This document provides an overview of chemical reactions, including the key parts of a chemical equation, types of reactions such as synthesis, decomposition, and combustion reactions, and how to write and balance chemical equations. It also discusses writing total ionic and net ionic equations by identifying soluble and insoluble reactants and products. The document aims to teach students the essential concepts and skills for understanding and working with chemical reactions.

1. Unit 9:
Chemical Reactions

2. Bellringer 10/15
• List 4 signs that let you know a
chemical reaction has occurred.

3. Objectives
• Follow directions
• Identify a cause and effect
relationship in balancing
chemical reactions
• Balance Chemical reactions
• List the signs of a chemical
reaction

4. Introduction
• Chemical reactions occur when bonds
between the outermost parts of 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
describe compounds, chemical equations
describe a chemical reaction

5. Parts of a Reaction
Equation
• Chemical equations show the conversion of
reactants (the molecules shown on the left of
the arrow) into products (the molecules shown
on the right of the arrow).
• A + sign separates molecules on the same
side
• The arrow is read as “yields”
• Example
C + O2  CO2
• This reads “carbon plus oxygen react to yield
carbon dioxide”

6. Chemical Equations
Their Job: Depict the kind of
reactants and products and their
relative amounts in a reaction.
4 Al (s) + 3 O2 (g) ---> 2 Al2O3 (s)
The numbers in the front are called
Coefficients
The letters (s), (g), and (l) are the
physical states of compounds.

7. Chemical Equations
Continued
• You can indicate the physical state
of a substance by putting a symbol
after each formula.
• (s) – solid
• (l) – liquid
• (g) – gas
• (aq) – aqueous solution (in water)
For example: K(s) + Cl(g) → KCl(s)

8. Conservation of Matter
• Matter cannot be created or
destroyed.
• It can change forms
• Total mass in reactants must
equal total mass of products

9. Terms with Equations
• Activation energy—minimum
energy colliding particles must
have in order to react
• Endothermic reaction—process
that absorbs heat from the
surroundings
• Exothermic reaction—process
that releases heat to the
surroundings

10. Preview/Predict
• Looking at the two rxns, label
one as endothermic & one as
exothermic. Explain why.
SP

11. Catalysts
• A substance
used to speed up
the rate of a
reaction.
• Neither a
product nor a
reactant.
• Written above
the arrow.

12. Symbols Used in
Equations
• Solid ___
• Liquid ___
• Gas ___
• Aqueous solution ___
H2SO4
• Catalyst
• Escaping gas ()
• Change of temperature ()

13. • The charcoal used in a grill is basically carbon. The
carbon reacts with oxygen to yield carbon dioxide. The
chemical equation for this reaction, C + O2  CO2,
contains the same information as the English
sentence but has quantitative meaning as well.

14. Chemical Equations
Because of the principle of the
conservation of matter,
an equation must be
balanced.
It must have the same
number of atoms of the
same kind on both sides.
Lavoisier, 1788

15. 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)

16. 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.

17. Chemical Equations
4 Al(s) + 3 O2(g)
---> 2 Al2O3(s)
This equation means
4 Al atoms + 3 O2 molecules
---produces--->
2 molecules of Al2O3
AND/OR
4 moles of Al + 3 moles of O2
---produces--->
2 moles of Al2O3

18. Steps to Balancing
Equations
There are four basic steps to balancing a chemical equation.
1. 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!
2. 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.
3. 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.
4. Check your answer to see if:
− The numbers of atoms on both sides of the equation are
now balanced.
− The coefficients are in the lowest possible whole
number ratios. (reduced)

19. Some Suggestions to Help You
Some Helpful Hints for balancing equations:
• Take one element at a time, working left to right
except for H and O. Save H for next to last, and O
until last.
• IF everything balances except for O, and there is no
way to balance O with a whole number, double all
the coefficients and try again. (Because O is
diatomic as an element)
• (Shortcut) Polyatomic ions that appear on both
sides of the equation should be balanced as
independent units

21. Balancing Equations
2 2
___ H2(g) + ___ O2(g) ---> ___ H2O(l)
What Happened to the Other Oxygen
Atom?????
This equation is not balanced!
Two hydrogen atoms from a hydrogen
molecule (H2) combines with one of the
oxygen atoms from an oxygen molecule
(O2) to form H2O. Then, the remaining
oxygen atom combines with two more
hydrogen atoms (from another H2 molecule)
to make a second H2O molecule.

