Chemical_Reactions.ppt

it's all about chemical reaction

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
H
2 gas
and
O
2gas
AFTER
H
2Oliquid

3. Parts of a Chemical Reaction
Reactants Products
Reactants: Substances that are destroyed by the
chemical change (bonds break).
Products: Substances created by the chemical
change (new bonds 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 H
2O)
•“+” 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 equationuses 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!

12. One more example…
•6Na (s) + Fe
2O
3(s) 3Na
2O (s) + 2Fe (s)
–The numbers preceding the chemical formulae are
coefficients. They are used to balancethe 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 neverbe read in terms of
grams, only in terms of particles or groups of particles
(moles).

13. Conservation of Mass
During a chemical reaction, atoms are neither
created nor destroyed (Conservation of
Mass).
Hydrogen and oxygen gas react to form
water:
H
2(g)+ O
2 (g)H
2O
(l)

14. Conservation of Mass
H
2(g)+ O
2 (g)H
2O
(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:
2H
2(g)+ 1O
2 (g)2H
2O
(l)
The redcoefficients represent the # of molecules
(or the # of moles) of each reactant or product.

15. 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

16. TYPES OF CHEMICAL
REACTIONS

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

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

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

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

21. 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 activitiesreplace
elements with lower activitiesduring a
single-replacement reaction, but notvice-
versa.

22. 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

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

24. 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.

25. 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

26. 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

27. 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 actuallyoccur.
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 aboveare produced and both products are
(aq), then there is no reaction (NR)!

28. 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
2O (l)
(water-forming, acid-base, neutralization)
CaCO
3(s) + 2HCl (aq) CaCl
2(aq) + H
2CO
3
(gas-forming)
H(OH)
H
2O (l) + CO
2 (g)

29. 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).

30. 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
3stays as a
solid, (s) or (ppt).
This is fortunate for many
sea-creatures!
Seashells are made of CaCO
3!

31. 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?

32. Solubility Rules Chart

33. 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.

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

35. 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
2O(l) Ca(OH)
2(s)

36. Synthesis Reactions (cont’d)
•Metallic and nonmetallic elements react to form ionic
compounds. The resultant compound should be charge
balancedby 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 Structurefor the product.
2H
2+ O
22H
2O
or
H
2+ O
2H
2O
2
But NOT
H
2+ O
22OH

37. 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

38. DECOMPOSITION REACTIONS
(Cont’d)
Decomposition of a compound produces two or
moreelements and/or compounds
The products are always simplerthan the
reactant.
Gasesare 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.

39. The Decomposition of Water by
Electrolysis
2H
2O 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!

40. 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).

41. 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
hydrocarbonalways produces CO
2
and H
2O
d)Incomplete combustion of a
hydrocarbonwill produce CO and
possibly C(black carbon soot) as
well
Ex: CH
4+ 2O
2=> CO
2+ 2H
2O (complete combustion –blue flame)
Ex: CH
4+ 1.5O
2=> CO + 2H
2O(incomplete combustion –yellow flame)
Ex: CH
4+ O
2=> C + 2H
2O (incomplete combustion –yellow flame, soot)

42. Combustion (cont’d)
•Anysynthesis reaction which involves O
2as a
reactant is also considered to be a combustion
reaction!
Ex. 2Mg + O
22MgO
(metal oxide)
This is called the combustionof magnesiumor
the synthesisof 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!

43. 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)

44. 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

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

46. Counting Atoms
SnO
2+ 2H
2→ Sn + 2H
2O
SUBSCRIPT COEFFICIENT

47. Rules for Counting Atoms
1)Coefficients propagate to the right through the
entire compound, whether or not parentheses
are present.
2) Subscripts affect only the element to the left of
the subscript, unless…
3) If a subscript occurs to the right of a
parentheses, the subscript propagates to the left
through the parentheses.
4) When a coefficient and subscript “meet”, you
must multiply the two.

