as possible. While searching for various atomic mass units, scientists initially took 1/ 16 of the mass of an atom of naturally occurring oxygen as the unit. This was considered relevant due to two reasons: • oxygen reacted with a large number of elements and formed compounds. • this atomic mass unit gave masses of most of the elements as whole numbers. However, in 1961 for a universally accepted atomic mass unit, carbon-12 isotope was chosen as the standard reference for measuring atomic masses. One atomic mass unit is a mass unit equal to exactly one-twelfth (1/12th) the mass of one atom of carbon-12. The relative atomic masses of all elements have been found with respect to an atom of carbon-12. Imagine a fruit seller selling fruits without any standard weight with him. He takes a watermelon and says, “this has a mass equal to 12 units” (12 watermelon units or 12 fruit mass units). He makes twelve equal pieces of the watermelon and finds the mass of each fruit he is selling, relative to the mass of one piece of the watermelon. Now he sells his fruits by relative fruit mass unit (fmu), as in Fig. 3.4. Fig. 3.4 : (a) Watermelon, (b) 12 pieces, (c) 1/12 of watermelon, (d) how the fruit seller can weigh the fruits using pieces of watermelon Similarly, the relative atomic mass of the atom of an element is defined as the average mass of the atom, as compared to 1/12th the mass of one carbon-12 atom. Table 3.2: Atomic masses of a few elements Element Atomic Mass (u) Hydrogen 1 Carbon 12 Nitrogen 14 Oxygen 16 Sodium 23 Magnesium 24 Sulphur 32 Chlorine 35.5 Calcium 40 3.2.3 HOWDOATOMSEXIST? Atoms of most elements are not able to exist independently. Atoms form molecules and ions. These molecules or ions aggregate in large numbers to form the matter that we can see, feel or touch. uestions Q 1. Define the atomic mass unit. 2. Why is it not possible to see an atom with naked eyes? 3.3 What is a Molecule? A molecule is in general a group of two or more atoms that are chemically bonded together, that is, tightly held together by attractive forces. A molecule can be defined as the smallest particle of an element or a compound that is capable of an independent existence and shows all the properties of that substance. Atoms of the same element or of different elements can join together to form molecules. ATOMSAND MOLECULES Element Ratio by mass Atomic mass (u) Mass ratio/ atomic mass Simplest ratio H O 1 8 1 16 1 1 =1 8 16 = 1 2 2 1 What you have learnt • During a chemical reaction, the sum of the masses of the reactants and products remains unchanged. This is known as the Law of Conservation of Mass. • In a pure chemical compound, elements are always present in a definite proportion by mass. This is known as the Law of Definite Proportions. • An atom is the smallest particle of the element that cannot usually exist independently and retain all its chemical properties. • A molecule is the smallest particle of an element or a compound capable of independent existence under ordinary conditions. It shows all the properties of the substance. • A chemical formula of a compound shows its constituent elements and the number of atoms of each combining element. • Clusters of atoms that act as an ion are called polyatomic ions. They carry a fixed charge on them. • The chemical formula of a molecular compound is determined by the valency of each element. • In ionic compounds, the charge on each ion is used to determine the chemical formula of the compound. • Scientists use the relative atomic mass scale to compare the masses of different atoms of elements. Atoms of carbon-12 isotopes are assigned a relative atomic mass of 12 and the relative masses of all other atoms are obtained by comparison with the mass of a carbon-12 atom. • The Avogadro constant 6.022 × 1023 is defined as the number of atoms in exactly 12 g of carbon-12. • The mole is the amount of substance that contains the same number of particles (atoms/ ions/ molecules/ formula units etc.) as there are atoms in exactly 12 g of carbon-12. • Mass of 1 mole of a substance is called its molar mass. Exercises 1. A 0.24 g sample of compound of oxygen and boron was found by analysis to contain 0.096 g of boron and 0.144 g of oxygen. Calculate the percentage composition of the compound by weight. 2. When 3.0 g of carbon is burnt in 8.00 g oxygen, 11.00 g of carbon dioxide is produced. What mass of carbon dioxide will ATOMSAND MOLECULES

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