1. Introduction to Reactions

2. Chemical Equation Reactants  Products Fe + O2  Fe2O3
A catalyst is a substance that speeds up the reaction but is not changed by it.
It is neither a reactant or a product.

3. Signs of a Reaction Release of a gas
CO2 is released when acid is placed in a solution containing CO32- ions
Formation of a solid (precipitate)
A solution containing Ag+ ions mixed with a solution containing Cl- ions
Heat is produced or absorbed
Acid and base are mixed together
Color changes

4. Common Symbols Symbol Meaning  forms, produces ↔ reversible reaction
(s) Solid state
(l) Liquid state; water only
(g) Gaseous state
(aq) aqueous state, all liquids besides water
heat/energy is supplied to the reaction
Catalyst is used, here platinum

5. Features of a Chemical Equation
Products and reactants must be specified using chemical symbols
Reactants – written on the left of arrow
Products – written on the right
 – energy is needed
Physical states are shown in parentheses

6. Writing Equations 2H2 (g) + O2(g)  2H2O(g)
Identify the substance involved
Coefficients - how many?
Chemical Formula – of what?
Physical State – in what state?
Remember Diatomic Elements
Magic Seven

7. Example
Two atoms of aluminum react with three units of aqueous copper (II) chloride to produce three atoms of copper and two units of aqueous aluminum chloride?
How many?
Of what?
What physical state?

8. Example
Two atoms of aluminum react with three units of aqueous copper (II) chloride to produce three atoms of copper and two units of aqueous aluminum chloride?
How many?
Of what?
What physical state?
2 Al(s) + 3 CuCl2(aq)  3 Cu(s) + 2AlCl3(aq)

9. Describing Equations Describing Coefficients: 3CO2  2Mg  4MgO 
individual atom = “atom”
covalent substance = “molecule”
ionic substance = “unit”
3CO2 
2Mg 
4MgO 

10. Describing Equations Describing Coefficients: 3CO2  2Mg  4MgO 
individual atom = “atom”
covalent substance = “molecule”
ionic substance = “unit”
3CO2 
2Mg 
4MgO 
3 molecules of carbon dioxide
2 atoms of magnesium
4 units of magnesium oxide

11. Zn(s) + 2HCl(aq)  ZnCl2(aq) + H2(g)
Describing Equations
Zn(s) + 2HCl(aq)  ZnCl2(aq) + H2(g)
How many?
Of what?
In what state?

12. Zn(s) + 2HCl(aq)  ZnCl2(aq) + H2(g)
Describing Equations
Zn(s) + 2HCl(aq)  ZnCl2(aq) + H2(g)
How many?
Of what?
In what state?
One atom of solid zinc reacts with two units of aqueous hydrochloric acid to produce one unit of aqueous zinc chloride and one molecule of hydrogen gas

13. Balancing Reaction

14. Balancing Reactions 4Fe + 3O2  2Fe2O3
Law of conservation of mass - matter cannot be created or destroyed
mass of the products = mass of the reactants
Coefficient: # of moles of products & reactants
4Fe + 3O2  2Fe2O3
Diatomic elements (The Magic 7)
H2, N2, O2, F2, Cl2, Br2, I2

15. Coefficient - how many of that substance are in the reaction
Balancing
Coefficient - how many of that substance are in the reaction
The equation must be balanced
All the atoms of every reactant must also appear in the products
Number of Hg on left? 2
on right 2
Number of O on left? 2

16. Examine the Equation H2 + O2  H2O
Is the law of conservation of mass obeyed as written? NO
You never change subscripts
WRONG: H2 + O2  H2O2

17. Steps in Equation Balancing
H2 + O2  H2O
The steps to balancing:
Step 1. Count the number of moles of atoms of each element on both product and reactant sides
Reactants Products
2 mol H mol H
2 mol O mol O

18. Step 2. Determine which elements are not balanced – Oxygen is not balanced
Step 3. Balance one element at a time by changing the coefficients
H2 + O2  2H2O
This balances oxygen, but is hydrogen still balanced?
2H2 + O2  2H2O
Step 4. Make sure the law of conservation of mass is obeyed
Reactants Products
4 mol H 4 mol H
2 mol O 2 mol O

19. Practice Equation Balancing
Balance the following equations: 1. C2H2 + O2  CO2 + H2O 2. AgNO3 + FeCl3  Fe(NO3)3 + AgCl 3. C2H6 + O2  CO2 + H2O 4. N2 + H2  NH3

