1. Unit 9 Chemical Equations
CP Chemistry

2. Describing Chemical Change
OBJECTIVES:
Write equations describing chemical reactions, using appropriate symbols

3. Describing Chemical Change
OBJECTIVES:
Write balanced chemical equations, when given the names or formulas of the reactants and products in a chemical reaction.

4. 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

5. Symbols in equations
the arrow separates the reactants from the products
Read “reacts to form” or “yields”
The plus sign + “and”
(s) after the formula = solid
(g) after the formula = gas
(l) after the formula = liquid

6. Symbols used in equations
(aq) after the formula - dissolved in water, an aqueous solution.
­ used after a product indicates a gas (same as (g))
¯ used after a product indicates a solid (same as (s))

7. Symbols used in equations
indicates a reversible reaction (more later)
shows that heat is supplied to the reaction
is used to indicate a catalyst is supplied, in this case, platinum.

8. What is a catalyst?
A substance that speeds up a reaction, without being changed or used up by the reaction.
Enzymes are biological or protein catalysts.

9. Exothermic and Endothermic
Exothermic Reactions give off heat
Ex: explosions
Endothermic Reactions absorb heat and feel cold
Ex: ice packs from the nurse

10. In a chemical reaction The way atoms are joined is changed
Atoms aren’t created or destroyed.
Can be described several ways:
1. In a sentence
Copper reacts with chlorine to form copper (II) chloride.
2. In a word equation
Copper + chlorine ® copper (II) chloride

11. DO NOT FORGET DIATOMIC MOLECULES!!!
BrINClHOF
Never travel alone!!

12. Convert these to equations
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.

13. Now, write these: Fe(s) + O2(g) ® Fe2O3(s)
Cu(s) + AgNO3(aq) ® Ag(s) + Cu(NO3)2(aq)
NO2 (g) N2(g) + O2(g)

14. Balancing Chemical Equations

15. Balanced Equation Atoms can’t be created or destroyed
All the 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.

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

17. 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
2 NaCl is okay, Na2Cl is not.

18. Example
H2 + O2
H2O
Make a table to keep track of where you
are at

19. Example
H2 + O2
H2O R P 2 H 2 2 O 1
Need twice as much O in the product

20. Example
H2 + O2 2
H2O R P 2 H 2 2 O 1
Changes the O

21. Example
H2 + O2 2
H2O R P 2 H 2 2 O 1 2
Also changes the H

22. Example
H2 + O2 2
H2O R P 2 H 2 4 2 O 1 2
Need twice as much H in the reactant

23. Example 2
H2 + O2 2
H2O R P 2 H 2 4 2 O 1 2
Recount

24. Example 2
H2 + O2 2
H2O R P 4 2 H 2 4 2 O 1 2
The equation is balanced, has the same
number of each kind of atom on both sides

25. Example R P 4 2 H 2 4 2 O 1 2 2 H2 + O2 ® 2 H2O This is the answer
Not this

26. Balancing Examples _Mg + _N2 ® _Mg3N2 _P + _O2 ® _P4O10
_AgNO3 + _Cu ® _Cu(NO3)2 + _Ag
_Mg + _N2 ® _Mg3N2
_P + _O2 ® _P4O10
_Na + _H2O ® _H2 + _NaOH
_CH4 + _O2 ® _CO2 + _H2O

27. Balancing Examples 2AgNO3 + 2Cu ® Cu(NO3)2 + 2Ag 3Mg +N2 ® Mg3N2
4P + 5O2 ® P4O10
2Na + 2H2O ® H2 + 2NaOH
CH4 + 2O2 ® CO2 + 2H2O

28. Types of Chemical Reactions
OBJECTIVES:
Identify a reaction as combination, decomposition, single-replacement, double-replacement, or combustion

29. Types of Chemical Reactions
OBJECTIVES:
Predict the products of combination, decomposition, single-replacement, double-replacement, and combustion reactions.

