1. Chemical Reactions

2. Effects of chemical reactions: Reactants  Products
Chemical reactions rearrange the atoms in the reactants to form new products.
The identities and properties of the products are completely different from that of the reactants!
Production of gases and color changes are signs (evidence) of chemical reactions.

3. Energy and Reactions Energy must be added to break bonds.
Energy is released when bonds are formed.
Chemical energy is CONSERVED in chemical reactions.

4. Exo- vs. Endo- EXOTHERMIC REACTIONS: release energy
More energy is released as the products form bonds than energy is absorbed to break the bonds in the reactants.
ENDOTHERMIC REACTIONS: absorb energy
More energy is absorbed to break the bonds in the reactants, than energy is released to bonds form in the products.

5. Chemical Equations 2H2 + O2 —> 2H2O
Chemical equations are used to represent or describe chemical reactions.
For example when hydrogen (H2) burns, it reacts with oxygen (O2) in the air to form water.
We write the chemical equation for this reaction as follows:
2H2 + O2 —> 2H2O

6. Chemical Equations An equation shows… Formulas of reactants
Formulas of products
Molar ratios of all compounds in the reaction.

7. Chemical Equations 2H2 + O2 → 2H2O mole ratio = 2:1:2
Here, we read the (+) sign as “reacts with” and the arrow (—>) as “produces” or “yields”.
2H2 + O2 → 2H2O
mole ratio = 2:1:2
Reactants
Products

8. To show physical states of each substance:
(s) or  solid
(l) liquid
(g) or  gas
(aq) aqueous (dissolved in water)

9. To show physical states of each substance:
Consider the reaction of iron (a solid) with oxygen (a gas) to form iron(III) oxide, or rust (a solid).
Fe(s) + O2(g)  Fe2O3(s) (unbalanced)

10. Coefficients & Subscripts
COEFFICIENTS: numbers in front of reactant or product that represents the number of moles
SUBSCRIPTS: lower, smaller numbers that define the formula of a compound
2H2SO4
Coefficient
Subscript

11. H2O One molecule of water
2H2O Two molecules of water
H2O2 One molecule of hydrogen peroxide

12. During a chemical reaction, atoms are rearranged, not created nor destroyed!
Chemical equations must be balanced to show the relative amounts of all substances.
Balanced: each side of the equations has the same number of atoms of each element.
CH4 + O2 —> H2O + CO2 Unbalanced
CH4 + 2O2 —> 2H2O + CO2 Balanced

13. In order to balance… Reactants  Products
Write correct formulas for all reactants and products
Count the number of atoms of each element in reactants & products.
Balance one at a time using coefficients.
Check for balance
Are the coefficients in the lowest possible ratio?

14. Balancing Equations
NOTE: When balancing equations, you may change coefficients as much as you need to, but you may never change subscripts because you can’t change what substances are involved.

15. Fe(s) + O2(g)  Fe2O3(s) (unbalanced)
4Fe(s) + 3O2(g)  2Fe2O3(s) (balanced)

16. Sample Problem 1 H2O  H2 + O2 2H2O  2H2 + O2
Water is decomposed (broken down) to form the gaseous products hydrogen (H2) and oxygen (O2). Write the balanced equation.
H2O  H2 + O2
2H + 1O  2H + 2O
O is not balanced
2H2O  2H2 + O2
4H + 2O  4H + 2O
The equation is balanced!

17. Sample Problem 2 Cl2 + KBr  KCl + Br2 Cl2 + 2KBr  2KCl + Br2
Chlorine gas (Cl2) reacts with potassium bromide (KBr) to form potassium chloride and bromine (Br2). Write the balanced equation.
Cl2 + KBr  KCl + Br2
2Cl + 1K + 1Br  1Cl + 1K +2Br
Cl and Br are not balanced
Cl2 + 2KBr  2KCl + Br2
2Cl + 2K + 2Br  2Cl + 2K +2Br The equation is balanced!

18. Balancing equations involves a great deal of “trial and error” at first,
but there are some tricks…

19. Na + H2O —> NaOH + H2 Let’s start with the even number two!
For example…..
Sodium metal reacts with water to produce sodium hydroxide and hydrogen gas.
Na + H2O —> NaOH + H2
Note that on the product side (right side) there are an odd number of hydrogens (3).
On the reactant side (left side) there is an even number (2).
This implies there must be an even coefficient in front of the NaOH.
Let’s start with the even number two!

20. _Na + _H2O —> 2NaOH + _H2
Now lets balance sodium; we need a 2 in front of the Na…
2Na + _H2O —> 2NaOH + _H2
Now consider hydrogen…
2Na + 2H2O —> 2NaOH + H2

21. the equation is balanced.
2Na + 2H2O —> 2NaOH + (1)H2
Check to see if it balances…
2 Na on the left 2 Na on the right
4 hydrogen = 4 hydrogen
2 oxygen 2 oxygen
the equation is balanced.

