Chemical reactions · Evidence for chemical reactions. The chemical equation ... atoms on left...

Chemical reactions

ClassificationsReactions in solution

Ionic equations

Learning objectives

Distinguish between chemical and physical change

Write and balance chemical equations

Describe concepts of oxidation and reduction

Classify reaction according to types of reactants and products

Distinguish among strong, weak and non-electrolytes

Identify common acids and bases by from chemical formula

Predict formation of precipitates by application of solubility rules

Write total and net ionic equations from balanced molecular equations

Chemical vs physical redux

Physical: No new substance!

Chemical: New substance formed!

Evidence for chemical reactions

The chemical equation

aA + bB = cC + dD

Law of Conservation of Matter states that matter is

neither created nor destroyed

Means: atoms on left equals atoms on right

Reactant

side

Product

sidecoefficientELEMENT or

COMPOUND

Chemical book-keeping

Keys to balancing equations:

“Have I gained or lost any atoms?”

Put down the correct formula for each

reactant or product

Formulas cannot be changed in order to

balance the equation

Reaction of hydrogen with oxygen to produce

water: reactants are H2 and O2, product is H2O

Count the atoms: 4 H and 2 O 4 H and 2 O

The big number

multiplies every

atom after it

The subscript

only multiplies

the atom before it

2 H2 + O2 → 2 H2O

Molecular representation of the

reaction

Balance the equationsA method of trial and error

CH4 + O2 = CO2 + H2O

Balance the equations

CH4 + O2 = CO2 + H2O– CH4 + 2O2 = CO2 + 2H2O

C3H8 + O2 = CO2 + H2O– C3H8 + 5O2 = 3CO2 + 4H2O

N2 + H2 = NH3

– N2 + 3H2 = 2NH3

Do balancing equation exercises

One approach to classification

Oxidation – reduction: focusing on

electrons

Oxidation is loss of electrons

Reduction is gain of electrons

Oxidation is always accompanied by

reductionThe total number of electrons is kept constant

Oxidizing agents oxidize and are

themselves reduced

Reducing agents reduce and are

themselves oxidized

Redox in chemistry

All reactions involve rearrangement of

atoms and molecules

Some reactions involve rearrangement of

atoms and molecules and electrons

– Photosynthesis, respiration, combustion...

These are called redox reactions

Any reaction involving elements must be

redox

Combination (synthesis)

reactionsElement + element compound (redox)– S + O2 → SO2

– Metal + nonmetal binary ionic compound

– Nonmetal + nonmetal binary covalent compound

Compound + element compound (redox)– CO + O2 → CO2

Compound + compound compound– SO2 + H2O →H2SO3

Decomposition reactions

Compound element +

element (redox)

– HgO → Hg + O2

Compound element +

compound (redox)

– PCl5 → PCl3 + Cl2

Compound compound

+ compound

– CaCO3 → CaO + CO2

Single replacement (displacement)

Element displaces another

element from compound

(redox)

Cu + 2 AgNO3 → Cu(NO3)2 + 2 Ag

Predicting single replacement

reactions: the activity series

Element higher in the

will displace one

lower in the series

The element higher is

a stronger reducing

agent

The element lower is

a stronger oxidizing

agent

Three types of double

displacement reaction

Compounds

exchanging partners

– Usually ionic

compounds in solution

Precipitation

Acid-base

neutralization

Gas formation

Precipitation

Identify ions and swap them

BaCl2 + Na2SO4 → BaSO4 + 2 NaCl

Acid – base neutralization:

special case of double replacement

KOH(aq) + HNO3(aq) = KNO3(aq) + H2O(l)

Product is liquid water not a solid

BASE ACID SALT WATER

Gas formationProduct is either a gas or is unstable and

decomposes to a gas

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

Writing balanced molecular equations

for double replacement reactions

Use correct formulae

– Metal ion charge predicted

from group number

– Use table for correct

formula and charge for

polyatomic ions

Identify as solid (s), gas

(g), liquid (l) or dissolved

(aq)

Balance: atoms (groups)

on left = atoms (groups)

on right

Balancing double replacement equations

It’s very much a matter of states – show

them!

