Our world is made up of chemical reactions. They happen all of the time, all around us…even inside of us. Chemical reactions occur as we digest our food and breathe in the air around us. We represent these chemical reactions by using chemical equations, which is a way to use shorthand notation to represent information about a chemical reaction.
We will start with word equations. One example of a word equation (for the rusting of iron) is provided below: iron + oxygen → iron(III)oxide Iron and oxygen are the reactants. Iron(III)oxide is the product of this reaction. While most reactions describe the formation of a product or products, some reactions may describe the decomposition, or the breaking down, of a compound. We will look at some of these later. Writing word equations can be tedious, so we use chemical equations as a form of word equation shorthand. The word equation above can be rewritten as a chemical equation (see below). This chemical equation is
Fe + O
2
→ Fe
2
O
3 also called a skeleton equation because it does not indicate the relative amounts of the reactants or the products. Notice, however, that oxygen is diatomic if it is alone, so we write “O
2
” instead of “O” in the equation. We can add more information to the equation by indicating the physical states (solid, liquid or gas) of the reactants and the products in the equation. We use s, l, and g for solids, liquids and gases. If a substance is dissolved in water (in an aqueous solution), we use (aq) . Using our equation above, we can now add more information to it. The unbalanced equation for the rusting of iron is:
Sometimes, a catalyst is used to speed up a reaction. A catalyst is not used in the chemical reaction itself so it is neither a reactant nor a product. The catalyst is written above the arrow (which means “yields”) in a chemical equation. The table below shows two examples of this. One in which “heat” is a catalyst and the other, where Pt (platinum) is a catalyst. The equations
Fe (s) + O
2
(g) → Fe
2
O
3
(s) that we have written so far are unbalanced.
They do not indicate anything about how much of the reactants are used or how much product would be formed. Balancing chemical equations is a skill that we will work on in class. There are some rules to follow when balancing chemical equations. First, you must make sure that each chemical formula (reactants and products) are written correctly. Second, you cannot change subscripts just to balance the equation. A change in a single subscript represents a different chemical compound.
Third, you can ONLY use coefficients to balance equations. A coefficient in algebra
(the “2” in 2x) is the same as a coefficient
In chemistry (the “2” in 2 H
2
O). We balance chemical equations because of the law of conservation of mass. This was part of
Dalton’s atomic theory. It says that when a chemical change occurs, the numbers of atoms (mass) that we start off with is the same as the number of atoms (mass) that we end up with. New atoms do not spontaneously appear. Likewise, atoms are not destroyed. The atoms are simply rearranged into other compounds. We will do many examples of balancing chemical reactions in class.
There are FIVE general types of chemical reactions. These are summarized in the table on the back of this paper. The first type is a combination or synthesis reaction. A synthesis reaction is one in which two or more substances react to form a single new substance. The product is often more complex than the reactants. A decomposition reaction is simply the opposite of a synthesis reaction. A single compound breaks down into two or more simpler products. Decomposition reactions are easy to pick out because the “yields” arrow is at the front of the equation. An example is on the table on the back. Another type of reaction is a single replacement reaction. This type of reaction occurs when one element replaces a second element in a compound. Both the reactants and the products consist of a single element + a compound. Again, look at the example provided in the table. We use data from the activity series of metals (on your polyatomic reference sheet) to determine if one metal will displace another metal in a reaction. The metals in an activity series table are usually given in order of decreasing reactivity. A metal will displace another metal ONLY if it is more active (higher on the list) than the other metal. For example, aluminum is higher than lead in the series. Aluminum will replace lead in a reaction but lead will NOT replace aluminum. Halogens are the same way, but the reactivity of halogens decrease as you move down the halogen family in the periodic table. The fourth type of chemical reaction is a double replacement reaction. This is also called an ionic reaction because it involves two aqueous ionic compounds combining and the ions rearranging to form either new

compounds or free floating ions in the aqueous solution. These reactions often form precipitates, gases, or water as products. We can predict the precipitates that may be formed by looking at the solubility rules (on your polyatomic reference sheet). These rules tell you which compounds are soluble (dissolve in water so NO precipitate) and which are insoluble (do not dissolve in water so form a precipitate). We will work more with both the activity series and the solubility rules in class. The final general type of reaction is a combustion reaction. A combustion reaction is one in which an element or compound reacts with oxygen, often producing energy in the form of heat and/or light. Oxygen is ALWAYS one of the reactants in a combustion reaction. The other reactant is often a hydrocarbon. Products often (but not always) include carbon dioxide (CO
(CO) may also be products of combustion reactions.
