TYPES OF CHEMICAL EQUATIONS
Equations in Chemistry
consist of two major 'divisions': the reactants side and the products side. The
reactants (those elements and compounds which are reacting) are written
symbolically on the left and the products (the results of the reactants
reaction together) are written on the left, like this:
A + B ----> C + D
A and B are reactants; C and D are products. It is read "A
reacts with B to yield (the arrow is read as 'yield' or 'yields") C and D.
Also, when writing the
reactants and products, other symbols are used to indicate what kind of
reactants and products are being used and being formed (whether they are solid,
liquid or gaseous). An up-pointing arrow means a gas is being formed and a
down-pointing arrow means a precipitate (a solid) is being formed.
A(s) + B(l) ---->
C(g) + D(s) + E(aq)
This means that a solid
element or compound A is reacting with liquid phase B to yield gaseous element
or compound C plus solid D and element or compound E which is dissolved in water
('aq' means aqueous, or in water).
Equations must represent
the facts of the chemical reaction, all elements and compounds must be written
correctly and the equation must be balanced to satisfy the Law of Conservation
of Atoms. No mass is gained or lost in a chemical reaction. The total mass of
reactants equals the total mass of the products.
Until an equation is
balanced, it does not truly represent a chemical reaction.
To write an equation:
(a) Represent the facts
(b) Write correctly the formulae of all elements and compounds involved
(c) Balance the equation so that equal numbers of each type of atom are present
on each side of the equation.
You might want to write
the equation in words and translate the words to symbols (if it IS not
already), proceeding as indicated above from there.
KEEP IN MIND THAT THERE ARE SEVEN DIATOMIC GASES WHOSE
FORMULAE ARE NOT WHAT THEY WOULD BE IF WRITTEN STRAIGHT FROM THE PERIODIC
TABLE.
To be able to write
equations as they should be, you must know the symbols of the
elements, know the usual
oxidation numbers of the elements, know the radicals and their charges and
their names, you have to know what is reacting and forming and whether it is a
gas, liquid or solid or dissolved in water, you have to make sure formulae for
compounds are written correctly and you must balance the equation correctly (so
that equal numbers of all and each type of atom appear on both sides of the
equation). You use coefficients in front only of the symbols of the elements
and compounds involved in the reaction. You never change the formula of a
compound or element to make it balance correctly. If you did that, you change
the name of the compound or element, as well as change the chemical nature
(physical and chemical properties) of that element. In other words, if I had
iron (II) oxide reacting and wrote the symbols for iron (III) oxide, I would
not be representing the facts for that part of the equation and the equation
would be inaccurate. YOU MUST REPRESENT THE FORMULAE FOR ALL REACTANTS AND
PRODUCTS CORRECTLY!
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GENERAL TYPES OF CHEMICAL REACTIONS:
Following is a list of
the general kinds of chemical reactions. After the list and a short explanation
of the reaction types, there will be a section devoted to each general type of
reaction, the specific subtypes and descriptions of each.
1. Composition (also called synthesis)
reactions.
These
are reactions in which two or simpler substances (elements or compounds) are
combined to form a more complex substance (a compound.) It has the following
general form:
A + X ----> AX
2. Decomposition reactions.
These
are reactions in which one substance breaks down, forming two or more simpler
substances. It is the reverse of the composition (synthesis) reaction. The
general form of the reaction is:
AX ----> A + X
3. Replacement (also called single
replacement) reactions.
In this type of reaction, one substance is
replaced by a more reactive material. Either the cation
(positive part) or the anion (negative part) undergoes replacement, but both do
not undergo replacement at the same time. The general forms are as follows:
A + BX ----> AX + B (cation
replacement)
Y + BX ----> BY + X (anion replacement)
The reason that this
happens is that one element (substance) is more reactive (greater electronegativity) than another.
4. Ionic (also called double replacement)
reactions.
In this
type of reaction, there is no electron transfer going on (as there must be in
single replacement for one substance to get to the free or uncombined state-
see equations in #3 for illustrations). Rather, substances which are
dissolved/ionized in solution exchange anions (or exchange cations,
depending on your point of view). Further, the new products are formed in such
a way that at least one of the products leaves the 'reaction environment' (the
surroundings where the reaction is taking place). Usually, that will mean that
a gas is evolved (given off) or a precipitate (insoluble -cannot remain
dissolved in the 'reaction environment'). These reactions are also called
exchange reactions.
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1. SYNTHESIS REACTIONS
(a) two
elements forming a more complex substance:
sodium + chlorine ----> sodium chloride
2Na + Cl2 ----> 2NaCl
(b) two
compounds forming a more complex substance:
hydrogen oxide + sulfur trioxide ----> hydrogen sulfate
H2O + SO3 ----> H2SO4
Combinations of the above
are possible. This would be like an element plus a compound forming a more complex
compound. In general, there will be only one product to this type of reaction.
2. DECOMPOSITION REACTIONS:
There are six types of
decomposition reactions dealt with here:
(a) Metallic carbonates
yield metallic oxides plus carbon dioxide.