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

23. Balancing
Equations
____C3H8(g) + _____ O2(g) ---->
_____CO2(g) + _____ H2O(g)

24. ____B4H10(g) + _____ O2(g) ---->
___ B2O3(g) + _____ H2O(g)

25. Balancing Equations
Sodium phosphate + iron (III) oxide 
sodium oxide + iron (III) phosphate
Na3PO4 + Fe2O3 ---->
Na2O + FePO4

26. Diatomic Elements
• Fluorine • F2
• Chlorine • Cl2
• Bromine • Br2
• Iodine • I2
• Hydrogen • H2
• Nitrogen • N2
• Oxygen • O2

27. Bellringer
• List and describe the five types
of chemical reactions listed in
your book.

28. Types of Reactions
• There are five types of chemical
reactions we will talk about:
1. Synthesis reactions
2. Decomposition reactions
3. Single displacement reactions
4. Double displacement reactions
5. Combustion reactions
• You need to be able to identify
the type of reaction and predict
the product(s)

29. Steps to Writing Reactions
• Some steps for doing reactions
1. Identify the type of reaction
2. Predict the product(s) using the type
of reaction as a model
3. Balance it
Don’t forget about the diatomic
elements! (BrINClHOF) 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 a compound!

30. 1. Combination or Synthesis
• Two or more substances
combine to form a single
substance.
• Forms only ONE PRODUCT!
R + S → RS
For Example: Al + N2 → AlN

31. 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) 

32. 2. Decomposition
• A single compound is broken
down into two or more
products.
• Has only ONE REACTANT!
RS → R + S

33. 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) 

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

35. Bellringer 10/19
• Predict the products of the
following single replacement
reaction, then balance the
reaction, and correctly name
each product:
NaCl(s) + F2(g) 
When you are done pick up your lab
book.

36. Free Response
• IN Your Lab Book:
• Write an outline for the following
prompt:
When exposed to the natural elements
Iron metal reacts with oxygen gas. In
this natural state Iron usually has an
oxidation state of +3. Tell what type
reaction this is, write a balanced
chemical equation, give the oxidation
state for oxygen, and identify what
substance is reduced in this reaction,
and tell why that substance is
reduced.

37. Bellringer
• Pick up a worksheet from the
front desk.
• Tear a sheet of paper in half,
widthwise, and label it January
PreTest
• Read the passages and answer
the questions in the order they
occur (#’s are messed up)

38. 3. Single-Replacement
Reactions
• An element replaces another
element in a compound.
• Whether one metal will displace
another metal is determined by the
activity series of metals chart.
• A reactive metal will replace any
metal listed below it in the activity
series. For example, Mg will replace
Zn.
T + RS → TS + R
When water splits it splits into H & OH

39. Single Replacement
Reactions
• Write and balance the following
single replacement reaction
equation:
• Sodium metal reacts with
aqueous hydrochloric acid
Na(s) + HCl(aq) 
Note: Sodium replaces the
hydrogen ion in the reaction

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

41. 4. Double-Replacement
• Two ionic compounds react by
exchanging cations to form two
different compounds.
• Again, whether one metal will
replace another depends on the
activity series of metals chart.
RS + TU → RU + TS

42. 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) 

43. More on dbl replacement
• Takes place in aqueous
solutions
• Usually produces a precipitate,
gas, or molecular compound
• Aqueous solutions often
disassociate (break apart) in
water

44. Practice
• Predict the products. Balance the
equation
1. HCl(aq) + AgNO3(aq) 
2. CaCl2(aq) + Na3PO4(aq) 
3. Pb(NO3)2(aq) + BaCl2(aq) 
4. FeCl3(aq) + NaOH(aq) 
5. H2SO4(aq) + NaOH(aq) 
6. KOH(aq) + CuSO4(aq) 

45. 5. Combustion
• A hydrocarbon reacts with
oxygen producing energy as
light and heat.
• Hydrocarbons are compounds
composed of C, H, and
sometimes O.
• Always forms carbon dioxide
and water!!
CHO + O2 → CO2 + H2O

