· Acids, Bases, and Salts • You should be able to • Understand the acid-base theories of Arrhenius, Brønsted-Lowry and Lewis • Identify strong acids and bases and calculate

Acids and

Bases

Mini Research

What is an acid?

What makes acids dangerous?

Is acid rain an issue for us?

What does pH balanced mean?

Table of Contents‘Acids, Bases, and Salts’

• Definitions• pH Scale• Common Acids• Common Bases• LeChatelier’s Principle• Conjugate Acid-Base Pairs• pKa• Concentration vs. Strength• pH Indicators• Buffers• Titration

Acids, Bases, and Salts

• You should be able to

• Understand the acid-base theories of Arrhenius, Brønsted-Lowry and Lewis

• Identify strong acids and bases and calculate their pH’s

• Calculate the pH of a weak acid or base

• Calculate the concentration of a strong or weak acid or base from its pH

• Calculate the pH and ion concentration in a polyprotic acid

• Predict the pH of a salt from its formula and then calculate the pH of the salt

• Be familiar with titration curves and selection of an acid-base indicator

Acids and Bases

• Acids and bases play an important role in our lives but are sometimes safe and sometimes dangerous.

• Demonstrations explain pH and how it is measured.

https://cvod.infobase.com/PortalPlaylists.aspx?wID=225368&xtid=35316

pH scale

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

ACID BASE

NEUTRAL

Each step on pH scale represents a factor of 10.

pH 5 vs. pH 6 (10X more acidic)pH 3 vs. pH 5 (100X different)

pH 8 vs. pH 13 (100,000X different)

: measures acidity/basicity

10x10x10x100x

pH scale

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

ACID BASE

NEUTRAL

Each step on pH scale represents a factor of 10.

pH 5 vs. pH 6 (10X more acidic)pH 3 vs. pH 5 (100X different)

pH 8 vs. pH 13 (100,000X different)

: measures acidity/basicity

10x10x10x100x

Søren Sorensen(1868 - 1939)

pH is defined as the negative base-10 logarithm of the hydrogen ion concentration

pH = – log [H+] or [H+] = 10-pH

Hydrogen ion concentration in pure water is 1 x 10-7 M at 25ºC

Example: the pH of pure water is – log [1.0 x 10-7] = 7.00

pH decreases with increasing [H+] — adding an acid to pure water increases the hydrogen ion concentration and decreases the hydroxide ion concentration.

Adding a base to pure water increases the hydroxide ion concentration and decreases the hydrogen ion concentration—pH increases with decreasing [H+].

pH scale runs from pH = 0 (corresponding to 1M H+) to pH 14 (corresponding to 1 M OH–).

Relationships between acidity, basicity, and pH:

If pH = 7.0, the solution is neutral

If pH < 7.0, the solution is acidic

If pH > 7.0, the solution is basic

Acid

Base

pH = -log [H1+]

pH = 7

Acidic Basic

Neutral

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

Acids and Bases

pH < 7 pH > 7

taste sour taste bitter

react w/bases react w/acids

proton (H1+) donor proton (H1+) acceptor

turn litmus red turn litmus blue

lots of H1+/H3O1+ lots of OH1–

react w/metals don’t react w/metals

Both are electrolytes.

Acid vs. Base

Acid

pH > 7

bitter taste

does not

react with

metals

pH < 7

sour taste

react with

metals

Alike Different

Related to

H+ (proton)

concentration

pH + pOH = 14

Affects pH

and

litmus paper

Base

Different

Topic Topic

Properties

electrolytes

turn litmus red

sour taste

react with metals to

form H2 gas

slippery feel

turn litmus blue

bitter taste

ChemASAP

vinegar, milk, soda,

apples, citrus fruits

ammonia, lye, antacid,

baking soda

electrolytes

Acid

Sour taste

Turns blue litmus red

Reacts with some metals to produce H2Dissolves carbonate salts, releasing CO2

Base

Bitter taste

Turns red litmus blue

Slippery to the touch

Common Acids and Bases

Strong Acids (strong electrolytes)

HCl hydrochloric acid

HNO3 nitric acid

HClO4 perchloric acid

H2SO4 sulfuric acid

Weak Acids (weak electrolytes)

CH3COOH acetic acid

H2CO3 carbonic

Strong Bases (strong electrolytes)

NaOH sodium hydroxide

KOH potassium hydroxide

Ca(OH)2 calcium hydroxide

Weak Base (weak electrolyte)

NH3 ammonia

Weak Base (weak electrolyte)

NH4OH ammonia

NH3 + H2O → NH4OH

Acid + Base → Salt + Water

• Orange juice + milk → bad taste

• Evergreen shrub + concrete → dead bush

• Under a pine tree + fertilizer → white powder

HCl + NaOH → NaCl + HOH

salt water

Acid-Base Neutralization

1+ 1-

+ +

Hydronium ion Hydroxide ion

H3O+ OH-

Water

H2O

Water

H2O

Water

H2O

Water

H2O

Acid-Base Neutralization

1+ 1-

+ +

Hydronium ion Hydroxide ion Water

H3O+ OH- H2O

Water

H2O

Common Acids

Common Acids

Sulfuric Acid H2SO4

Nitric Acid HNO3

Phosphoric Acid H3PO4

Hydrochloric Acid HCl

Acetic Acid CH3COOH

Carbonic Acid H2CO3

Battery acid

Used to make fertilizers

and explosives

Food flavoring

Stomach acid

Vinegar

Carbonated water

http://upload.wikimedia.org/wikipedia/commons/0/08/Phosphoric-acid-3D-vdW.pnghttp://upload.wikimedia.org/wikipedia/commons/2/24/Sulfuric-acid-3D-vdW.pnghttp://upload.wikimedia.org/wikipedia/commons/e/ef/Acetic-acid-3D-balls.pnghttp://upload.wikimedia.org/wikipedia/commons/6/69/Nitric-acid-3D-balls-B.pnghttp://upload.wikimedia.org/wikipedia/commons/4/48/Hydrogen-chloride-3D-vdW-labelled.pnghttp://upload.wikimedia.org/wikipedia/commons/8/81/Carbonic-acid-3D-vdW.png

Common Acids

Formula Name of Acid Name of Negative Ion of Salt

HF hydrofluoric fluoride

HBr hydrobromic bromide

HI hydroiodic iodide

HCl hydrochloric chloride

HClO hypochlorous hypochlorite

HClO2 chlorous chlorite

HClO3 chloric chlorate

HClO4 perchloric perchlorate

H2S hydrosulfuric sulfide

H2SO3 sulfurous sulfite

H2SO4 sulfuric sulfate

HNO2 nitrous nitrite

HNO3 nitric nitrate

H2CO3 carbonic carbonate

H3PO3 phosphorous phosphite

H3PO4 phosphoric phosphate

Formation of Hydronium Ions

1+

hydronium ion

H3O+

+

hydrogen ion

H+

water

H2O

1+

(a proton)

1+

Sulfuric Acid, H2SO4

Sulfuric acid is the most commonly produced industrial chemical in the world.

Uses: petroleum refining, metallurgy, manufacture of fertilizer,

many industrial processes: metals, paper, paint, dyes, detergents

Sulfuric acid is used in

automobile batteries.

H2SO4“oil of vitriol”

http://upload.wikimedia.org/wikipedia/commons/2/24/Sulfuric-acid-3D-vdW.png

Nitric Acid, HNO3

Nitric acid stains proteins yellow (like your skin).

Uses: make explosives, fertilizers, rubber, plastics, dyes, and pharmaceuticals.

HNO3

“aqua fortis”

O

OO

N H

http://upload.wikimedia.org/wikipedia/commons/6/69/Nitric-acid-3D-balls-B.png

Hydrochloric Acid, HCl

The stomach produces HCl to aid in the digestion of food.

Uses: For ‘pickling’ iron and steel.

Pickling is the immersion of metals in acid solution to remove

surface impurities.

A dilute solution of HCl is called muriatic acid (available in many hardware

stores). Muriatic acid is commonly used to adjust pH in swimming pools

and in the cleaning of masonry.

HCl(g) + H2O(l) HCl(aq)hydrogen chloride water hydrochloric acid

http://upload.wikimedia.org/wikipedia/commons/8/88/HCl_molecule_model-VdW_surface.svg

Common Bases

Common Bases

Sodium hydroxide NaOH lye or caustic soda

Potassium hydroxide KOH lye or caustic potash

Magnesium hydroxide Mg(OH)2 milk of magnesia

Calcium hydroxide Ca(OH) 2 slaked lime

Ammonia water NH3 H2O household ammonia

Name Formula Common Name

.NH4OH

NH41+ + OH1-

ammonium hydroxide

hydroxide

ion

OH1-

http://upload.wikimedia.org/wikipedia/commons/5/5a/Hydroxide-3D-vdW.pnghttp://upload.wikimedia.org/wikipedia/commons/5/5a/Hydroxide-3D-vdW.pnghttp://upload.wikimedia.org/wikipedia/commons/5/5a/Hydroxide-3D-vdW.pnghttp://upload.wikimedia.org/wikipedia/commons/5/5a/Hydroxide-3D-vdW.pnghttp://upload.wikimedia.org/wikipedia/commons/5/5a/Hydroxide-3D-vdW.pnghttp://upload.wikimedia.org/wikipedia/commons/5/5a/Hydroxide-3D-vdW.png

