Read Section 8.1
Information available from a chemical equation:
Reactants --------> Products
__________________ - starting substances in a chemical reaction
__________________ - substances formed in a chemical reaction
----------> means “yields”
Example:
Word equation: hydrogen plus oxygen yields water
Chemical equation: H2
+ O2
--------> H2O
The Law of Conservation of Mass tells us that we can neither create
nor destroy matter in a chemical reaction, ...
H2
+ O2
---------> H2O
But you can change their partners.
2 H2 +
O2
----------> 2 H2O
2 molecules H2 + 1 molecules O2 -----------> 2 molecules
H2O
2 moles H2 +
1 mole O2 -----------> 2 moles H2O
4.04 g H2
+ 32.00 g O2 ----------->
36.04 g
H2O
Symbols used in writing chemical equations
+
used to separate two reactants or two products
--------> “yields”, separates reactants from products
--------> heat applied to the reaction
catalyst
--------> a catalyst is required (gets reaction started faster)
(s) _________________
(l) _________________
(g) __________
(aq) _______________; substance is dissolved in water
gas being produced by the reaction
solid being produced by the reaction
Writing and Balancing Chemical Equations
1. Make sure you have the correct formulas for all the reactants and products in the reaction. Are elements present in the reaction that should be written as diatomic molecules?
2. Write the reactants on the left and the products on the right with a yield sign in between. If there are two or more reactants or products involved, use plus signs to separate their formulas.
3. Compare the number of atoms of each element in the reactants with the numbers of atoms of each element in the products. If the numbers on both sides are equal you are finished; the equation is balanced.
4. Balance the elements on each side of the equation by placing a coefficient, a whole number, in front of the chemical formulas as appropriate. Start with elements that appear only once on each side of the equation. Never change subscripts to balance the equation. That would alter the chemical formula of a substance, making it another substance altogether.
5. Double-check each kind of atom to be sure that the equation is balanced.
6. Check to see that the coefficients are in the lowest possible ratio. For example, 12:6:4 should be reduced to 6:3:2.
Writing and Balancing Chemical Equations
Strategies for balancing equations:
1. EVEN/ODD
Example: Write the equation for the reaction of iron and oxygen to form iron (III) oxide
1. Make sure you have the correct formula for all reactants and
products.
iron
oxygen
iron (III) oxide
2. Write the reactants on the left and the products on the right
with a yield sign in between.
Fe +
O2 ------------> ______________
3. Equation is not balanced with an even number of oxygen atoms on the left of the equation and odd number of O atoms on the right.
4. Balance using coefficients
In the reaction,
Fe +
________ --------->
_________
the number of oxygen atoms on the left must be even. Begin by making the number of oxygen atoms on the right even, then balance the iron atoms.
5. double-check
6. lowest possible ratio
Example: Write an equation for the formation of ammonia from the reaction of nitrogen and hydrogen.
Substances:
nitrogen _______
hydrogen _______
ammonia ____________
Skeleton equation - chemical equation that does not indicate the
relative amount of reactants and products
_____________ + ____________
--> _____________
Equation is not balanced
Draw a box around reactants to remind yourself that you cannot change the subscripts; the boxes are off-limits - you cannot change anything inside the box
__________________________________________
use the even/odd strategy
__________________________________________
Double-check number of atoms
Check that coefficients are in the lowest possible ratio
Example: Write the equation showing that sodium sulfide is formed from sodium and sulfur.
Substances:
sodium ______ sulfur
______
sodium sulfide
Skeleton equation - chemical equation that does not indicate
the relative amount of reactants and products
_______ + ________ ------->
____________
Equation is not balanced
________ + ________ --------->
_________
Double check number of atoms on both sides of the equation
Check that coefficients are in lowest possible ratio
2. INTACT POLYATOMIC ION
Example: Write the equation for the formation of iron (III) acetate and lead from iron and lead (II) acetate
Substances:
1) lead ______
iron _______
iron (III) acetate
lead (II) acetate
Skeleton equation:
Equation not balanced
the atoms in the polyatomic ion, acetate, can be balanced
as a group
also use even/odd strategy
Is the equation balanced?
Are the coefficients in the lowest possible ratio?
Example: Sodium chloride and potassium nitrate react to form
sodium nitrate and potassium chloride.
Substances:
sodium chloride
_______________
sodium nitrate ________________
potassium nitrate
_______________
potassium chloride _____________
Equation:
3. COMBUSTION OF HYDROCARBONS
The number of hydrogen atoms in the reactants must be divisible by four.
