Structure of Atom Class 9 PDF

CLASS-IX CHAPTER-4 STRUCTURE OF ATOM Essentially, the structure of an atom comprises of protons, neutrons and electrons. These basic components provide the mass and charge of the atoms. The nucleus comprises of proton and neutron, with the electron orbiting around...

CLASS-IX CHAPTER-4 STRUCTURE OF ATOM Essentially, the structure of an atom comprises of protons, neutrons and electrons. These basic components provide the mass and charge of the atoms. The nucleus comprises of proton and neutron, with the electron orbiting around that. Atoms  Atoms are the basic building blocks of matter.  Different kinds of matter exist because there are different kinds of atoms present in them. Charged Particles in Matter  Whenever we rub two objects together, they become electrically charged. This is because atoms contain charged particles in them. Therefore, atoms can be divided further into particles i.e. proton, electron and neutron. It was known by 1900 that the atom was indivisible particle but contained at least one sub-atomic particle – the electron identified by J.J. Thomson. Even before the electron was identified, E. Goldstein in 1886 discovered the presence of new radiations in a gas discharge and called them canal rays. Canal rays : were positively charged radiations which ultimately led to the discovery of another sub-atomic particle proton.  Canal rays/Protons were discovered by E. Goldstein.  Nucleus of an atom was discovered by Ernest Rutherford.  Electrons were discovered by J.J. Thomson, in his cathode ray tube experiment.  Neutrons were discovered by James Chadwick. Properties of electrons, protons, and neutrons  Atoms consist of protons and electrons in a balanced proportion.  Protons exist in the interiors of the atom and electrons exist in the exteriors of the atom. Therefore, electrons can be removed from an atom. Limitations of Dalton’s Atomic Theory: 1. Dalton suggested that atom was indivisible and indestructible. But the discovery of two fundamental particles (electrons and protons) inside the atom, led to failure of this aspect of Dalton’s atomic theory. 2. The atoms of same element are similar in all respects, but isotopes of same element have different mass. Thomson’s Model of an Atom  According to Thomson, (i) An atom consists of a positively charged sphere, and the electrons are embedded in it. (ii) The negative and positive charges are equal in magnitude. So, the atom as a whole is electrically neutral  The first model of an atom to be put forward and taken into consideration.  He proposed a model of the atom be similar to that of a Christmas pudding/watermelon.  The red edible part of the watermelon is compared with the positive charge in the atom.  The black seeds in the watermelon are compared with the electrons which are embedded on it. Rutherford’s Model of an Atom Ernest Rutherford was interested in knowing how the electrons are arranged within an atom. Rutherford designed an experiment for this. In this experiment, fast moving alpha (α)- particles were made to fall on a thin gold foil. - He selected a gold foil because he wanted as thin a layer as possible. This gold foil was about 1000 atoms thick. - α -particles are doubly-charged helium ions. Since they have a mass of 4 u, the fast-moving alpha particles have a considerable amount of energy. - It was expected that alpha-particles would be deflected by the sub-atomic particles in the gold atoms. Since the alpha particles were much heavier than the protons, he did not expect to see large deflections. Rutherford’s Experiment  He experimented with thin gold foil by passing fast moving alpha particles on it.  He expected that the gold atoms will deflect the Alpha particles. Rutherford’s observations: But, the a-particle scattering experiment gave totally unexpected results. The following observations were made: (i) Most of the fast moving a-particles passed straight through the gold foil. (ii) Some of the α -particles were deflected by the foil by small angles. (iii) Surprisingly one out of every 12000 particles appeared to rebound. In the words of Rutherford, “This result was almost as incredible as if you fire a 15-inch shell at a piece of tissue paper and it comes back and hits you”.  Inferences /Conclusion made out of these observations are as follows: Rutherford concluded from the a-particle scattering experiment that– (i) Most of the space inside the atom is empty because most of the a-particles passed through the gold foil without getting deflected. (ii) Very few particles were deflected from their path, indicating that the positive charge of the atom occupies very little space. (iii) A very small fraction of a-particles were deflected by 1800, indicating that all the positive charge and mass of the gold atom were concentrated in a very small volume within the atom. From the data he also calculated that the radius of the nucleus is about 10 5 times less than the radius of the atom. On the basis of his experiment, Rutherford put forward the nuclear model of an atom, which had the following features: There is a positively charged centre in an atom called the nucleus. Nearly all the mass of an atom resides in the nucleus. (i) The electrons revolve around the nucleus in circular paths. (ii) The size of the nucleus is very small as compared to the size of the atom. Drawbacks /Limitations of Rutherford’s model  He explained that the electrons in an atom revolve around the nucleus in well- defined orbits. Particles in a circular orbit would experience acceleration.  