# Writing and Balancing Chemical Equations

####
**Chemical
Reactions and Chemical Equations**

** A

**chemical reaction**is a process in which a substance (or substances) is changed into one or more new substances.
** To
communicate with one another about chemical reactions, chemists have devised a
standard way to represent them using chemical equations.

** A

**chemical equation**uses chemical symbols to show what happens during a chemical reaction.
** In this subject
we will learn how to write chemical equations and balance them.

####
**Writing
Chemical Equations**

###
**Reaction
between Hydrogen and Oxygen **

** Consider
what happens when hydrogen gas (H

_{2}) burns in air (which contains oxygen, O_{2}) to form water (H_{2}O). This reaction can be represented by the chemical equation.
** where the “plus”
sign means “reacts with” and the arrow means “to yield.” Thus, this symbolic
expression can be read: “Molecular hydrogen reacts with molecular oxygen to
yield water.” The reaction is assumed to proceed from left to right as the
arrow indicates.

** Equation (1)
is not complete, however, because there are twice as many oxygen atoms on the left
side of the arrow (two) as on the right side (one). To conform with the law of conservation
of mass, there must be the same number of each type of atom on both sides of
the arrow; that is, we must have as many atoms after the reaction ends as we
did before it started.

** We can balance
Equation (1) by placing the appropriate coefficient (2 in this case) in front
of H

_{2}and H_{2}O:
** This

**balanced chemical equation**shows that “two hydrogen molecules can combine or react with one oxygen molecule to form two water molecules” (Figure).
Because the
ratio of the number of molecules is equal to the ratio of the number of moles,
the equation can also be read as (2 moles of hydrogen molecules react with 1
mole of oxygen molecules to produce 2 moles of water molecules.)

** We know the
mass of a mole of each of these substances, so we can also interpret the equation
as (4.04 g of H

_{2}react with 32.00 g of O_{2}to give 36.04 g of H_{2}O.)
** These three
ways of reading the equation are summarized in the following Table:

####
**Important
Notes for Writing Chemical Equations**

**(1)**

**Reactants**are the starting materials in a chemical reaction. H

_{2}and O

_{2}are Reactants in Equation (1)

**(2)**

**Products**are the substance formed as a result of a chemical reaction. Water is the

**Product**in Equation (1)

**(3)**A

**chemical equation**is just the chemist’s shorthand description of a reaction.

**(4)**In a chemical equation the reactants are conventionally written on the left and the products on the right of the arrow:

**(5)**To provide additional information, chemists often indicate the physical states of the reactants and products by using the letters g, l, and s to denote gas, liquid, and solid, respectively.

For example:

To represent
what happens when sodium chloride (NaCl) is added to water, we write

where (aq) denotes
the aqueous (that is, water) environment. Writing H

_{2}O above the arrow symbolizes the physical process of dissolving a substance in water, although it is sometimes left out for simplicity####
**Balancing
Chemical Equations**

** Suppose we
want to write an equation to describe a chemical reaction that we have just carried
out in the laboratory. How should we go about doing it? Because we know the
identities of the reactants, we can write their chemical formulas. The
identities of products are more difficult to establish. For simple reactions,
it is often possible to guess the product(s). For more complicated reactions
involving three or more products, chemists may need to perform further tests to
establish the presence of specific compounds.

** Once we have
identified all the reactants and products and have written the correct formulas
for them, we assemble them in the conventional sequence— reactants on the left
separated by an arrow from products on the right. The equation written at this
point is likely to be unbalanced; that is, the number of each type of atom on
one side of the arrow differs from the number on the other side.

###
**Rules
for balancing chemical equation **

In general, we
can balance a chemical equation by the following steps

**(1)**Identify all reactants and products and write their correct formulas on the left side and right side of the equation, respectively.

**(2)**Begin balancing the equation by trying different coefficients to make the number of atoms of each element the same on both sides of the equation. We can change the coefficients (the numbers preceding the formulas) but not the subscripts (the numbers within formulas).

Changing the
subscripts would change the identity of the substance. For example, 2NO

_{2}means “two molecules of nitrogen dioxide,” but if we double the subscripts, we have N_{2}O_{4}, which is the formula of dinitrogen tetroxide, a completely different compound.**(3)**First, look for elements that appear only once on each side of the equation with the same number of atoms on each side. The formulas containing these elements must have the same coefficient. Therefore, there is no need to adjust the coefficients of these elements at this point.

