This document provides an overview of chemical reactions, defining key concepts such as reactants, products, and various types of reactions including single replacement, double replacement, synthesis, decomposition, and combustion. It explains how to write word equations and skeleton equations, and emphasizes the conservation of mass in reactions, along with guidelines for balancing them. The document also discusses evidence for chemical reactions and the activity series that determines whether certain reactions occur.

1. Introduction to Chemical
Reactions

2. What is a Chemical Reaction?
It is a chemical change in which one or
more substances are destroyed and one
or more new substances are created.
BEFORE
H2 gas
and
O2 gas
AFTER
H2O liquid

3. Parts of a Chemical Reaction
Reactants  Products
Reactants: Substances that are destroyed by the
chemical change (bonds break
break).
Products: Substances created by the chemical
change (new bonds form
form).
The arrow () is read as “yields”.

4. Other symbols in chemical
reactions
• (s) = solid
• (l) = liquid
• (g) = gas
• (aq) = aqueous solution (the substance is
dissolved in H2O)
• “+” separates two or more reactants or
products
• “” yield sign separates reactants from
products

5. Evidence for a Chemical Reaction
1) Evolution of light or heat.

6. Evidence for a Chemical Reaction
2) Temperature change (increase or
decrease) to the surroundings.

7. Evidence for a Chemical Reaction
3) Formation of a gas (bubbling or an odor)
other than boiling.

8. Evidence for a Chemical Reaction
4) Color change (due to the formation of a
new substance).

9. Evidence for a Chemical Reaction
5) Formation of a precipitate (a new solid
forms) from the reaction of two aqueous
solutions.

10. Word Equations
• Statements that indicate the reactants and
products in a chemical reaction.
• Ex. Iron (s) + chlorine (g)  iron (III) chloride (s)
• This is read as:
“Solid iron and chlorine gas react (combine) to produce
solid iron (III) chloride”

11. Translating Word Equations to
Skeleton Equations
• A skeleton equation uses chemical formulas
rather than words to identify the reactants and
products of a chemical reaction.
• The word equation
Iron (s) + chlorine (g)  iron (III) chloride (s)
• The skeleton equation
Fe(s) + Cl2(g)  FeCl3 (s)
A skeleton equation is not yet “balanced” by coefficients!

12. One more example…
• 6 Na (s) + Fe2O3 (s)  3 Na2O (s) + 2 Fe (s)
– The numbers preceding the chemical formulae are
coefficients. They are used to balance the reaction.
– The numbers within the chemical formulae are
subscripts.
– You can read the above balanced reaction as:
• “6 atoms of solid sodium plus 1 formula unit of solid
iron (III) oxide yields 3 formula units of solid sodium
oxide and 2 atoms of solid iron” or…
• “6 moles of solid sodium plus 1 mole of solid iron (III)
oxide yields 3 moles of solid sodium oxide plus 2
moles of solid iron”
• Chemical reactions can never be read in terms of
grams, only in terms of particles or groups of particles
(moles).

13. TYPES OF CHEMICAL
REACTIONS

14. There are 5 basic types….
• Single Replacement (Displacement)
(Redox)
• Double Replacement (Displacement)
(Metathesis)
• Synthesis (Combination)
• Decomposition
• Combustion

15. A single uncombined
element replaces
another element in
an ionic compound.
There are two
reactants and two
products.
1) SINGLE REPLACEMENT
REACTION
Ex: Zn + CuSO4  ZnSO4 + Cu

16. Single Replacement Reactions
Single replacement reactions have the
general form, A + BC  AC + B.
Question: Do all single replacement
reactions actually occur?
Answer: Not necessarily…

17. Single Replacement Reactions
Examine the reaction:
Zn + CuSO4  ZnSO4 + Cu
This reaction does occur!’
Now let’s try:
Cu + ZnSO4  No Reaction
Conclusion: Zn will replace Cu in
solution, but not vice versa!

18. Single Replacement Reactions
How do we know which reactions will occur
and which ones will not?
We look at the “activity series”.
Elements with higher activities replace
elements with lower activities during a
single-replacement reaction, but not vice-
versa.

19. HIGHEST ACTIVITY
Li
Rb
K
Ba
Ca
Na
Mg
Al
Mn
Zn
Cr
Fe
Ni
Sn
Pb
H
Cu
Hg
Ag
Pt
Au
LOWEST ACTIVITY
Activity Series for
Metals

20. Activity Series for Nonmetals
Highest Activity
F
Cl
Br
I
Lowest Activity

21. Parts of two
aqueous ionic
compounds switch
places to form two
new compounds.
There are two
reactants and two
products.
2) DOUBLE REPLACEMENT
REACTION
Example:
AgNO3 + NaCl 
AgCl + NaNO3

22. Double Replacement Reactions
The general form of a double replacement reaction is:
AB + CD  AD + CB
Just like single replacement reactions, not all double
replacement reactions actually occur.
We can experimentally attempt a D.R. reaction. The
reaction occurs if:
1) A solid precipitate is produced, or
2) A gas is produced, or
3) Water is produced.
If none of the above are produced and both products are
(aq), then there is no reaction (NR)!

