1. Ch. 9 Chemical Reactions & Equations
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Ch. 9 Chemical Reactions & Equations
Zn + I2
Reactants
Zn I2
Product

2. Introduction
Chemical reactions occur when bonds between atoms are formed or broken
Chemical reactions involve changes in matter, the making of new materials with new properties, and energy changes.
Symbols represent elements,
Formulas represent compounds,
Chemical equations represent chemical reactions

3. Combustion of Methane CH4 + 2O2  CO2 + 2H2O
Atoms are rearranged!

4. Chemical Equations
Yields or produces!
Their Job: Depict the kind of reactants and products and their relative amounts in a reaction.
4 Al (s) O2 (g) > 2 Al2O3 (s)
The numbers in the front of formulas are called COEFFICIENTS
The letters (s), (g), and (l) are the physical states of compounds.
Starting Materials
What’s created

5. The charcoal used in a grill is basically carbon
The charcoal used in a grill is basically carbon. The carbon reacts with oxygen to yield carbon dioxide.

6. This is the “Word Equation” for that reaction
Word Equations: show the names of the reactants and
the products
Lavoisier, 1788
This is the “Word Equation” for that reaction
carbon + oxygen carbon dioxide
A skeleton equation does NOT indicate the
relative amounts of reactants and products
(no coefficients!)
The skeleton equation for that reaction is:
C + O2  CO2
Chemical Equations must be balanced in order to conform to the Law of Conservation of Mass - same # & type of atoms on each side of the yield arrow.

7. Skeleton equation: CH4 + O2 CO2 + H2O
Word Equation:
Methane + Oxygen gas carbon dioxide + water
Skeleton equation: CH4 + O CO2 + H2O

8. Symbols Used in Equations
Solid (s)
Liquid (l)
Gas (g)
Aqueous solution (aq) (dissolved in water)
Catalyst H2SO4
Escaping gas ()
Change of temperature ()
Precipitate ( )

9. Balancing Equations
When balancing a chemical reaction you may add coefficients in front of the compounds to balance the reaction, but you may not change the subscripts.
Changing the subscripts changes the compound. Subscripts are determined by the valence electrons (charges for ionic or sharing for covalent)
Never put a coefficient in the middle of a formula
2 NaCl is ok Na2Cl is not.

10. Subscripts vs. Coefficients
The subscripts tell you how many atoms of a particular element are in a compound. The coefficient tells you about the quantity, or number, of molecules of the compound.

11. There are a number of ways to interpret balanced equations
2 H2(s) + O2(g) ---> 2 H2O(s)
This equation means:
2 molecules H2 + 1 molecules O2 ---->2 molecules H2O(g)

12. Steps to Balancing Equations
There are four basic steps to balancing a chemical equation.
Write the correct formula for the reactants and the products. DO NOT TRY TO BALANCE IT YET! You must write the correct formulas first. And most importantly, once you write them correctly DO NOT CHANGE THE FORMULAS!
Find the number of atoms for each element on the left side. Compare those against the number of the atoms of the same element on the right side.
Determine where to place coefficients in front of formulas so that the left side has the same number of atoms as the right side for EACH element in order to balance the equation.
Check your answer to see if:
The numbers of atoms on both sides of the equation are equal.
The coefficients are in the lowest possible whole number ratios. (reduced)

13. Some Suggestions to Help You
Helpful Hints for balancing equations:
Take one element at a time.
Save pure elements for last
IF everything balances except for that last pure element, and there is no way to balance it with a whole number, double all the coefficients and try again.
(Shortcut) Polyatomic ions that appear on both sides of the equation should be balanced as independent units
That is, don’t separate them into individual atoms!

14. Example Make a table to keep track of atoms H2 + O2  H2O R P 2 H O 1
Need another O on the product side

15. Example H2 + O2  2H2O Place a coefficient of 2 in front of H2O R P 21
Changes the O

16. Example
H2 + O2  2H2O R P 2 H O 1 2
This also changes the H

17. Example
H2 + O2  2H2O R P 2 H O 1 4 2
Now we need twice as much H in the reactant

18. Example
2H2 + O2  2H2O R P 2 H O 1 4 2
Add a coefficient of 2 in front of H2

19. Example
2H2 + O2  2H2O
Your answer R P 2 H O 1 4 4 2
Recount to check

20. Balancing Equations Na3PO4 + Fe2O3 ---> Na2O + FePO4
Sodium phosphate + iron (III) oxide  sodium oxide + iron (III) phosphate
Na3PO Fe2O3 ---> Na2O FePO4 3 2 2 R L 2 1 3 1 2 6 2 3 Na
PO4 Fe O 6 2

