CHEMICAL REACTIONS
I. A CHEMICAL REACTION is the process by which a chemical change is brought about.
A.
The Law of Conservation of Matter states that matter cannot be created nor destroyed during a
chemical change. Matter is only rearranged. The same atoms that were present before the reaction
are present after the reaction.
B.
Word equations tell the chemicals used but does not tell how much.
1. Sodium and chlorine produce sodium chloride.
2. Hydrogen and oxygen yield water.
C.
Formula equations tell the chemicals used and also tell how much: qualitative –chemicals that are
present, quantitative  how much (mass) of a chemical is present.
1. EXAMPLES:
a.
2 Na + Cl2  2 NaCl
b.
2 H2 + O2  2 H2O
c.
2Ag + S
 Ag2S
2. An arrow  stands for yield or forms or produces.
3. Reactants are chemicals on the left side of the arrow.
4. Products are chemicals on the right side of the arrow.
5. Plus signs ( + ) means and.
Understanding Chemical Equations
Let’s use the following reaction as an example…
HCl + NaOH → H2O + NaCl
The reactants are your starting chemicals…
HCl & NaOH
The products are what you end up with after the reactants interact. Here the
H2O & NaCl
If you were going to read the above reaction out loud you would say…
HCl plus NaOH yields H2O plus NaCl
products are…
D.
Rules for writing and balancing an equation.
1. The total number of atoms must be the same on both sides of the equation.
a. H2 + O2  H2O
this equation is not balanced
2H
2O
2H and 1O
the number of atoms on each side are not the same
therefore, use coefficients to balance the equation.
balanced
2 H2 + O2  2 H2O
b.
Zn + HCl  ZnCl2 + H2 not balanced
c.
Zn + 2HCl  ZnCl2 + H2
coefficients are added to balance the equation
AgNO3 + MgCl2  AgCl + Mg(NO3)2
not balanced
2AgNO3 + MgCl2  2AgCl + Mg(NO3)2
balanced
2. Read the statement. Write out symbols of elements with correct formulas.
a. 7 diatomic elements have to be written with the subscript 2.
H2
N2
O2
F2
Cl2
Br2
I2 BrINClHOF
b.
Formulas are written and balanced before being put in the equation.
c.
Once the formula is balanced with subscripts, DO NOT CHANGE IT.
3. Make the number of atoms equal on each side of the arrow, by adding a coefficient in front of
each formula that needs one. DO NOT EVER CHANGE THE FORMULA!
4. Count the number of atoms on each side of the equation. If they are equal the equation is
balanced. If they are not equal, recheck the charges of the elements, the subscripts of the
formulas and the coefficients of the compounds.
Understanding Balancing Chemical Equations
Definition – An equation is balanced when there are equal
numbers of all atoms on each side of the
Example…
Na + Cl2 → NaCl
This reaction shows sodium reacting with chlorine…
Notice that there are 2Cl atoms on the left of the
arrow and only one on the right.
Na + Cl2 → 2NaCl
So the equation is not balanced!
But what if we add some coefficients?
2Na + Cl2 → 2NaCl
Now we have 2 Na’s on each side..
And 2 Cl’s on each side.
This equation is now balanced!
equation.
E.
EXAMPLES:
1. Magnesium and hydrogen chloride produce magnesium chloride and hydrogen.
2. Sodium and oxygen yield sodium oxide.
3. Hydrogen and iodine forms hydrogen iodide.
4. Potassium and chlorine make potassium chloride.
3 POINTS TO REMEMBER WHEN WRITING & BALANCING EQUATIONS:

You cannot balance an equation using subscripts, this would change the formulas of the

compounds.

You cannot insert a coefficient between elements in a chemical formula.

Write the chemical formula for H 2O as HOH.
3 SUGGESTIONS FOR BALANCING:
 Balance metal and nonmetal atoms first.
 Balance hydrogen and oxygen atoms last.
 Balance polyatomic ions as a group if they appear unchanged on both sides of the equation.
II. TYPES OF REACTIONS:
A.
COMPOSITION (a.k.a. synthesis) reactions occur when two or more reactants combine to create
one product.
1. General form:
A + B  AB
2. EXAMPLES: (Balance them)
a. C + O2
 CO2
B.
b.
Mg + O2

MgO
c.
K

KCl
+ Cl2
DECOMPOSITION reactions occur when one reactant breaks down to form two or more products.
There are three types. You will need to memorize the specific products they always form.
1. General form:
AB  A + B
2. Metal Carbonates: Metal carbonates decomposes into metal oxide and carbon dioxide.
a. K2CO3  K2O + CO2
b. CaCO3  CaO + CO2
3. Metal hydroxides: Metal hydroxides decompose into metal oxide and water.
a. Mg(OH)2 
MgO + HOH
b. 2NaOH

Na2O + HOH
4. Metal chlorates: Metal chlorates decompose into metal chlorides and oxygen.
a. 2 KClO3 
2 KCl + 3 O2
b. Ca(ClO3)2 
CaCl2 + 3 O2
5. Reversible reactions-some reactions can go either way depending on the conditions.
C.
SINGLE REPLACEMENT reactions occur when the reactant are one element and a compound.
The element replaces one of the elements of the compound. A cation replaces a cation and an anion
replaces an anion.
1. General form:
cation replaces cation
AB + X  XB + A
anion replaces anion
AB + Y  AY + B
2. EXAMPLES: (Balance them)
a.
Zn
+
HCl

ZnCl2
b.
Ca
+
HOH

Ca(OH)2 + H2
c.
Zn
+ H2SO4
d.
D.

