Transcript Chapter 11 Chemical Reactions

Chapter 11
“Chemical
Reactions”
1
Section 11.1
p. 321
Describing Chemical
Reactions
2
All chemical reactions…



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have two parts:
1. Reactants = substances you
start with
2. Products = end with
reactants turn into products
Reactants  Products
- Page 321
Products
Reactants
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In a chem rxn
Atoms not created or destroyed (Law of
Conservation of Mass)
 rxn described in a:

#1. sentence every item is a word
Copper reacts with chlorine to form copper (II)
chloride.
#2. word
equation some symbols used
Copper + chlorine  copper (II) chloride
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Symbols in equations? – Text page 323
 arrow
(→) separates reactants from
products (points to products)
–Read as: “reacts to form” or yields
 plus sign = “and”
 (s) after formula = solid:
Fe(s)
 (g) = gas:
CO2(g)
 (l) = liquid:
H2O(l)
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Symbols used in equations
 (aq) after formula = dissolved in
water, aqueous solution: NaCl(aq)
is salt water solution
 used after product - indicates gas
produced: H2↑
 used after product - indicates
solid produced: PbI2↓
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Symbols used in equations
■
double arrow indicates a
reversible reaction (more later)

heat
■  
shows that
 ,    
heat supplied to rxn
Pt
■   indicates catalyst
supplied (here, platinum is
catalyst)
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What is a catalyst?
substance that speeds up
rxn, w/o being changed or
used up in rxn
 Enzymes - biological or
protein catalysts in your body

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#3. The Skeleton Equation
 Uses formulas and symbols to
describe rxn
–but doesn’t indicate how many;
means they’re NOT balanced
 All chem equations are description
of rxn
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Write a skeleton equation for:
1.
2.
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Solid iron (III) sulfide reacts with
gaseous hydrogen chloride to form
iron (III) chloride and hydrogen
sulfide gas.
Nitric acid dissolved in water reacts
with solid sodium carbonate to form
liquid water and carbon dioxide gas
and sodium nitrate dissolved in
water.
Now, read these equations:
Fe(s) + O2(g)  Fe2O3(s)
Cu(s) + AgNO3(aq)  Ag(s) + Cu(NO3)2(aq)
Pt
NO2(g)  
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N2(g) + O2(g)
#4. Balanced Chemical Equations
 Atoms
can’t be created or destroyed
in an ordinary reaction:
–All atoms we start with we must
end up with (balanced!)
 balanced equation has same # of
each element on both sides of
equation
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Rules for balancing:
1) Assemble correct formulas for all
reactants and products, using “+” and “→”
2) Count # of atoms of each type on both sides
3) Balance elements one at a time by adding
coefficients (the numbers in front) where you
need more - save balancing the H and O
until LAST!
(hint: I prefer to save O until the very last)
4) Double-Check make sure balanced
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Never change subscript (only change
coefficients)
– changing subscript (formula) describes
different chemical
– H2O different than H2O2
 Never put coefficient in middle of formula;
only in front

2NaCl is okay, but Na2Cl is not.
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Practice Balancing Examples
 _AgNO
2
3
 _Mg
3
 _P
4
+ _N2  _Mg3N2
+ _O
5
2  _P4O10
 _Na
2
+ _H
2
2 2O  _H2 + _NaOH
 _CH4
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+ _Cu  _Cu(NO3)2 + 2_Ag
+ _O
2
2 2O
2  _CO2 + _H
Balancing Equations
Balancing Chemical Reactions
Mark Rosengarten – 8:21
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Section 11.2
p. 330
Types of Chemical Reactions
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Types of Reactions
5
major types.
predict the products
 predict whether or not they will happen at all

 How? We recognize them by their reactants
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#1 - Combination Reactions
 Combine
= put together
 2 substances combine to make one
cmpd (also called “synthesis”)
 Ca + O2 CaO
 SO3 + H2O  H2SO4
 predict products, especially if reactants
are 2 elements
Mg3N2 (symbols, charges, cross)
 Mg + N2 _______
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Complete and balance:
+ Cl2 
 Fe + O2  (assume iron (II) oxide is the product)
 Al + O2 
 Ca
first step…write correct
formulas – you can still change
subscripts at this point, but not while
balancing!
 Then balance by changing just
coefficients only
 Remember
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#1 – Combination Reactions
 Additional
Notes:
a) Some nonmetal oxides react
with H2O - produces acid:
SO2 + H2O  H2SO3
(how “acid rain” forms)
b) Some metallic oxides react with
H2O - produces base:
CaO + H2O  Ca(OH)2
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#2 - Decomposition Reactions
 decompose
= fall apart
 one reactant breaks apart into 2 or
more elements or cmpds
electricity
 NaCl   
Na + Cl2

 CaCO3  
CaO + CO2
 Note:
energy (heat, sunlight,
electricity, etc.) usually required
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#2 - Decomposition Reactions
 predict
products if binary cmpd
(made of 2 elements)
–It breaks apart into the elements:
electricity
 H2O   
 HgO  

