Atoms And Molecules
Class 10 Unit 7 Science
UNIT-7
ATOMS AND MOLECULES
Learning Objectives
At the end of the lesson the students will be able to:
- acquire the ability to learn about the atoms and molecules.
- comprehend atomic mass and molecular mass.
- have information about gram atomic mass and gram molecular mass.
- perceive the intended meaning of Avogadro's hypothesis of gases.
- interpret the application of Avogadro's hypothesis.
- determine the atomicity of a molecule.
- interpret the relation between vapour density and relative molecular mass.
- have the facts about the relationship between the volume of a gas and the number of molecules present in it.
- grasp the idea of mole concept and solve many problems using it.
- Calculate the percentage of composition of a compound.
INTRODUCTION
You have learnt, in your lower classes that matter is around us everywhere. Matter is made of atoms. Curiously the idea of atom was first proposed by the Greek philosophers in the fifth century BC (BCE). But their theory was more philosophical than scientific. The first scientific theory of the atom was proposed by John Dalton. Few of the postulates of Dalton's theory about an atom were found incorrect by the later on studies made by J.J. Thomson, Rutherford, Neils Bohr and Schrodinger. In the light of the result of the research most of the limitations of the Dalton's theory were removed and a new theory known as the modern atomic theory was put forward.
The main postulates of modern atomic theory are as follows:
- An atom is no longer indivisible (after the discovery of the electron, proton, and neutron).
- Atoms of the same element may have different atomic mass. (Discovery of isotopes 17Cl35, 17Cl37).
- Atoms of different elements may have same atomic masses (discovery of Isobars 18Ar40, 20Ca40).
- Atoms of one element can be transmuted into atoms of other elements. In other words, atom is no longer indestructible (discovery of artificial transmutation).
- Atoms may not always combine in a simple whole number ratio (Example Glucose C6H12O6 C is to H is to O = 6 is to 12 is to 6 or 1 is to 2 is to 1 and Sucrose C12H22O11 C is to H is to O = 12 is to 22 is to 11).
- Atom is the smallest particle that takes part in a chemical reaction.
- The mass of an atom can be converted into energy (E = mc²).
The modern atomic theory is the basis for all the studies of chemical and physical processes that involve atoms. You have studied the most fundamental ideas about an atom in your lower classes. Let us discuss some more concepts about atoms in this lesson.
7.1 ATOM AND ATOMIC MASS
As you know, anything that has mass and occupies space is called matter. Atoms are the building blocks of matter. Since matter has mass, it must be due to its atoms. According to the modern atomic theory, an atom contains subatomic particles such as protons, neutrons and electrons. Protons and neutrons have considerable mass, but electrons don't have such a considerable mass. Thus, the mass of an atom is mainly contributed by its protons and neutrons and hence the sum of the number of protons and neutrons of an atom is called its mass number.
Individual atoms are very small, and it is difficult to measure their masses. You can measure the mass of macroscopic materials in gram or kilogram. The mass of an atom is measured in atomic mass unit (amu). Atomic mass unit is one-twelfth of the mass of a carbon-12 atom; an isotope of carbon, which contains 6 protons and 6 neutrons. (Note: The symbol 'amu' is no longer used in the modern system and instead, it uses the symbol 'u' to denote unified atomic mass. The mass of a proton or neutron is approximately 1 amu).
7.1.1 Relative Atomic Mass (RAM)
As an atom is very small, its absolute mass cannot be determined directly. The early pioneers of chemistry used to measure the atomic mass of an atom relative to an atom of another element. They measured the masses of equal number of atoms of two or more elements at a time, to determine their relative masses. They established one element as a standard, gave it an arbitrary value of atomic mass and using this value they measured the relative mass of other elements. The mass obtained by this way is called relative atomic mass.
In the beginning, the mass of hydrogen atom was chosen as a standard and masses of other atoms were compared with it, because of the existence of isotopic character of hydrogen (1H1, 1H2, 1H3). Later hydrogen atom was replaced by oxygen atom as the standard. Now, the stable isotope of carbon (C-12) with atomic mass 12 is used as the standard for measuring the relative atomic mass of an element.
Relative atomic mass of an element is the ratio between the average mass of its isotopes to 1/12th part of the mass of a carbon-12 atom. It is denoted as Ar. It is otherwise called "Standard Atomic Weight".
