Chapter 11Chemical Reactions

Section 11. 1Describing Chemical Reactions

All chemical reactions… Have two parts:

Reactants - the substances you start with.

Products - the substances you end up with.

The reactants turn into the products.

Reactants Products

A reaction can be described several ways:

In a word equation (some symbols used)

Copper + chlorine copper (II) chloride

2

To write a word equation, write the names of the reactants to the left of the arrow separated by plus signs and write the names of the products to the right of the arrow separated by plus signs.e.g. Hydrogen peroxide decomposes to form water and oxygen gas. Write the word equation of this reaction.

hydrogen peroxide Water + oxygen(Reactants) (Products)

e.g. the burning of methane (combining with oxygen) produces carbon dioxide and water. Write the word equation of this reaction.

Methane + oxygen carbon dioxide + water(Reactants) (Products)

3

But it is easier to use the formulas for the reactants and products to describe the chemical reactions.

Chemical equation: is a representation of a chemical reaction by using the formulas of the reactants (on the left) followed by an arrow then the formulas of the products (on the right).

4

5

used after a product indicates a solid has

been produced as precipitate: PbI2 ↓

used after a product indicates a gas has

been produced (evolved) : H2 ↑Catalyst: is a substance that speeds up the reaction

but is not used up in the reaction.6

The Skeleton EquationThe Skeleton EquationUses formulas and symbols to describe a reaction but doesn’t indicate the relative amounts of the reactants and products.

All chemical equations are a description that describe reactions.

Write a skeleton equation for:

1. Solid iron (III) sulfide reacts with gaseous hydrogen chloride to form iron (III) chloride and hydrogen sulfide gas.

2. Nitric acid dissolved in water reacts with solid sodium carbonate to form liquid water and carbon dioxide gas and sodium nitrate dissolved in water.

7

Write the word equation of the following:

Fe(s) + O2(g) Fe2O3(s)

Cu(s) + AgNO3(aq) Ag(s) + Cu(NO3)2(aq)

NO2 (g) N2(g) + O2(g)

8

Law of Conservation of MLaw of Conservation of Matteratter

A natural law describing the fact that matter is neither created nor destroyed in any process

The amount of matter that you start with has to equal to the amount of matter that you end with

Atoms can’t be created or destroyed in an ordinary reaction:

All the number of atoms we start with ,

we must end up with

A balanced equation has the same number of each element on both sides of the equation.

9

For Chemical Reactions This Means

• The amount of reactants has to equal the amount of products.

• Matter cannot be created or destroyed through a chemical reaction.

• Chemical equations have to be balanced.

10

Rules for Balancing:Rules for Balancing:

1. Assemble the correct formulas for all the reactants and products, use + and →

2. Count the number of atoms of each type appearing on both sides

3. Balance the elements one at a time by adding coefficients where needed (the numbers in front) - save balancing the H and O until LAST!

4. Check to make sure it is balanced.11

Never change a subscript to balance an equation.

– If you change the formula you are describing a different reaction.

H2O is a different compound than H2O2

Never put a coefficient in the middle of a formula

2NaCl is okay, but Na2Cl is not.

12

Balancing Chemical Equations

Example:

HCl + NaOH NaCl + H2O

H=2 H=2

Cl=1 Cl=1

Na=1 Na=1

O=1 O=1

The equation is balanced because the number of atoms in the reactants are equal to the number of atoms in the products.

13

Balancing Chemical Equations

Example: H2 + O2 H2O

H=2 O=2 H=2 O=1

H2 + O2 2 H2O

H=2 O=2 H=4 O=2

2H2 + O2 2 H2O

H=4 O=2 H=4 O=214

Balancing Chemical Equations

Example:

Cu + AgNO3 Cu(NO3)2 + Ag

Cu=1 Ag=1 N=1 O=3 Cu=1 Ag=1 N=2 O=6

Cu + 2AgNO3 Cu(NO3)2 + Ag

Cu=1 Ag=2 N=2 O=6 Cu=1 Ag=1 N=2 O=6

Cu + 2AgNO3 Cu(NO3)2 + 2Ag

Cu=1 Ag=2 N=2 O=6 Cu=1 Ag=2 N=2 O=6

15

Balancing Chemical Equations NaHCO3 + H3C6H5O7 CO2 + H2O + Na3C6H5O7

Na=1 H=9 C=7 O=10 Na=3 H=7 C=7 O=10

3NaHCO3 + H3C6H5O7 CO2 + H2O + Na3C6H5O7

Na=3 H=11 C=9 O=16 Na=3 H=7 C=7 O=10

3NaHCO3 + H3C6H5O7 3CO2 + H2O + Na3C6H5O7

Na=3 H=11 C=9 O=16 Na=3 H=7 C=9 O=14

3NaHCO3 + H3C6H5O7 3CO2 + 3H2O + Na3C6H5O7

Na=3 H=11 C=9 O=16 Na=3 H=11 C=9 O=1616

Practice Balancing Examples

…AgNO3 + …Cu …Cu(NO3)2 + …Ag

…Mg + …N2 …Mg3N2

…P + …O2 …P4O10

…Na + …H2O …H2 + …NaOH

…CH4 + …O2 …CO2 + …H2O

17

End of Section 11.1

Section 11.2

Types of Chemical Reactions

Types of ReactionsThere are 5 major types of chemical reactions

1. Combination reaction or Synthesis reaction

2. Decomposition reaction

3. Single Replacement reaction

4. Double Replacement reaction

5. Combustion reaction

Not all reactions fit into only one category

Patterns of chemical reactions will help you predict the products of the reaction

20

Combination Reactions• Combine = put together

• 2 substances combine to make one compound.