48. Examples of Counting Atoms
SnO
2+ 2H
2→ Sn + 2H
2O
2 C
4H
10+ 13 O
2→ 8 CO
2+ 10 H
2O
Cu + 2AgNO
3→ Cu(NO
3)
2+ 2Ag
3Pb(NO
3)
2+ 2AlCl
3→ 3PbCl
2+ 2Al(NO
3)
3

49. Classwork
Complete “The Count” worksheet
on counting atoms in chemical
reactions.

50. Warm-Up
2Ca
3(PO
4)
2+ 6 SiO
2+ 10C 
6 CaSiO
3+ P
4+10CO
Atom # Atoms on
Left Side
# Atoms on
Right Side
Ca
P
O
Si
C

51. Rules for Balancing
Chemical Reactions
__H
2+ __ O
2__H
2O
Balancing is about finding the
right coefficients!

52. Rules for Balancing
Chemical Reactions
1)You canchange the coefficients, but
NEVER the subscripts!
__H
2+ __ O
2__H
2O
Off Limits!

53. Rules for Balancing
Chemical Reactions
2) The coefficients must reduced to
represent the lowest possible numbers.
4H
2+ 2O
24H
2O

54. 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).
H
2+ ½ O
2H
2O

55. 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

56. Rules for Balancing
Chemical Reactions
5) Do a final balance check for each
element!
2H
2+ O
22H
2O

57. Practice
1)K + Br KBr
2)HgO Hg + O
2
3) Na + H
2O NaOH + H
2

58. Practice
4) CaO + H
2O Ca(OH)
2
5) Al + HCl AlCl
3+ H
2

59. Energy Changes Accompanying
Chemical Reactions
All chemical reactions involve a net release or absorption of
energy. Therefore, heat energy moves between the
chemical system and the surroundings. This exchange of
heat can be monitored by keeping track of changes in
temperature of the surroundings (calorimetry).
Remember, q = mc
pT
where q = change in heat (in Joules)
m = mass of H
2O (in grams)
c
p= specific heat capacity of
H
2O (J/g ◦C )
T = change in temperature
of H
2O (in ◦C)

60. Where does the energy come from
during a chemical reaction?
•During chemical reactions, bonds are broken and new bonds
are formed.
•The heat energy that moves between the system and
surroundings during chemical reactions is basically the energy
that is used to break bonds and the energy that is released
when bonds form. (i.e. bond energy)
•The energy changethat accompanies any chemical reaction is
called the enthalpy (heat) of reaction or H
0
rxn.
H
0
rxn= H
final–H
initial
•H
0
simply means that the energy changes during chemical
reactions are generally measured at “standard state” conditions
of 298 K (25
◦
C) and 1 atm pressure.
•It is important to note that absolute amounts of energy
within a chemical system cannot be measured. We can
only measure changes in energy within a chemical system.
Hence we use the “” sign.

61. Exothermic Reactions
A chemical reaction is exothermicif energy is given off by the system to
the surroundings (the energy exits):
Reactants Products + Energy Released
The temperature of the surroundings (including the temperature probe)
increases during exothermic reactions because the system releases
energy. The H
0
rxnis negativebecause H
finalis less than H
initial. In other
words, the system lost energy.(sign goes with the system)
The majorityof chemical reactions are exothermic because nature favors a
low chemical potential energy.
System
Surroundings
reactants
products
Chemical Potential
Energy (H)
Reaction progress
H
rxnis (-)

62. Example: An Exothermic Reaction
The “Smashing” Thermite Reaction:
2Al(s) + Fe
2O
3(s) 2Fe (s) + Al
2O
3(s)
Reaction Progress
Chemical Potential
Energy (H)

63. Endothermic Reactions
A chemical reaction is endothermicif energy is absorbed by the system
from the surroundings (the energy enters):
Reactants + Energy Absorbed Products
The temperature of the surroundings (including the temperature probe)
decreases during endothermic reactions because the system absorbs
energy. The H
0
rxnis positivebecause H
finalis more than H
initial. In other
words, the system gained energy. (sign goes with the system)
Endothermic chemical reactions are generally unfavorable but may occur
only if they are accompanied by an increase in entropyor disorder of the
system (due to more particles formed, liquids/gases formed, mixtures
formed, volume of gas increases).
System
Surroundings
reactants
products
Chemical Potential
Energy (H)
Reaction progress
H
rxnis (+)

64. Example: An Endothermic Reaction
Ba(OH)
28H
2O (s) + 2NH
4(NO
3) (s) 
Ba(NO
3)
2(aq) + 2NH
3(g) + 10 H
2O (l)
Reaction Progress
Chemical Potential
Energy (H)

65. Do you have to actually perform and
observe a chemical reaction to know if it is
exothermic or endothermic?
•No –you can calculate H
0
rxnfrom data that has
already been measured and tabulated by
thermo-chemists (see handout).
•H
0
f= standard heat of formation for a compound
(in kJ/mol). It is determined by forming the
compound from its elements in their stable forms
at conditions of 298K and 1 atm of pressure
inside of a calorimeter.
•For most compounds, H
0
fis negative because
bond formation is exothermic!
•H
0
fof an element is always 0 kJ/mol by def.

