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Acids and Bases

Properties,

Reactions, pH, and

Titration

C-19

2017

Properties of acids

1. Taste Sour (don’t try this except with foods).

2. Are electrolytes (conduct electricity).

– Some are strong, some are weak.

3. Change indicator colors. (litmus →red).

4. React with metals to form hydrogen gas.

5. React with hydroxides to form water and a

salt.

Acid’s Reaction with Metals

Metals: Dissolves; Problem: bridges, cars,

buildings

– Magnesium:

2HCl + Mg ---> H2 + MgCl2

– Iron:

2HCl + Fe ---> H2 + FeCl2

– Copper:

2HCl + Cu ---> H2 + CuCl2

Common Acids

Fruits – citric acid

Milk – lactic acid

Vinegar – acetic acid

Soda pop – carbonic and

phosphoric acid

And lots more!!!!

Properties of Bases

1. React with acids to form water and a salt.

2. Taste bitter.(Don’t try this)

3. Feel slippery (Don’t try this either).

4. Can be strong or weak electrolytes.

5. Change indicators (litmus → blue).

Common Bases

Windex – ammonia

Baking soda – sodium

bicarbonate

Drain cleaner – NaOH

Milk of Magnesia –

Mg(OH)2

And more…..

Organic acids

found in living things (fruits, etc)

contain -COOH a carboxyl group

weak acids are only slightly ionized to -COO-

Called carboxylic acids

Mineral acids

from inorganic materials (rocks)

traditional acids - used industrially

Common Industrial Acids

Sulfuric acid - H2SO4 – petroleum, fertilizer, metallurgy, paper, paints,batteries, etc

Nitric acid – HNO3 – explosives, rubber, plastics, pharmaceuticals, etc.

Phosphoric – H3PO4 – fertilizer, flavoring agent, detergents, etc.

Hydrochloric – HCl – pickling metal, cleaning, chlorination (pools)

Acetic Acid – CH3COOH – plastics, food supplements, etc.

Nomenclature

Two basic types of acids: binary and oxyacids

1. Binary acids – 2 elements only

hydro + stem + ic acid

– HCl – hydrochloric acid

– HI – hydroiodic acid

Nomenclature

2. Oxyacid names –

anion stem + ous (ite anions) NO2-1 (nitrite)

– HNO2 nitrous acid

Or anion stem + ic (ate anions) NO3-1 (nitrate)

– HNO3 nitric acid

– HClO4 perchloric acid

Nature: Electrolytes are classified as Acids, Bases, or Salts

Acids - react with H2O and produce H+

– The H+ ion combines with water and forms H3O+

– called the “hydronium ion”

Bases – dissociate with H2O and produce OH-

Salts - Ionic combinations of metal/nonmetal

ions.

Strong vs. Weak

STRONG electrolytes show complete

ionization in water (all ions); good conductors

– Soluble salts, SA, SB

– NaCl → Na+(aq) + Cl-(aq)

WEAK electrolytes show partial ionization in

water (mostly molecules); poor conductors

– WA, WB

– NH3 + H+ NH4+

Aqueous acids

Arrhenius definition: acids ionize in water to form H+ ions -

- are polar covalent compounds and all have H.

- may ionize in more than 1 step. (ex H2SO4)

Strong acids show complete ionization (100%)

HA → H+1 + A-1

Weak acids produce few ions (less than 5%); are

dissolved intact as molecules.

HA H+1 + A-1

Arrhenius Base

Bases dissociate and produce OH- ions.

Strong bases – 100% dissociation

– Group I and II hydroxides

Weak bases – less than 5% dissociation

– Ammonia, aniline, carbonates are not included.

– All other hydroxides are.

Memorize the Strong Acids

HCl - hydrochloric

HBr – hydrobromic

HI - hydroiodic

H2SO4 - sulfuric

HClO4 – perchloric

HNO3 - nitric

Memorize the Strong Bases

NaOH - sodium hydroxide

KOH - potassium hydroxide

LiOH – lithium hydroxide

RbOH - rubidium hydroxide

Ba(OH)2 – barium hydroxide

Sr(OH)2 – strontium hydroxide

Ca(OH)2 - calcium hydroxide

Mg(OH)2 – magnesium hydroxide

Acid definitions

Bronsted Lowry

– Acids are proton donors

– Bases are proton acceptors

Acids and bases occur in conjugate pairs

Come in Pairs

General equation

HA(aq) + H2O(l) H3O+(aq) + A-(aq)

Acid + Base Conjugate + Conjugate

acid base

Conjugate pairs

This is an equilibrium.

B(aq) + H2O(l) BH+(aq) + OH-(aq)

Base + Acid Conjugate acid +Conjugate base

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

In Bronsted-Lowry theory, bases do not require OH-

Bases are able to accept protons

Allows ammonia and carbonate ions to be

considered bases, others as well.