Relative Atomic Mass = Average mass of its isotopes / (1/12) mass of C-12 atom
Modern methods of determination of atomic mass by Mass Spectrometry uses C-12 as standard. For most of the elements, the relative atomic mass is very closer to a whole number, and it is rounded off to a whole number, to make calculations easier.
Table 7.1 Relative atomic mass of elements (C-12 Scale) | Element | Symbol | Relative atomic mass |
|---|
| Hydrogen | H | 1 |
| Carbon | C | 12 |
| Nitrogen | N | 14 |
| Oxygen | O | 16 |
| Sodium | Na | 23 |
| Magnesium | Mg | 24 |
| Sulphur | S | 32 |
Do You Know?
Relative Atomic Mass is only a ratio, so it has no unit.
If the atomic mass of an element is expressed in grams, it is called as Gram Atomic Mass
- Gram Atomic Mass of hydrogen = 1 gram
- Gram Atomic Mass of carbon = 12 gram
- Gram Atomic Mass of nitrogen = 14 gram
- Gram Atomic Mass of oxygen = 16 gram
7.0.2 Average Atomic Mass (AAM)
How can one measure the atomic mass of an element? It is somewhat more complicated because most of the naturally occurring elements exist as a mixture of isotopes, each of which has its own mass. Thus, it is essential to consider this isotopic mixture while calculating the atomic mass of an element. The average atomic mass of an element is the weighted average of the masses of its naturally occurring isotopes. But, the abundance of isotopes of each element may differ. So, the abundancy of all these isotopes is taken into consideration while calculating the atomic mass.
Then, what do we mean by a weighted average? Let us consider an element which exists as a mixture of 50% of an isotope having a mass of 9 amu, and 50% of another isotope having a mass of 10 amu. Then, its average atomic mass is calculated by the following equation:
Average atomic mass = (Mass of 1st isotope × % abundance of 1st isotope) + (Mass of 2nd isotope × % abundance of 2nd isotope)
Thus, for the given element the average atomic mass = (9 × 0.50) + (10 × 0.50) = 4.5 + 5.0 = 9.5 amu
(Note: In the calculations involving percentages, you need to convert percentage abundance into fractional abundance. For example, 50 percent is converted into 50/100 or 0.50 as shown in a foresaid calculation.)
The atomic masses of elements, given in the periodic table, are average atomic masses. Sometimes, the term atomic weight is used to mean average atomic mass. It is observed, from the periodic table that atomic masses of most of the elements are not whole numbers. For instance, the atomic mass of carbon given in the periodic table is 12.01 amu, not 12.00 amu. The reason is that while calculating the atomic mass of carbon, both of its natural isotopes such as carbon-12 and carbon-13 are considered. The natural abundance of C-12 and C-13 are 98.90% and 1.10% respectively. The average of the atomic mass of carbon is calculated as follows:
Average atomic mass of carbon = (12 × 0.9890) + (13 × 0.0110) = 11.868 + 0.143 = 12.011 amu
So, it is important to understand that if it is mentioned that the atomic mass of carbon is 12 amu, it refers to the average atomic mass of the carbon isotopes, not the mass of the individual atoms of carbon.
Table 7.2 Atomic mass of some elements | Atomic Number | Name | Symbol | Atomic Mass (amu) |
|---|
| 1 | Hydrogen | H | 1.008 |
| 2 | Helium | He | 4.003 |
| 3 | Lithium | Li | 6.941 |
| 4 | Beryllium | Be | 9.012 |
| 5 | Boron | B | 10.811 |
Calculation of average atomic mass – Solved Examples
Example 1: Oxygen is the most abundant element in both the Earth's crust and the human body. It exists as a mixture of three stable isotopes in nature as shown in Table 7.3:
Table 7.3 Isotopes of oxygen | Isotope | Mass (Atomic Mass Unit) | % abundance |
|---|
| 8O16 | 15.9949 | 99.757 |
| 8O17 | 16.9991 | 0.038 |
| 8O18 | 17.9992 | 0.205 |
The atomic mass of Oxygen = (15.9949×0.99757) + (16.9991× 0.00038) + (17.9992×0.00205) = 15.999 amu.