Common Bases

Sodium hydroxide NaOH lye or caustic soda

Potassium hydroxide KOH lye or caustic potash

Magnesium hydroxide Mg(OH)2 milk of magnesia

Calcium hydroxide Ca(OH) 2 slaked lime

Ammonia water NH3 H2O household ammonia

Name Formula Common Name

.NH4OH

NH41+ + OH1-

ammonium hydroxide

hydroxide

ion

OH1-

http://upload.wikimedia.org/wikipedia/commons/5/5a/Hydroxide-3D-vdW.png

Relative Strengths of Acids and Bases

perchloric HClO4hydrogen chloride HCl

nitric HNO3sulfuric H2SO4hydronium ion H3O

+

hydrogen sulfate ion HSO4-

phosphoric H3PO4acetic HC2H3O2carbonic H2CO3hydrogen sulfide H2S

ammonium ion NH4+

hydrogen carbonate ion HCO3-

water H2O

ammonia NH3hydrogen H2

Decre

asin

g A

cid

Str

ength

perchlorate ion ClO4-

chloride ion Cl-

nitrate ion NO3-

hydrogen sulfate ion HSO4-

water H2O

sulfate ion SO42-

dihydrogen phosphate ion H2PO4-

acetate ion C2H3O2-

hydrogen carbonate ion HCO3-

hydro sulfide ion HS-

ammonia NH3carbonate ion CO3

2-

hydroxide ion OH-

amide ion NH2-

hydride ion H-

Decre

asin

g B

ase S

trength

Acid Formula Conjugate base Formula

Metcalfe, Williams, Catska, Modern Chemistry 1966, page 229 acid conjugate base + H+

Binary Hydrogen Compounds

Oxysalts + H2O → Oxyacids

Binary Hydrogen Compoundsof Nonmetals When Dissolved in Water

(These compounds are commonly called acids.)

The prefix hydro- is used to represent hydrogen, followed by the name

of the nonmetal with its ending replaced by the suffix –ic and the word

acid added.

Examples:

*HCl

HBr

*The name of this compound would be hydrogen chloride if it was NOT dissolved in water.

Hydrochloric acid

Hydrobromic acid

Naming Simple Chemical Compounds

Ionic (metal and nonmetal) Covalent (2 nonmetals)

Metal

Forms

only one

positive

ion

Forms

more than

one positive

ion

Nonmetal

Use the

name of

element

Use element

name followed

by a Roman

numeral to

show the charge

First

nonmetal

Second

nonmetal

Before

element name

use a prefix

to match

subscript

Use a prefix

before

element name

and end

with ide

Single

Negative

Ion

Polyatomic

Ion

Use the name

of the

element, but

end with ide

Use the

name of

polyatomic

ion (ate or

Ite)

Naming Ternary Compounds from Oxyacids

The following table lists the most common families of oxy acids.

one more

oxygen atom

most

“common”

one less

oxygen

two less

oxygen

HClO4perchloric acid

HClO3chloric acid

HClO2chlorous acid

HClO

hypochlorous acid

H2SO4sulfuric acid

H2SO3sulfurous acid

H3PO4phosphoric acid

H3PO3phosphorous acid

H3PO2hypophosphorous acid

HNO3nitric acid

HNO2nitrous acid

(HNO)2hyponitrous acid

An acid with a

name ending in

A salt with a

name ending in

-ic

-ous

-ate

-iteforms

forms

Oxyacids → Oxysalts

If you replace hydrogen with a metal, you have formed an oxysalt.

A salt is a compound consisting of a metal and a non-metal. If the

salt consists of a metal, a nonmetal, and oxygen it is called an

oxysalt. NaClO4, sodium perchlorate, is an oxysalt.

HClO4perchloric acid

HClO3chloric acid

HClO2chlorous acid

HClO

hypochlorous acid

NaClO4sodium perchlorate

NaClO3sodium chlorate

NaClO2sodium chlorite

NaClO

sodium hypochlorite

OXYACID OXYSALT

ACID SALT

per stem ic changes to per stem ate

stem ic changes to stem ate

stem ous changes to stem ite

hyper stem ous changes to hypo stem ite

HClO3 + Na1+ NaClO3 + H

1+

acid cation salt

Definitions

Arrhenius Acids and Bases

Acids release hydrogen ions in water.

Bases release hydroxide ions in water.

An acid is a substance that produces hydronium ions, H3O+,

when dissolved in water.

Lewis Definitions

A Lewis acid is a substance than can accept (and share) an electron pair.

A Lewis base is a substance than can donate (and share) an electron pair.

Lewis Acid

Brønsted-Lowry Definitions

A Brønsted-Lowry acid is a proton donor; it donates a hydrogen ion, H+.

A Brønsted-Lowry base is a proton acceptor; it accepts a hydrogen ion, H+.

Brønsted-Lowry

Arrhenius

acids

Acid Definitions

Acid Definitions

Lewis acids

Brønsted-Lowry

Arrhenius

acids

The Arrhenius model of acids

and bases was broadened by

the Brønsted-Lowry model.

The Lewis acid-base model is

the most general in scope.

The Lewis definition of an acid

includes any substance that

is an electron pair acceptor;

a Lewis base is any substance

that can act as an electron pair

donor.

Lewis acids

Brønsted-Lowry

Arrhenius

acids

The Arrhenius model of acids

and bases was broadened by

the Brønsted-Lowry model.

The Lewis acid-base model is

the most general in scope.

The Lewis definition of an acid

includes any substance that

is an electron pair acceptor;

a Lewis base is any substance

that can act as an electron pair

donor.

Acid Definitions

Acid – Base Systems

Type Acid Base

Arrhenius H+ or H3O +

producer

OH - producer

Brønsted-

Lowry

Proton (H +)

donor

Proton (H +)

acceptor

Lewis Electron-pair

acceptor

Electron-pair

donor

Neutralization

Neutralization is a chemical reaction between an acid and a base

to produce a salt (an ionic compound) and water.

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

base acid salt water

Some neutralization reactions:

H2SO4(aq) + NaOH(aq) Na2SO4 + HOH

sulfuric acid sodium hydroxide sodium sulfate water

HC2H3O2(aq) + Ca(OH)2(aq) Ca(C2H3O2)2 + HOH

acetic acid calcium hydroxide calcium acetate water

2 2

2 2

ACID + BASE → SALT + WATER

HCl + NaOH → NaCl + H2O

HC2H

3O

2+ NaOH → NaC

2H

3O

2+ H

2O

• Salts can be neutral, acidic, or basic.

• Neutralization does not mean pH = 7.

weak

strong strong

strong

neutral

basic

Keys to Success

• You must recognize that while each classification has a specific definition, any given molecule can fall into more than one category• Some fall into all 3

categories.

Arrhenius Acids

• An Arrhenius acid is any species that increases the concentration of [H3O+] ions—or protons—in aqueous solution.

• Acid: a substance that produces H3O+ ions in aq solution.

In other words…• An Arrhenius acid is a

molecule that when dissolved in H2O will donate an H+ in solution

• This is known as a proton donor

The trick:

Look for a molecule that starts with an H and typically contains an O or a Halogen

Common Examples of Arrhenius Acids Include:

HCl

HNO3H2SO4

HCH3CO2

• An acid dissociating in water does not form a free-floating proton

• One of the water molecules in solution will grab the H+ yielding a hydronium or H3O

+ ion

• Example: Here’s what happens when nitric acid dissociates in water.

Example: Under the Arrhenius definition, HCl is an acid because it produces H3O

+ ions in solution.

HCl (aq) + H2O→ H+ + Cl-

HCl (aq) + H2O → H3O+ (aq) + Cl- (aq)

The process that converts a molecule such as HCl into ions is

called ionization. Ionization is the production of ions from molecular

compounds. No ions were initially present.

But what if the acid is not dissolved in water?

Bronsted-Lowry Acid

• A Bronsted-Lowry acid, like an Arrhenius acid, is a compound that breaks down to give an H+ in solution

• The only difference is that the solution does not have to be water

• We will still use the same acids list but our solvent is going to change to ammonia, alcohol, or anything else

What does that look like?

• Now let’s see what happens when an acid dissolves in ammonia (NH3)

HNO3 + NH3→ NH4+ + NO3

-

•NH3 picked up the free floating H+

Let’s see what happens when an acid dissolves in methanol (CH3OH)

HNO3 + CH3OH → CH3OH2+ + NO3

-

Bronsted-Lowry Bases

• This definition focuses on the transfer of H+ ions in an acid-base reaction—this definition focuses on the idea of a proton donor and a proton acceptor. • Remember that an acid is a proton (H+ ion)

donor• A Bronsted-Lowry Base is a proton (H+ ion)

acceptor

Examples of Bronsted-Lowry Base & Acid

• HCl is an acid because it donates a proton to water.

HCl (aq) + H2O (l) → H3O+ (aq) + Cl- (aq)

• NH3 is a base because it accepts a proton from water.

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

Identifying Bronsted-Lowry Acids & Bases

• All acids have a conjugate base, and all bases have a conjugate acid.

• In an acid-base reaction, • A base accepts a proton and becomes a

conjugate acid.• An acid donates a proton and becomes a

conjugate base.

NH3(aq) + H2O(l) → NH4+(aq) + OH-(aq)

Acid Conj. baseBase Conj. acid

Example

• In the following reaction, identify the Bronsted-Lowry acid, the Bronsted-Lowry base, conjugate acid, and conjugate base.

H2SO4(aq) + H2O(l) → HSO4-(aq) + H3O(aq)

Step #1: Identify your acid & base. Remember that:1. Acid: proton (H+

ion) donor.2. Base: proton (H+

ion) acceptor.

Step #2: Identify your conjugate acid & base

Check Your Understanding

In the following reaction, identify the Bronsted-Lowry acid, the Bronsted-Lowry base, conjugate acid, and conjugate base.

