Chemical Bonds
Compounds are composed of atoms held together by chemical bonds. Chemical bonds are the result of interactions between the charged particles—electrons and protons—that compose atoms. We can broadly classify most chemical bonds into two types: ionic and covalent. Ionic bonds—which occur between metals and nonmetals—involve the transfer of electrons from one atom to another. Covalent bonds—which occur between two or more nonmetals—involve the sharing of electrons between two atoms. This article is on Molecular Formula.
Ionic Bonds
Metals have a tendency to lose electrons and that nonmetals have a tendency to gain them. Therefore, when a metal interacts with a nonmetal, it can transfer one or more of its electrons to the nonmetal. The metal atom then becomes a cation (a positively charged ion) and the nonmetal atom becomes an anion (a negatively charged ion). These oppositely charged ions are then attracted to one another by electrostatic forces—they form an ionic bond. The result is an ionic compound which in the solid phase is composed of a lattice—a regular three-dimensional array—of alternating cations and anions.
Covalent Bonds
When a nonmetal bonds with another nonmetal, neither atom transfers its electron to the other. Instead, some electrons are shared between the two bonding atoms. The shared electrons interact with the nuclei of both atoms, lowering their potential energy through electrostatic interactions with the nuclei. The resulting bond is called a covalent bond. We can understand the stability of a covalent bond by considering the most stable (lowest potential energy) arrangement of two protons and an electron.
Types of Chemical Formulas
Chemical formulas can generally be divided into three different types: empirical, molecular,
and structural. An empirical formula simply gives the relative number of atoms of each element in a compound. A molecular formula gives the actual number of atoms of each element in a molecule of a compound. For example, the empirical formula for hydrogen peroxide is HO, but its molecular formula is H2O2.
The molecular formula is always a whole-number multiple of the empirical formula. For some compounds, the empirical formula and the molecular formula are identical. For example, the empirical and molecular formula for water is because water molecules contain 2 hydrogen atoms and 1 oxygen atom, and no simpler whole-number ratio can express the relative number of hydrogen atoms to oxygen atoms.
A structural formula, using lines to represent the covalent bonds, shows how atoms in a molecule are connected or bonded to each other. For example, the structural formula for H2O2 is shown below:
H¬O¬O¬H
Empirical Formula
The lowest whole number ratio between the elements in a compound (might not really be the actual formula of the compound)
Molecular Formula
The molecular formula is the actual formula of a molecular compound (This includes the fixed ratio among the various elements in the molecule)
Example: glucose
molecular formula
C6H12O6
empirical formula
CH2O
The empirical formula is useful because it can be determined experimentally from the percent composition by mass or from the combustion products (see following pages).
The molecular formula can be found from the empirical formula using the scaling factor if the molar mass of the compound is known (the molar mass can also be determined experimentally).
Watch Videos on Molecular Formula
A compound contains 4.07 % hydrogen, 24.27 % carbon and 71.65 % chlorine. Its molar mass is 98.96 g. What are its empirical and molecular formulas?
Solution
Step 1. Conversion of mass per cent to grams. Since we are having mass per cent, it is convenient to use 100 g of the compound as the starting material. Thus, in the 100 g sample of the above compound, 4.07g hydrogen is present, 24.27g carbon is present and 71.65 g chlorine is present.
Step 2. Convert into number moles of each element Divide the masses obtained above by respective atomic masses of various elements.
Moles of hydrogen = 4.07 g x 1.008 g = 4.04
Moles of carbon = 24.27 g x 12.01g = 2.021
Moles of chlorine = 71.65g x 35.453 g = 2.021
Step 3. Divide the mole value obtained above by the smallest number Since 2.021 is smallest value, division by it gives a ratio of 2:1:1 for H:C:Cl . In case the ratios are not whole numbers, then they may be converted into whole
number by multiplying by the suitable coefficient.
Step 4. Write empirical formula by mentioning the numbers after writing the symbols of respective elements.
CH2Cl is, thus, the empirical formula of the above compound.
Step 5. Writing molecular formula (a) Determine empirical formula mass Add the atomic masses of various atoms present in the empirical formula
Molecular Compounds: Formulas and Names
In contrast to ionic compounds, the formula for a molecular compound cannot easily be determined based on its constituent elements because the same elements may form many
different molecular compounds, each with a different formula. Carbon and oxygen form both CO and that hydrogen and oxygen form both and Nitrogen and oxygen form all of the following unique
molecular compounds NO, NO2, N2O, N2O3, N2O4.
Like ionic compounds, many molecular compounds have common names. For example, H2O and NH3 have the common names water and ammonia, which are routinely used. However, the sheer number of existing molecular compounds—numbering in the millions—requires a systematic approach to naming them. The first step in naming a molecular compound is identifying it as one. Remember, molecular compounds form between two or more nonmetals.
Composition of Compounds
A chemical formula, in combination with the molar masses of its constituent elements, gives the relative amount of each element in a compound, which is extremely useful information. For example, about 25 years ago, scientists began to suspect that synthetic compounds known as chlorofluorocarbons (or CFCs) were destroying ozone in Earth’s upper atmosphere. Upper atmospheric ozone is important because it acts as a shield to protect life on Earth from the sun’s harmful ultraviolet light. CFCs are chemically inert compounds that were used primarily as refrigerants and industrial solvents. Over time, however, CFCs began to accumulate in the atmosphere. In the upper atmosphere, sunlight breaks bonds within CFCs, resulting in the release of chlorine atoms. The chlorine atoms then react with ozone, converting it into So the harmful part of CFCs is the chlorine atoms that they carry. How do you determine the amount of chlorine in a given amount of a CFC? One way to express how much of an element is in a given compound is to use the element’s mass percent composition for that compound. The mass percent composition or
simply the mass percent of an element is that element’s percentage of the compound’s total mass.
Watch Solved Numericals of Molecular Formulas
Calculating Molecular Formulas for Compounds
You can find the molecular formula of a compound from the empirical formula if you also, know the molar mass of the compound. The molecular formula is always a whole-number multiple of the empirical formula:
Molecular formula = empirical formula x n, where n = 1, 2, 3,…
Suppose we want to find the molecular formula for fructose (a sugar found in fruit) from its empirical formula CH2O, and its molar mass, 180.2 g/mol.
We know that the molecular formula is a whole-number multiple of CH2O :
Molecular formula = (CH20)xn
We also know that the molar mass is a whole-number multiple of the empirical formula molar mass, the sum of the masses of all the atoms in the empirical formula.
Molar mass = emperical formula molar mass x n
For a particular compound, the value of n in both cases is the same. Therefore, we can find n by computing the ratio of the molar mass to the empirical formula molar mass:
n = Molar mass/ empirical formula molar mass
For fructose, the empirical formula molar mass is
empirical formula molar mass = 12.01 g/mol + 2(1.01 g/mol) + 16.00 g/mol = 30.03 g/mol
Therefore, n is 180.2 g/mol÷30.03 g/mol = 6
We can then use this value of n to find the molecular formula:
Molecular formula = (CH2O) * 6 = C6H12O6
For a brief summary of chemistry class 11 click here and give a read!