MCO3 ----> MO + CO2 (g)
(b) Many (most, but not
all) metallic hydroxides form metallic oxides and water when heated.
MOH ----> MO + H2O [(l) or (g), depending on
temperature]
(c) Metallic chlorates
form metallic chlorides and oxygen when heated.
MClO3 ----> MCl + O2 (g)
(d) Acids (formed of
nonmetallic oxides' and water's reaction and having hydrogen as their first
element in the compound) decompose into nonmetallic oxides and water when
heated.
HNmO ----> NmO + H2O [(l) or (g) depending on the temperature]
(e) Some oxides, when
they are heated sufficiently, form simpler substances (usually the metal and
oxygen form. Metallic oxides are the ones for whom this reaction is true).
2MO ----> M + O2
(f) Electricity, by
supplying electrons in a reaction, can cause decomposition.
Electricity……………..
2H2O ---------------> 2H2 (g) + O2 (g)
3. SINGLE REPLACEMENT REACTIONS:
There are four types of
single replacement (also called replacement or displacement) reactions. These
are always redox reactions. (Read much further on for
an idea of what redox reactions are).
(a) Replacement of a
metal in a compound by a more active metal.
Remember the general
rule? Activity increases over and up the Periodic Table.
Zn + CuSO4 ----> ZnSO4 + Cu
There is a list called
the Activity Series of the Elements which shows the more active (or usual)
metals involved in chemical reactions and the order in which they are able to
replace one another in compounds. In this list (which will come later on) any
element (metal) will displace any other metal below it from the less active
(lower) element's (metal's) compound. The further apart the metals are the
faster and more readily the reaction will take place.
(b) Replacement of
hydrogen in water by metals.
Any metal above hydrogen
in the Activity Series will displace hydrogen from
water, forming a new compound plus hydrogen gas.
Metal + H2O ----> MetalOH + H2
Ca + H2O ----> Ca(OH)2 + H2
(c) Replacement of hydrogen in acids by metals.
Hydrogen may be replaced from acids (hydrogen, if it is the first element in a
compound, is in a compound type called an acid) by metals above it in the
Activity Series.
Metal + HX(acid) ----> MetalX + H2
Zn + HCl ----> ZnCl2 + H2
(d) Replacement by
halogens.
Group VIIA on the
Periodic Table has another name: the Halogen Family. (Halogen means
'salt-former'.) They are similar in properties (which is reasonable since they
have similar outer shell arrangements- remember: all elements in the same 'A'
column have the same number of electrons in the outer (highest) energy level).
The higher up in the column you go, the move reactive the halogen. (The most
active element is fluorine; it is the most active halogen). Therefore, a
halogen may be replaced from its compound by a halogen which lies above it on
the Periodic Table.
X2 + MetalX' ----> MetalX + X2'
Cl2 + 2NaI ----> 2NaCl + I2
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Many times, chemical
reactions can be reversed, using the reverse process which formed the new
products. That is, you can re-form the reactants. When this happens without any
outside help, and all the products and reactants are still present in the same
'reaction environment', it happens until a condition known as equilibrium
occurs. Equilibrium means the products are forming the reactants as fast as the
reactants are forming the products.
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4. IONIC REACTIONS
In ionic reactions, there
is no electron exchange. Instead, the anions of the compounds 'exchange' (swap)
and products are formed which leave the reaction environment.
AB (aq) + CD (aq)
----> AD + CB
One or both of the
products AD or CB may be solid or gaseous or some combination. A specific
example follows:
CaCl2 (aq) + 2AgNO3 (aq) ----> Ca(NO3)2 (aq) + 2AgCl (s)
While describing this type
reaction, it would be useful to note that when a solid forms, this type of
reaction can be referred to as a precipitation reaction.
Also, reactions in which two compounds react to form two new compounds and
there is no change in oxidation number are called metathesis reactions. If any
of the compounds involved change oxidation numbers on any element contained in
the compound, the reaction is referred to as an oxidation-reduction (redox) reaction. One element or compound serves as reducing
agent (loses electrons) and one element or compound serves as the oxidizing
agent gains electrons).
Examples:
a. Metathesis Reaction:
Pb(NO3)2 + K2CrO4
-----> PbCrO4 + 2KNO3
None of the involved
elements change oxidation number.
b. Precipitation reaction
AgNO3 (aq) + NaCl
(aq) -----> AgCl (s) +
NaNO3 (aq)
Both silver nitrate and
sodium chloride are dissolved in water. They are mixed, with the precipitate
silver chloride being formed and sodium nitrate remaining as the other reaction
product. The sodium nitrate is dissolved in the water.
c. Reactions Involving
Oxidation Number Changes
1. Oxidation-Reduction (Redox)
SnCl2 (aq) + 2FeCl3 (aq)-----> SnCl4 (aq) + 2FeCl2 (aq)
The reactants tin (II)
chloride and iron (III) chloride are altered in oxidation number on the
respective metal ions. Tin (II) goes to tin (IV) on the product side,
indicating
that it has lost electrons (and is therefore called the
reducing agent - the one
responsible for reducing the charge on another ion). Iron (III) goes
from 3+ on the reactant side to 2+ (iron (II)) on the product side. It has been
reduced in oxidation number and is called the oxidizing agent (the one
responsible for causing another ion to increase in charge).