21. 4 10 2 11
__B4H ___O >___ B2O ___H2O
__C3H8 + __O > __CO2 + __ H2O 5 3 4 3
Balancing Equations

22. Types of Reactions
There are millions of chemical reactions. The only way to be sure what your products will be is to carry them out in the lab!
Not very practical – or cost effective, however…
There are five types of chemical reactions we will talk about:
Synthesis reactions
Decomposition reactions
Single Replacement reactions
Double Repalcement reactions
Combustion reactions
You need to be able to identify the type of reaction and predict the product(s)

23. 1. Synthesis reactions
Synthesis reactions occur when two or more substances (generally elements) combine to form a single compound.
(Sometimes these are called combination reactions.)
reactant + reactant  1 product
Basically: A + B  AB
Example: 2H2 + O2  2H2O
Example: C + O2  CO2

24. Synthesis Reactions
Here is another example of a synthesis reaction

25. Balancing Equations 2 3
___ Al(s) + ___ Br2(l) ---> Al2Br6(s)

26. Practice
Predict the products. Write and balance the following synthesis reaction equations.
Sodium metal reacts with chlorine gas
Na(s) + Cl2(g) 
Solid Magnesium reacts with fluorine gas
Mg(s) + F2(g) 
Aluminum metal reacts with fluorine gas
Al(s) + F2(g) 

27. 2. Decomposition Reactions
Decomposition reactions occur when a single compound breaks up into one or more elements or simpler compounds
1 Reactant  Product + Product
In general: AB  A + B
Example: 2 H2O  2H2 + O2
Example: 2 HgO  2Hg + O2

28. Decomposition Reactions
Another view of a decomposition reaction:

29. Example: Decomposition of hydrogen peroxide in the presence of a catalyst.
Word equation:
hydrogen peroxide  water + oxygen

30. Example: Decomposition of hydrogen peroxide in the presence of a catalyst.
Skeleton equation:
H2O2  H2O O2

31. Example: Decomposition of hydrogen peroxide in the presence of a catalyst.
Balanced Equation:
2H2O2  H2O O2

32. Example: Decomposition of hydrogen peroxide in the presence of a catalyst.
Balanced equation showing
the catalyst (MnO2)
the state of the reactants and products:
MnO2
2H2O2(l)  H2O (l) O2 (g)

33. Decomposition Exceptions
Carbonates and chlorates are special case decomposition reactions that do not go to the elements.
Carbonates (CO32-) decompose to carbon dioxide and a metal oxide
Example: CaCO3  CO2 + CaO
Chlorates (ClO3-) decompose to oxygen gas and a metal chloride
Example: 2 Al(ClO3)3  2 AlCl3 + 9 O2
There are other special cases, but we will not explore those in Chemistry I

34. Practice N2(g) + O2(g)  BaCO3(s)  Co(s)+ S(s)  Nitrogen monoxide
Identify the type of reaction for each of the following synthesis or decomposition reactions, and write the balanced equation:
N2(g) + O2(g) 
BaCO3(s) 
Co(s)+ S(s) 
NH3(g) + H2CO3(aq) 
NI3(s) 
Nitrogen monoxide
(make Co be +3)

35. 3. Single Replacement Reactions
Single Replacement Reactions occur when one element replaces another in a compound.
A metal can replace a metal (+) OR a nonmetal can replace a nonmetal (-).
element + compound element + product
A + BC  AC + B (if A is a metal) OR
A + BC  BA + C (if A is a nonmetal)
(remember the cation always goes first!)
Use Activity Series of Metals
to see if reaction works.

36. We have looked at several reactions: Fe + CuSO4  Cu + Fe2(SO4)3
Li + H2O  LiOH + H2
Such experiments reveal trends. The activity series ranks the relative reactivity of metals.
It allows us to predict if certain chemicals will undergo single displacement reactions when mixed: metals near the top are most reactive and will displacing metals near the bottom.
Q: Which of these will react?
Fe + CuSO4 
Ni + NaCl 
Li + ZnCO3 
Al + CuCl2  K Na Li Ca Mg Al Zn Fe Ni Sn Pb H Cu Hg Ag Au
No, Ni is below Na
Yes, Li is above Zn
Yes, Al is above Cu
Yes, Fe is above Cu
Cu + Fe2(SO4)3
NR (no reaction)
Zn + Li2CO3
Cu + AlCl3