Example 2: Boron naturally occurs as a mixture of boron-10 (5 protons + 5 neutrons) and boron-11 (5 protons + 6 neutrons) isotopes. The percentage abundance of B-10 is 20 and that of B-11 is 80. Then, the atomic mass of boron is calculated as follows:
Atomic mass of boron = (10 × 20/100) + (11 × 80/100) = (10 × 0.20) + (11 × 0.80) = 2 + 8.8 = 10.8 amu
7.2 MOLECULE AND MOLECULAR MASS
Except noble gases, atoms of most of the elements are found in the combined form with itself or atoms of other elements. It is called as a molecule. A molecule is a combination of two or more atoms held together by strong chemical forces of attraction, that is chemical bonds.
Activity 7.1
Complete the following table by filling the appropriate values/terms
| Element | Number of Protons | Number of Neutrons | Mass Number | Stable Isotopes (abundance) | Atomic Mass (amu) |
|---|
| Nitrogen | 7 | | | N-14(99.6%) N-15(0.4%) | |
| Silicon | 14 | | 28 | Si-28(92.2%) Si-29(4.7%) Si-30(3.1%) | |
| Chlorine | | | | Cl-35(75%) Cl-37(25%) | |
7.2.1 Classification of molecules
A molecule may contain atoms of the same element or may contain atoms of two or more elements joined in a fixed ratio, in accordance with the law of definite proportions. Thus, a molecule may be an element or a compound.
If the molecule is made of similar kind of atoms, then it is called homoatomic molecule. The molecule that consists of atoms of different elements is called heteroatomic molecule. A compound is a heteroatomic molecule. The number of atoms present in the molecule is called its 'atomicity'.
Table 7.4 Classification of molecules | Atomicity | Number of atoms present | Name |
|---|
| 1 | 1 | Monoatomic |
| 2 | 2 | Diatomic |
| 3 | 3 | Triatomic |
| >3 | >3 | Polyatomic |
Let us consider oxygen. Oxygen gas exists in two allotropic forms: Oxygen (O2) and Ozone (O3). In oxygen molecule, there are two oxygen atoms. So, its atomicity is two. Since both the atoms are similar, oxygen (O2) is a homo diatomic molecule. Other elements that exist as diatomic molecules are hydrogen (H2), nitrogen (N2) and halogens: fluorine (F2), chlorine (Cl2), bromine (Br2) and iodine (I2).
Ozone (O3) contains three oxygen atoms and hence it is called homo triatomic molecule. If a molecule contains more than three atoms, then it is called polyatomic molecule.
Consider hydrogen chloride. It consists of two atoms, but of different elements, that is hydrogen and chlorine. So, its atomicity is two. It is a hetero diatomic molecule. Similarly, the water molecule contains two hydrogen atoms and one oxygen atom. So, its atomicity is three. It is a hetero triatomic molecule.
Activity 7.2
Classify the following molecules based on their atomicity and fill in the table:
Fluorine(F2), Carbon dioxide(CO2), Phosphorous(P4), Sulphur(S8), Ammonia(NH3), Hydrogen iodide(HI), Sulphuric acid(H2SO4), Methane(CH4), Glucose(C6H12O6), Carbon monoxide(CO)
| Molecule | Diatomic | Triatomic | Polyatomic |
|---|
| Homo | Hetero | Homo | Hetero | Homo | Hetero |
|---|
| | | | | | |
7.2.2 Relative Molecular Mass (RMM)
As the molecules are made of atoms, they also have their own mass. The mass of the molecule of an element or compound is measured in the C-12 scale and hence called relative molecular mass.
The Relative Molecular Mass of a molecule is the ratio between the mass of one molecule of the substance to 1/12th mass of an atom of Carbon 12
Do You Know?
Relative Molecular Mass is only a ratio. So, it has no unit.
If the molecular mass of a compound is expressed in grams, it is called Gram Molecular Mass.
- Gram Molecular Mass of water = 18 gram
- Gram Molecular Mass of carbon dioxide = 44 gram
- Gram Molecular Mass of ammonia = 17 gram
- Gram Molecular Mass of HCl = 36.5 gram
The relative molecular mass is obtained by adding together the relative atomic masses of all the atoms present in a molecule.