66. H
0
rxn= nH
0
f(products) -nH
0
f(reactants)
•Not as hard as it looks 
•Basically, you just
1) multiply the coefficient of each product times its
standard heat of formation and add together for all
products
2) multiply the coefficient of each reactant times its
standard heat of formation and add together for all
reactants
3) take the difference of 1 and 2
(always products -reactants)
4) If the difference is (-) the reaction is exothermic;
if the difference is (+) the reaction is endothermic.

67. Try this…
•Calculate the H
0
rxnfor the thermite reaction
using tabulated data (see handout):
2Al (s) + Fe
2O
3(s) 2Fe (s) + Al
2O
3(s)
H
0
rxn= nH
0
f(products) -nH
0
f(reactants)

68. Try this…
•Calculate the H
0
rxnfor this reaction based on
tabulated data:
Ba(OH)
28H
2O (s) + 2NH
4(NO
3) (s) 
Ba(NO
3)
2(aq) + 2NH
3(g) + 10 H
2O (l)
H
0
rxn= nH
0
f(products) -nH
0
f(reactants)
Compound H
0
f
(kcal/mol)
NH
4(NO
3) (s) -87.73
Ba(OH)
28H
2O (s) -798.8
Ba(NO
3)
2(aq) -227.62
NH
3(g) -11.02
H
2O (l) -68.32
1 kcal = 4.184 kJ

69. Summarizing H
0
rxn
•If H
0
rxn is (-) the reaction is exothermic and the
bonds formedare strongerand more stable
than the bonds broken.
•If H
0
rxn is (+) the reaction is endothermic and
the bonds formedare weakerand less stable
than the bonds broken. However, if the entropy
of the system has increased to sufficiently to
counteract this increase in enthalpy, then the
reaction can still occur.

70. Bond Enthalpies
•Another way to determining an enthalpy
change (H
0
rxn) for a chemical reaction is to
compute the difference in bond enthalpies
between reactants and products
•The energy to required to break a covalent
bond in the gaseous phase is called a bond
enthalpy (bond dissociation energy).
•Bond enthalpy tables give the average
energy to break a chemical bond. Actually
there are slight variations depending on the
environment in which the chemical bond is
located

71. Bond Enthalpy Table
The average bond enthalpies for several types of
chemical bonds are shown in the table below:

72. Bond Enthalpies
•Bond enthalpies can be used to calculate the
enthalpy change (H
0
rxn)for a chemical
reaction.
•Energy is required to break chemical bonds.
Therefore when a chemical bond is broken
its enthalpy changecarries a positive sign.
•Energy is released when chemical bonds
form. When a chemical bond is formed its
enthalpy changeis expressed as a negative
value.
•By combining the enthalpy required and the
enthalpy released for the breaking and
forming chemical bonds, one can calculate
the overall enthalpy change for a chemical
reaction.

73. Bond Enthalpy Calculations
Example : Calculate the enthalpy change (H
0
rxn)
for the reaction N
2+ 3 H
22 NH
3
Bonds broken (energy in)
1N≡N: = 945
3H-H: 3(435) = 1305
Total = 2250 kJ/mol
Bonds formed (energy out)
2x3 = 6 N-H: 6 (390) = -2340 kJ/mol
H
0
rxn= [energy used for breaking bonds] + [energy released in forming bonds]
Net enthalpy change (H
0
rxn)
= + 2250 + (-2340) = -90 kJ/mol (exothermic reaction)
H -H
H -H
H -H
You may have
to draw a
Lewis
Structure to
know what
type of bonds
are present!

74. Another Way to Think About It
Chemical Potential Energy (H) of System
Start
+ 2250 kJ/mol
(energy in
when bonds
break)
-2340 kJ/mol
(energy out
when bonds
form)
H
0
rxn= -90 kJ/mol (net)
released by the system to the
surroundings