Example: Write an equation for the combustion of acetylene to form carbon dioxide and water.
Substances:
CO2
O2
C2H2
H2O
Skeleton equation:
__________ + _______ ---------> ________ + _________
Balance: 1) Make the number of hydrogen atoms on the reactant side - a
number divisible by 4
2) Balance carbon atoms
3) Balance oxygen atoms
Begin by making the number of hydogen atoms on the left 4 by choosing a coefficient of 2 for
__________; the equation then balances easy
___________ + _________ -------> _________ + __________
Is the equation balanced?
Are the coefficients in the lowest possible ratio?
When the number of hydrogen atoms is already divisible by 4, this technique is not needed, as in:
__________ + _________--------> _________ + _________
Read Section 8.2
Types of Chemical Reactions
I Combination
II Decomposition
III Single-Replacement
IV Double-Replacement
V Combustion
I. Combination Reaction
Two or more substances combine to form a single product
Reactants are elements or compounds
A. Group A metal cations combine with non-metal anions to form an ionic
compound
Examples:
2 Ca (s) + O2 (g) ------------> 2 CaO(s)
4 Al (s) + 3 O2 (g) ------------> 2 Al2O3 (s)
Practice:
zinc + nitrogen -->
sodium + chloride -->
B. Two non-metals react with each other - or -
a transition metal reacts with a non-metal
Example: non-metal and non-metal
2 Cl2 (g) + O2
-------> 2 Cl2O (g)
dichlorine monoxide
Cl2 (g) + 2 O2 (g)
------> 2 ClO2 (g) chlorine dioxide
Example: transition metal and non-metal
2 Cu (s) + O2 (g) ----------> 2 CuO (s) copper (II) oxide
4 Cu (s) + O2 (g) ----------> 2 Cu2O (s) copper (I) oxide
C. Non-metal oxides react with water to form acids
(compounds that produce hydrogen ions, H+, when
combined with water)
Example: CO2 (g) + H2O (l) ---------> H2CO3 (aq) carbonic acid
Practice: SO3
+ H2O
--->
N2O5
+ H2O --->
D. Metallic oxides react with water to form bases
(compounds which contain hydroxide ions, OH-)
Example: Na2O (s) + H2O (l) ---------->
2 NaOH (aq)
Practice: Li2O
+ H2O -->
BaO +
H2O -->
II. Decomposition Reactions
In decomposition reactions, it is more difficult to determine the products. Compounds that contain only 2 kinds of elements (binary compounds) will separate into those two elements.
A. Binary compounds separate into two elements
Example: 2 Ag2O (s)
--> 4 Ag (s)
+ O2 (g)
Practice:
B. Some acids will decompose into water and a nonmetal oxide
(if heated)
Example: H2CO3 (aq)
-----> CO2 +
H2O
Practice:
C. Some bases will decompose into water and a metal oxide (if
heated)
Example: Ba(OH)2 (aq)
--> BaO (s)
+ H2O (l)
Practice:
Compounds containing three or more elements are less predictable.
III. Single-Replacement Reactions
One element replaces a second element in a compound; the element that begins the reaction alone replaces another
element that begins the reaction as part of a compound; the element
that ends up alone in the products is the one
that was originally part of the reactants
A. An active metal will replace the metallic ion in a compound of
a less active metal from a
compound
Example: Cu (s)
+ AgNO3 (aq)
---> Ag (s)
+ CuNO3 (aq)
Practice: Fe (s) + CuSO4 (aq)
--->
B. Active metals such as Zn, Pb, and Al will replace H in acids to
give a salt (ionic compound) and hydrogen gas
Example:
C. Halogens will replace less active halogens.
Example:
IV. Double-Replacement Reactions
Positive ions exchange places between two compounds
Three indications of reaction
1) precipitate forms
2) gas given off
3) water produced when acid mixed with base
AX + BY --------> AY + BX
Example:
_________ (aq) + __________ (aq) --------> __________ (s) + 2 NaCl (aq)
Practice:
2 KI (aq) + _________ (aq) --------> 2 ________ (aq) + _________ (s)
V. Combustion
Substance combines with oxygen
Products of complete combustion are carbon dioxide and water
Exothermic
CxHy + _______ -------> ________+ ________
Example: the combustion of octane in gasoline
2 _________ (l) + 25 _______ (g) ----> 16 ______ (g) + 18 _____ (g)
Example: the combustion of heptane: _________(g) + _____ (g) ------>
_________ (g) + ________ (g) -------> ________ (g) + _______(g)
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