Calculation of relative molecular mass – Solved examples:
Example 1: Relative molecular mass of sulphuric acid (H2SO4) is calculated as follows:
Sulphuric acid contains 2 atoms of hydrogen, 1 atom of sulphur and 4 atoms of oxygen.
Therefore, Relative molecular mass of sulphuric acid = (2 × mass of hydrogen) + (1 × mass of sulphur) + (4 × mass of oxygen) = (2 × 1) + (1 × 32) + (4 × 16) = 98
That is one molecule of H2SO4 is 98 times as heavy as 1/12th of the mass of a carbon 12.
Example 2: Relative molecular mass of water (H2O) is calculated as follows:
A water molecule is made of 2 atoms of hydrogen and one atom of oxygen.
So, the relative molecular mass of water = (2 × mass of hydrogen) + (1 × mass of oxygen) = (2 × 1) + (1 × 16) = 18
That is one molecule of H2O is 18 times as heavy as 1/12th of the mass of a carbon 12.
7.3 DIFFERENCE BETWEEN ATOMS AND MOLECULES
Even though atoms are the basic components of molecules, they differ in many aspects when compared to the molecules. Table 7.5 consolidates the major difference between atoms and molecules.
Table 7.5 Difference between atoms and molecules | Atom | Molecule |
|---|
| An atom is the smallest particle of an element | A molecule is the smallest particle of an element or compound. |
| Atom does not exist in free state except in a noble gas | Molecule exists in free state |
| Except some of noble gas, other atoms are highly reactive | Molecules are less reactive |
| Atom does not have a chemical bond | Atoms in a molecule are held by chemical bonds |
7.4 MOLE CONCEPT
So far, we discussed about matters in terms of individual atoms and molecules. Atomic mass units provide a relative scale for the masses of the elements. Since the atoms have such small masses, no usable scale can be devised to weigh them in the calibrated units of atomic mass units. In any real situation, we deal with macroscopic samples containing enormous number of atoms. Therefore, it is convenient to have a special unit to describe a very large number of atoms.
The idea of a 'unit' to denote a particular number of objects is not new. For example, the pair (2 items) and the dozen (12 items), are all familiar units. Chemists measure atoms and molecules in moles. So, you can now understand that 'mole' denotes a number of particles.
In the SI system, the mole (mol) is the amount of a substance that contains as many elementary entities (atoms, molecules, or other particles) as there are atoms in exactly 12 gram (or 0.012 kg) of the carbon 12 isotope.
The actual number of atoms in 12 gram of carbon 12 is determined experimentally. This is called Avogadro's Number (NA), named after an Italian scientist Amedeo Avogadro who proposed its significance. Its value is 6.023 × 10²³. So, one mole of a substance contains 6.023 × 10²³ entities. Thus, 5 moles of oxygen molecules contain 5 × 6.023 × 10²³ molecules.
Mole Concept:
The study of the collection of particles by using mole as the counting unit, in order to express the mass and volume of such unit particles in a bulk of matter is known as mole concept.
The number of moles of a substance can be calculated by various means depending on the data available, as follows:
- Number of moles of molecules.
- Number of moles of atoms.
- Number of moles of a gas (Standard molar volume at STP = 22.4 litre).
- Number of moles of ions.
Note: STP-Standard Temperature and Pressure (273.15 Kelvin, 1.00 standard atmosphere pressure)
Mole of atoms:
One mole of an element contains 6.023 × 10²³ atoms and it is equal to its gram atomic mass. That is, one mole of oxygen contains 6.023 × 10²³ atoms of oxygen, and its gram atomic mass is 16 gram.
Mole of molecules:
One mole of matter contains 6.023 × 10²³ molecules, and it is equal to its gram molecular mass. That is, one mole of oxygen contains 6.023 × 10²³ molecules of oxygen and its gram molecular mass is 32 gram.
Molar volume:
One mole of any gas occupies 22.4 litre or 22400 ml at S.T.P. This volume is called as molar volume.
Calculation of number of moles by Different modes
Number of moles = Mass/Atomic Mass = Mass/Molecular mass = Number of Atoms/(6.023 × 10²³) = Number of Molecules/(6.023 × 10²³)
7.5 PERCENT COMPOSITION
So, for, we were dealing with the number of entities present in a given substance. But many times, the information regarding the percentage of a particular element present in a compound is required. The percentage composition of a compound represents the mass of each element present in 100 gram of the compound. Let us understand the percentage composition of oxygen and hydrogen by taking the example of H2O. It can be calculated using the formula:
% of an element = (Mass of element in the compound/Molecular mass of compound) × 100
This percentage composition is useful to determine the empirical formula and molecular formula.