HCO3-(aq) + H2O(l) → H2CO3(aq) + OH

-


F -

H2PO4-

H2O

HF

H3PO4

H3O+

◆ Give the conjugate base for each of the following:


Br -

HSO4-

CO32-

HBr

H2SO4

HCO3-

◆ Give the conjugate acid for each of the following:

Lewis Acids

• The Lewis definition for acids is the most extreme because it’s not dealing with protons specifically

• Instead the Lewis definition deals with the movement of electrons → picks up an e- pair

• The atom getting attacked or accepting those electrons is the Lewis acid in that reaction

Common Lewis Acid Examples

• All cations are Lewis acids since they are able to accept electrons. (e.g., Cu2+, Fe2+, Fe3+)

• An atom, ion, or molecule with an incomplete octet of electrons can act as an Lewis acid (e.g., BF3, AlF3).

• Molecules where the central atom can have more than 8 valence shell electrons can be electron acceptors, and thus are classified as Lewis acids (e.g., SiBr4, SiF4).

• Molecules that have multiple bonds between two atoms of different electronegativities (e.g., CO2, SO2)

Example: Formation of Sulfuric Acid

SO3 + H2O → H2SO4

Difference between

Lewis Acids & Bases

• Lewis Acids• Are electrophilic; e- attracting

• Lewis Bases• Are nucleophilic; “attack” a positive

charge with a lone pair

Lewis Bases

• An atom, ion, or molecule with a lone-pair of electrons can thus be a Lewis base.

Each of the following anions can "give up" their electrons to an acid, e.g., OH− , CN− , CH3COO− , :NH3 , H2O: , CO: .

Warm-up• For each

molecule or ion in the table, identify whether it can act as an acid or a base and put a checkmark under each theory or theories that describe it.

Molecule/Ion

Acid or Base

ArrheniusBronsted-

LowryLewis

Br-

CN-

H2CO3

NH3

HNO2

Ba(OH)2

HCl

AlCl3

Cl-

KOH

IO3

CH3COOH

Quiz Time!

pH Scale

Acid Base

0

7

14

Zumdahl, Zumdahl, DeCoste, World of Chemistry 2002, page 515

[H+] pH

10-14 14

10-13 13

10-12 12

10-11 11

10-10 10

10-9 9

10-8 8

10-7 7

10-6 6

10-5 5

10-4 4

10-3 3

10-2 2

10-1 1

100 0

1 M NaOH

Ammonia

(household

cleaner)

Blood

Pure water

Milk

Vinegar

Lemon juice

Stomach acid

1 M HCl

Acid

ic

N

eutr

al

Basic

pH of Common Substances

Timberlake, Chemistry 7th Edition, page 335

1.0 MHCl0

gastricjuice1.6

vinegar2.8

carbonated beverage3.0

orange3.5

apple juice3.8

tomato4.2

lemonjuice2.2 coffee

5.0

bread5.5

soil5.5

potato5.8

urine6.0

milk6.4

water (pure)7.0

drinking water7.2

blood7.4

detergents8.0 - 9.0

bile8.0

seawater8.5

milk of magnesia10.5

ammonia11.0

bleach12.0

1.0 MNaOH(lye)14.0

8 9 10 11 12 14133 4 5 621 70

acidic neutral basic

[H+] = [OH-]

pH of Common Substance

14 1 x 10-14 1 x 10-0 0

13 1 x 10-13 1 x 10-1 1

12 1 x 10-12 1 x 10-2 2

11 1 x 10-11 1 x 10-3 3

10 1 x 10-10 1 x 10-4 4

9 1 x 10-9 1 x 10-5 5

8 1 x 10-8 1 x 10-6 6

6 1 x 10-6 1 x 10-8 8

5 1 x 10-5 1 x 10-9 9

4 1 x 10-4 1 x 10-10 10

3 1 x 10-3 1 x 10-11 11

2 1 x 10-2 1 x 10-12 12

1 1 x 10-1 1 x 10-13 13

0 1 x 100 1 x 10-14 14

NaOH, 0.1 M

Household bleach

Household ammonia

Lime water

Milk of magnesia

Borax

Baking soda

Egg white, seawater

Human blood, tears

Milk

Saliva

Rain

Black coffee

Banana

Tomatoes

Wine

Cola, vinegar

Lemon juice

Gastric juice

More

basic

More

acid

icpH [H1+] [OH1-] pOH

7 1 x 10-7 1 x 10-7

7

Acid – Base Concentrations

pH = 3

pH = 7

pH = 11

OH-

H3O+OH-

OH-H3O+

H3O+

[H3O+] = [OH-][H3O

+] > [OH-] [H3O+] < [OH-]

acidic

solution

neutral

solutionbasic

solution

co

nc

en

tra

tio

n (

mo

les

/L)

10-14

10-7

10-1


pH

pH = -log [H1+]

Kelter, Carr, Scott, Chemistry A World of Choices 1999, page 285

pH Calculations

pH

pOH

[H3O+]

[OH-]

pH + pOH = 14

pH = -log[H3O+]

[H3O+] = 10-pH

pOH = -log[OH-]

[OH-] = 10-pOH

[H3O+] [OH-] = 1 x10-14

pH = - log [H+]

pH = 4.6

pH = - log [H+]

4.6 = - log [H+]

- 4.6 = log [H+]

- 4.6 = log [H+]

Given:

2nd log

10x

antilog

multiply both sides by -1

substitute pH value in equation

take antilog of both sides

determine the [hydronium ion]

choose proper equation

[H+] = 2.51x10-5 M

You can check your answer by working backwards.

pH = - log [H+]

pH = - log [2.51x10-5 M]

pH = 4.6

Recall, [H+] = [H3O+]

Acid Dissociation

monoprotic

diprotic

polyprotic

HA(aq) H1+(aq) + A1-(aq)

0.03 M 0.03 M 0.03 M

pH = - log [H+]

pH = - log [0.03M]

pH = 1.52

e.g. HCl, HNO3

H2A(aq) 2 H1+(aq) + A2-(aq)

0.3 M 0.6 M 0.3 M

pH = - log [H+]

pH = - log [0.6M]

pH = 0.22

e.g. H2SO4

Given: pH = 2.1

find [H3PO4]

assume 100%

dissociation

e.g. H3PO4

H3PO4(aq) 3 H1+(aq) + PO4

3-(aq)

? M x M

pH = ?

Given: pH = 2.1

find [H3PO4]

assume 100%

dissociation

H3PO4(aq) 3 H1+(aq) + PO4

3-(aq)

X M 0.00794 M

Step 1) Write the dissociation of phosphoric acid

Step 2) Calculate the [H+] concentration pH = - log [H+]

2.1 = - log [H+]

- 2.1 = log [H+]

2nd log - 2.1 = log [H+]2nd log

[H+] = 10-pH

[H+] = 10-2.1

[H+] = 0.00794 M

[H+] = 7.94 x10-3 M7.94 x10-3 M

Step 3) Calculate [H3PO4] concentration

Note: coefficients (1:3) for (H3PO4 : H+)

7.94 x10-3 M3

= 0.00265 M H3PO4

How many grams of magnesium hydroxide are needed to add to 500 mL of H2O

to yield a pH of 10.0?

Step 1) Write out the dissociation of magnesium hydroxide Mg2+ OH1-

Mg(OH)2Mg(OH)2(aq) Mg2+(aq) 2 OH1-(aq)+

Step 2) Calculate the pOH pH + pOH = 14

10.0 + pOH = 14

pOH = 4.0

Step 3) Calculate the [OH1-] pOH = - log [OH1-]

[OH1-] = 10-OH

[OH1-] = 1 x10-4 M

1 x10-4 M0.5 x10-4 M5 x10-5 M

Step 4) Solve for moles of Mg(OH)2

L

mol M =

L 0.5

molx M x105 5- = x = 2.5 x 10-5 mol Mg(OH)2

Step 5) Solve for grams of Mg(OH)2

x g Mg(OH)2 = 2.5 x 10-5 mol Mg(OH)2 1 mol Mg(OH)2

= 0.00145 g Mg(OH)258 g Mg(OH)2

Equilibrium

• LeChatelier’s Principle

CO2 + CaO CaCO3“chicken

breath”“food” “egg shell”

I WISH I HAD

SWEAT GLANDS.

As temperature increases, chickens “pant” more.

This stresses the system and shifts the equilibrium to the LEFT

This makes the egg shells THIN and fragile.

[ CaO ] , shift

[ CO2 ] , shift

-- shift ; eggshells are thinner

In a chicken… CaO + CO2 CaCO3(eggshells)

In summer, [ CO2 ] in a chicken’s blood due to panting.

How could we increase eggshell thickness in summer?

-- give chickens carbonated water

-- put CaO additives in chicken feed

-- air condition the chicken house TOO much $$$

-- pump CO2 gas into the chicken house

would kill all the chickens!

I wish I had

sweat glands.

LeChatelier’s Principle

N2 + 3 H2 2 NH3 + heat

Raising the temperature……favors the endothermic reaction (the reverse

reaction) in which the rise in temperature is

counteracted by the absorption of heat.

Increasing the pressure……favors the forward reaction in which 4 mol

of gas molecules is converted to 2 mol.

Decreasing the concentration

of NH3…

…favors the forward reaction in order to

replace the NH3 that has been removed.

Dorin, Demmin, Gabel, Chemistry The Study of Matter 3rd Edition, page 532 Animation by Raymond Chang

All rights reserved

http://www.mhhe.com/physsci/chemistry/essentialchemistry/flash/lechv17.swf

Equilibrium Expression

322

2

3eq

HN

NHK = reactantsproducts

Keq =

N2 + 3 H2 2 NH3 + heat

Dorin, Demmin, Gabel, Chemistry The Study of Matter 3rd Edition, page 532

Haber Process

reversible reaction:

H2SO4 2 H1+ + SO4

2–

Acid dissociation is a reversible reaction.