37. A: No Mg + H2O  No Reaction
H is the only nonmetal listed. H2 may be displaced from acids or can be given off when a metal reacts with H2O (producing H2 + metal hydroxide). The reaction with H2O depends on metal reactivity.
Q: will Mg react with H2O? K Na Li Ca Mg Al Zn Fe Ni Sn Pb H Cu Hg Ag Au
A: No Mg + H2O  No Reaction
Complete these reactions:
Al + HCl
Cu + Sn2S 
Ca + Cu2SO4 
Na + PbO 
No Reaction NR
2 Cu + CaSO4 2
Pb + Na2O

38. Single Replacement Reactions
Another view:

39. 2 Al + Fe2O3
2 Fe + Al2O3
Here is a mixture of aluminum powder and iron (III) oxide
powder.
Once ignited the aluminum rips the oxygen off of the iron
oxide resulting in molten iron, which falls down to the
pot of sand below. One time someone used wet sand,
which was a mistake because the molten iron turned
the water to steam instantly. This sprayed molten iron
all over the front two rows of students.
Write the balanced equation for this reaction.

40. Single Replacement Reactions
Write and balance the following single replacement reaction equation:
Zinc metal reacts with aqueous hydrochloric acid
Zn(s) HCl(aq)  ZnCl2 + H2(g)
Note: Zinc replaces the hydrogen ion in the reaction 2

41. Single Replacement Reactions
Sodium chloride solid reacts with fluorine gas
NaCl(s) + F2(g)  NaF(s) + Cl2(g)
Note that fluorine replaces chlorine in the compound
Aluminum metal reacts with aqueous copper (II) nitrate
Al(s)+ Cu(NO3)2(aq) 2 2

42. 4. Double Replacement Reactions
Double Replacement Reactions occur when a metal in one compound replaces a metal in another compound.
compound + compound  compound + compound
AB + CD  AD + CB

43. Double Replacement Reactions
Think about it like “foil”ing in algebra, first and last ions go together + inside ions go together
Example:
AgNO3(aq) + NaCl(s)  AgCl(s) + NaNO3(aq)
Another example:
K2SO4(aq) + Ba(NO3)2(aq)  KNO3(aq) + BaSO4(s) 2

45. Practice
Predict the products. Then, write and balance the following decomposition reaction equations:
Solid Lead (IV) oxide decomposes PbO2(s) 
Aluminum nitride decomposes
AlN(s) 

46. MnO2 + CO --> Mn2O3 + CO2 Carbon monoxide is commonly used to strip off
oxygen atoms from metals.
This reaction shows the first step. If more CO is present, eventually all oxygen atoms will be grabbed by CO and manganese (Mn) metal will be left.
This is how metal ores get converted to metals.
2 MnO2 + CO --> Mn2O3 + CO2
Balance this equation!

47. Practice Predict the products. Balance the equation
HCl(aq) + AgNO3(aq) 
CaCl2(aq) + Na3PO4(aq) 
Pb(NO3)2(aq) + BaCl2(aq) 
FeCl3(aq) + NaOH(aq) 
H2SO4(aq) + NaOH(aq) 
KOH(aq) + CuSO4(aq) 

48. 5. Combustion Reactions
Combustion reactions occur whenever something reacts with oxygen gas.
This is also called burning!!! In order to burn something you need the 3 things in the “fire triangle”: 1) A Fuel (hydrocarbon) 2) Oxygen to burn it with 3) Something to ignite the reaction (spark)

49. Complete Combustion of Hydrocarbons
In general: CxHy + O2  CO2 + H2O
Products are ALWAYS carbon dioxide and water. (although incomplete burning does cause some by-products like carbon monoxide)
Combustion is used to heat homes and run automobiles (octane, as in gasoline, is C8H18)

50. Combustion
Example
C5H O2  CO2 + H2O
Write the products and balance the following combustion reaction:
C10H O2  5 8 6

51. Combustion Reactions
Edgar Allen Poe’s drooping eyes and mouth are potential signs of CO poisoning.

52. Mixed Practice
State the type, predict the products, and balance the following reactions:
BaCl2 + H2SO4 
C6H12 + O2 
Zn + CuSO4 
Cs + Br2 
FeCO3 

53. Some steps for Writing Reactions
Identify the type of reaction
Predict the product(s) using the type of reaction as a model
Remember! If you write a formula – You MUST balance it!
3. Balance the equation (using coefficients)
Don’t forget about the diatomic elements! (BrINClHOF or HON 7) For example, Oxygen is O2 as an element.
In a compound, it can’t be a diatomic element because it’s not an element anymore, it’s part of a compound!