22. Examples CuCl2(aq) + Al(s)  Cu(s) +AlCl3(aq)
(3:2:3:2)

23. Balance C – then H – then O
Examples
Propane (C3H8) burns in oxygen (O2) to form carbon dioxide and water.
C3H8 + O2 CO2 + H2O
Balance C – then H – then O
C3H8 + 5O2 3CO2 + 4H2O
(1:5:3:4)

24. Balance C – then H – then O
Examples
Pentane (C5H12) burns in oxygen (O2) to form carbon dioxide and water.
C5H12 + O2 CO2 + H2O
Balance C – then H – then O
C5H12 + 8O2 5CO2 + 6H2O
(1:8:5:6)

25. Examples
Silver nitrate reacts with copper to produce silver and copper(II) nitrate.
AgNO3 + Cu Ag + Cu(NO3)2
2AgNO3 + Cu 2Ag + Cu(NO3)2
(2:1:2:1)

26. Examples
Phosphorus reacts with oxygen gas to produce diphosphorus pentoxide.
P + O2  P2O5
4P + 5O2  2P2O5
(4:5:2)

27. Balance C – then H – then O
Examples
C7H14 + O2 CO2 + H2O
Balance C – then H – then O
C7H ½O2 7CO2 + 7H2O
2C7H O2 14CO2 + 14H2O
(2:21:14:14)

28. Types of
Reactions

29. Types of Chemical Reactions
Synthesis / Combination
Decomposition
Single Replacement
Double Replacement
Combustion

30. Synthesis / Combination Reactions
Definition: Reaction where two or more substances react to form a single substance.
A + B  AB
Examples:
2K(s) + Cl2(g)  2KCl(s)
SO2(g) + H2O(l)  H2SO3(aq)

31. Decomposition Reactions
Definition: Reaction where a single compound is broken down into two or more products.
AB  A + B
Examples:
2H2O(l)  2H2(g) + O2(g)
CaCO3  CaO + CO2

32. Single-Replacement Reactions
Definition: Reaction where atoms of one element replace atoms of a second element in a compound.
XA + B  BA + X
Note: A reactive metal will replace any metal listed below it in the activity series. Generally, nonmetal replacement is limited to the halogens. The activity of the halogens decreases as you go down Group 7A of the periodic table. See handout.
Examples:
2AgNO3 + Mg  Mg(NO3)2+2Ag
Mg+LiNO3  no reaction

33. Any element will replace any element below it.Li K Ca Na Mg Al Zn Fe Pb
(H)* Cu Hg Ag
Activity Series
Increasing Activity
Any element will replace any element below it.
*Metals from Li to Na will replace H from acids and water; from Mg to Pb they will replace H from acids only

34. For Example… Ca + MgO  CaO + Mg
The Ca will replace the Mg because Ca is more active than Mg.
That is to say:
Ca is above Mg on the activity list.
For

35. Double-Replacement Reactions
Definition: Reaction that involves an exchange of positive ions between two compounds.
XA + BY  BA + XY
Note: These reactions generally take place between two ionic compounds in aqueous solution, and are often characterized by one of the products coming out of solution in some way. Examples:
2NaCN(aq)+H2SO4(aq)  2HCN(g)+Na2SO4(aq)
Na2S(aq)+Cd(NO3)2(aq)  CdS(s)+2NaNO3(aq)

36. CH4+2O2  CO2+2H2O + heat + light
Combustion Reactions
Definition: Reaction where an element or compound reacts with oxygen, often producing energy in the form of heat and light.
Examples:
CH4+2O2  CO2+2H2O + heat + light
2Mg(s)+O2(g)  2MgO(s)

37. Combustion of Hydrocarbons
If the reactant is a hydrocarbon, the products are always carbon dioxide and water.
CH4 + 2O2  CO H2O

38. Ionic Equations An aqueous solution is ions dissolved in water
When a soluble substance is dissolved in water, the substance often breaks into ions. This solution is said to be an aqueous solution.
Pb(NO3)2(aq)  Pb2+ + 2NO3-
NaI(aq)  Na+ + I-
An aqueous solution is ions dissolved in water

39. Ionic Equations Consider the reaction…
Pb(NO3)2(aq) + NaI(aq)  PbI2(s) + NaNO3(aq)
What is really going on is…
Pb2+ + NO3- + Na+ + I-  PbI2(s) + Na+ + NO3-
Note that the Na+ ion and the NO3- ion are not reacting. They are said to be spectator ions.

40. Net Ionic Equations
It is often useful to write an equation showing only the species that are actually reacting. This is called a net ionic equation. It does not show the spectator ions.
Pb2+ + NO3- + Na+ + 2I-  PbI2(s) + Na+ + NO3-
becomes….
Pb I-  PbI2(s)