Combination reaction: is a chemical change in which two or more substances react to form a single new substance.

• Ca +O2 CaO (2 elements form 1 compound)

• SO3 + H2O H2SO4 (2 compounds form another)

• When 2 non metals react (or a transition metal and a non metal) in a combination reaction, often more than one product is possible.

S(s) + O2 (g) SO2 (g)

2S(s) + 3O2 (g) 2SO3 (g)

21

Complete and balance• Ca + Cl2

• Fe + O2

• Al + O2

• Remember that the first step is to write the correct formulas – you can still change the subscripts at this point, but not later!

• Then balance by using the coefficients only

22

#2 - Decomposition Reactions

• decompose = fall apart

• one reactant breaks apart into two or more elements or compounds.

• NaCl Na + Cl2

• CaCO3 CaO + CO2

• Note that energy (heat, sunlight, electricity, etc.) is usually required

electricity

23

• Can predict the products if it is a binary compound-Made up of only two elements

– breaks apart into its elements:

• H2O

• HgO

electricity

H2 + O2

Hg + O2

24

#3 - Single Replacement

• One element replaces another

• Reactants must be an element and a compound.

• Products will be a different element and a different compound.

• Na + KCl No reaction

• F2 + LiCl LiF + Cl2

25

• Metals replace other metals (and they can also replace hydrogen)

• K + AlN • Zn + HCl • Think of water as: HOH

– Metals replace one of the H, and then

combine with the hydroxide.• Na + HOH

26

• We can even tell whether or not a single replacement reaction will happen:– Some chemicals are more “active” than others– More active replaces less active

• There is a list on page 333 - called the Activity Series of Metals

Higher on the list replaces lower

27

The Activity Series of the Metals Lithium Potassium Calcium Sodium Magnesium Aluminum Zinc Chromium Iron Nickel Lead HydrogenHydrogen Bismuth Copper Mercury Silver Platinum Gold

• Group 1, 2, & 3 Metals are more active than Hydrogen and any other metals (transition metals).

So Group 1, 2, & 3 Metals can

replace Hydrogen and any other metals (transition metals).

Higher activity

Lower activity

28

Practice:Practice:

• Fe + CaSO4

• Pb + KCl

• Al + HCl

No Reaction

No Reaction

AlCl3 + H2

29

The Activity Series of the Halogens

Fluorine Chlorine Bromine Iodine

Halogens can replace other halogens in compounds, provided that they are above the halogen that they are trying to replace.

2NaCl(s) + F2(g) 2NaF(s) + Cl2(g)

MgCl2(s) + Br2(g) ???No ReactionNo Reaction

???

Higher Activity

Lower Activity

30

#4 - Double Replacement

• Two things replace each other.–Reactants must be two ionic compounds–Usually in aqueous solution

• NaOH + FeCl3

–The positive ions change place.

• NaOH + FeCl3 Fe+3 OH- + Na+1 Cl-1

• NaOH + FeCl3 Fe(OH)3 + NaCl31

Complete and balance:

• assume all of the following reactions actually take place:CaCl2 + NaOH

CuCl2 + K2S

KOH + Fe(NO3)3

(NH4)2SO4 + BaF2

32

Practice Examples:Practice Examples:

H2 + O2

H2O

Zn + H2SO4

HgO

KBr + Cl2

AgNO3 + NaCl

Mg(OH)2 + H2SO3

33

#5 - Combustion• Means “add oxygen”

• Normally, a compound composed of only C, H, (and maybe O) is reacted with oxygen – usually called “burning”

• If the combustion is complete, the products will be CO2 and H2O

• If the combustion is incomplete, the products will be CO (or possibly just C) and H2O.

34

Combustion Examples:

C4H10 + O2

C3H8 + O2

C6H12O6 + O2

C8H8 + O2

35

(assume complete)

SUMMARY: an equation...

• Describes a reaction

• Must be balanced in order to follow the Law of Conservation of Mass

• Can only be balanced by changing the coefficients.

• Has special symbols to indicate physical state, if a catalyst or energy is required, etc.

36

How to Recognize which type:

Look at the reactants:

A + B =AB (Combination)

AB = A + B (Decomposition)

A + BC = AC + B (Single replacement)

AB + CD = AD + CB (Double replacement)

A + O2 = (Combustion)

37

End of Section 11.2

Section 11.3

Reactions in Aqueous Solution

Predicting the formation of a precipitate

Some combination of solutions produce precipitates, while others do not, whether or not a precipitate forms depends upon the solubility of the new compounds that form.

You can predict the formation of a precipitate by using the general rules for solubility of ionic compounds.

These rules are shown in the following table:

40

Solubility Rules for Ionic Compounds

Compounds Solubility

Sodium, potassium, and ammonium salts Soluble

All nitrates and chlorates salts Soluble

All chlorides except silver chloride and lead chloride

Soluble

All sulfates except, silver sulfate, lead sulfate, and barium sulfate

Soluble

All carbonates, phosphates, hydroxides, sulfides and chromates salts except with sodium, potassium and ammonium

Insoluble

Insoluble salt = Precipitate

41

Example:

CaCl2(s) + Pb(NO3)2(aq) PbCl2(s) + Ca(NO3)2(aq)

42