Class 11 CBSE Ch 1 Some basic concept of chemistry

4. Recapitulate
 In the previous part we have learnt about Atomic and Molecular
Mass

=322 u

5.  To analyze the transformations that occur between individual atoms or
molecules in a chemical reaction, it is therefore absolutely essential for
chemists to know how many atoms or molecules are contained in a measurable
quantity in the laboratory—a given mass of sample. The unit that provides this
link is the mole (mol), from the Latin moles, meaning “pile” or “heap”.
 Many familiar items are sold in numerical quantities that have unusual names.
For example, cans of soda come in a six-pack, eggs are sold by the dozen (12),
and pencils often come in a gross (12 dozen, or 144). Sheets of printer paper are
packaged in reams of 500, a seemingly large number.
 Atoms are so small, however, that even 500 atoms are too small to see or
measure by most common techniques. Any readily measurable mass of an
element or compound contains an extraordinarily large number of atoms,
molecules, or ions, so an extraordinarily large numerical unit is needed to count
them. The mole is used for this purpose.

6.  A mole is defined as the amount of a substance that contains the number of carbon
atoms in exactly 12 g of isotopically pure carbon-12. According to the most recent
experimental measurements, this mass of carbon-12 contains 6.022142 × 1023 atoms
 Just as 1 mol of atoms contains 6.022 × 1023 atoms, 1 mol of eggs contains
6.022 × 1023 eggs.
 The number in a mole is called Avogadro’s number (NA), after the 19th-century Italian
scientist who first proposed a relationship between the volumes of gases and the
numbers of particles they contain.
 Obviously, the mole is not a term we need for most things in daily life. Instead of being
used for things we encounter in daily life, the mole is used by scientists when talking
about enormous numbers of particles like atoms, molecules, and electrons
Figure : Carbon-12, with6 protonsand 6 neutrons, is the isotope
that used to define one mole.

7.  The important point is that 1 mol of carbon—or of anything else,
whether atoms, compact discs, or houses—always has the same number
of objects: 6.022 × 1023.
 The mole is so large that it is useful only for measuring very small
objects, such as atoms.
 Using the concept of the mole, we can now restate Dalton’s theory: 1
mol of a compound is formed by combining elements in amounts whose
mole ratios are small whole numbers. For example, 1 mol of water (H2O)
has 2 mol of hydrogen atoms and 1 mol of oxygen atoms.

8. Mole example

9.  The molar mass of a chemical compound is defined as the mass of a sample of that compound
divided by the amount of substance in that sample, measured in moles. The molar mass is a
bulk, not molecular, property of a substance. The molar mass is an average of many instances
of the compound, which often vary in mass due to the presence of isotopes.
 The molecular weight is very commonly used as a synonym of molar mass, particularly for
molecular compounds. The formula weight is a synonym of molar mass that is frequently used
for non-molecular compounds, such as ionic salts.
 The molar mass is an intensive property of the substance, that does not depend on the size of
the sample. In the International System of Units (SI), the base unit of molar mass is kg/mol.
However, for historical reasons, molar masses are almost always expressed in g/mol.
 For example, the average mass of a molecule of water is about 18.0153 Daltons, and the molar
mass of water is about 18.0153 g/mol.
 For chemical elements without isolated molecules, such as carbon and metals, the molar mass
is computed dividing by the number of moles of atoms instead. Thus, for example, the molar
mass of iron is about 55.845 g/mol.

10.  The molar mass of atoms of an element is given
Molar mass
Common
symbols
M
SI unit kg/mol
Other units g/mol
M(H) = 1.00797(7) × 1.000000 g/mol = 1.00797(7) g/mol
M(S) = 32.065(5) × 1.000000 g/mol = 32.065(5) g/mol
M(Cl) = 35.453(2) × 1.000000 g/mol = 35.453(2) g/mol
M(Fe) = 55.845(2) × 1.000000 g/mol = 55.845(2) g/mol.
Some elements are usually encountered as molecules, e.g. hydrogen (H2), sulfur (S8),
chlorine (Cl2). The molar mass of molecules of these elements is the molar mass of
the atoms multiplied by the number of atoms in each molecule:
M(H2) = 2 × 1.007 97(7) × 1.000000 g/mol = 2.01588(14) g/mol
M(S8) = 8 × 32.065(5) × 1.000000 g/mol = 256.52(4) g/mol
M(Cl2) = 2 × 35.453(2) × 1.000000 g/mol = 70.906(4) g/mol.

11. Example  Along with telling us the mass of one mole of
an element, molar mass also acts as a
conversion factor between the mass of a
sample and the number moles in that
sample.
 For example, 24 grams of 12C atoms would be
equal to two moles since 24 grams divided
by the mass of one mole (12) equals 2.
 Further, Avogadro’s number acts as the
conversion factor for converting between
the number of moles in a sample and the
actual number of atoms or molecules in that
sample.
 Extending our example, two moles of 12C
atoms contains 2 times 6.02 x 1023 atoms,
which equals 12.04 x 1023 atoms, which can
be written as 1.204 x 1024 atoms.

12. Watch small clip on Molar mass

13. Formula and conversion

14. How to Convert Moles to Mass?
Using the formula
m= n x M

15. The Mole andMolar Mass

17. Q. Calculate number of moles in
392g of H2SO4
Put the value in given formula
n = 392/ 98
= 4 mol
Given: mass = 392 g moles = ?

18. Q. Calculate number of atoms of the
constituent Element in 53 g of Na2CO3
First, we need to find out no. of moles then no. of molecules and finally no. of
atoms
For that first we need to find out GMM of Sodium carbonate i.e. 106 g
Continue to next slide

19. Now, if one mole = 6.022×10²³ mol⁻¹ then
0.5 mole will be,

21. Criss Cross method for making formula