NH3 + H+ → NH4+

Base + H+ → Conjugate acid

Most accepted theory

Acid & Base Reactions

Neutralization Reaction:

– Acid + Base → “salt” + H2O (usually)

– “Salt” = general term for an ionic compound

– Example:

HCl + NaOH → NaCl + H2O

Acid-Base reactions

Are equilibrium reactions (reversible)

Compare strength of the two acids (charts)

Equilib. shifts away from the stronger acid.

HClO4 + H2O ⇆ H3O+ + ClO4

-

Acid + base ⇆ cong.acid + cong. Base

HClO4 is a stronger acid than H3O+ so….

Equilibrium shifts to the right →

– away from HClO4

Protons are Hydrogen ions

Monoprotic acids have one proton to donate

ex. HCl

Diprotic acids have two protons to donate

ex. H2SO4 (one step at a time)

Polyprotic – two or more protons to donate

ex. H3PO4

Amphoteric substances

Substances which can either accept or donate

a proton.

Water is an example

H2O + H+ → H3O+ (water as a base)

H2O → H+ + OH- (water as an acid)

Other examples are NH3 and HSO4-

Lewis Theory

Lewis Acid – accepts an electron pair

Lewis Base – donates an electron pair

Not frequently used for chemists

Most general definition

(same G. Lewis that made e-dot diagrams)

Lewis Acids and Bases

Lewis Acid

– A species (atom, ion or molecule) that is an

electron pair acceptor.

Lewis Base

– A species that is an electron pair donor.

base acid adduct

Showing Electron Movement

Focus On Acid Rain

CO2 + H2O → H2CO3

H2CO3 + H2O → HCO3- + H3O

+

HCO3- + H3O

+ → CO2 + H2O

3 NO2 + H2O → 2 HNO3 + NO

or

SO2 + H2O → H2SO3

Acid Rain

Acid rain

Gases like sulfur dioxide and nitrogen dioxide are produced from burning coal, oil, and other fuels.

These gases react with water vapor in the atmosphere to form acids.

Acid rain can be stopped with govt. regulations.

Less in US/Canada now, but more in China/India

Acid/Base Titration - a lab process

Basic Concepts:

– 1. Acids & bases neutralize each other

– 2. From the balanced equation, the number of

moles needed of the “known” reactant & the

“unknown” reactant are given.

– 3. An indicator is selected based on the strength of

the “known” reactant.

– 4. The indicator will change color when the “known”

reactant equals the “unknown”.

– 5. Concentration of the “unknown” is calculated.

7 Steps

1. Fill Burette with NaOH (known)

2. Place 20ml HCl in flask (unknown)

– The amount may be different, but record

3. Place indicator in HCl

4. Slowly add NaOH until the endpoint is reached (color change).

5. Record amount of NaOH used (let’s pretend 19.9ml)

6. Use the factor label method to find the number of moles of

NaOH.

7. Look at the balanced equation to determine the ratio of moles

between the “Known” NaOH & “unknown” HCl.

Titration calculation

Use the equation: Ma x Va = Mb x Vb

Example: 25 ml of HCl is neutralized by 20 ml of 0.5 M NaOH. Find conc. of HCl.

Solution: Ma = Mb x Vb / Va

Ma = 0.5 M x 20mL / 25mL = 0.4 M HCl

Water

• Self ionization of water. (very small amount)

• H2O H+ + OH- • [H+ ] = [OH-] = 1 x 10-7M

• A neutral solution.

• In water: Kw = [H+ ] x [OH-] = 1 x 10-14

• Kw is called the ion product constant.

Ion Product Constant

• H2O H+ + OH-

• Kw is constant 1 x 10-14

• If [H+] > 10-7 then [OH-] < 10-7 (acidic)

• If [H+] < 10-7 then [OH-] > 10-7 (basic)

• If we know one, we can determine the other.

• If [H+] = 1x 10-3 Find [OH-] • Kw/ [H+] = [OH-]

• 1 x 10-14/1 x 10-3 = [OH-] = 1 x 10-11

Logarithms

• Powers of ten. A shorthand form

• pH = -log[H+]

• in neutral pH = - log(1 x 10-7) = 7

• in acidic solution [H+] > 10-7

• pH < -log(10-7)

• pH < 7

• in base pH > 7

pH and pOH equations

• pH = -log[H+]

• pOH = - log [OH-]

• [H+] x [OH-] = 1 x 10-14

• pH + pOH = 14

0 1 3 5 7 9 11 13 14

0 1 3 5 7 9 11 13 14

Basic

100 10-1 10-3 10-5 10-7 10-9 10-11 10-13 10-14

Basic 100 10-1 10-3 10-5 10-7 10-9 10-11 10-13 10-14

Acidic Neutral

[OH-]

pH

[H+]

pOH