Rate at which

R → P

Rate at which

P → R=

looks like nothing is happening, however…

system is dynamic, NOT static

equilibrium:

Reactant → Product and P → RReactant Product

Remove NH3…………………..

“ “ NH3…………………

“ “ H2…………………..

Add more N2…………………..

Le Chatelier’s principle

N2(g) + 3 H2(g) 2 NH3(g)

Le Chatelier’s principle:

Disturbance Equilibrium Shift

no shift

When a system at equilibrium is disturbed, it shifts to a

new equilibrium that counteracts the disturbance.

Add a catalyst…………………

Increase pressure…………….Fritz Haber

shift to a new equilibrium:

Then go inside…

shift to a new equilibrium:

Light-Darkening Eyeglasses

AgCl + energy Ago + Clo

“energy”

Go outside… Sunlight more intense than inside light;

GLASSES DARKEN

(clear) (dark)

“energy”

GLASSES LIGHTEN

Sensitive Sunglasses

Oxidation-reduction reactions are the basis for many interesting and useful applications of technology.

One such application is photochromic glass, which is used for the lenses in light sensitive glasses.

Lenses manufactured by the Corning Glass Company can change from transmitting 85% of light to only

transmitting 22% of light when exposed to bright sunlight.

Photochromic glass is composed of linked tetrahedrons of silicon and oxygen atoms jumbled together

in a disorderly array, with crystals of silver chloride caught in between the silica tetrahedrons. When the

glass is clear, the visible light passes right through the molecules. The glass absorbs ultraviolet light,

however, and this energy triggers an oxidation-reduction reaction

between Ag+ and Cl-:

Ag+ + Cl- --> Ag0 + Cl0

To prevent the reaction from reversing itself immediately, a few ions of Cu+ are incorporated into the

silver chloride crystal. These Cu+ ions react with the newly formed chlorine atoms:

Cu+ + Cl0 --> Cu2+ + Cl-

The silver atoms move to the surface of the crystal and form small colloidal clusters of silver metal.

This metallic silver absorbs visible light, making the lens appear dark (colored).

As the glass s removed from the light, the Cu2+ ions slowly move to the surface of the crystal where

they interact with

the silver metal:

Cu2+ + Ag0 --> Cu+ + Ag+

The glass clears as the silver ions rejoin chloride ions in the crystals.

Maintaining Blood pH

Acid entering the blood stream Carbon dioxide is exhaled

HCO31- + H+ H2CO3 H2O + CO2

Bicarbonate ion circulates in the blood stream where it is in equilibrium with H+ and OH-.

In the lungs, bicarbonate ions combine with a hydrogen ion and lose a water molecule

to form carbon dioxide, which is exhaled.


Alkalosis

If our breathing becomes too fast (hyperventilation)…

Carbon dioxide is removed from the blood too quickly.

This accelerates the rate of degradation of carbonic acid into carbon dioxide and water.

The lower level of carbonic acid encourages the combination of hydrogen ions and

bicarbonate ions to make more carbonic acid. The final result is a fall in blood H1+

levels that raises blood pH which can result in over-excitability or death.


Acidosis

If breathing becomes too slow (hypoventilation)…

…free up acid, pH of blood drops, with associated health risks such as depression

of the central nervous system or death.

The normal pH of blood is between 7.2 – 7.4.

This pH is maintained by the bicarbonate ion and other buffers.


H+ A - H+ A - HA

A - H+ A - H+ A –

H+ A - H+ A - H+

A - HA H+ A -

H+ A - H+ A - H+

HA HA HA HA

HA HA HA

H+ A - HA HA

HA HA H + A –

HA H + A – HA HA

H+ A- H+ A- H+ A- H+ A- HA

A- H+ A- H+ A- H+ A- H+ A -

H+ A- HA H+ A- H+ A- H+ A-

A- H+ A- H+ A- H+ A- H+ A- H+

H+ A - H + A - H + A - HA H + A -

A- H+ A- H+ A- H+ A- H+ A–

H+ A- H+ A- H+ A- H+ A- H+

A- H+ A- H+ A- H+ A- H+ A-

HA A- H+ A- H+ A- H+ A- H+

HA HA H+ A- HA HA HA

HA HA HA HA HA H+ A-

H+ A- HA HA HA HA HA


HA HA HA H+ A- HA HA


HA HA HA H+ A- HA HA



DILUTECONCENTRATED

ST

RO

NG

WE

AK

STRONG ACIDS

Dissociate nearly 100%

HA H1+ + A-

WEAK ACIDS

Dissociate very little

HA H1+ + A-

Acids: Concentration vs. Strength

Comparison of Strong and Weak Acids

Type of acid, HA Reversibility

of reactionKa value

Ions existing when acid,

HA, dissociates in H2O

StrongNot

reversibleKa value very large

H+ and A-, only.

No HA present.

Weak reversible Ka is small H+, A-, and HA

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

The equilibrium expression for the reaction is

Ka = [H3O

+] [A-]

[HA]Note: H3O

+ = H+

Strong vs. Weak Acid


Concentrated vs. Dilute

0.3 M HCl

2.0 M HCl

12.0 M HCl

10.0 M CH3COOHDilute, strong acid

Concentrated, strong acidOR Dilute, strong, acid

Concentrated, strong acid

Concentrated, weak acid

Naming Acids

_________ ide

(chloride, Cl1-)

_________ite

(chlorite, ClO2-)

(hypochlorite, ClO-)

_________ ate

(chlorate, ClO3-)

(perchlorate, ClO4-)

Hydro____ ic acid

(hydrochloric acid, HCl)

_________ic acid

(chloric acid, HClO3)

(perchloric acid, HClO4)

______ous acid

(chlorous acid, HClO2)

(hypochlorous acid, HClO)

Anion Acid

add H+

add H+

add H+

ions

ions

ions

4A

Group

5A 6A 7A

Period 2CH4

No acid or

base

properties

NH3

Weak base

H2O

---

HF

Weak acid

Period 3SiH4

No acid or

base

properties

PH3

Weak base

H2S

Weak acid

HCl

Strong acid

Increasing acid Strength

Increasing base Strength

Incre

asin

g a

cid

Str

ength

Incre

asin

g b

ase S

trength

Brown, LeMay, Bursten, Chemistry 2000, page 625

[H3O+]

Equilibrium and pH Calculations

HA + H2O H3O+ + A-

Weak acid

HA + H2O H3O+ + A-

Strong acid

acid-dissociation

constant calculations

Ka = [A-] [H3O

+]

[HA]

[HA] = [H3O+]

+

pH0 7 14

antilog(-pH)

-log [H3O+] [OH-]

-

1 x 10-14

[OH-]=

1 x 10-14

[H3O+]

=

Tocci, Viehland, Holt Chemistry Visualizing Matter 1996, page 525

HA H+ + A-

Kw = [H3O+][OH-]

1 x 10-14 = [H3O+][OH-]

[H3O+]

Equilibrium and pH Calculations

HA + H2O H3O+ + A-

Weak acid

HA + H2O H3O+ + A-

Strong acid

acid-dissociation

constant calculations

Ka = [A-] [H3O

+]

[HA]

[HA] = [H3O+]

+

pH0 7 14

antilog(-pH)

-log [H3O+] = [OH-]

-

1 x 10-14

[OH-]=

1 x 10-14

[H3O+]

=

Tocci, Viehland, Holt Chemistry Visualizing Matter 1996, page 525

Strengths of Conjugate Acid-Base Pairs

strong medium weak very weak

Acid strength increases

HCl H2SO4 HNO3 H3O+ HSO4

- H3PO4 HC2H3O2 H2CO3 H2S H2PO4- NH4

+ HCO3- HPO4

2- H2O

negligible very weak weak medium strong

Base strength increases

Cl- HSO4- NO3 H2O SO4

2- H2PO4- C2H3O2

- HCO3- HS- HPO4

2- NH3 CO32- PO4

3- OH-

Kw = [H3O+][OH-]

1 x 10-14 = [H3O+][OH-]

Keqequilibrium constant

Kwwater dissociation

constant

Kaacid dissociation

constant

Kbbase dissociation

constant

H+ + NH3 NH4+NH4

+ H+ + NH3acid CB

CAbase

HA H+ + A-

HA H+ + A-

strong acid

weak acid

0.1 M 0.1 M 0.1 M

0.1 M ? M

Conjugate Acid Strength

Very

strong

Strong

Weak

Very

weak

Relative

acid

strength

Relative

conjugate

base

strength

Very

weak

Very

strong

Weak

Strong

HA H+ + A-

pKa =[H+] [A-]

[HA]


Solutions of Acids and Bases: The Leveling Effect

• No acid stronger than H3O+ and no base stronger than OH– can exist

in aqueous solution, leading to the phenomenon known as the leveling effect.

• Any species that is a stronger acid than the conjugate acid of water (H3O

+) is leveled to the strength of H3O+ in aqueous solution

because H3O+ is the strongest acid that can exist in equilibrium with

water.