20. Practice Equation Balancing
Balance the following equations: 1. 2C2H2 + 5O2  4CO2 + 2H2O 2. 3AgNO3 + FeCl3  Fe(NO3)3 + 3AgCl 3. 2C2H6 + 5O2  4CO2 + 6H2O 4. N2 + 3H2  2NH3

21. Types of Reaction

22. Combination Reactions
Synthesis reactions
The joining of two or more elements or compounds, producing a product of different composition
Examples:
metal + nonmetal  salt: 2Na(s) + Cl2(g)  2NaCl(s)
H + Cl  HCl
MgO(s) + CO2(g)  MgCO3(s)
A + B  AB

23. Decomposition Reactions
Produce two or more products from a single reactant
Reverse of a combination reaction
Examples:
2HgO(s)  2Hg(l) + O2(g)
CaCO3(s)  CaO(s) + CO2(g)
Removal of water from a hydrated material
AB  A + B

24. Replacement Reactions
Single-replacement
One atom replaces another in the compound producing a new compound
Examples:
Cu(s)+2AgNO3(aq)  2Ag(s)+Cu(NO3)2(aq)
2Na(s) + 2H2O(l)  2NaOH(aq) + H2(g)
A + BC  B + AC

25. Single Replacement Rxn.
Activity Series – lists metals in order of decreasing reactivity (p.333)
Reactive metals will replace any metal listed below it in the activity series
If the metal is below, no reaction occurs
Halogen(7A) can replace other halogens that are below it in the periodic table

26. Activity Series

27. Single Replacement Rxn.
2K(s) + 2H2O(l) 
Zn(s) + Cu(NO3)2(aq) 
Cu(s) + Al2O3(aq) 
Br2(aq) + 2NaI(aq) 
Br2(aq) + NaCl 

28. Single Replacement Rxn.
2K(s) + 2H2O(l)  2KOH(aq) + H2(g)
Zn(s) + Cu(NO3)2(aq)  Cu(s) + Zn(NO3)2(aq)
Cu(s) + Al2O3(aq)  No reaction
Br2(aq) + 2NaI(aq)  2NaBr(aq) + I2(aq)
Br2(aq) + NaCl  No reaction

29. Two compounds undergo a “change of partners”
Double Replacement
Two compounds undergo a “change of partners”
Two compounds react by exchanging atoms to produce two new compounds
AB + CD  AD + CB

30. Double Replacement Rxn.
Double-displacement reaction
Exchange of positive ions
Occur in aqueous solution
To occur:
One of the products is slightly soluble and a precipitates forms
One product is a gas
One of the products is a molecular compound, like water

31. Types of Double-Replacement
Acid + base  water and salt
HCl(aq)+NaOH(aq) NaCl(aq)+H2O(l)
Formation of solid lead chloride from lead nitrate and sodium chloride
Pb(NO3)2(aq) + 2NaCl(aq)  PbCl2(s) + 2NaNO3(aq)
AB + CD  AD + CB

32. Precipitation Reactions
Chemical change in a solution that results in one or more insoluble products
Solubility Rules (p.344)
1. salts of alkali metals and ammonia  soluble
2. nitrate salts and chlorate salts  soluble
3. sulfate salts, except compounds with Pb, Ag, Hg, Ba, Sr, and Ca  soluble
4. Chloride salts, except with Ag, Pb, and Hg soluble
5. carbonates, phosphates, chromates, sulfides, and hydroxides  most are insoluble

33. Predicting Whether Precipitation
Will Occur
Recombine the ionic compounds to have them exchange partners
Examine the new compounds formed and determine if any are insoluble
Any insoluble salt will be the precipitate
Pb(NO3)2(aq) + NaCl(aq) 
PbCl2 (?) + NaNO3 ( ?)
(s)
(aq)

34. Precipitates Predict Whether These Reactions Form Precipitates
Potassium chloride and silver nitrate
Potassium acetate and silver nitrate

35. Precipitates Predict Whether These Reactions Form Precipitates
Potassium chloride and silver nitrate
KCl(aq) + AgNO3(aq)  KNO3(aq) + AgCl (s)
Potassium acetate and silver nitrate
KC2H3O2 + AgNO3(aq)  KNO3(aq) + AgC2H3O2(s)

36. Reactions with Oxygen
Reactions with oxygen generally release energy in the form of light or heat
Combustion
Reactants: Oxygen and a hydrocarbons
Products: CO2 and H2O
Combustion of natural gas
CH4+2O2CO2+2H2O
Rusting or corrosion of iron
4Fe + 3O2  2Fe2O3