30. Types of Reactions There are millions of reactions.
Can’t remember them all
Fall into several categories.
We will learn 5 major types.
Will be able to predict the products.
For some, we will be able to predict whether they will happen at all.
Will recognize them by the reactants

31. #1 - Combination (SYNTHESIS)
Combine - put together
2 substances combine to make one compound.
Ca +O2 ® CaO
SO3 + H2O ® H2SO4
We can predict the products if they are two elements.
Mg + N2 ® __?__

32. Write and balance Ca + Cl2 ® Fe + O2 ® Al + O2 ®
Remember that the first step is to write the correct formulas
Then balance by using coefficients only

33. #2 - Decomposition Reactions
decompose = fall apart
one reactant falls apart into two or more elements or compounds.
NaCl Na + Cl2
CaCO CaO + CO2
Note that energy is usually required to decompose

34. #2 - Decomposition Reactions
Can predict the products if it is a binary compound
Made up of only two elements
Falls apart into its elements
H2O
HgO

35. #2 - Decomposition Reactions
If the compound has more than two elements you will only be asked to balance them
__ NiCO3 ® ___ NiO + ___ CO2
__KClO3(aq) __ KCl +__ O2

36. #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.
K + NaCl ® Na + KCl
F2 + LiCl ® LiF + Cl2

37. #3 Single Replacement Metals replace other metals (and hydrogen)
K + AlN ®
Zn + HCl ®
Think of water as HOH
Metals replace one of the H, combine with hydroxide.
Na + HOH ®

38. #3 Single Replacement We can tell whether a reaction will happen
Some chemicals are more “active” than others
More active replaces less active
There is a list on the last page of your packet - called the Activity Series of Metals
Higher on the list replaces lower.
If something isn’t on the activity series you have, assume the one with the higher atomic number is higher on the list

39. #3 Single Replacement Note the * concerning Hydrogen
H can be replaced in acids by everything higher
Li, K, Ba, Ca, & Na replace H from acids and water
Al + HCl ®
Fe + CuSO4 ®
Pb + KCl ®

40. #3 - Single Replacement Nonmetals can replace other nonmetals
Limited to F2 , Cl2 , Br2 , I2 (halogens)
Higher replaces lower.
F2 + HCl ®
Br2 + KCl ®

41. #4 - Double Replacement Two things replace each other.
Reactants must be two ionic compounds or acids.
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

42. #4 - Double Replacement Has certain “driving forces”
Will only happen if one of the products:
doesn’t dissolve in water and forms a solid (a “precipitate”), or
is a gas that bubbles out, or
is a covalent compound (usually water).

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

44. How to recognize which type
Look at the reactants:
E + E = Combination
C = Decomposition
E + C = Single replacement
C + C = Double replacement

45. Examples H2 + O2 ® H2O ® Zn + H2SO4 ® HgO ® KBr +Cl2 ® AgNO3 + NaCl ®
Mg(OH)2 + H2SO3 ®

46. #5 - Combustion Means “add oxygen”
A compound composed of only C, H, and maybe O is reacted with oxygen
If the combustion is complete, the products will be CO2 and H2O.
If the combustion is incomplete, the products will be CO (possibly just C) and H2O.

47. Examples
C4H10 + O2 ®
C6H12O6 + O2 ®
C8H8 +O2 ®

48. 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, and if a catalyst or energy is required.

49. Reactions Come in 5 major types.
Can tell what type they are by the reactants.
Single Replacement happens based on the activity series
Double Replacement happens if the product is a solid, water, or a gas.

50. Net Ionic Equations
Many reactions occur in water- that is, in aqueous solution
Many ionic compounds “dissociate”, or separate, into cations and anions when dissolved in water
Now we can write a complete ionic equation

51. Net Ionic Equations Example: AgNO3 + NaCl  AgCl + NaNO3
1. this is the full equation
2. now write it as an ionic equation
3. can be simplified by eliminating ions not directly involved (spectator ions) = net ionic equation

52. Predicting the Precipitate
Insoluble salt = a precipitate - note Figure 8.13, p.227
General rules: Table 8.3, p. 227, Reference p.7 (back of textbook), and in Lab manual p.338
Sample problem 8-11, p.228