• In aqueous solution, any base stronger than OH– is leveled to the strength of OH– because OH– is the strongest base that can exist in equilibrium with water

• Any substance whose anion is the conjugate base of a compound that is a weaker acid than water is a strong base that reacts quantitatively with water to form hydroxide ion

[C] [D]

[Products]

A(g) + 2 B(g) 3 C(g) + D(g)

Weak Acids (pKa)

Weak Acids – dissociate incompletely (~20%)

Strong Acids – dissociate completely (~100%)

Equilibrium constant (Keq) =

Keq = LeChatelier’s Principle(lu-SHAT-el-YAY’s)

[Reactants]

[A][B]

3

2

H+(aq) + C2H3O21-

(aq)CH3COOH(aq)HC2H3O2(aq)

[Reactant]

[Product]Equilibrium constant Keq =

= Ka = Acid dissociation constant

Ka = 1.8 x 10-5 @ 25 oC for acetic acid

[H+][C2H3O21-]

[HC2H3O2]

[H+][C2H3O21-]

[HC2H3O2]Ka =

[H+][C2H3O21-]

[HC2H3O2]=1.8 x 10-5

Assume we begin with 0.1 M acetic acid.

[0.1 M ]

[X ][X ]

X2 = 1.8 x 10-6 M

= 1.34 x 10-3 M[H+]X

pH = -log[H+]

pH = -log[1.34 x10-3]

pH = 2.87

http://upload.wikimedia.org/wikipedia/commons/e/ef/Acetic-acid-3D-balls.png

HC2H3O2 H+ + C2H3O2

1-

HCl H+ + Cl1- very large

HNO3 H+ + NO3

1- very large

H2SO4 H+ + HSO4

1- large

1.8 x 10-5

H2S H+ + HS1- 9.5 x 10

-8

Ionization Constants for Acids

Ka

Ionization of Acids

Acid Ionization Equation Ionization Constant, pKa

Hydrochloric HCl H1+ + Cl1- very large

Sulfuric H2SO4 H1+ + HSO4

1- large

Acetic HC2H3O2 H1+ + C2H3O2

1- 1.8 x 10-5

Formula Name Value of Ka*

Values of Ka for Some Common Monoprotic Acids

HSO4- hydrogen sulfate ion 1.2 x 10-2

HClO2 chlorous acid 1.2 x 10-2

HC2H2ClO2 monochloracetic acid 1.35 x 10-3

HF hydrofluoric acid 7.2 x 10-4

HNO2 nitrous acid 4.0 x 10-4

HC2H3O2 acetic acid 1.8 x 10-5

HOCl hypochlorous acid 3.5 x 10-8

HCN hydrocyanic acid 6.2 x 10-10

NH4+ ammonium ion 5.6 x 10-10

HOC6H5 phenol 1.6 x 10-10

*The units of Ka are mol/L but are customarily omitted.

Incre

asin

g a

cid

str

ength

H2SO4 2 H+ + SO4

2- in dilute solutions...occurs ~100%

H2SO4 H+ + HSO4

1- & HSO41- H+ + SO4

2-

One gram of concentrated sulfuric acid (H2SO4) is diluted to a 1.0 dm3 volume

with water. What is the molar concentration of the hydrogen ion in this solution?

What is the pH?

x mol H2SO4 = 1 g H2SO4

Solution)

First determine the number of moles of H2SO4

Sample 1)

= 0.010 mol H2SO4

OVERALL:

pH = - log [H+]

pH = 1.69

0.010 M 0.020 M

substitute into equation pH = - log [0.020 M]

98 g H2SO4

1 mol H2SO4

A volume of 5.71 cm3 of pure acetic acid, HC2H3O2, is diluted with water at

25 oC to form a solution with a volume of 1.0 dm3.

Step 2) Find the number of moles of acid.

x mol acetic acid = 6.00 g HC2H3O2 = 0.10 mol acetic acid (in 1 L)

M = 0.1 molar HC2H3O2Step 3) Find the [H+]

Ka =

Step 1) Find the mass of the acid

Mass of acid = density of acid x volume of acid

= 1.05 g/cm3 x 5.71 cm3

= 6.00 g

Molarity: M = mol / L

Substitute into equation M = 0.10 mol / 1 L

What is the molar concentration of the hydrogen ion, H+, in this solution?

(The density of pure acetic acid is 1.05 g/cm3.)

(From the formula of acetic acid,

you can calculate that the molar mass of acetic acid is 60 g / mol).

60 g HC2H3O2

1 mol HC2H3O2

Step 3) Find the [H+]

H C H O

HC H O

1

2 3 2

2 3 2

[ ][ ]

[ ]

−+

1.8 x 10-5 =

Ka = 1.8 x 10-5 @ 25 oC for acetic acid

H C H O

HC H O

1

2 3 2

2 3 2

[ ][ ]

[ ]

−+

Ka =

Substitute into equation: ]OH[HC

[x][x] 10 x 1.8

232

5- =

]M [0.10

x 10 x 1.8

25- =

x2 = 1.8 x 10-6 M

x = 1.3 x 10-3 molar = [H+]

HC2H3O2 H+ + C2H3O2

1-

0.1 M

pH = - log[H+]

pH = - log [1.3 x10-3 M]

pH = 2.9

?0.1 Mweak acid

How do the concentrations of

H+ and C2H3O21- compare?

Moles of Acid used to make

1 L of solutionH+ pH

0.010 mol H2SO4 Strong acid

0.100 mol HC2H3O2 Weak acid

Note: although the sulfuric acid is 10x less

concentrated than the acetic acid...

…it produces > 10x more H+

H+ Concentrations…Strong vs. Weak Acid

pH = - log[H+]

1.7

2.9

0.0200 M

0.0013 M

1a) What is the molar hydrogen ion concentration in a 2.00 dm3 solution

of hydrogen chloride in which 3.65 g of HCl is dissolved?

1b) pH

2a) What is the molar concentration of hydrogen ions in a solution

containing 3.20 g of HNO3 in 250 cm3 of solution?

2b) pH

3a) An acetic acid solution is 0.25 M. What is its molar concentration of

hydrogen ions?

3b) pH

4) A solution of acetic acid contains 12.0 g of HC2H3O2 in 500 cm3

of solution. What is the molar concentration of hydrogen ions?

1a) 0.0500 M 2a) 0.203 M 3a) 2.1 x 10-3 M 4) 2.7 x 10-3 M

1b) pH = 1.3 2b) pH = 0.7 3b) pH = 2.7

Practice Problems:

Weak Acids

Cyanic acid is a weak monoprotic acid. If the pH of 0.150 M cyanic acid is 2.32.

calculate Ka for cyanic acid.

HCN(aq) H+(aq) + CN1-(aq)

H3O+(aq)

0.150 M 4.8 x 10-3 M

Ka = [Products]

[Reactants]Ka =

[H3O+]

[HCN]

[CN1-]

Ka = [4.8 x 10-3 M]

[0.150 M]

[CN1-][4.8 x 10-3 M]

Ka = 1.54 x 10-4

4.8 x 10-3 M

pH = -log[H3O+]

10-pH = [H3O+]

10-2.32 = [H3O+]

4.8 x10-3 M = [H3O+]

Weak Acids

Cyanic acid is a weak monoprotic acid. If the initial concentration of cyanic

acid is 0.150 M and the equilibrium concentration of H3O+ is 4.8 x 10-3 M,

calculate Ka for cyanic acid.

HCN(aq) H+(aq) + CN1-(aq)

H3O+(aq)

0.150 M 4.8 x 10-3 M

Ka = [Products]

[Reactants]Ka =

[H3O+]

[HCN]

[CN1-]

Ka = [4.8 x 10-3 M]

[0.150 M]

[CN1-][4.8 x 10-3 M]Ka = 1.54 x 10

-4

How is [H3O+] determined?

4.8 x 10-3 M

Measure pH of solution and work backwards

Acid Dissociation


HCl

Conjugate baseAcid

Conjugate pair

+

1-

Cl

H

Acid Dissociation


HCl

Conjugate baseAcid

Conjugate pair

+

1-

Cl

H

http://upload.wikimedia.org/wikipedia/commons/f/f9/3D_model_hydrogen_bonds_in_water.jpghttp://upload.wikimedia.org/wikipedia/commons/f/f9/3D_model_hydrogen_bonds_in_water.jpghttp://upload.wikimedia.org/wikipedia/commons/f/f9/3D_model_hydrogen_bonds_in_water.jpghttp://upload.wikimedia.org/wikipedia/commons/f/f9/3D_model_hydrogen_bonds_in_water.jpghttp://upload.wikimedia.org/wikipedia/commons/f/f9/3D_model_hydrogen_bonds_in_water.jpghttp://upload.wikimedia.org/wikipedia/commons/f/f9/3D_model_hydrogen_bonds_in_water.jpghttp://upload.wikimedia.org/wikipedia/commons/f/f9/3D_model_hydrogen_bonds_in_water.jpghttp://upload.wikimedia.org/wikipedia/commons/f/f9/3D_model_hydrogen_bonds_in_water.jpghttp://upload.wikimedia.org/wikipedia/commons/f/f9/3D_model_hydrogen_bonds_in_water.jpg

Conjugate Acid-Base Pairs

HCl + H2O H3O+ + Cl-

acid base

base acid

conjugates

conjugates


acid base CA CB


NH3 + H2O NH41+ + OH-

base acid

acid base

conjugates

conjugates

base acid CA CB

NH3 + H2O NH41+ + OH-

Water is Amphoteric

base acid CA CB

NH3 + H2O NH41+ + OH-


acid base CA CB

Amphoteric or Amphiprotic substances:

Substances which can act as either proton donors (acids) or

proton acceptors (bases) depending on what substances are present.

Amphoteric

1-

+ +

sulfuric acid

H2SO4water

H2O

hydrogen sulfate

ion

HSO4-

hydronium ion

A substance that can act as either an acid or a base.