Example 1: Find the mass percentage composition of methane (CH4).
Molecular mass of CH4 = 12 + (4 × 1) = 16
% of Carbon = (12/16) × 100 = 75%
% of Hydrogen = (4/16) × 100 = 25%
7.6 AVOGADRO HYPOTHESIS
In 1811 Avogadro framed a hypothesis based on the relationship between the number of molecules present in equal volumes of gases in different conditions. The Avogadro's law states that "equal volumes of all gases under similar conditions of temperature and pressure contain equal number of molecules"
It follows that the volume of any given gas must be proportional to the number of molecules in it. If 'V' is the volume and 'n' is the number of molecules of a gas, then Avogadro law is represented, mathematically, as follows:
V ∝ n
V = constant × n
Thus, one litre (1 dm³) of hydrogen contains the same number of molecules as in one litre of oxygen, that is, the volume of the gas is directly proportional to the number of molecules of the gas.
Explanation
Let us consider the reaction between hydrogen and chlorine to form hydrogen chloride gas
H2 + Cl2 → 2HCl
1 volume + 1 volume → 2 volumes
According to Avogadro's law 1 volume of any gas is occupied by "n" number of molecules.
n molecules + n molecules → 2n molecules
If n=1 then 1 molecule + 1 molecule → 2 molecules.
½ molecule + ½ molecule → 1 molecule
1 molecule of hydrogen chloride gas is made up of ½ molecule of hydrogen and ½ molecule of chlorine. Hence, the molecules can be subdivided. This law is in agreement with Dalton's atomic theory.
Activity 7.3
Under same conditions of temperature and pressure if you collect 3 litres of O2, 5 litres of Cl2 and 6 litres of H2, Which has the highest number of molecules? Which has the lowest number of molecules?
7.7 APPLICATIONS OF AVOGADRO'S LAW
- It explains Gay-Lussac's law.
- It helps in the determination of atomicity of gases.
- Molecular formula of gases can be derived using Avogadro's law
- It determines the relation between molecular mass and vapour density.
- It helps to determine gram molar volume of all gases (i.e., 22.4 litre at S.T.P)
7.8 RELATIONSHIP BETWEEN VAPOUR DENSITY AND RELATIVE MOLECULAR MASS
1. Relative molecular mass: (Hydrogen scale)
The Relative Molecular Mass of a gas or vapour is the ratio between the mass of one molecule of the gas or vapour to mass of one atom of Hydrogen.
2. Vapour Density:
Vapour density is the ratio of the mass of a certain volume of a gas or vapour, to the mass of an equal volume of hydrogen, measured under the same conditions of temperature and pressure.
Vapour Density (V.D.) = Mass of certain volume of gas or vapour / Mass of equal volume of hydrogen
According to Avogadro's law, equal volumes of all gases contain equal number of molecules. Thus, let the number of molecules in one volume = n, then
Vapour density = (n × molecular mass of gas) / (n × molecular mass of hydrogen)
Cancelling 'n' which is common, you get
Vapour density = Molecular mass of gas / Molecular mass of hydrogen
However, since hydrogen is diatomic
Molecular mass of hydrogen = 2
Vapour density = Molecular mass of gas / 2
When you compare the formula of vapour density with relative molecular mass, they can be represented as
Vapour density = Molecular mass of gas / 2 (Equation 7.1)
Relative molecular mass (hydrogen scale) = Molecular mass of gas / 1 (Equation 7.2)
You can therefore substitute the above equation to an Equation 7.1 and arrive at the following formula
Vapour density = Relative molecular mass / 2
Now on cross multiplication, you have
2 × vapour density = Relative molecular mass of a gas
(Or) Relative molecular mass = 2 × Vapour density
7.9 SOLVED PROBLEMS
7.10 I. Calculation of molecular mass
Calculate the gram molecular mass of the following.
