

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