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Summary Stoichiometry:Calculations with chemical formulas and equations
- Course
- Introduction to chemistry
- Institution
- University Of South Africa (Unisa)
Stoichiometry: Calculations with chemical formulas and equations
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1Week 5 : Stoichiometry: Calculations with chemical
formulas and equations
Learning Objectives:
1. Balance chemical equations.
2. Write a chemical equation from a word reaction and convert a word reaction into a
chemical equation.
3. Classify a reaction as a combination, decomposition, combustion, single or double
replacement reaction.
4. Complete a chemical equation and assign physical states to reactants and
products.
5. Calculate the formula weight (molar mass) of a molecule/compound.
6. Calculate percentage compositions for elements in a molecule/compound and
elements in
a sample.
7. Convert between the number of moles, mass and molar mass for a substance.
8. Convert between the number of moles and the number of units
(particles/molecules/ions/atoms)
using Avogadro’s constant.
9. Calculate the mass/moles/number of units of an element or ion in a chemical
formula using
stoichiometry.
10. Use stoichiometry in a balanced equation to convert quantitative information of
one
substance into quantitative information for another substance in the reaction.
11. Derive the empirical formula and molecular formula for a given compound.
12. Determine the limiting reagent and excess reagent in a chemical reaction.
13. Determine the theoretical and percentage yield of a product.
How to balance chemical equations
We represent chemical reactions by chemical equations. When the gas hydrogen
1H22
burns, for example, it reacts with oxygen 1O22 in the air to form water 1H2O2. We
write
the chemical equation for this reaction as
2 H2 + O2 ¡ 2 H2O [3.1]
We read the + sign as “reacts with” and the arrow as “produces.” The chemical
formulas to the left of the arrow represent the starting substances, called reactants.
The
, 2
chemical formulas to the right of the arrow represent substances produced in the
reaction, called products. The numbers in front of the formulas, called coefficients,
indicate
the relative numbers of molecules of each kind involved in the reaction. (As in
algebraic
equations, the coefficient 1 is usually not written.)
Because atoms are neither created nor destroyed in any reaction, a balanced
chemical
equation must have an equal number of atoms of each element on each side of the
arrow.
On the right side of Equation 3.1, for example, there are two molecules of H2O, each
composed of two atoms of hydrogen and one atom of oxygen (Figure 3.2). Thus, 2
H2O (read
“two molecules of water”) contains 2 * 2 = 4 H atoms and 2 * 1 = 2 O atoms. Notice
that the number of atoms is obtained by multiplying each subscript in a chemical
formula by
the coefficient for the formula. Because there are four H atoms and two O atoms on
each
side of the equation, the equation is balanced.
Chemists write unbalanced equations all the time in order to identify the reactants
and
products in a reaction. But to determine the amount of product that can be made, or
the
amount of a reactant that is required, the chemical equation needs to be balanced.
To construct a balanced chemical equation we start by writing the formulas for the
reactants on the left-hand side of the arrow and the products on the right-hand side.
Next we balance the equation by determining the coefficients that provide equal
numbers of each type of atom on both sides of the equation. For most purposes, a
balanced
equation should contain the smallest possible whole-number coefficients.
In balancing an equation, you need to understand the difference between
coefficients and subscripts. As Figure 3.3 illustrates, changing a subscript in a
formula—
from H2O to H2O2, for example—changes the identity of the substance. The
substance
H2O2, hydrogen peroxide, is quite different from the substance H2O, water. Never
change
subscripts when balancing an equation. In contrast, placing a coefficient in front of a
formula changes only the amount of the substance and not its identity. Thus, 2 H2O
means
two molecules of water, 3 H2O means three molecules of water, and so forth.
Steps to Balancing Chemical Equations
, 3
1. Count the number of atoms of each element in the reactants and the
products. List each element and how many atoms it has underneath
the reactants and products.
Example:
HCl + Na2S H2S + NaCl
H–1H–2
Cl – 1 Cl – 1
Na – 2 Na – 1
S–1S–1
2. To balance out the atoms add in coefficients.
Example:
2 HCl + Na2S H2S + NaCl
H–12H–2
Cl – 1 2 Cl – 1
Na – 2 Na – 1
S–1S–1
Example:
2 HCl + Na2S H2S + 2 NaCl
H–12H–2
Cl – 1 2 Cl – 1 2
Na – 2 Na – 1 2
S–1S–1
3. Make sure that there are equal numbers of atoms for each element on
both sides of the equation.
Chemical Equation
A chemical equation shows the overall change of reactants to products in a chemical
reaction.
Sometimes, state symbols are required to indicate the physical states of the
substances in a chemical reaction.
The following table gives the physical states and the state symbols used in chemical
equations: solid, liquid, gas, aqueous.
, 4
The following table gives the valency of some common ions. This table can be used
to help you work out the chemical formula of the reactants and products.
Here are some simple covalent formulas that you will find useful to remember:
● Water H2O,
● Carbon Dioxide CO2,
● Ammonia NH3,
● Hydrogen H2,
● Oxygen O2,
● Nitrogen N2,
● Sulfur Dioxide or Sulphur Dioxide SO2,
● Methane CH4
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