1) H2O
Atomic masses of H = 1, O = 16
Gram molecular mass of H2O = (1×2) + (16×1) = 2+16
Gram molecular mass of H2O = 18 gram
2) CO2
Atomic masses of C=12, O=16
Gram molecular mass of CO2 = (12×1) + (16×2) = 12+32
Gram molecular mass of CO2 = 44 gram
3) Ca3(PO4)2
Atomic masses of Ca=40, P=30, O=16.
Gram molecular mass of Ca3(PO4)2 = (40×3) + [30+(16×4)] × 2 = 120+(94×2) = 120+188
Gram molecular mass of Ca3(PO4)2 = 308 gram
2. Calculation based on number of moles from mass and volume
1) Calculate the number of moles in 46 g of sodium?
Number of moles = 46/23 = 2 moles of sodium
2) 5.6 litre of oxygen at S.T.P
Number of moles of oxygen = 5.6/22.4 = 0.25 moles of oxygen
3) Calculate the number of moles of a sample that contains 12.046 × 10²³ atoms of iron?
Number of moles = 12.046 × 10²³ / 6.023× 10²³ = 2 moles of iron
3. Calculation of mass from mole
Calculate the mass of the following
1) 0.3 mole of Aluminium (Atomic mass of Aluminium is 27)
Mass = No. of moles × atomic mass
So, mass of Al = 0.3 × 27 = 8.1 gram
2) 2.24 litre of SO2 gas at S.T.P
Molecular mass of SO2 = 32 + (16×2) = 32 + 32 = 64
Number of moles of SO2 = Given volume of SO2 at S.T.P / Molar volume SO2 at S.T.P
Number of moles of SO2 = 2.24/22.4 = 0.1 mole
Mass = No. of moles × molecular mass
Mass = 0.1 × 64 = 6.4 gram
3) 1.51 × 10²³ molecules of water
Molecular mass of H2O = 18
Number of moles = 1.51 × 10²³ / 6.023 × 10²³ = 1/4 = 0.25 mole
0.25 = mass/18
Mass = 0.25 × 18 = 4.5 gram
4) 5 × 10²³ molecules of glucose?
Molecular mass of glucose = 180
Mass = (180 × 5 × 10²³) / (6.023 × 10²³) = 149.43 gram
IV. Calculation based on number of atoms or molecules.
1) Calculate the number of molecules in 11.2 litre of CO2 at S.T.P
Number of moles = 11.2/22.4 = 0.5 mole
Number of molecules of CO2 = number of moles of CO2 × Avogadro's number
= 0.5 × 6.023 × 10²³ = 3.011 × 10²³ molecules of CO2
2) Calculate the number of atoms present in 1 gram of gold (Atomic mass of Gold=198)
Number of atoms of Gold = (1/198) × 6.023 × 10²³
Number of atoms of Gold = 3.042 × 10²¹ atoms
3) Calculate the number of molecules in 54 gm of H2O?
Number of molecules of water = (6.023 × 10²³ × 54)/18 = 18.069 × 10²³ molecules
4) Calculate the number of atoms of oxygen and carbon in 5 moles of CO2.
1 mole of CO2 contains 2 moles of oxygen
5 moles of CO2 contain 10 moles of oxygen
Number of atoms of oxygen = Number of moles of oxygen × Avogadro's number
= 10 × 6.023 × 10²³ = 6.023 × 10²⁴ atoms of Oxygen
1 mole of CO2 contains 1 mole of carbon
5 moles of CO2 contain 5 moles of carbon
No. of atoms of carbon = No. of moles of carbon × Avogadro's number
= 5 × 6.023 × 10²³ = 3.011 × 10²⁴ atoms of Carbon
5. Calculation based on molar volume
Calculate the volume occupied by:
1) 2.5 mole of CO2 at S.T.P
Volume of CO2 at S.T.P = 22.4 × 2.5 = 56 litres.
2) 12.046 × 10²³ of ammonia gas molecules
Number of moles = Number of molecules/Avogadro's number
= 12.046 × 10²³ / 6.023 × 10²³ = 2 moles
Volume occupied by NH3 = number of moles × molar volume = 2 × 22.4 = 44.8 litres at S.T.P
3) 14 gram nitrogen gas
Number of moles = 14/28 = 0.5 mole
Volume occupied by N2 at S.T.P = number of moles × molar volume = 0.5 × 22.4 = 11.2 litres.