KBr + Cl2
+ H2
ZnSO4
+ H2
KCl
+ Br2

DOUBLE REPLACEMENT reactions occur when the reactants are two compounds. The
compounds “switch partners”. The cation of each compound switch places. The cation is always
written first.
1. General form:
AB + CD  AD + CB
2. EXAMPLES: (balance them)
a.
H2SO4
b.
Al2(CO3)3
c.
HCl
+ CaCl2
 CaSO4
+ Ca(OH)2 
+ Ca(OH)2

+
HCl
CaCO3 +
CaCl2
+
Al(OH)3
HOH
III. Energy Transfers in Chemical Reactions:
A. Law of Conservation of Energy states that energy is not created nor destroyed. It is only changed
from one form to another.
1. Energy can be stored in bonds and released when those bonds are broken.
2. Energy can be absorbed into bonds from outside.
3. Energy can be in different forms
a. Light
b. Heat-measured in calories or joules.
B.
2 Different Types of Energy Reactions
1. Exothermic reactions involve heat being released (exits). The reaction vessel gets hot.
2. Endothermic reactions involve heat being absorbed (pulled in from the environment). The
reaction vessel gets cold.
IV. Reaction Rate
A. Kinetics is the study of rates of chemical reactions or how fast a reaction occurs.
Some factors that affect the reaction rate:
1. Concentration-defines how much of the substance is present in a given space.
In general, the more concentrated, the faster the rate.
2. Surface area-describes the contact area the two reacting substances have.
Example: stomach vs. small intestine
3. Temperature-the measure of the average kinetic energy of molecules. In general, the more
kinetic energy, the faster the rate.
4. Catalysts – substances that speed up a reaction without becoming part of the reaction or being
changed.
Name_________________________________________Date_________
BALANCING CHEMICAL EQUATIONS ASSIGNMENT UNIT 6
A. Fill in the appropriate coefficient to balance the following equations.
 ____CaCl2 + ____H2 + _____CO2
1. ____HCl + ____CaCO3
2. _____C3H8 + _____O2  _____CO2 + _____H2O
3. _____AlBr3 + _____K2SO4
 _____KBr + _____ Al2(SO4 )3
4. _____H2O2  _____H2O + _____O2
5. _____N2 + _____H2
 _____NH3
B. Rewrite the following word equations as balanced chemical equations. On the left write C for
composition, D for decomposition, SR for single replacement, DR for double replacement.
6. sodium chlorate when heated yields sodium chloride and oxygen gas.
D
2 NaClO3  2 NaCl + 3 O2
7. hydrogen and nitrogen monoxide yields water and nitrogen gas.
8. aluminum and copper (II) chloride yield aluminum chloride and copper.
9. potassium phosphate and calcium chloride yield calcium phosphate and
potassium chloride.
10. iron (III) oxide and carbon monoxide yield iron and carbon dioxide.
C. Finish the word reaction and predict the products of the reactions below. To the left write the type
of reaction.
11. silver nitrate and zinc chloride yields silver chloride and zinc nitrate
DR
2AgNO3 + ZnCl2  Zn(NO3)2 + 2 AgCl
12. aluminum and iron (III) oxide yields __________________________________
13. magnesium bromide and chlorine yields ____________________________________
14. hydrogen and chlorine yields ____________________________
15. calcium and oxygen yields ______________________________
16. sodium and iodine produces ______________________________
17. iron (II) sulfate and ammonium sulfide forms ______________________________________
18. potassium iodide and chlorine yield ___________________________________________
19. barium chloride and sodium sulfate yield ______________________________________
20. zinc chloride and ammonium sulfide yields ___________________________________
21. sodium and water yields ______________________________________________
22. silver and sulfur yields ______________________________________
23. zinc and lead (II) acetate yields __________________________________________
24. lead (II) acetate and hydrogen sulfide produces ______________________________
25. sodium chlorate decomposes into ________________________________________
26. magnesium chlorate forms _____________________________________________
27. aluminum chlorate yields ___________________________________________
28. iron (III) hydroxide yields _______________________________________
29. barium hydroxide yields _____________________________________________
30. lithium hydroxide yields ____________________________________________
31. magnesium carbonate yields ___________________________________________
32. sodium carbonate yields ________________________________________________
33. barium carbonate yields ________________________________________________
34. hydrogen carbonate yields ______________________________________________