46. Combustion Reactions
• Products in combustion
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)

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

48. Bellringer 10/18
• Identify the following reactions
as decomposition, combustion,
single replacement, double
replacement, or combination.
1. T + RS → TS + R
2. CHO + O2 → CO2 + H2O
3. RS → R + S
4. Al + N2 → AlN

49. Bellringer
• give the products
1. T+ + RS →
2. CHO + O2 →
3. RS →
4. Al + N2 →

50. Total Ionic Equations
• Once you write the molecular equation
(synthesis, decomposition, etc.), you
should check for reactants and
products that are soluble or insoluble.
• We usually assume the reaction is in
water
• We can use a solubility table to tell us
what compounds dissolve in water.
• If the compound is soluble (does
dissolve in water), then splits the
compound into its component ions
• If the compound is insoluble (does NOT
dissolve in water), then it remains as a
compound

51. Solubility Table

52. Solubilities Not on the Table!
• Gases only slightly dissolve in water
• Strong acids and bases dissolve in
water
• Hydrochloric, Hydrobromic, Hydroiodic,
Nitric, Sulfuric, Perchloric Acids
• Group I hydroxides (should be on your chart
anyway)
• Water slightly dissolves in water! (H+
and OH-)
• For the homework… SrSO4 is insoluble;
BeI2 and the products are soluble
• There are other tables and rules that
cover more compounds than your table!

53. Total Ionic Equations
Molecular Equation:
K2CrO4 + Pb(NO3)2  PbCrO4 + 2 KNO3
Soluble Soluble Insoluble Soluble
Total Ionic Equation:
2 K+ + CrO4 -2 + Pb+2 + 2 NO3- 
PbCrO4 (s) + 2 K+ + 2 NO3-

54. Net Ionic Equations
• These are the same as total
ionic equations, but you should
cancel out ions that appear on
BOTH sides of the equation
Total Ionic Equation:
2 K+ + CrO4 -2 + Pb+2 + 2 NO3- 
PbCrO4 (s) + 2 K+ + 2 NO3-
Net Ionic Equation:
CrO4 -2 + Pb+2  PbCrO4 (s)

55. Net Ionic Equations
• Shows the molecules broken
apart
AgNO3(aq) + NaCl(aq)  AgCl(s) + NaNO3(aq)
vs.
Ag+ + NO3- + Na+ + Cl-  AgCl(s) + Na+ + NO3-
Which ions are the same on both sides of the
reaction?
These are spectator ions! Na+ NO3-
So, the Net equation is: Ag+ + Cl-  AgCl

56. Net Ionic Equations
• Try this one! Write the molecular, total
ionic, and net ionic equations for this
reaction: Silver nitrate reacts with Lead (II)
Chloride in hot water.
Molecular:
Total Ionic:
Net Ionic:

57. Compounds Solubility
Salts of alkali Soluble
metals &
ammonia Solubility
Nitrate salts Soluble
and chlorate
salts • Precipitates form
Sulfate salts, Soluble when an insoluble
except
compounds solid is formed.
with Pb2+, Ag+,
Hg22+, Ba2+, • This can be
Sr2+, and Ca2+
Chloride salts, Soluble
predicted by using
except the solubility table
compounds
with Pb2+, Ag+ on pg 344!
and, Hg 2+
• Not soluble =
2
Carbonates, Most are
phosphates,
chromates,
insoluble precipitate
sulfides, and
hydroxides

58. Will a Precipitate Form?
•2AgNO3 + H2S  Ag2S + 2 HNO3
(s)
• Yes salts with Ag are not soluble!
• The solid Ag2S will precipitate out!

59. Formation of a precipitate

60. Bellringer
• Using complete sentences tell
the difference between a net
ionic equation and a complete
ionic equation, then give an
example of each.

61. Recall!
1. Write 1 characteristic about
each of the 5 reactions
discussed in class to help you
remember that reaction.

62. Recall
1. Comb/synth: two (single
substances) reactants become one
product
2. Decomp: one reactant breaks into
two products
3. Combust: combines with O2
4. SR: Has 1 single ion plus a
compound on reactants side
5. DR: Has 2 compounds on reactants
side

63. Exit
• What type of practice has
helped you best understand
writing and balancing
equations?