23. Examples of Double Replacement
Reactions:
Pb(NO3)2 (aq) + 2NaI (aq)  PbI2 (s) + 2NaNO3 (aq)
(precipitate forming)
HCl (aq) + NaOH (aq)  NaCl (aq) + H2O (l)
(water-forming, acid-base, neutralization)
CaCO3 (s) + 2HCl (aq)  CaCl2 (aq) + H2CO3
(gas-forming)
H(OH)
H2O (l) + CO2 (g)

24. Two or more simple substances
(the reactants) combine to form
a more complex substance (the
product).
3) SYNTHESIS REACTION
Ex: 2Mg + O2
 2MgO

25. SYNTHESIS REACTION
Types of synthesis:
a)Element A + Element B Compound
Na(s) + Cl2 (g)  2NaCl(s)
a)Element + Compound A  Compound B
O2(g) + 2SO2(g)  2SO3(g)
a)Compound A + Compound B  Compound C
CaO(s) + H2O(l)  Ca(OH)2 (s)

26. A more complex substance (the
reactant) breaks down into two
or more simple parts (products).
Synthesis and decomposition
reactions are opposites.
4) DECOMPOSITION REACTION
Ex: 2H2O  2H2 + O2
Electrolysis of
Water

27. DECOMPOSITION REACTIONS
(Cont’d)
Decomposition of a compound produces two or
more elements and/or compounds
The products are always simpler than the
reactant.
Gases are often produced (H2, N2, O2, CO2, etc.)
in the decomposition of covalent compounds.
Ionic compounds may be decomposed into pure
elements by using electricity (electrolysis). This is
how pure metals are obtained from salts.

28. 5) COMBUSTION REACTIONS
a) All involve oxygen (O2) as a reactant,
combining with another substance
b) All combustion reactions are are
exothermic
c) Complete combustion of a
hydrocarbon always produces CO2
and H2O
d) Incomplete combustion of a
hydrocarbon will produce CO and
possibly C (black carbon soot) as well
Ex: CH4 + 2O2 => CO2 + 2H2O (complete combustion – blue flame)
Ex: CH4 + 1.5O2 => CO + 2H2O (incomplete combustion – yellow flame)
Ex: CH4 + O2 => C + 2H2O (incomplete combustion – yellow flame, soot)

29. Combustion (cont’d)
• Any synthesis reaction which involves O2 as a
reactant is also considered to be a combustion
reaction!
Ex. 2Mg + O2  2MgO
(metal oxide)
This is called the combustion of magnesium or
the synthesis of magnesium oxide. The
combustion of a metal always produces a metal
oxide (in this case, magnesium oxide). Make
sure the metal product is criss-crossed
correctly!

30. TRY TO CLASSIFY THESE:
1) C4H8 + 6O2  4CO2 + 4H2O
2) HCl + NaOH  H2O + NaCl
3) 2KNO3(s)  2KNO2(s) + O2(g)

31. TRY TO CLASSIFY THESE:
4) 2Ag + S  Ag2S
5) MgCO3(s)  MgO(s) + CO2(g)
6) Cl2 + 2KBr  2KCl + Br2

32. Check Your Answers…
1) Combustion (of a hydrocarbon)
2) Double replacement (water forming)
3) Decomposition
4) Synthesis
5) Decomposition
6) Single Replacement

33. Conservation of Mass
During a chemical reaction, atoms are neither
created nor destroyed (Conservation of
Mass).
Hydrogen and oxygen gas react to form
water:
H2 (g) + O2 (g)  H2O (l)

34. Conservation of Mass
H2 (g) + O2 (g)  H2O (l)
What is wrong with this equation above? Doesn’t
it appear that one oxygen atom “went missing”?
According to conservation of mass, the proper way
to write this reaction is:
2H2 (g) + 1O2 (g)  2H2O (l)
The red coefficients represent the # of molecules
(or the # of moles) of each reactant or product.

35. Not All Properties are Conserved
During Chemical Reactions!
CONSERVED NOT CONSERVED
Mass
Types of atoms
Number of each atom
Color
Physical state (solid,
liquid, gas)
Volume
Number of moles of
reactants/products

36. Counting Atoms
SnO2 + 2H2 → Sn + 2H2O
SUBSCRIPT COEFFICIENT

37. Rules for Balancing
Chemical Reactions
__H2 + __ O2  __H2O
Balancing is about finding the
right coefficients!

38. Rules for Balancing
Chemical Reactions
1) You can change the coefficients, but
NEVER the subscripts!
__H2 + __ O2  __H2O
Off Limits!

39. Rules for Balancing
Chemical Reactions
2) The coefficients must reduced to
represent the lowest possible numbers.
4H2 + 2 O2  4H2O

40. Rules for Balancing
Chemical Reactions
3) It is OK to use fraction coefficients, but
you must get rid of them in the end
(multiply through by denominator).
H2 + ½ O2  H2O

41. Rules for Balancing
Chemical Reactions
4) Often, it is helpful to save the following
elements until the end (do other
elements first):
H, C, O

42. Rules for Balancing
Chemical Reactions
5) Do a final balance check for each
element!
2H2 + O2  2H2O

43. Practice
1) K + Br  KBr
2) HgO  Hg + O2
3) Na + H2O  NaOH + H2

44. Practice
4) CaO + H2O  Ca(OH)2
5) Al + HCl  AlCl3 + H2