2
) and water. Elemental carbon and/or carbon monoxide
Examples (unbalanced) Reaction Type Key Parts combination or synthesis forms a single new substance decompostion arrow at front of equation single replacement element + compound
Format
A + B → AB
AB → A + B
A + BC → B + AC
Mg + O
HgO
Zn + Cu(NO s
)
2
→
2
→ MgO
Hg + O
2
→ Cu + Zn(NO
3
)
2 double replacement or ionic combustion reactants are aqueous ionic compounds oxygen always a reactant
AB + CD → AD + CB
N/A
NaCl + CuSO
C
8
H
18
4
→ CuCl
+ O
2
→ CO
2
2
+ Na
+ H
2
2
SO
O
Notice in the table above that the synthesis reaction between magnesium and oxygen could also be called a combustion reaction.
4
When we were working with ionic bonding, you completed a lab involving double displacement (ionic) reactions. In that lab, you wrote your first unbalanced chemical equations and you predicted the products of the reaction. Remember that when an ionic compound is dissolved in water, the cations and anions dissociate (separate). If a precipitate is formed, then the remaining cation and anion are floating freely in the water – which is why it would conduct electricity. One such reaction is given below:
AgNO
3
(aq) + NaCl (aq) → AgCl ↓ + NaNO
3
(aq)
So, we know that when sodium chloride is dissolved in water, it separates sodium ions (Na+(aq)) and chloride ions (Cl-(aq)). Silver nitrate, when dissolved in water, separates into silver ions (Ag+(aq)) and nitrate ions (NO
3
(aq)) . We can then use these ions to write a complete ionic equation showing the dissolve ionic compounds as dissociated free ions. We also know that it is the silver chloride that forms a precipitate because the solubility rules tell us that silver combined with any group 17 halogen is insoluble. This means that the
Na+ ions and the NO
3
- ions are still floating in the aqueous solution. We can use a ↓ or (s) to show this.
COMPLETE IONIC EQUATION
Ag+ (aq) + NO
3
- (aq) + Na+ (aq) + Cl- (aq) → AgCl (s) + Na+ (aq) + NO
3
- (aq)
Notice that the nitrate (NO
3
-) ion and the sodium (Na+) ion are unchanged on both sides of the equation. These ions do not participate
in the reaction. They are called spectator ions. If we rewrite the complete ionic equation and leave out the spectator ions, we have a simpler version of the equation called a net ionic equation. A net ionic equation shows ONLY the ions directly involved in the chemical change.
NET IONIC EQUATION
Ag+ (aq) + Cl- (aq) → AgCl (s)
Notice that the original equation above between silver nitrate and sodium chloride is already balanced. Therefore the complete and net ionic equations are also already balanced.
Let’s look at another example in which the net ionic equation needs to be balanced with respect to the mass and the charges.
Original Equation (balanced): 3 MgCl
2
(aq) + 2 Na
3
PO
4
(aq) → Mg
3
(PO
4
)
2
(s) + 6 NaCl
Complete Ionic Equation:
3 Mg+ 2 (aq) + 6 Cl- (aq) + 6 Na+ (aq) + 2 PO
4
3 (aq) → Mg
3
(PO
4
Net Ionic Equation: 3 Mg+ 2 (aq) + 2 PO
4
3 (aq) → Mg
3
(PO
4
)
2
(s)
)
2
(s) + 6 Na+ (aq) + 6 Cl- (aq)