H3O+

1+

1-

+ +

sulfate ion

SO42-

water

H2O

hydrogen sulfate

ion

HSO4-

hydroxide ion

OH-

1-

2-

1-

+ +

1+

sulfuric acid

H2SO4

water

H2O

hydrogen sulfate

ion

HSO4-

hydronium ion

H3O+

(HSO4- as a base)

Amphoteric


Amphoteric


1-

+

hydrogen sulfate

ion

HSO4-

hydroxide ion

OH-

1-

+

sulfate ion

SO42-

water

H2O

2-

(HSO4- as an acid)


HC2H3O2 + H2O H3O1+ + C2H3O2

-

acid1 base1

base2 acid2

conjugates

conjugates

acid base CA CB

HC2H3O2 + H2O H3O1+ + C2H3O2

-

The reaction proceeds in the direction such that the stronger acid

donates its proton to the stronger base.

Litmus Paper

Copyright © 2007 Pearson Benjamin Cummings. All rights reserved.

pH Paper

pH 0 1 2 3 4 5 6

pH 7 8 9 10 11 12 13

Desired Features of Sensors

pH paper

1904

Detection limit

Low deflection

High sensitivity

High selectivity

Wide dynamic

range

Simple to use

Cost-effective

pH 0 1 2 3 4 5 6

pH 7 8 9 10 11 12 13

Range and Color Changes of Some

Common Acid-Base Indicators

Indicators

pH Scale

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

Methyl orange red 3.1 – 4.4 yellow

Methyl red red 4.4 6.2 yellow

Bromthymol blue yellow 6.2 7.6 blue

Neutral red red 6.8 8.0 yellow

Phenolphthalein colorless 8.0 10.0 red colorless beyond 13.0

Bromthymol blue indicator would be used in titrating a strong acid with a strong base.

Phenolpthalein indicator would be used in titrating a weak acid with a strong base.

Methyl orange indicator would be used in titrating a strong acid with a weak base.


pH

Indicator Acid color Transition color Base color

Litmus

Bromthymol blue

STRONG ACID – STRONG BASE

pH

2 3 4 5 6 7 8 9 10 11 12

INDICATOR COLORS IN TITRATION

2 3 4 5 6 7 8 9 10 11 12


Phenolphthalein

Phenol red

WEAK ACID – STRONG BASE

pH


2 3 4 5 6 7 8 9 10 11 12


Methyl orange

Bromphenol blue

STRONG ACID – WEAK BASE

pH


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

Phenolphthalein

Methyl Red

Orange IV

Colorless Pink Red

Red Orange Yellow

Orange Peach Yellow

pH

phenolphthalein methyl red methyl orange

http://upload.wikimedia.org/wikipedia/commons/5/54/Methyl-orange-2D-skeletal.pnghttp://upload.wikimedia.org/wikipedia/commons/d/d1/Methyl-orange-3D-vdW.pnghttp://upload.wikimedia.org/wikipedia/commons/3/38/Methyl_red.pnghttp://upload.wikimedia.org/wikipedia/commons/5/50/Phenolphthalein.pnghttp://upload.wikimedia.org/wikipedia/commons/f/f8/Phenolphthalein-at-pH-9.jpghttp://upload.wikimedia.org/wikipedia/commons/1/11/Phenolphthalein-in-conc-sulfuric-acid.jpg

Common pH Indicators


Edible Acid-Base IndicatorsCOLOR CHANGES AS A FUNCTION OF pH

INDICATOR pH 2 3 4 5 6 7 8 9 10 11 12

RED APPLE SKIN

BEETS

BLUEBERRIES

RED CABBAGE

CHERRIES

GRAPE JUICE

RED ONION

YELLOW ONION

PEACH SKIN

PEAR SKIN

PLUM SKIN

RADISH SKIN

RHUBARB SKIN

TOMATO

TURNIP SKIN

*

*YELLOW at pH 12 and above

Red Cabbage IndicatorCopyright © 2007 Pearson Benjamin Cummings. All rights reserved.

http://images.google.com/imgres?imgurl=http://www.funsci.com/fun3_en/acids/acids_01.jpg&imgrefurl=http://www.funsci.com/fun3_en/acids/acids.htm&h=294&w=304&sz=17&hl=en&start=1&tbnid=PW71Yfkt-yvLCM:&tbnh=112&tbnw=116&prev=/images%3Fq%3Dacids%26gbv%3D2%26hl%25

Phenolphthalein Indicator

Colorless = Acidic pH

Pink = Basic pH

H+

-OO

C

C

O

O-

(Colorless acid form, HIn) (Pink base form, In-)

OH

OH

HO

C

C

O

O-

Aspirin Synthesis

Preparation of an Ester Acetylsalicylic Acid (Aspirin)

OBJECTIVE: To become familiar with the techniques and principle of esterification. DISCUSSION:

Aspirin is a drug widely used as an antipyretic agent (to reduce fever), as an analgesic agent (to reduce pain), and/or as an anti-inflammatory agent (to reduce redness, heat or swelling in tissues). Chemically, aspirin is an ester. Esters are the products of reaction of acids with alcohols, as shown in the following equation using type formulas:

R – C – OH + R’ – OH R – C – O – R’ + H2O ACID ALCOHOL ESTER WATER

The symbol R refers to the hydrocarbon portion (radical) of the molecules aside from the O functional group. In an organic acid, R – C – OH, the functional group is the carboxyl O group (-COOH) or –C-OH. The type of formula for an alcohol is R-OH, where the functional group is the hydroxyl group (-OH). The symbol R’ indicates that the two R-groups in the ester formula need not be the same. It has been shown by radioactive tracer methods that in the mechanism of the esterification reaction, the –OH group is split from the acid and the –H from the alcohol.

Aspirin can be made as follows:

C – OH C – OH CH3 – C – OH + HO – CH3 – C – O– + H2O Acetic acid Salicylic acid Aspirin (containing an (acetylsalicylic acid, -OH group) an ester)

The use of acetic anhydride instead of acetic acid, however, is a better preparative method, because the anhydride with the water to form acetic acid tends to drive the reaction to the right as shown below. An acid catalyst also is used to speed up the reaction.

C – OH C – OH + HO – CH3 – C – O– + CH3 – C – OH

Acetic anhydride Salicylic acid Aspirin Acetic acid

(138.12 g/mol) (180.15 g/mol)

O C

CH3

CH3

C O

O

O O

14.55 mL

23

24

How to read a buret volume

23.45 mL

(not 24.55 mL)

24.55 mL?

Titration

• Titration

• Analytical method in which a standard solution is used to determine the concentration of an unknown solution.

standard solution

unknown solutionCourtesy Christy Johannesson www.nisd.net/communicationsarts/pages/chem

Titration

• Equivalence point (endpoint)

• Point at which equal amounts of H3O+

and OH- have been added.• Determined by…

• indicator color change

• dramatic change in pH

Courtesy Christy Johannesson www.nisd.net/communicationsarts/pages/chem

Titration

? M of HCl 30.0 mL of 2.0 M of NaOH

If it requires 10.5 mL of ? M HCl to titrate 30.0 mL of 2.0 M NaOH to its endpoint:

what is the concentration of the HCl?

M1V1 = M2V2

M V = M VH+ H+ OH- OH- HCl(aq) → H+(aq) + Cl-(aq)

0.1 M 0.1 M 0.1 M

H2SO4(aq)→ 2 H+(aq) + SO4

2-(aq)

0.1 M “0.2 M” 0.1 M

proper term is Normality (N)

M V n = M V nH+ H+ OH- OH-

Al(OH)3(aq) → Al3+(aq) + 3 OH-(aq)

10.5 mL

HCl must be ~ __x

more concentrated

than the NaOH.

6

(x M)(10.5 mL) = (2.0 M)(30.0 mL)

X = 5.7 M

30.0 mL of NaOH with bromthymol blue indicator

muriatic acid

sunnyside

0.1 molar H2SO4 is 0.2 normal

10.5 mL of HCl

Endpoint of titration is reached…color change.

Titration

moles H3O+ = moles OH-

MVn = MVn

M: Molarity

V: volume

n: # of H+ ions in the acidor OH- ions in the base


Solution

of NaOH

Solution

of KOH

Solution

of H2SO4

50.0 mL

Titration

42.5 mL of 1.3M KOH are required to neutralize 50.0 mL of H2SO4

Find the molarity of H2SO4.

H3O+

M = ?

V = 50.0 mL

n = 2

OH-

M = 1.3M

V = 42.5 mL

n = 1

MV# = MV#

M(50.0mL)(2)

=(1.3M)(42.5mL)(1)

M = 0.55M H2SO4


Acid-Base Titration

Calibration Curve

Acid (mL)

Base (

mL)

0.10 M HCl ? M NaOH

0.00 mL

1.00 mL

2.00 mL

4.00 mL

9.00 mL

17.00 mL

27.00 mL

42.00 mL

1.00 mL

1.00 mL

2.00 mL

5.00 mL

8.00 mL

10.0 mL

15.0 mL

1) Create calibration curve of six data points

2) Using [HCl], determine concentration of NH33) Determine vinegar concentration using [NaOH]

determined earlier in lab

Solution

of NaOH

Solution

of NaOH

Solution

of HCl

5 mL

Data Table

Titration Curve

indicator -changes color

to indicate pH change

e.g. phenolpthalein is colorless in acid

and pink in basic solution

Pirate…”Walk the plank”

once in water, shark eats and

water changes to pink color

pH

endpoint

equivalence

point

base

7

pink

Titration

Calibration Curve

Acid (mL)

Base (

mL)

pH

endpoint

equivalence

point

indicator

base

7

pink

- changes color to indicate pH change

e.g. phenolphthalein is colorless in acid

and pink in basic solution

Pirate…”Walk the plank”

once in water, shark eats and

water changes to pink color

Titration Curve


equivalence point

14.0

12.0

10.0

8.0

6.0

4.0

2.0

0.00.0 10.0 20.0 30.0 40.0

pH

Volume of 0.100 M NaOH added

(mL)

Titration of a Strong Acid With a Strong Base

Solution

of NaOH

Solution

of NaOH

Solution

of HClH+

H+H+

H+

Cl

Cl-

Cl-

Cl-

Na+

Na+

Na+

Na+

OH-

OH-OH-

OH-

Acid-Base Titrations

Adding NaOH from the buret to hydrochloric acid in the flask,

a strong acid. In the beginning the pH increases very slowly.

