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!
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.
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.
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.
M + H2O ----> MOH + 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.
M + HX ----> MX + 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 + MX' ----> MX + X2'
Cl2 + 2NaI ----> 2NaCl + I2
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.
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).
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
2. Reactions of Oxygen
a. Formation of peroxides and superoxides.
Oxygen reacts with metals of Group IA to produce metal oxides, peroxides, or superoxides. It can react with lithium to form an oxide, with sodium to form peroxide (in the presence of excess oxygen) and with the heavier members of Group IA to form superoxides. In oxides the oxidation number of oxygen is 2-, in peroxides two oxygen atoms are joined, sharing 2 gained electrons for an overall 2- charge, and in superoxides two oxygen atoms are joined with an overall 1- charge (which corresponds to a charge of -1/2 on each atom of oxygen involved).
2Li (s) + O2 (g) -----> 2Li2O
2Na (s) + O2 (g) -----> Na2O2 (g)
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)
Following is a brief listing of the Activity Series of the Elements:
Activity Series of the Elements
lithium ' fluorine
potassium ' chlorine
calcium ' bromine
sodium ' iodine
iron ' DECREASING ACTIVITY DOWN
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.