6. Calculation based on % composition
Calculate % of Sulpher in H2SO4
Molecular mass of H2SO4 = (1 × 2) + (32 × 1) + (16 × 4) = 2 + 32 + 64 = 98 gram
% of Sulphur in H2SO4 = (mass of sulphur/molecular mass of H2SO4) × 100
= (32/98) × 100 = 32.65%
Points to Remember
- Two or more forms of an element having the same atomic number, but different mass number are called Isotopes (17Cl35, 17Cl37).
- Atoms of different elements having the same mass number, but different atomic numbers are called Isobars (18Ar40, 20Ca40).
- Atoms of different elements having the same number of neutrons, but different atomic number and different mass number are called Isotones (6C13, 7N14).
- Relative atomic mass of an element is the ratio between the mass of one atom of the element to 1/12th of the mass of the atom of carbon -12.
- Average atomic mass of an element is calculated by adding the masses of its isotopes, each multiplied by their natural abundance on the Earth.
- Relative molecular mass of a molecule is the ratio between the mass of one molecule of the substance to 1/12th of the mass of the atom of carbon – 12.
- The Avogadro's law states that "equal volumes of all gases under similar conditions of temperature and pressure contain equal number of molecules".
- The vapour density is defined as "the ratio between the masses of equal volumes of a gas (or a vapour) and hydrogen under the same condition".
- Atomicity of a monoatomic element = Molecular mass/Atomic Mass.
- Molecular mass = 2 × Vapour density.
TEXTBOOK EVALUATION
Exercise 1. Choose the best answer.
1. Which of the following has the smallest mass?
a. 6.023 × 10²³ atoms of He b. 1 atom of He c. 2 g of He d. 1 mole atoms of He
2. Which of the following is a triatomic molecule?
a. Glucose b. Helium c. Carbon dioxide d. Hydrogen
3. The volume occupied by 4.4 gram of CO2 at S.T.P
a. 22.4 litre b. 2.24 litre c. 0.24 litre d. 0.1 litre
4. Mass of 1 mole of Nitrogen atom is
a. 28 atomic mass unit b. 14 atomic mass unit c. 28 gram d. 14 gram
5. Which of the following represents 1 amu?
a. Mass of a C-12 atom b. Mass of a hydrogen atom c. 1/12th of the mass of a C-12 atom d. Mass of O-16 atom
6. Which of the following statement is incorrect?
a. 12 gram of C – 12 contains Avogadro's number of atoms. b. One mole of oxygen gas contains Avogadro's number of molecules. c. One mole of hydrogen gas contains Avogadro's number of atoms. d. One mole of electrons stands for 6.023 × 10²³ electrons.
7. The volume occupied by 1 mole of a diatomic gas at S.T.P is
a. 11.2 litre b. 5.6 litre c. 22.4 litre d. 44.8 litre
8. In the nucleus of 20Ca40, there are
a. 20 protons and 40 neutrons b. 20 protons and 20 neutrons c. 20 protons and 40 electrons d. 40 protons and 20 electrons
9. The gram molecular mass of oxygen molecule is
a. 16 g b. 18 g c. 32 g d. 17 g
10. 1 mole of any substance contains _____ molecules.
a. 6.023 × 10²³ b. 6.023 × 10⁻²³ c. 3.0115 × 10²³ d. 12.046 × 10²³ Exercise II. Fill in the blanks
1. Atoms of different elements having mass number, but atomic numbers are called isobars.
2. Atoms of one element can be transmitted into atoms of other elements by .
3. The sum of the numbers of protons and neutrons of an atom is called its .
4. Relative atomic mass is otherwise known as .
5. The average atomic mass of hydrogen is amu.
6. If a molecule is made of similar kind of atoms, then it is called atomic molecule.
7. The number of atoms present in a molecule is called its .
8. One mole of any gas occupies milli litre at S.T.P
9. Atomicity of phosphorous is .
Exercise III. Match the following
| 1. 8 gram of O2 | |
| 2. 4 gram of H2 | |
| 3. 52 gram of He | |
| 4. 112 gram of N2 | |
| 5. 35.5 gram of Cl2 | |
Exercise IV. True or False: (If false give the correct statement) 1. Two elements sometimes can form more than one compound. 2. Noble gases are Diatomic 3. The gram atomic mass of an element has no unit 4. 1 mole of Gold and Silver contain same number of atoms 5. Molar mass of CO2 is 42g. Exercise V. Assertion and Reason:
Answer the following questions using the data given below:
- A and R are correct, R explains A.
