1. Chapter 11 Chemical Reactions
Section 11. 1
Describing Chemical Reactions

2. 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

3. 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)

4. 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).

6. 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.

7. The Skeleton Equation
Uses 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:
Solid iron (III) sulfide reacts with gaseous hydrogen chloride to form iron (III) chloride and hydrogen sulfide gas.
Nitric acid dissolved in water reacts with solid sodium carbonate to form liquid water and carbon dioxide gas and sodium nitrate dissolved in water.

8. 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)

9. Law of Conservation of Matter
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.

10. 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.

11. Rules for Balancing:
Assemble the correct formulas for all the reactants and products, use + and →
Count the number of atoms of each type appearing on both sides
Balance the elements one at a time by adding coefficients where needed (the numbers in front) - save balancing the H and O until LAST!
Check to make sure it is balanced.

12. Never 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.

13. Balancing Chemical Equations
Example:
HCl + NaOH NaCl + H2O
H=2 H=2
Cl= Cl=1
Na=1 Na=1
O= O=1
The equation is balanced because the number of atoms in the reactants are equal to the number of atoms in the products.

14. Balancing Chemical Equations
Example:
H2 + O H2O
H=2 O= H=2 O=1
H2 + O H2O
H=2 O= H= O=2
2H2 + O H2O
H=4 O= H=4 O=2

15. Balancing Chemical Equations
Example:
Cu + AgNO Cu(NO3)2 + Ag
Cu=1 Ag=1 N=1 O= Cu=1 Ag=1 N=2 O=6
Cu + 2AgNO Cu(NO3)2 + Ag
Cu=1 Ag=2 N=2 O= Cu=1 Ag=1 N=2 O=6
Cu + 2AgNO Cu(NO3)2 + 2Ag
Cu=1 Ag=2 N=2 O= Cu=1 Ag=2 N=2 O=6

16. Balancing Chemical Equations
NaHCO3 + H3C6H5O CO2 + H2O + Na3C6H5O7
Na=1 H=9 C=7 O= Na=3 H=7 C=7 O=10
3NaHCO3 + H3C6H5O CO2 + H2O + Na3C6H5O7
Na=3 H=11 C=9 O= Na=3 H=7 C=7 O=10
3NaHCO3 + H3C6H5O CO2 + H2O + Na3C6H5O7
Na=3 H=11 C=9 O= Na=3 H=7 C=9 O=14
3NaHCO3 + H3C6H5O CO2 + 3H2O + Na3C6H5O7
Na=3 H=11 C=9 O= Na=3 H=11 C=9 O=16

17. Practice Balancing Examples
…AgNO3 + …Cu ® …Cu(NO3)2 + …Ag
…Mg + …N2 ® …Mg3N2
…P + …O2 ® …P4O10
…Na + …H2O ® …H2 + …NaOH
…CH4 + …O2 ® …CO2 + …H2O

18. End of Section 11.1

19. Types of Chemical Reactions
Section 11.2
Types of Chemical Reactions

20. Types of Reactions There are 5 major types of chemical reactions
Combination reaction or Synthesis reaction
Decomposition reaction
Single Replacement reaction
Double Replacement reaction
Combustion reaction
Not all reactions fit into only one category
Patterns of chemical reactions will help you predict the products of the reaction

21. 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)

22. 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

23. #2 - Decomposition Reactions
decompose = fall apart
one reactant breaks apart into two or more elements or compounds.
NaCl Na + Cl2
CaCO CaO + CO2
Note that energy (heat, sunlight, electricity, etc.) is usually required

24. Can predict the products if it is a binary compound-Made up of only two elements
breaks apart into its elements:
H2O
HgO
H2 + O2
Hg + O2

25. #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

26. 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 ®

27. Higher on the list replaces lower
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 called the
Activity Series of Metals
Higher on the list replaces lower

28. The Activity Series of the Metals
Higher activity
Lithium
Potassium
Calcium
Sodium
Magnesium
Aluminum
Zinc
Chromium
Iron
Nickel
Lead
Hydrogen
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).
Lower activity

29. Practice: Fe + CaSO4 Pb + KCl Al + HCl No Reaction No Reaction
AlCl H2

30. The Activity Series of the Halogens
Higher Activity
Halogens can replace other
halogens in compounds, provided that they are above the halogen that they are trying to replace.
Fluorine
Chlorine
Bromine
Iodine
Lower Activity
2NaCl(s) + F2(g) 
2NaF(s) + Cl2(g)
???
MgCl2(s) + Br2(g) 
No Reaction
???

31. #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 + NaCl

32. Complete and balance:
assume all of the following reactions actually take place:
CaCl2 + NaOH
CuCl2 + K2S
KOH + Fe(NO3)3
(NH4)2SO4 + BaF2

33. Practice Examples:
H2 + O2
H2O
Zn + H2SO4
HgO
KBr + Cl2
AgNO3 + NaCl
Mg(OH)2 + H2SO3

34. #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.

35. Combustion Examples: C4H10 + O2 C3H8 + O2 C6H12O6 + O2 C8H8 + O2
(assume complete)
C4H10 + O2
C3H8 + O2
C6H12O6 + O2
C8H8 + O2

36. 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.

37. 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)

38. End of Section 11.2

39. Section 11.3 Reactions in Aqueous Solution

40. 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:

41. Solubility Rules for Ionic Compounds
Insoluble salt = Precipitate
Solubility Rules for Ionic Compounds
Compounds
Solubility
Sodium, potassium, and ammonium salts
Soluble
All nitrates and chlorates salts
All chlorides except silver chloride and lead chloride
All sulfates except, silver sulfate, lead sulfate, and barium sulfate
All carbonates, phosphates, hydroxides, sulfides and chromates salts except with sodium, potassium and ammonium
Insoluble

42. Example:
CaCl2(s) + Pb(NO3)2(aq)
PbCl2(s) + Ca(NO3)2(aq)