Adding additional NaOH is added. pH rises as

the equivalence point is approached.

Additional NaOH is added. pH increases and then levels off as

NaOH is added beyond the equivalence point.

equivalence point

14.0

12.0

10.0

8.0

6.0

4.0

2.0

0.00.0 10.0 20.0 30.0 40.0

pH


(mL)


0.00 1.00

10.00 1.37

20.00 1.95

22.00 2.19

24.00 2.70

25.00 7.00

26.00 11.30

28.00 11.75

30.00 11.96

40.00 12.36

50.00 12.52

NaOH added

(mL) pH

Titration Data

Solution

of NaOH

Solution

of NaOH

Solution

of HClH+

H+H+

H+

Cl-

Cl-

Cl-

Cl-

Na+

Na+

Na+

Na+

OH-

OH-OH-

OH-

25 mL

Bromthymol blue is best indicator: pH change 6.0 - 7.6

Yellow Blue


equivalence

point

14.0

12.0

10.0

8.0

6.0

4.0

2.0

0.00.0 10.0 20.0 30.0

pH


(mL)

Color change

methyl violet

Color change

bromphenol blue

Color change

bromthymol blue

Color change

phenolpthalein

Color change

alizarin yellow R

(20.00 mL of 0.500 M HCl by 0.500 M NaOH)

Hill, Petrucci, General Chemistry An Integrated Approach 2nd Edition, page 680

equivalence point

14.0

12.0

10.0

8.0

6.0

4.0

2.0

0.00.0 10.0 20.0 30.0 40.0

pH


(mL)

Titration of a Weak Acid With

a Strong Base

0.00 2.89

5.00 4.14

10.00 4.57

12.50 4.74

15.00 4.92

20.00 5.35

24.00 6.12

25.00 8.72

26.00 11.30

30.00 11.96

40.00 12.36

NaOH added

(mL) pH

Titration Data

Titration of a Weak Acid With a Strong Base

Phenolphthalein is best indicator: pH change 8.0 - 9.6

equivalence point

14.0

12.0

10.0

8.0

6.0

4.0

2.0

0.00.0 10.0 20.0 30.0 40.0

pH

Volume of 0.100 M HCl added

(mL)

Titration of a Weak Base With a Strong Acid

0.00 11.24

10.00 9.91

20.00 9.47

30.00 8.93

40.00 8.61

45.00 8.30

47.00 7.92

48.00 7.70

49.00 7.47

50.00 5.85

51.00 3.34

HCl added

(mL) pH

Titration Data

Titration of a Weak Base With a Strong Acid

50.0

7. What is the pH of a solution made by dissolving 2.5 g NaOH in 400 mL water?

Determine number of moles of NaOH

x mol NaOH = 2.5 g NaOH =

NaOH g 40

NaOH mol 10.0625 mol NaOH

Calculate the molarity of the solution

L

mol M =

L 0.4

NaOH mol 0.0625 [Recall 1000 mL = 1 L]

MNaOH = 0.15625 molar

NaOH Na1+ + OH1-

0.15625 molar 0.15625 molar0.15625 molar

pOH = -log [OH-]

pOH = -log [0.15625 M]

pOH = 0.8

pOH + pH = 14

or kW = [H+] [OH-]

1 x 10-14 = [H+] [0.15625 M]

[H+] = 6.4 x 10-14 M

pH = -log [H+]

pH = 13.2 pH = -log [6.4 x 10-14 M]0.8 + pH = 14

What volume of 0.5 M HCl is required to titrate 100 mL of 3.0 M Ca(OH)2?

x = 600 mL of 0.5 M HCl

HCl H1+ + Cl1-

0.3 mol 0.3 mol0.3 mol

HCl + Ca(OH)2 CaCl2 + HOH 22

x mL

0.5 M

100 mL

3.0 M

M1V1 = M2V2(0.5 M) (x mL) = (3.0 M) (100 mL)

x = 1200 mL of 0.5 M HCl

M1V1 = M2V2(0.5 M) (x mL) = (6.0 M) (100 mL)

Ca(OH)2 Ca2+ + 2OH1-

0.3 mol 0.6 mol0.3 mol

M

mol

L

HCl

molHCl = M x L

mol = (0.5 M)(0.6 L)

mol = 0.3 mol HCl

Ca(OH)2

mol = (3.0 M)(0.1 L)

mol = 0.3 mol Ca(OH)2

mol = M x LCa(OH)2

[H+] = [OH-]

"6.0 M"

6. 10.0 grams vinegar

M

mol

L

NaOH

molNaOH = M x L

mol = (0.150 M)(0.0654 L)

mol = 0.00981 mol NaOH

titrated with 65.40 mL of 0.150 M NaOH

(acetic acid + water)

moles NaOHmoles HC2H3O2 =

therefore, you have ...

0.00981 mol HC2H3O2

B)

A)

x g HC2H3O2 = 0.00981 mol HC2H3O2 =

232

232

OHHC mol 1

OHHC g 600.59 g HC2H3O2

C) % = 100% x whole

part

% = 100% x vinegar g 10.0

acidacetic g 0.59

% = 5.9 % acetic acid

Commercial vinegar is sold as 3 - 5 % acetic acid

H

OC

O

C

H

H

H

H

H

O

49 mL 0.2 M HCl + 50 mL 0.2 M NaOH

A) molHCl = M . L

molHCl = (0.2 M) . (0.049 L)

molHCl = 0.0098 mol

B) molNaOH = M . L

molNaOH = (0.2 M) . (0.05 L)

molNaOH = 0.010 mol

49 mL

0.2 M HCl50 mL

0.2 M NaOH

99 mL H2O

1 mL of 0.2 M NaOH

0.010 mol OH1-

0.0098 mol H1+-

0.0002 mol OH1-“net”

1) What is the pH of a solution made by combining 49 mL of 0.2 M HCl

with 50 mL of 0.2 M NaOH?

1) What is the pH of a solution made by adding 1mL

of 0.2 M NaOH with 99 mL H2O?

HCl + NaOH H2O + NaCl

M

mol

L

49 mL

0.2 M HCl50 mL

0.2 M NaOH

99 mL H2O

1 mL of 0.2 M NaOH

0.010 mol OH1-

0.0098 mol H1+-

0.0002 mol OH1-“net”

1) What is the pH of a solution made by adding

1mL of 0.2 M NaOH with 99 mL H2O?

NaOH → Na1+ + OH1-

Calculate the molarity of the solution

[Recall 1000 mL = 1 L]

MNaOH = 0.002020 molar

NaOH Na1+ + OH1-

0.002020 molar 0.002020 molar0.002020 molar

pOH = -log [OH-]

pOH = -log [0.002020 M]

pOH = 2.7

pOH + pH = 14

or kW = [H+] [OH-]

1 x 10-14 = [H+] [0.002020 M]

[H+] = 4.95 x 10-12 M

pH = -log [H+]

pH = 11.3 pH = -log [4.95 x 10-12 M]2.7 + pH = 14

M = mol

LM =

0.0002 mol NaOH

0.0099 L

Carboxylic Acid

HC2H3O2

CH3COOH

C2H4O2

R - COOH

H C C

H

H

O

O

H

carboxylic acid

H+

= acetic acid

1-

Lactic Acid

H3C C CO2H

H

OH

Lactic acidC3H6O3

http://upload.wikimedia.org/wikipedia/commons/5/59/Lactic-acid-3D-balls.png

Titration

? M NaOH1.0 M HCl titrate with

1.00 mL 2.00 mL

M1 V1 = M2 V2(1.0 M)(1.00 mL) = (x M)(2.00 mL)

X = 0.5 M NaOH

? M NaOH1.0 M H2SO4 titrate with

1.00 mL 2.00 mL

M1 V1 = M2 V2(1.0 M)(1.00 mL) = (x M)(2.00 mL)

X = 0.5 M NaOH

2.0 M H1+

?

Calibration Curve

Vinegar Ammonia

1 mL

3 mL

5 mL

10 mL

15 mL

Acid Basevinegar ammonia

vinegar

am

mo

nia

Using 3 mL vinegar… titrate with 0.130 M NaOH solution.

M

mol

L

NaOH

molNaOH = M x L

mol = (0.130 M)(0.0196 L)

mol = 0.002548 mol NaOH

moles NaOHmoles HC2H3O2 =

therefore, you have ...

0.002548 mol HC2H3O2

B)

A)

x g HC2H3O2 = 0.002548 mol HC2H3O2 =

232

232

OHHC mol 1

OHHC g 600.153 g HC2H3O2

C) % = 100% x whole

part

% = 100% x vinegar g 3.0

acidacetic g 0.1529

% = 5.1 % acetic acid Commercial vinegar is sold

as 3 - 5 % acetic acid

Calculate molarity (M) of acetic acid. M1V1 = M2V2Calculate % acetic acid in vinegar. % = part/whole x100

http://upload.wikimedia.org/wikipedia/commons/2/2f/Red_pog2.svghttp://upload.wikimedia.org/wikipedia/commons/2/2f/Red_pog2.svghttp://upload.wikimedia.org/wikipedia/commons/2/2f/Red_pog2.svghttp://upload.wikimedia.org/wikipedia/commons/2/2f/Red_pog2.svghttp://upload.wikimedia.org/wikipedia/commons/2/2f/Red_pog2.svg

Calibration Curve

Vinegar Ammonia

1 mL

3 mL

5 mL

10 mL

15 mL

vinegar

am

mo

nia

Using 3 mL vinegar… titrate with 0.130 M NaOH solution.