Next, look for
elements that appear only once on each side of the equation but in unequal numbers
of atoms. Balance these elements. Finally, balance elements that appear in two
or more formulas on the same side of the equation.

**(4)**Check your balanced equation to be sure that you have the same total number of each type of atoms on both sides of the equation arrow.

####
**Examples
on Writing and Balancing Chemical Equations**

###
**Example
(1): Heating of potassium Chlorate**

** Let’s
consider a specific example. In the laboratory, small amounts of oxygen gas can
be prepared by heating potassium chlorate (KClO

_{3}). The products are oxygen gas (O_{2}) and potassium chloride (KCl). From this information, we write:
(For
simplicity, we omit the physical states of reactants and products.)

** All three
elements (K, Cl, and O) appear only once on each side of the equation, but only
for K and Cl do we have equal numbers of atoms on both sides. Thus, KClO

_{3}and KCl must have the same coefficient. The next step is to make the number of O atoms the same on both sides of the equation. Because there are three O atoms on the left and two O atoms on the right of the equation, we can balance the O atoms by placing a 2 in front of KClO_{3}and a 3 in front of O_{2}.
** Finally, we
balance the K and Cl atoms by placing a 2 in front of KCl:

** As a final
check, we can draw up a balance sheet for the reactants and products where the number
in parentheses indicates the number of atoms of each element:

** Note that
this equation could also be balanced with coefficients that are multiples of 2
(for KClO

_{3}), 2 (for KCl), and 3 (for O_{2}); for example:
** However, it
is common practice to use the simplest possible set of whole-number coefficients
to balance the equation. Equation (2) conforms to this convention.

###
**Example
(2): Combustion of ethane**

** Now let us
consider the combustion (that is, burning) of the natural gas component ethane (C

_{2}H_{6}) in oxygen or air, which yields carbon dioxide (CO_{2}) and water. The unbalanced equation is:
** We see that
the number of atoms is not the same on both sides of the equation for any of
the elements (C, H, and O). In addition, C and H appear only once on each side
of the equation; O appears in two compounds on the right side (CO

_{2}and H_{2}O).
** To balance
the C atoms, we place a 2 in front of CO

_{2}:
** To balance
the H atoms, we place a 3 in front of H

_{2}O:
** At this
stage, the C and H atoms are balanced, but the O atoms are not because there are
seven O atoms on the right-hand side and only two O atoms on the left-hand side
of the equation. This inequality of O atoms can be eliminated by writing 7/2 in
front of the O

_{2}on the left-hand side:
** The “logic”
for using 7/2 as a coefficient is that there were seven oxygen atoms on the right
hand side of the equation, but only a pair of oxygen atoms (O

_{2}) on the left. To balance them we ask how many pairs of oxygen atoms are needed to equal seven oxygen atoms. Just as 3.5 pairs of shoes equal seven shoes, 7/2 O_{2}molecules equal seven O atoms. As the following tally shows, the equation is now balanced:
** However, we
normally prefer to express the coefficients as whole numbers rather than as
fractions. Therefore, we multiply the entire equation by 2 to convert 7/2 to 7:

The final tally
is:

** Note that
the coefficients used in balancing the last equation are the smallest possible set
of whole numbers.

###
**Example
(3): Formation of aluminum oxide**

** When
aluminum metal is exposed to air, a protective layer of aluminum oxide (Al

_{2}O_{3}) forms on its surface. The unbalanced equation is:
** In a
balanced equation, the number and types of atoms on each side of the equation
must be the same. We see that there is one Al atom on the reactants side and
there are two Al atoms on the product side. We can balance the Al atoms by
placing a coefficient of 2 in front of Al on the reactants side.

** There are
two O atoms on the reactants side, and three O atoms on the product side of the
equation. We can balance the O atoms by placing a coefficient of 3/2 in front
of O

_{2}on the reactants side
** This is a
balanced equation. However, equations are normally balanced with the smallest set
of whole number coefficients. Multiplying both sides of the equation by 2 gives
whole number coefficients

** Check For an
equation to be balanced, the number and types of atoms on each side of the equation
must be the same. The final tally is

The equation is
balanced.

**Reference:***General Chemistry: The Essential Concepts / Raymond Chang , Jason Overby. (sixth edition)**.*

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