ACID-BASE REACTIONSACID-BASE REACTIONS

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ACIDSACIDSAcids are substances that ionize in

aqueous solutions to form hydrogen ions, thereby increasing the concentration of H+ ions.

Because hydrogen atom consists of a proton and an electron, H+ is simply a proton.

Thus, acids are often called proton donors.

Molecules of different acids can ionize

to form different numbers of H+ ions.

Both hydrochloric acid and nitric acid

are monoprotic acids, which yield one

per molecule of acid.

Sulfuric acid is a diprotic acid, one that

yields two H+ per molecule of acid.

BASESBASES

Bases are substances that accept H+

ions.

Bases produce hydroxide ions when

they dissolve in water.

When dissolved in water, they

dissociate into their component ions,

introducing OH- ions into the solution.

COMMON PROPERTIES OF COMMON PROPERTIES OF ACIDS AND BASES:ACIDS AND BASES:

ACID BASE

Sour taste Bitter taste

Neutralizes bases Neutralizes acids

Turns litmus paper blue to red Turns litmus paper red to blue

Soapy and slippery feeling

Indicator Color in strongly

acidic solution

pH at which color changes

Color in strongly alkaline solution

Methyl orange

Red 4 Yellow

Litmus Red 7 Blue

Phenolphthalein

Colorless 9 Red

Screened Methyl Orange

Red 4 Green

STRONG AND WEAK ACIDS STRONG AND WEAK ACIDS AND BASESAND BASES

Acids and bases that are strong

electrolytes (completely ionized in

solution) are called strong acids and

strong bases.

Those that are weak electrolytes

(partly ionized) are called weak

acids and weak bases.

COMMON STRONG ACIDS COMMON STRONG ACIDS AND BASESAND BASES

STRONG ACIDS STRONG BASES

Hydrochloric acid Lithium hydroxide

Hydrobromic acid Sodium hydroxide

Hydroiodic acid Potassium hydroxide

Chloric acid Rubidium hydroxide

Perchloric acid Cesium hydroxide

Nitric acid Calcium hydroxide

Sulfuric acid Strontium hydroxideBarium hydroxide

IDENTIFYING STRONG AND IDENTIFYING STRONG AND WEAK ELECTROLYTESWEAK ELECTROLYTES

To classify a soluble substance as a

strong electrolyte, weak electrolyte, or

nonelectrolyte, we simply use the

following table:

SUMMARY OF THE ELECTROLYTIC SUMMARY OF THE ELECTROLYTIC BEHAVIOR OF A COMMON SOLUBLE BEHAVIOR OF A COMMON SOLUBLE

IONIC AND MOLECULAR COMPOUNDSIONIC AND MOLECULAR COMPOUNDS

Strong Electrolyte

Weak Electrolyte

Nonelectrolyte

Ionic All None None

Molecular Strong acids Weak acids

Weak bases All other compounds

If an acid is not listed, it is probably a weak electrolyte.

NH3 is only a weak base that we consider.

NEUTRALIZATION NEUTRALIZATION REACTIONS AND SALTSREACTIONS AND SALTS

When a solution of an acid and that

of a base are mixed, a

neutralization reaction occurs.

The products of the reaction have

none of the characteristics properties

of either the acidic and the basic

solutions.

HCl + NaOH HHCl + NaOH H22O + O + NaClNaCl

By analogy to this reaction, the term

salt has come to mean any ionic

compound whose cation comes from a

base and whose anion comes from an

acid.

A neutralization reaction between an

acid and a metal hydroxide produces

water and salt.

(acid) (base) (water) (acid) (base) (water) (salt)(salt)

ACID-BASE THEORIESACID-BASE THEORIES

1. Arrhenius Acids and Bases

2. Bronsted-Lowry Acids and Bases

3. Lewis Acids and Bases

ARRHENIUS ACIDS AND ARRHENIUS ACIDS AND BASESBASESSwedish chemist Svante Arrhenius (1859-1927)

proposed a revolutionary way of defining and

thinking about acids and bases.

He said that acids are hydrogen-containing

compounds that ionize to yield hydrogen ions

(H+) in aqueous solutions.

He also said that bases are compounds that

ionize to yield hydroxide ions (OH-) in aqueous

solutions.

Acids that contain one ionizable hydrogen,

such as nitric acid, are called monoprotic

acids.

Acids that contain two ionizable hydrogens,

such as sulfuric acid, are called diprotic acids.

Acids that contain three ionizable hydrogens,

such as phosphoric acid, are called triprotic

acids.

BRONSTED-LOWRY ACIDS BRONSTED-LOWRY ACIDS AND BASESAND BASES In 1923, the Danish chemist Johannes

Bronsted and the English chemist Thomas

Lowry independently proposed a new

definition.

Defines an acid as a hydrogen-ion donor.

Defines a base as a hydrogen-ion

acceptor.

A conjugate base is the particle that remains

when an acid has donated a hydrogen ion.

A conjugate base is the particle that remains

when an acid has donated a hydrogen ion.

Conjugate acids and bases are always paired with

a base or an acid, respectively.

A conjugate acid-base pair consists of two

substances related by the loss or gain of a single

hydrogen ion.

LEWIS ACIDS AND BASESLEWIS ACIDS AND BASES

The third theory of acids and bases was

proposed by Gilbert Lewis.

Lewis focused on the donation or acceptance

of a pair of electrons during a reaction.

This concept is more general than either the

Arrhenius theory or the Bronsted-Lowry

theory.

A Lewis acid is a substance that can accept a

pair of electrons to form a covalent bond.

A Lewis base is a substance that can donate a

pair of electrons to form a covalent bond.

A hydrogen ion (Bronsted-Lowry acid) can accept

apair of electrons in forming a bond.

A hydrogen ion, therefore, is also a Lewis acid.

A Bronsted-Lowry base, or a substance that

accepts a hydrogen ion, must have a pair of

electrons available and is also a Lewis base.

pH ConceptpH Concept

A widely used system for expressing [H+] is

the pH scale, proposed in 1909 by the Danish

scientist Soren Sorenson.

It ranges from 0-14, neutral solutions have a

pH of 7.

A pH of 10 is strongly basic.

The pH of a solution is the negative logarithm

of the hydrogen-ion concentration.

The pH may be represented

mathematically using the

following equation:

pH = - log [H+]

Similarly, the pOH of a solution

equals the negative logarithm of

the hydroxide-ion concentration.

pOH = - log [OH-]

A neutral solution has a pOH of 7.

A solution with a pOH less than 7 is basic.

A solution with a pOH greater than 7 is acidic.

A simple relationship between pH and pOH makes

it easy to find either one when the other is known.

pH + pOH = 14

pH = 14 – pOH

pOH = 14 - pH

pHpH

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

10-1 10-2 10-3 10-4 10-5 10-6 10-7 10-8 10-9 10-10 10-1110-12 10-

1310-14 N

EU

TR

AL

[H[H++]]

Increasing BasicityIncreasing Acidity

PRACTICE EXERCISE:PRACTICE EXERCISE:What is the pH of a solution with

a hydrogen-ion concentration of 1.0 x 10-10 M?

Find the pH of each solution:◦[H+] = 1.0 x 10 -4 M◦[H+] = 0.0010 M◦[H+] = 1.0 x 10 -9 M◦[H+] = 1.0 x 10 -12 M◦[H+] = 0.010 M