Chemical Reaction (Reactivity Series).ppt

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1. Chemical Reactions

2. Starter
Lets revise on writing a word equation
Potassium hydroxide and sulfuric acid as
reacted. To form potassium sulfate and
water
write the word equation for the reaction

3. 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) + Cl
2
(g)  FeCl
3
(s)
A skeleton equation is not yet “balanced” by coefficients!

4. One more example…
•6 Na (s) + Fe
2O
3 (s)  3 Na
2O (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).

5. TYPES OF CHEMICAL
REACTIONS

6. 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

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

8. Predicting the Products of Single
Replacement Reactions
1) Write the reactants.
2) Identify the cation and anion of the reactant
that is a compound.
3) Use the activity series to see if the single
element will replace one of the elements in
the compound. If no reaction will occur,
just write “NR” for the products and you
are done.
4) Identify the reactant that is the element.
Determine its charge when it becomes an
ion.
5) Perform criss-cross to predict the new
compound on the products side of the
reaction.
6) Write both new products.
7) Balance the reaction.

9. Single Replacement Between
Metals and Water
•Some metals have a higher activity than hydrogen
and can replace it in a single replacement
reaction. In these reactions, you may think of
water (H
2O) as H(OH).
•Ex: Na + H
2O  ?
Na + HOH  ?
Na + H
+
OH
-
 Na
+
OH
-
+ H

2Na + 2H
2O  2NaOH + H
2

10. 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:
AgNO
3
+ NaCl 
AgCl + NaNO
3

11. 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)!

12. Examples of Double Replacement
Reactions:
Pb(NO
3)
2 (aq) + 2NaI (aq)  PbI
2 (s) + 2NaNO
3 (aq)
(precipitate forming)
HCl (aq) + NaOH (aq)  NaCl (aq) + H
2
O (l)
(water-forming, acid-base, neutralization)
CaCO
3
(s) + 2HCl (aq)  CaCl
2
(aq) + H
2
CO
3
(gas-forming)

H(OH)
H
2
O (l) + CO
2
(g)

13. How do you determine if one of the products
of a double replacement reaction will be a
precipitate?
•Use the solubility rules….
Soluble compounds
These compounds break down when put in water.
Example: In water, NaCl  Na
1+
and Cl
1-
.
We say that NaCl…
 has dissolved.
 is soluble.
 forms an aqueous solution (aq).

14. The Solubility Rules
Insoluble compounds
These compounds do NOT
break down when put in
water.
Example: In water, CaCO
3
does NOT break down
into Ca
2+
and CO
3
2-
ions.
The CaCO
3
stays as a
solid, (s) or (ppt).
This is fortunate for many
sea-creatures!
Seashells are made of CaCO
3
!

15. The Solubility Rules
You do not have to memorize these rules,
but you do have to know how to use them
to determine if a product is a precipitate.
See the chart on the next slide…..
Let’s check NaCl and CaCO
3… Are these
compounds soluble or insoluble in
aqueous solution?

16. Solubility Rules Chart

17. Predicting the Products of Double
Replacement Reactions…
Step Example
1) Write the two reactants (both are ionic
compounds)
2) Identify the cations and anions in both of the
compound reactants
3) Pair up each cation with the anion from the
other compound
(i.e. – switch the cations)
4) Write the formula for each product using the
criss-cross method
5) Write the complete equation for the double
replacement reaction
6) Balance the equation.
7) Use the solubility rules chart to figure out which
product is a precipitate (s) and which product
is an aqueous solution (aq). If both products
are (aq) it is really not a reaction.

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

19. SYNTHESIS REACTION
Types of synthesis:
a)Element A + Element B Compound
Na(s) + Cl
2 (g)  2NaCl(s)
a)Element + Compound A  Compound B
O
2(g) + 2SO
2(g)  2SO
3(g)
a)Compound A + Compound B  Compound C
CaO(s) + H
2
O(l)  Ca(OH)
2
(s)

20. Synthesis Reactions (cont’d)
•Metallic and nonmetallic elements react to form ionic
compounds. The resultant compound should be charge
balanced by the criss-cross method.
Ex. 4Li + O
2  2Li
2O
•Nonmetals react with each other to form covalent (molecular)
compounds. You should be able to draw a valid Lewis
Structure for the product.
2H
2
+ O
2
 2H
2
O
or
H
2
+ O
2
 H
2
O
2
But NOT
H
2 + O
2  2OH

21. 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: 2H
2O  2H
2 + O
2
Electrolysis of
Water

22. 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 (H
2
, N
2
, O
2
, CO
2
, 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.

23. The Decomposition of Water by
Electrolysis
2H
2
O  2H
2
+ O
2
An electrical
current can be
used to chemically
separate water into
oxygen gas and
hydrogen gas.
Notice that twice
as much hydrogen
is produced
compared to
oxygen!

24. Electrolysis of Molten Sodium
Chloride Many pure metals are
obtained by using
electrolysis to separate
metallic salts (ex. NaCl
is used to obtain pure
Na).

25. 5) COMBUSTION REACTIONS
a)All involve oxygen (O
2
) as a reactant,
combining with another substance
b)All combustion reactions are are
exothermic
c)Complete combustion of a
hydrocarbon always produces CO
2
and H
2O
d)Incomplete combustion of a
hydrocarbon will produce CO and
possibly C
(black carbon soot) as well
Ex: CH
4
+ 2O
2
=> CO
2
+ 2H
2
O (complete combustion – blue flame)
Ex: CH
4
+ 1.5O
2
=> CO + 2H
2
O (incomplete combustion – yellow flame)
Ex: CH
4
+ O
2
=> C + 2H
2
O (incomplete combustion – yellow flame, soot)

26. Combustion (cont’d)
•Any synthesis reaction which involves O
2
as a
reactant is also considered to be a combustion
reaction!
Ex. 2Mg + O
2
 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!

27. TRY TO CLASSIFY THESE:
1) C
4H
8 + 6O
2  4CO
2 + 4H
2O
2) HCl + NaOH  H
2O + NaCl
3) 2KNO
3(s)  2KNO
2(s) + O
2(g)

28. TRY TO CLASSIFY THESE:
4) 2Ag + S  Ag
2S
5) MgCO
3(s)  MgO(s) + CO
2(g)
6) Cl
2
+ 2KBr  2KCl + Br
2

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