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Chapter 4
Chapter 4
Chemical Reactions
Chemical reactions describe processes involving
chemical change
The chemical change involves rearranging matter
Converting one or more pure substances into new
pure substances
Reactants
Substances combined in the reaction
Products
Chemical Equations
Chemical equations are used to describe chemical
reactions
The chemical symbols for the reactants are
shown on the left
The chemical symbols for the products are shown
on the right
An arrow (→) is used to indicate that reactants are
converting to products
A plus sign (+) is used to separate individual
reactants and products
Substances produced in the reaction
Chapter 4
Chapter 4
Chemical Equations
Chemical Equations
“hydrogen and oxygen react to form water”
Chemical equation describing the action:
2 H2(g) +
O2(g) → 2 H2O(l)
Subscripts?
2 H2(g) + O2(g) → 2 H2O(l)
Or,
P4(s) + 6Cl2 → 4PCl3
The equation accomplishes several things:
1.
2.
Coefficients?
Letters in parentheses?
3. Demonstrates the fundamental law of conservation of
matter
Atoms are neither created nor destroyed in chemical
reactions, they are only rearranged
Chapter 4
Balancing Chemical Equations
Stoichiometry
Pronounced as stoy-key-AHM-uh-tree
For example, in the equation describing the formation
of liquid water from hydrogen gas and oxygen gas
2 H2(g) + O2(g) → 2 H2O(l)
There are four hydrogen atoms on both the left
and right sides of the equation
There are two oxygen atoms on both the left and
right sides of the equation
Therefore the equation is balanced
“The relationship between the quantities of
reactants and products”
Chapter 4
Chapter 4
Balancing Chemical Equations
Other examples
NO(g) + O2(g) → NO2(g)
Is it balanced? NO(g) + O(g) → NO2(g)
Is this OK?
How to Balance a Chemical Equation?
NH3(g) + O2(g) → NO(g) +H2O (g)
First step: balance nitrogen atoms on both sides
NH3(g) + O2(g) → NO(g) +H2O (g)
2nd step: Then hydrogen atoms
2NH3(g) + O2(g) → NO(g) +3H2O (g)
But I will need to balance the N atoms again
2NH3(g) + O2(g) → 2NO(g) +3H2O (g)
3rd step: Balance oxygens
2NH3(g) + (5/2) O2(g) → 2NO(g) +3H2O (g)
Is it balanced? NO(g) + ½ O2(g) → NO2(g)
Is this OK?
4th and final step: Just remove the fractions
4NH3(g) + 5 O2(g) → 4NO(g) +6H2O (g)
Now try
NO(g) + O2(g) → NO2(g)
Chapter 4
Chapter 4
Balancing Chemical Equations
H2S(aq) + I2(aq) → HI(aq) + S(s)
Fe(s) + O2 → Fe2O3
KClO3(s) → KCl(s) + O2(g)
Mg + O2(g) →MgO
Ca(OH)2(s) + H3PO4(aq) → Ca3(PO4)2(s) +H2O(l)
P4(s) +O2 (g) →P4O10(s)
Ba(ClO3)2(aq) + H2SO4(aq) → HClO3(aq) + BaSO4(s)
NH3(g)+O2(g) →NO(g) + H2O(g)
C3H8(g) + O2(g) → CO2(g) + H2O(g)
C8H18(l) + O2(g) → CO2(g) + H2O(g)
Chapter 4
Balancing Chemical Equations
H2S(aq) + I2(aq) → 2HI(aq) + S(s)
4Fe(s) + 3O2 → 2Fe2O3
2KClO3(s) → 2KCl(s) + 3O2(g)
2Mg + O2(g) →2MgO
3Ca(OH)2(s) + 2H3PO4(aq) → Ca3(PO4)2(s) +6H2O(l)
P4(s) +5O2 (g) →P4O10(s)
Ba(ClO3)2(aq) + H2SO4(aq) → 2HClO3(aq) + BaSO4(s)
4NH3(g)+5O2(g) →4NO(g) + 6H2O(g)
C3H8(g) + 5O2(g) → 3CO2(g) + 4H2O(g)
2C8H18(l) + 25O2(g) → 16CO2(g) + 18H2O(g)
Chapter 5
More Chemical Equations
Balance the following equations
SO2(g) + O2(g) → SO3(g)
Types of Reactions
Reaction Types
Combination
Decomposition
Double Replacement
Precipitation
Acid-Base
Single Replacement
Oxidation-Reduction
Gas Forming
N2(g) + O2(g) → N2O(g)
C6H14(l) + O2(g) → CO2(g) + H2O(l)
HOMEWORK
Chapter 5
Chapter 5
Combination Reactions
Combination reactions have the form
A+B→C
Combination Reactions
Other combination reactions:
2 Na(s) + S(s) → Na2S(s)
Two or more reactants produce a single product
SO3(g) + H2O(l) → H2SO4(aq)
2 H 2 + O2
=H
2 H2O
2Na (s) + Cl2 (g) → 2NaCl (s)
=O
Chapter 5
Chapter 5
Decomposition Reactions
Decomposition reactions have the form
A→B+C
Single reactant breaks down into two or more
products
2 H2O2
2 H2O + O2
=H
=O
Decomposition Reactions
Further examples:
2 HgO(s) → 2 Hg(l) + O2(g)
This reaction was used by Joseph Priestley in
the discovery of oxygen in 1774
CaCO3(s) → CaO(s) + CO2(g)
This reaction is used industrially to produce
