Transcript Chapter 8 Chemical Reactions
1 Chapter 7 Chemical Reactions
2 Section 7.1
Describing Chemical Change
OBJECTIVES:
–
Write equations describing chemical reactions, using appropriate symbols
3 Section 7.1
Describing Chemical Change
OBJECTIVES:
–
Write balanced chemical equations, when given the names or formulas of the reactants and products in a chemical reaction.
4 All chemical reactions
have two parts:
–
Reactants - the substances you start with
–
Products - the substances you end up with
The reactants turn into the products.
Reactants
Products
5 In a chemical reaction
The way atoms are joined is changed
Atoms aren’t created of destroyed.
Can be described several ways: 1. In a sentence Copper reacts with silver nitrate to form silver and copper (II) nitrate .
2. In a word equation Copper + silver nitrate
silver + copper (II) nitrate
6 Or a skeleton equation
Cu ( ) + AgNO 3 ( )
Ag ( ) + Cu(NO 3 ) 2 ( )
reactants Or a balanced equation products
7 Symbols in equations-p.144
the arrow separates the reactants from the products
Read “reacts to form”
The plus sign = “and”
(s) after the formula = solid
(g) after the formula = gas
(l) after the formula = liquid
8 Symbols used in equations
(aq) after the formula - dissolved in water, an aqueous solution.
used after a product indicates a gas (same as (g))
used after a product indicates a solid (same as (s))
9 Symbols used in equations
indicates a reversible reaction (more later)
, shows that heat is supplied to the reaction
is used to indicate a catalyst is supplied, in this case, platinum.
10 What is a catalyst?
A substance that speeds up a reaction, without being changed or used up by the reaction.
Enzymes are biological or protein catalysts.
11 Skeleton Equation
Uses formulas and symbols to describe a reaction
doesn’t indicate how many.
All chemical equations are sentences that describe reactions.
12 Convert these to equations
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.
13 Now, read these:
Fe(s) + O 2 (g)
Fe 2 O 3 (s)
Cu(s) + AgNO 3 (aq)
Ag(s) + Cu(NO 3 ) 2 (aq)
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NO 2 (g) N 2 (g) + O 2 (g)
15 Balancing Chemical Equations
16 Balanced Equation
Atoms can’t be created or destroyed
All the atoms we start with we must end up with
A balanced equation has the same number of each element on both sides of the equation.
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C + O O
O C
C + O 2
CO 2
This equation is already balanced
What if it isn’t?
O
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C + O O
C O
C + O 2
CO
We need one more oxygen in the products.
Can’t change the formula, because it describes what it is (carbon monoxide in this example)
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C + O O
C O C O
Must be used to make another CO
But where did the other C come from?
C C + O O
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Must have started with two C
2 C + O 2
2 CO C O C O
21 Rules for balancing:
Assemble, write the correct formulas for all the reactants and products
Count the number of atoms of each type appearing on both sides Balance the elements one at a time by adding coefficients (the numbers in front) - save H and O until LAST!
E lement C arbon H ydrogen O xygen
Check to make sure it is balanced.
Never change a subscript to balance an equation.
–
If you change the formula ( subscripts ) you are describing a different reaction.
–
H 2 O is a different compound than H 2 O 2 Never put a coefficient in the middle of a formula
–
2 NaCl is okay, Na 2 Cl is not.
22
H 2 +
Example
O 2
H 2 O
23 Make a table to keep track of where you are at
H 2 +
Example
O 2
H 2 O R 2 2 H O P 2 1
Need twice as much O in the product 24
Changes the O 25
H 2 +
Example
O 2
2 H 2 O R 2 2 H O P 2 1
H 2 +
Example
O 2
2 H 2 O R 2 2 H O P 2 1 2
Also changes the H 26
H 2 +
Example
O 2
2 H 2 O R 2 2 H O P 2 1 4 2
Need twice as much H in the reactant 27
Recount 28
2 H 2 +
Example
O 2
2 H 2 O R 2 2 H O P 2 1 4 2
2 H 2 +
Example
O 2
2 H 2 O 4 R 2 2 H O P 2 1 4 2
29 The equation is balanced, has the same number of each kind of atom on both sides
30
2 H 2 +
Example
O 2
2 H 2 O 4 R 2 2 H O P 2 1 4 2
This is the answer Not this
31 Balancing Examples
_ AgNO 3 + _Cu
_Cu(NO 3 ) 2 _Mg + _N 2 _P + _O 2
_Mg _P 4 O 3 10 N 2 _Na + _H 2 O
_CH 4 + _O 2
_H 2 _CO 2 + _NaOH + _H 2 O + _Ag
32 Section 7.2
Types of Chemical Reactions
OBJECTIVES:
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Identify a reaction as combination, decomposition, single-replacement, double-replacement, or combustion
33 Section 7.2
Types of Chemical Reactions
OBJECTIVES:
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Predict the products of combination, decomposition, single-replacement, double-replacement, and combustion reactions.
Types of Reactions
There are millions of reactions.
Can’t remember them all
Fall into several categories.
We will learn 5 major types.
Will be able to predict the products.
For some, we will be able to predict whether they will happen at all.
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Will recognize them by the reactants
35 #1 - Combination Reactions
Combine - put together (synthesis)
2 substances combine to make one compound.
Ca +O 2
CaO SO 3 + H 2 O
H 2 SO 4 We can predict the products if they are two elements.
Mg + N 2
36 Write and balance
Ca + Cl 2 Fe + O 2
Al + O 2
iron (II) oxide
Remember that the first step is to write the correct formulas
Then balance by using coefficients only
37 #2 - Decomposition Reactions
decompose = fall apart
one reactant falls apart into two or more elements or compounds.