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#2 - Decomposition Reactions
 If
cmpd has > 2 elements you must
be given one of products
–other product from the missing
pieces
 NiCO3  
CO2 + ___
 H2CO3(aq) CO2 + ___
heat
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#3 - Single Replacement Reactions
 One
element replaces another (new
dance partner)
 Reactants
must be an element &
cmpd
 Products will be a different element
and different cmpd
 Na + KCl  K + NaCl (Cations switched)
 F2 + LiCl  LiF + Cl2 (Anions switched)
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#3 Single Replacement Reactions
Metals replace other metals (they can
also replace H)
 K + AlN 
 Zn + HCl 
 Think of water as: HOH
– Metals replace first H, then combines
w/ hydroxide (OH).
 Na + HOH 

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#3 Single Replacement Reactions
 can
even tell whether or not single
replacement rxn will happen:
–b/c some chemicals more “active” than
others
–More active replaces less active
 list
– p. 333 - Activity
Series of
Metals
 Higher
28
on list replaces lower
The “Activity Series” of Metals
Higher
activity
Lower
activity
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Lithium
Potassium
Calcium
Sodium
Magnesium
Aluminum
Zinc
Chromium
Iron
Nickel
Lead
Hydrogen
Bismuth
Copper
Mercury
Silver
Platinum
Gold
1) Metals can replace other
metals, if they are above
metal trying to replace
(i.e. Zn will replace Pb)
2) Metals above H can replace
H in acids.
3) Metals from Na upward can
replace hydrogen in H2O
The “Activity Series” of Halogens
Higher Activity
Fluorine
Chlorine
Bromine
Iodine
Halogens can replace other
halogens in compounds, if
they are above halogen they
are replacing
Lower Activity
2NaCl(s) + F2(g) 
MgCl2(s) + Br2(g) 
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2NaF
??? (s) + Cl2(g)
???Reaction!
No
#3 Single Replacement Reactions
Practice:

Fe + CuSO4 

Pb + KCl 

Al + HCl 
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#4 - Double Replacement Reactions

Two things replace each other.
– Reactants must be two ionic
compounds, in aqueous solution
NaOH + FeCl3 
– positive ions change place (dance partners)
 NaOH + FeCl3 Fe+3 OH- + Na+1 Cl-1
= NaOH + FeCl3 Fe(OH)3 + NaCl

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#4 - Double Replacement Reactions
 Have
certain “driving forces”, or reasons
–only happens if one product:
a) doesn’t dissolve in water & forms
solid (a “precipitate”), or
b) is gas that bubbles out, or
c) is molecular compound (usually
water)
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Complete and balance:
 assume all of the following
reactions actually take place:
CaCl2 + NaOH 
CuCl2 + K2S 
KOH + Fe(NO3)3 
(NH4)2SO4 + BaF2 
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How to recognize which type?
 Look at the reactants:
E + E = Combination
C
= Decomposition
E + C = Single replacement
C + C = Double replacement
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Practice Examples:
+ O2 
 H2O 
 Zn + H2SO4 
 HgO 
 KBr + Cl2 
 AgNO3 + NaCl 
 Mg(OH)2 + H2SO3 
 H2
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#5 – Combustion Reactions
 Combustion
means “add oxygen”
 Normally, a cmpd composed of only
C, H, (and maybe O) is reacted with
oxygen – called “burning”
 Complete combustion, products are
CO2 and H2O
 If incomplete, products are CO (or
possibly just C) and H2O
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Combustion Reaction Examples:
 C4H10
+ O2  (assume complete)
 C4H10
+ O2  (incomplete)
 C6H12O6
 C8H8
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+ O2  (complete)
+ O2  (incomplete)
SUMMARY: An equation...
 Describes
a rxn
 Must be balanced (follows the Law of
Conservation of Mass)
 only balance by changing coefficients
 special symbols to indicate physical
state, catalyst or energy required, etc.
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Reactions
5
major types
 We can tell what type they are by
looking at reactants
 Single Replacement happens based on
the Activity Series
 Double Replacement happens if one
product is: 1) a precipitate (an insoluble
solid), 2) water (a molecular compound), or 3) a gas
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Section 11.3
p. 342
Reactions in Aqueous Solution
NiCl2
Co(NO3)2
K2Cr2O7
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K2CrO4
CuSO4
KMnO4
Net Ionic Equations
 Many
reactions occur in water- that
is, in aqueous solution
 When dissolved in water, many
ionic cmpds “dissociate”, or
separate, into cations & anions
 Now write ionic equation
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Net Ionic Equations

Example (needs to be a double replacement reaction)
AgNO3 + NaCl  AgCl + NaNO3
1. this is the full balanced equation
2. next, write it as ionic equation by
splitting the cmpds into their ions:
Ag1+ + NO31- + Na1+ + Cl1- 
AgCl + Na1+ + NO31Note that the AgCl did not ionize, because it is a
“precipitate” (Table 11.3 p. 344)
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Net Ionic Equations
3. simplify by crossing out ions not
directly involved (called spectator ions)
Ag1+ + Cl1-  AgCl
This is called the net ionic equation
Let’s talk about precipitates before we
do some other examples
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Predicting the Precipitate


Insoluble salt is a precipitate
[note Figure 11.11, p.342 (AgCl)]
General solubility rules are found:
a) Table 11.3, p. 344 in textbook
b) Reference section - page R54
(Table B.9)
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Let’s do some examples
together of net ionic
equations, starting with
these reactants:
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BaCl2 + AgNO3 →
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NaCl + Ba(NO3)2 →
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Pb(NO3)2(aq) + H2SO4(aq) 
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