both lime (CaO) and CO2 from limestone
(CaCO3)
Air bags:
2 NaN3 (s) 2 Na (s) + 3 N2 (g)
Chapter 5
Chapter 5
Double Replacement
Double Replacement
Double replacement reactions are also called
“metathesis” reactions or “partner swapping”
reactions
They have the form AX + BY → BX + AY
Double replacement reactions often take place in
water and are of two basic types:
Acid base neutralization reactions:
HCl(aq) +
NaOH(aq) → NaCl(aq) + H2O(l)
(acid)
(base)
(salt)
(water)
H2SO4(aq) + 2 NaOH(aq) → Na2SO4(aq) + H2O(l)
+
HF
+
=H
NaOH
=O
→
NaF
= Na
+ H2O
Precipitation reactions (solid forms):
Ba(NO3)2(aq) + Na2SO4(aq) → BaSO4(s) + 2 NaNO3(aq)
Here the barium sulfate precipitates out of
solution
=F
Chapter 5
Chapter 5
Mixtures and Solutions
Solute
the component of a solution that is dissolved in another
substance
Mixtures and Solutions
Electrolyte
All ionic compounds that are soluble in water and conduct
electricity
Solvent
the medium in which a solute dissolved to form a solution
Strong electrolyte
Solution
a homogenous mixture in which the components are
evenly distributed in each other
Weak electrolyte
Aqueous
Any solution in which water is the solvent
Precipitate
Insoluble product of a reactant
Non-electrolyte
Chapter 5
Chapter 5
Solubility Rules
Mixtures and Solutions
Compounds Containing
+
+
Group IA (Na , K ) and NH4
Nitrates
Soluble
Soluble
Perchlorates (ClO4−)
Soluble
Chlorides, Bromides, Iodides
Soluble
CH3COOH(aq)
CH3COO- (aq)+ H+ (aq)
<5% ionized (weak electrolyte)
Chapter 5
Exceptions
Soluble
Acetates (CH3COO−)
CuCl2 (s) → Cu2+ (aq) + 2Cl- (aq)
100 % dissociation (strong electrolyte)
Figure 5.2
Solubility
+
Ag+, Pb2+, Hg+
2+
Fluorides
Soluble
Mg , Ca2+, Ag+, Pb2+, Sr2+
Sulfates
Soluble
Ag+, Ba2+, Sr2+, Pb2+, Hg+
Sulfides
Insoluble
Group IA and NH4+
Carbonates
Insoluble
Group IA and NH4+
Phosphates
Insoluble
Group IA and NH4+
Hydroxides
Insoluble
Group IA and NH4+
Chapter 5
Precipitation Reactions
Precipitation reactions
K2SO4 (aq) + Pb(NO3)2 (aq) →
Ionic Equations
When an ion appears on both sides of a chemical
equation it can be canceled out
“Spectator Ions”
Writing Equations
Molecular
Net Ionic Equation
Total ionic
Net ionic
http://www.dlt.ncssm.edu/TIGER/Flash/moles/DoubleDisp_Reaction-Precipitation.html
Chapter 5
Chapter 5
Ionic Equations
Ionic Equations
Mg(ClO4)2(aq) + K2CO3(aq) → MgCO3(s) + KClO4(aq)
Total:
Cancel:
KNO3(aq) + CaCl2(aq) →
Total:
Net Ionic:
Ba(NO3)2(aq) + Na2SO4(aq) → BaSO4(s) +
NaNO3(aq)
Cancel:
Net Ionic:
Total:
Cancel:
Net ionic:
Chapter 5
Chapter 5
Acids and Bases
Acid
a substance that, when dissolved in water, increases the concentration of
hydrogen ions (H+) in solution
Base
a substance that, when dissolved in water, increases the concentration of
hydroxide ions (OH-) in solution
Strong acid
Strong base
Strong electrolyte
Acids and Bases
Chapter 5
Chapter 5
Acid-Base Reactions
Acid-Base Reactions
HNO3(aq) + NaOH(aq) →
H2SO4(aq) + KOH(aq) →
Total:
Total:
Cancel:
Cancel:
Net Ionic:
Net Ionic:
Chapter 5
Chapter 5
Single Replacement
Single replacement reactions are also called
substitution reactions
They have the form A + BX → B + AX, where A
and B are elements and BX and AX are
compounds
Single Replacement
Other single replacement reactions
3 C(s) + 2 Fe2O3(s) → 4 Fe(s) + 3 CO2(g)
Cu(NO3)2(aq) + Zn(s) → Zn(NO3)2(aq) + Cu(s)
Single replacement reactions are also oxidationreduction (REDOX) reactions.
Oxidation
+
Reduction
H2
+
CuO
=H
→
=O
H2O
+ Cu
= Cu
OXIDATION NUMBER
Chapter 5
Chapter 5
Redox Reactions and Electron
Transfer
The oxidation number of an atom is the charge that atom
would have if the compound was composed of ions.
Oxidation of Mg (0) to Mg (II):
2Mg (s) + O2 (g) → 2MgO (s)
Reduction of Fe (III) to Fe(0):
Fe2O3 (s) + 3CO (g) →2Fe (s) + 3 CO2 (g)
Reduction of Ag (I) to Ag and oxidation of Cu (0) to
Cu (II):
2Ag+ (aq) + Cu(s) → 2 Ag (s) + Cu2+ (aq)
Chapter 5
Gas Forming Reactions
Gas Forming Reactions
Reactions leading to the formation of an insoluble
gas
Examples
Na2CO3 (aq) + 2HCl (aq) → 2NaCl (aq) + CO2 (g) + H2O (l)