Calculate molarity (M) of acetic acid. M1V1 = M2V2

M1 V1 = M2 V2

(Macetic acid)(3.0 mL) = (0.130 MNaOH )(19.6 mL)

Macetic acid = 0.8493 molar

It required 19.6 mL of NaOH to reach the endpoint.

Acid Basevinegar ammonia

http://upload.wikimedia.org/wikipedia/commons/2/2f/Red_pog2.svghttp://upload.wikimedia.org/wikipedia/commons/2/2f/Red_pog2.svghttp://upload.wikimedia.org/wikipedia/commons/2/2f/Red_pog2.svghttp://upload.wikimedia.org/wikipedia/commons/2/2f/Red_pog2.svghttp://upload.wikimedia.org/wikipedia/commons/2/2f/Red_pog2.svg

HCl

H2SO4

H3PO4

HNO3

CH3COOH

HF

Hydrochloric acid

Sulfuric acid

Phosphoric acid

Nitric acid

Acetic acid

Hydrofluoric acid

stomach acid, pickling metal

battery acid, # 1 selling chemical

food flavoring

fertilizer, explosives

vinegar

etch glass

NaOH Ca(OH)2 NH4OH

sodium hydroxide calcium hydroxide ammonium hydroxide

pH scale0

7

14

acid

neutral

base

[H+] = [OH-]

Soren Sorenson developed pH scale

pOH = -log [OH-]

kW = [H+] [OH-]

pH = -log [H+]

pH + pOH = 14

(alkalinity)

Arnold Beckman invented the pH meter

H+ + H2O H3O+

proton hydronium ion

kw = 1 x 10-14

Concentrated vs. Dilute

Concentration: Molarity molality Normality

M = mol

L m = mol

kg

H2SO4 2 H1+ + SO4

2-

3 M “6 M”

Strong / Weak Acid

Strong HA H+ + A- (~100% dissociation)

Weak HA H+ + A- (~20% dissociation)

Ka = [Product]

[Reactant]

acid dissociation constant

Ka

0.8 H3PO4 3H+ + PO4

3-

0.0021 HF H+ + F-

H2A 2 H+ + A-

Ka = [H+]2 [A-]

[H2A]

Acid + Base Salt + Water

How would you make calcium sulfate in the lab?

+ CaSO4

ACID

Sour taste, litmus red

Arrhenius – H+ as only ion in water

Brønsted-Lowry – proton donor

BASE

bitter taste, litmus blue

Arrhenius – OH- as only ion in water

Brønsted-Lowry – proton acceptor

H2SO4 Ca(OH)2 + 2 H2O? ?

phenolphthalein colorless pink

acid baseweak strong

bromthymol blue yellow blue

acid basestrong strong

universal indicator R O Y G B I V

pH 4 7 12

litmus paper & pH paper

Indicators

Buffers - salts of weak acids and weak bases that maintain a pH

LeChatelier’s Principle

- acidosis & alkalosis (bicarbonate ion acts as buffer)

- darkening glasses

- egg shells thinner in summer (warm)

e.g. Aspirin (acetyl salicylic acid) vs. Bufferin

low pH upsets stomach

Acid – Base Concentrations

pH = 3

pH = 7

pH = 11

OH-

H3O+OH-

OH-H3O+

H3O+

[H3O+] = [OH-][H3O

+] > [OH-] [H3O+] < [OH-]

acidic

solution

neutral

solutionbasic

solution

co

nc

en

tra

tio

n (

mo

les

/L)

10-14

10-7

10-1


Amino Acids – Functional Groups

Amine Carboxylic AcidBase Pair

NH21- R- COOH

NH3NH21- NH4

1+

amine ammonia ammonium ion

NH

HH

:

NH

HH

H :

1+

NH

H

: 1-

H+lose H+


Amine Carboxylic AcidBase Pair

NH21- R- COOH

O2-H+

H+ d−d+

http://fig.cox.miami.edu/~cmallery/150/chemistry/sf3x17b.jpg


Water – Amphiprotic

H2OOH1- H3O

1+

hydroxide water hydronium ion

NH

HH

:

NH

HH

H :

1+

NH

H

:

1-

H+lose H+

Water – Also Amphiprotic

H2OOH1- H3O

1+


H+lose H+

Amphiprotic – Act as an acid (proton donor) or base (proton acceptor)

O2-H+

H+ d−d+


H2OOH1- H3O

1+


H+lose H+

Amphiprotic – Act as an acid (proton donor) or base (proton acceptor)

O2-H+

H+ d−d+



Range and Color Changes of Some

Common Acid-Base Indicators

Indicators

pH Scale

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

Methyl orange red 3.1 – 4.4 yellow

Methyl red red 4.4 6.2 yellow

Bromthymol blue yellow 6.2 7.6 blue

Neutral red red 6.8 8.0 yellow

Phenolphthalein colorless 8.0 10.0 red colorless beyond 13.0

pH Paper

pH 0 1 2 3 4 5 6

pH 7 8 9 10 11 12 13

pH Paper

pH 0 1 2 3 4 5 6

pH 7 8 9 10 11 12 13

pH Paper

pH 0 1 2 3 4 5 6

pH 7 8 9 10 11 12 13

pH Paper

pH 0 1 2 3 4 5 6

pH 7 8 9 10 11 12 13

Neutralization of Bug Bites

Wasp - stings with base

(neutralize with lemon juice or vinegar)

Red Ant - bites with acid

(neutralize with baking soda)

Strength

• Strong Acid/Base• 100% ionized in water• strong electrolyte

- +

HCl

HNO3

H2SO4

HBr

HI

HClO4

NaOH

KOH

Ca(OH)2

Ba(OH)2


Strength

◆Weak Acid/Base

• does not ionize completely

• weak electrolyte

- +

HF

CH3COOH

H3PO4

H2CO3

HCN

NH3


Ionization of Water

H2O + H2O H3O+ + OH-

Kw = [H3O+][OH-] = 1.0 10-14


Why is pure water pH = 7?

• 1 in 500,000,000 water molecules will autoionize.

• H2O + H2O → H3O+ + OH1-

• This yields a hydronium ion concentration of 1 x 10-7 M H3O+ per liter of

solution

• pH = -log[H3O+]

• pH = -log[1 x 10-7] or pH = 7

H

OH

H

HO

H1-

1+

HO

H

H

O H

H

OH

H O

H

O

H

H

H

O

H

Ionization of Water

• Find the hydroxide ion concentration of 3.0 10-2 M HCl.

[H3O+][OH-] = 1.0 10-14

[3.0 10-2][OH-] = 1.0 10-14

[OH-] = 3.3 10-13 M

Acidic or basic? Acidic


pH = -log[H3O+]

pH Scale

0

7INCREASING

ACIDITYNEUTRAL

INCREASING

BASICITY

14

pouvoir hydrogène (Fr.)

“hydrogen power”Courtesy Christy Johannesson www.nisd.net/communicationsarts/pages/chem

pH Scale

pH = -log[H3O+]

pOH = -log[OH-]

pH + pOH = 14


pH Scale

• What is the pH of 0.050 M HNO3?

pH = -log[H3O+]

pH = -log[0.050]

pH = 1.3

Acidic or basic? Acidic


pH Scale

• What is the molarity of HBr in a solution that has a pOH of 9.6?

pH + pOH = 14

pH + 9.6 = 14

pH = 4.4

Acidic

pH = -log[H3O+]

4.4 = -log[H3O+]

-4.4 = log[H3O+]

[H3O+] = 4.0 10-5 M HBr


Resources - Acids and Bases

Objectives

Worksheet - vocabulary

Worksheet - pH and pOH calculations

Worksheet - practice problems (key)

Textbook - text ?'s chemical equilibrium

Worksheet - weak acid, pKa

Article - aspirin

Lab - synthesis of aspirin

Worksheet - aqueous acids and bases titration

Outline (general)

Video (VHS) - future of the past

Outline -

Textbook - Ch 15 Modern Chemistry

Worksheet -

Worksheet -

Worksheet -

Lab - titration

Textbook - questions

Episode 16 – The Proton in Chemistry
../../Objectives - Chemistry/AcidBaseTO.doc../Acid Word/1vocabacidbase.doc../Acid Word/1phpohws.doc../Acid Word/1Practice Problems.doc../Acid Word/1eqtextqs.doc../LABS & Activities/Aspirin Synthesis/Aspirin synthesis.doc../LABS & Activities/Ester Identification Activity/Salicylic Acid Derivatives.doc../Acid Word/1titration.doc../Outlines/u11ohnotes18f2005.doc../Outlines/Student Notes/u11lectout.doc../Acid Word/1futurepast.doc../Acid Word/abtextqs[1].doc../LABS & Activities/LABS/Acid Base Titration.doc../Outlines/Text Notes - General/u11textnotes.doc../Video/World of Chemistry Notes/Episode 16 - The Proton in Chemistry.dochttp://www.backflip.com/perl/go.pl?url=18424641

Documents

· Acids, Bases, and Salts • You should be able to • Understand the acid-base theories of Arrhenius, Brønsted-Lowry and Lewis • Identify strong acids and bases and calculate