Pb(NO3)2(aq) + 2KI(aq) = 2KNO3(aq) + PbI2(s)

Balance polyatomic ions as units:

– Pb2+, K+, I-, NO3-

Left hand side Right hand side1 Pb2+ 1 Pb2+

2 NO3- 2 NO3

-

2 K+ 2 K+

2 I- 2 I-

Molecular equation for reaction of

Na2SO4 + Ba(NO3)2

Combustion

Element or compound

reacting with oxygen

(redox)

– CH4 + O2 → CO2 + H2O

Associated with

production of heat and

light

Involves hydrocarbons

(CxHy), nonmetals (S) and

metals (Mg)

Sorting solution reactions:

dissolved species

Electrolytes:

– Ionic compounds produce ions in solution

(NaCl, NH4NO3 etc.)

Non-electrolytes:

– Covalent compounds do not produce ions in

solution (CH3OH, C6H12O6 etc.)

Electrolytes: distinguishing by

strength

All soluble substances that produce ions are

electrolytes

Strong electrolytes are characterized by

complete dissociation in water

Weak electrolytes dissociate to a much smaller

extent.

Strong, weak or non electrolyte?

All soluble salts are strong electrolytes

Strong acids and bases are strong

electrolytes

Weak acids and bases are weak

electrolytes

Insoluble compounds are non-electrolytes

Molecular compounds are non-electrolytes

Classification of electrolytes

Strong

electrolytes

Weak

electrolytes

Non-

electrolytes

ACIDS:

HCl, HBr, HI

HClO4, HNO3, H2SO4

ACIDS:

HF, H3PO4,

CH3CO2H

Molecular

covalent

compounds:

H2O,

CH3OH,

C12H22O11

(sucrose)

Most organic

compounds

and

INSOLUBLE

salts

SALTS:

KBr, Na3PO4

SALTS:

None

BASES:

NaOH, Ba(OH)2

BASES:

NH3

Flow chart for determining type of

electrolyte

3. Ionic or covalent?

1. Soluble in H2O?

2. Acid or base?

Nonelectrolyte

3. Weak or strong?

Strong

electrolyteWeak

electrolyte

Yes No

NoYes

Weak CovStrong Ionic

Recognizing acids and basesAcids usually have H at the beginning of the

formula – HCl

Bases usually have OH in the formula – NaOH

– Not in organic compounds though - CH3OH

Acid formula Name Base formula Name

HCl Hydrochloric acid NaOH Sodium hydroxide

H2SO4 Sulfuric acid KOH Potassium

hydroxide

H3PO4 Phosphoric acid Ba(OH)2 Barium hydroxide

HNO3 Nitric acid NH3 Ammonia

HClO4 Perchloric acid (CH3)3N Trimethylamine

CH3CO2H Acetic acid

HCO2H Formic acid

Citric acid

The strong acids and bases

Strong acids

(Only six)

Strong bases

(g1A and g2A)

HCl Hydrochloric acid LiOH Lithium hydroxide

HBr Hydrobromic acid NaOH Sodium hydroxide

HI Hydroiodic acid KOH Potassium

hydroxide

HNO3 Nitric acid Ca(OH)2 Calcium

hydroxide

H2SO4 Sulfuric acid Sr(OH)2 Strontium

hydroxide

HClO4 Perchloric acid Ba(OH)2 Barium hydroxide

Solubility roolsGroup I and ammonium (NH4

+) compounds soluble

Nitrates (NO3-), acetates (CH3CO2

-) soluble

Chlorides, bromides and iodides generally soluble

{except Pb(II), Ag(I) and Hg(I)}

Sulphates (SO42-) generally soluble (except g2A and

Pb2+)

Carbonates (CO32-), phosphates (PO4

3-) generally

insoluble (except gIA)

Hydroxides (OH-), sulphides (S2-) generally insoluble

(except gIA and gIIA)

Total ionic equations

Pb(NO3)2(aq) + K2CrO4(aq) = 2KNO3(aq) + PbCrO4(s)

Total ionic equation

Dissolved substances:

– Strong electrolytes show as ions

– Weak or non- electrolytes show as molecular

formula

Pb2+(aq) + 2NO3-(aq) + 2K+(aq) + CrO4

2-(aq) =

2K+(aq) + 2NO3-(aq) + PbCrO4(s)

Net ionic equations

Spectator ions are those ions that do not

undergo a change; they do not participate

in the chemical change and are the same

on both sides of the equation

Remove all spectator ions from the

equationPb2+(aq) + 2NO3

-(aq) + 2K+(aq) + CrO42-(aq) =

2K+(aq) + 2NO3-(aq) + PbCrO4(s)

Net ionic equations

Pb2+(aq) + CrO42-(aq) = PbCrO4(s)

Mass and charge must still balance, although overall charge may not be neutral in a net ionic equation

Net ionic equation for reaction of

Na2SO4 + Pb(NO3)2