NaCl Na + Cl 2
CaCO 3
CaO + CO 2
Note that energy is usually required to decompose
38 #2 - Decomposition Reactions
Can predict the products if it is a binary compound
Made up of only two elements
Falls apart into its elements
H 2 O
HgO
39 #2 - Decomposition Reactions
If the compound has more than two elements you must be given one of the products
The other product will be from the missing pieces
NiCO 3 CO 2 + ? H 2 CO 3 (aq)
CO 2 + ?
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The Activity Series of the Metals Lithium Potassium Calcium Sodium Magnesium Aluminum Zinc Chromium Iron Nickel Lead Hydrogen Bismuth Copper Mercury Silver Platinum Gold Metals can replace other metals provided that they are above the metal that they are trying to replace. Ex. Metals above hydrogen can replace hydrogen in acids.
Cu + AgNO Zn + NaCl 3 CuNO
Copy list to your periodic table
3 + Ag
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The Activity Series of the Halogens Fluorine Chlorine Bromine Iodine Halogens can replace other halogens in compounds, provided that they are above the halogen that they are trying to replace.
2NaCl(s) + F 2 (g) MgCl 2 (s) + Br 2 (g) 2NaF(s) + Cl 2 (g)
42 #3 - Single Replacement
One element replaces another
Reactants must be an element and a compound.
Products will be a different element and a different compound.
Na + KCl
F 2 + LiCl
K + NaCl LiF + Cl 2
#3 Single Replacement
Metals replace other metals (and hydrogen)
K + AlN
Zn + HCl
Think of water as HOH
Metals replace one of the H, combine with hydroxide.
Na + HOH
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#3 Single Replacement
We can tell whether a reaction will happen
Some chemicals are more “active” than others
More active replaces less active
There is a list on page 155 - called the Activity Series of Metals
Higher on the list replaces lower.
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#3 Single Replacement
Note the * concerning Hydrogen
H can be replaced in acids by everything higher
Li, K, Ba, Ca, & Na replace H from acids and water
Fe + CuSO 4 Pb + KCl
Al + HCl
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46 #3 - Single Replacement
What does it mean that Hg and Ag are on the bottom of the list?
Nonmetals can replace other nonmetals
Limited to F 2 , Cl 2 , Br 2 , I 2 (halogens)
Higher replaces lower.
F 2 + HCl
Br 2 + KCl
#4 - Double Replacement
Two things replace each other.
Reactants must be two ionic compounds or acids.
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Usually in aqueous solution NaOH + FeCl 3
The positive ions change place.
NaOH + NaOH + FeCl FeCl 3 3
Fe +3 OH
Fe( OH) 3 + Na +1 + Na Cl Cl -1
48 #4 - Double Replacement
Has certain “driving forces”
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Will only happen if one of the products:
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doesn’t dissolve in water and forms a solid (a “precipitate”), or
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is a gas that bubbles out, or
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is a covalent compound (usually water).
49 Complete and balance
assume all of the following reactions take place: CaCl 2 CuCl 2 + NaOH + K 2 S
KOH + Fe(NO 3 ) 3
(NH 4 ) 2 SO 4 + BaF 2
How to recognize which type
Look at the reactants : Combination Element + Element Compound (s,l or g) (s,l or g) (s,l or g)
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51 Decomposition
Compound Element + Element (s,l or g) (s,l or g) (s,l or g)
Single Displacement
Element + Compound (s,l or g) Element + Compound (aq) (s,l or g) (aq or s)
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Double Displacement
Compound + Compound Compound + Compound (aq) (aq) (aq or s) (aq or s)
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54
C x H y + O 2 (g)
Combustion
CO 2(g) + H 2 O (g)
55 Examples
H 2 + O 2 H 2 O
Zn + H 2 SO 4 HgO
KBr +Cl 2 AgNO 3
+ NaCl
Mg(OH) 2 + H 2 SO 3
56 #5 - Combustion
Means “add oxygen”
A compound composed of only C, H, and maybe O is reacted with oxygen
If the combustion is complete, the products
will be CO 2 and H 2 O.
If the combustion is incomplete, the products will be CO (possibly just C) and H 2 O.
57 Examples
C 4 H 10 + O 2
(assume complete) C 4 H 10 + O 2
(incomplete) C C 6 8 H H 12 8 O 6 +O 2 + O 2
(complete) (incomplete)
58 An equation...
Describes a reaction
Must be balanced in order to follow the Law of Conservation of Mass
Can only be balanced by changing the coefficients.
Has special symbols to indicate physical state, and if a catalyst or energy is required.
59 Reactions
Come in 5 major types.
Can tell what type they are by the reactants.
Single Replacement happens based on the activity series
Double Replacement happens if the product is a solid, water, or a gas.
60 Section 7.3
Reactions in Aqueous Solution
OBJECTIVES:
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Write and balance net ionic equations.
61 Section 7.3
Reactions in Aqueous Solution
OBJECTIVES:
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Use solubility rules to predict the precipitate formed in double replacement reactions.
Net Ionic Equations
Many reactions occur in water- that is, in
aqueous solution
Many ionic compounds “dissociate”, or separate, into cations and anions when dissolved in water
Now we can write a complete ionic equation
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Net Ionic Equations
Example:
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AgNO 3(aq) + NaCl(aq)
AgCl(s)+ NaNO 3(aq) 1. this is the full equation 2. now write it as an ionic equation 3. can be simplified by eliminating ions not directly involved (spectator ions) = net ionic equation
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Predicting the Precipitate
Insoluble salt = a precipitate - note Figure 7.13, p.156
General rules: Table 7.3, p. 161, Reference p.708 (back of textbook)
Sample problem 7-9, p.156
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