- A is correct, R is wrong.
- A is wrong, R is correct.
- A and R are correct, R doesn't explain A.
1. Assertion: Atomic mass of aluminium is 27
Reason: An atom of aluminium is 27 times heavier than 1/12th of the mass of the C – 12 atoms.
A and R are correct, R explains A. A is correct, R is wrong. A is wrong, R is correct. A and R are correct, R doesn't explain A.
2. Assertion: The Relative Molecular Mass of Chlorine is 35.5 atomic mass unit
Reason: The natural abundance of Chlorine isotopes is not equal.
A and R are correct, R explains A. A is correct, R is wrong. A is wrong, R is correct. A and R are correct, R doesn't explain A. Exercise VI. Short answer questions
1. Define: Relative atomic mass.
2. Write the different types of isotopes of oxygen and its percentage abundance.
3. Define: Atomicity
4. Give any two examples for hetero diatomic molecules.
5. What is Molar volume of a gas?
6. Find the percentage of nitrogen in ammonia.
Exercise VII. Long answer questions
1. Calculate the number of water molecule present in one drop of water which weighs 0.18 g.
2. N2 + 3H2 → 2NH3 (The atomic mass of nitrogen is 14, and that of hydrogen is 1)
1 mole of nitrogen ( gram) + 3 moles of hydrogen ( gram) gives 2 moles of ammonia ( gram).
3. Calculate the number of moles in i) 27 gram of Al ii) 1.51 × 10²³ molecules of NH4Cl
4. Give the salient features of "Modern atomic theory".
5. Derive the relationship between Relative molecular mass and Vapour density.
Exercise VIII. HOT question
1. Calcium carbonate is decomposed on heating in the following reaction CaCO3 → CaO + CO2
i. How many moles of Calcium carbonate are involved in this reaction?
ii. Calculate the gram molecular mass of calcium carbonate involved in this reaction
iii. How many moles of CO2 are there in this equation?
Exercise IX. Solve the following problems
1. How many grams are there in the following?
i. 2 moles of hydrogen molecule, H2
ii. 3 moles of chlorine molecule, Cl2
iii. 5 moles of sulphur molecule, S8
iv. 4 moles of phosphorous molecule, P4
2. Calculate the % of each element in calcium carbonate. (Atomic mass: C-12, O-16, Ca-40)
3. Calculate the % of oxygen in Al2(SO4)3. (Atomic mass: Al-27, O-16, S-32)
4. Calculate the % relative abundance of B-10 and B-11, if its average atomic mass is 10.804 atomic mass unit.
REFERENCE BOOKS
- Petrucci, Ralph H et.al. General Chemistry: Principles & Modern Applications (9th Edition). Upper Saddle River, NJ: Pearson Prentice Hall, 2007. Print.
- Raymond Chang. (2010). Chemistry. New York, NY: The Tata McGraw Hill Companies. Inc.
- Julia Burdge. (2011). Chemistry. New York, NY: The Tata McGraw Hill Companies. Inc.
INTERNET RESOURCES
- https://www2.estrellamountain.edu/faculty/farabee/biobk/BioBookCHEM1.html
- https://www.toppr.com/guides/chemistry/atoms-and-molecules/
📝 Summary
- Matter is made of atoms, which are the building blocks of all matter, and the modern atomic theory corrected several limitations of Dalton's original atomic theory through discoveries by scientists like J.J. Thomson, Rutherford, and others.
- The modern atomic theory states that atoms are no longer indivisible, atoms of the same element may have different atomic masses (isotopes), and atoms can be transmuted into other elements.
- An atom's mass is mainly contributed by protons and neutrons, with the mass number being the sum of protons and neutrons, while electrons have negligible mass contribution.
- Atomic mass is measured in atomic mass units (amu), where one atomic mass unit is defined as one-twelfth of the mass of a carbon-12 atom containing 6 protons and 6 neutrons.
- Since atoms are extremely small, their absolute mass cannot be measured directly, so scientists measure relative atomic mass by comparing the masses of equal numbers of atoms of different elements.