2. Disproportionation
Reactions
3NO2 (g) + H2O (l)-----> 2HNO3
(l) + NO (g)
Notice that the nitrogen
dioxide contains nitrogen, as does the nitric acid (HNO3) and the nitrogen
monoxide. Three compounds with nitrogen. Nitrogen undergoes a change in
oxidation number from 4+ in the nitrogen dioxide to 5+ in the nitric acid and
2+ in the NO. The nitrogen dioxide is both reduced and oxidized. It is both the
reducing and oxidizing agents.
OTHER REACTION
TYPES/CATEGORIES:
1. Reversible Reactions
In these type reactions,
the reactants form the product. At some point the product begins to decompose
back into the reactants. This may or may not happen at equilibrium.
2. Reactions of Hydrogen
Hydrogen forms two types
of hydrides (binary - 'hydrogen-other element') compounds.
a. Ionic hydrides, in
which hydrogen assumes a 1- oxidation number. This is generally with the active
metals of groups IA and IIA.
Li + H2 (g) -----> 2LiH
b. Covalent hydrides, in
which hydrogen assumes a 1+ oxidation number. This is in a binary combination
with a nonmetal. In combinations with elements in column VIIA, the compounds
are called hydrogen halides.
Covalent hydride
(example):
2H2 + O2 -----> 2H2O
Hydrogen halide
(example):
H2 + F2 -----> 2HF
Most (perhaps many would
be a better way to say it) of the covalent hydrides are
acidic.
2. Reactions of Oxygen
a. Formation of peroxides
and superoxides.
Oxygen reacts with metals
of
Oxide formation:
2Li (s) + O2 (g) -----> 2Li2O
Peroxide formation:
2Na (s) + O2 (g) -----> Na2O2 (g)
Superoxide formation:
K (s) + O2 (g) -----> KO2 (s)
b. Formation of metal
oxides
Oxygen reacts with metals
(other than those described above) to form oxides in which the oxidation number
of oxygen is 2-.
4Fe (s) + 3O2 (g) -----> 2Fe2O3 (s)
c. Metal oxides react
with water to form metal hydroxides. Metal oxides are sometimes called basic
anhydrides for this reason, since hydroxides are bases and when the water is
removed from their chemical structure, they are "anhydrides" (without
water).
ZnO (s) + H2O (l)
-----> Zn(OH)2 (aq)
d. Many nonmetals react with
oxygen to form nonmetal oxides. These are covalent compounds.
2S (s) + 3O2 (g) -----> 2SO3 (g)
e. Nonmetal oxides react
with water to form acids. For this reason, nonmetal oxides
are often called acidic anhydrides for reasons analogous to
basic anhydrides.
SO3 (g) + H2O (l) -----> H2SO4 (l)
f. Metal oxides and
nonmetal oxides react to form salts. This occurs with no change in oxidation
number for any involved element.
CO2 (g) + CaO (s) -----> CaCO3 (s)
g. Combustion reactions
involve the use of oxygen to oxidize materials. For this reason, these are redox reactions also. Most frequently it is thought of in
such situations as the burning of wood and fossil fuels (hydrocarbons).
CH4 (g) + 2O2 (g) -----> CO2 (g) + 2H2O(l)
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Following is a brief listing of the Activity Series of the Elements:
Activity Series of the Elements
Metals Nonmetals
lithium ' fluorine
potassium ' chlorine
calcium ' bromine
sodium ' iodine
magnesium '
aluminum '
zinc '
chromium '
iron ' DECREASING ACTIVITY DOWN
nickel '
tin \'/
lead '
HYDROGEN
copper
mercury
silver
platinum
gold
RULES REGARDING THE
ACTIVITY SERIES OF THE ELEMENTS:
1. Each element in the list
displaces from a compound any of the elements below it.
The
larger the interval between elements in the Series, the more vigorous the
action.
2. All metals above
hydrogen displace hydrogen from hydrochloric acid (HCl)
or dilute sulfuric acid (H2SO4).
3. Metals above magnesium
vigorously displace hydrogen from water. Magnesium displaces hydrogen from
steam.
4. Metals above silver combine directly with oxygen; those near the top do so
rapidly.
5. Metals below mercury
form oxides only indirectly.
6. Oxides of metals below
mercury decompose with mild heating.
7. Oxides of metals below
chromium easily undergo reduction to metals by heating with hydrogen.
8. Oxides of metals above
iron resist reduction by heating with hydrogen.
9. Elements near the top
of the Series are never found free in nature.
10. Elements near the
bottom of the list are often found free in nature.