Thus, the revolving electron would lose energy and finally fall into the nucleus.  But this cannot take place as the atom would be unstable and matter would not exist in the form we know.  Bohr's Model of an Atom Bohr came up with these postulates to overcome the objections raised against Rutherford’s model: Bohr suggested that –  Electrons spin around the nucleus in an individualized separate path or discrete orbits.  The electrons do not emit any energy while moving in these discrete orbits.  These orbits are also called as Energy Levels.  Electrons revolve around the nucleus in stable orbits without emission of radiant energy. Each orbit has a definite energy and is called an energy shell or energy level.  They are represented using letters or numbers as shown in the figure below – The Neutrons In 1932, J. Chadwick discovered that there is another sub-atomic particle present in the atom. This particle carries no charge and mass nearly equal to that of proton. It is known as a Neutron. It is present inside neutron along with proton. Therefore, we can conclude that atom consists of three types of particles - Electrons which carry a negative charge Protons which carry a positive charge Neutrons they are neutral The distribution of electrons in different shells or orbits was suggested by BOHR AND BURY.  If Orbit number = n 2  Then number of electrons present in an Orbit = 2n  So, for n =1 2  Maximum electrons present in shell – K = 2 * (1) = 2  The outermost shell can contain at most 8 electrons even if it has a capacity to accommodate more electrons. This is a very imp. Rule and is called Octet Rule. The presence of 8 electrons in the outermost shell makes the atom very stable.  The shells are always filled in a step-wise manner from the lower to higher energy levels. Electrons are not filled in the next shell unless previous shells are filled.  Thus, until the inner shells of an atom are filled completely the outer shells cannot contain any electrons.  ATOMIC STRUCTURE OF FIRST EIGHTEEN ELEMENTS VALENCY  The electrons present in the outermost shell of an atom are known as the valence electrons.  The combining capacity of the atoms or their tendency to react and form molecules with atoms of the same or different elements is known as valency of the atom.  Atoms of elements, having a completely filled outermost shell show little chemical activity.  Their combining capacity or valency is zero.  For example, we know that the number of electrons in the outermost shell of hydrogen is 1, and in magnesium, it is 2.  Therefore the valency of hydrogen is 1 as it can easily lose 1 electron and become stable.  On the other hand, that of magnesium is 2 as it can lose 2 electrons easily and also attain stability. What happens when the outermost shell contains a number of electrons that are close to its maximum capacity? (when outer shell electrons are more than 4) Valency in such cases is generated by subtracting the number of electrons present in the outermost orbit from octet (8). For example, oxygen contains 6 electrons in its outermost shell. Its valency is calculated as: 8 – 6 = 2. This means oxygen needs two electrons to form a bond with another element. ## Add complete handwritten big table of first twenty elements from Ch-3 Atomic Number of an Element Atomic Number (Z) = Number of protons in an atom Mass Number of an Element Mass Number = Number of protons + Number of neutrons What are nucleons? – Protons and Neutrons are collectively called as Nucleons.  Isotopes  The atoms of an element can exist in several forms having similar atomic numbers but varying mass numbers.  Isotopes are pure substances.  Isotopes have similar chemical properties.  Isotopes have different physical properties. Where can we use Isotopes? 1. The fuel of Nuclear Reactor – Isotope of Uranium 2. Treatment of Cancer – Isotope of Cobalt 3. Treatment of Goiter – Isotope of Iodine Example: Consider two atomic species namely U and V. Are they isotopes? U V Protons 5 5 Neutrons 5 6 Mass Number 5 + 5 = 10 5 + 6 = 11 Atomic Number 5 5 From the above example, we can infer that U and V are isotopes because their atomic number is the same. Calculation of average atomic mass/mass number for isotopic elements When an element has an isotope, the mass number can be calculated by the different proportions it exists in. Average atomic mass= (Atomic mass of isotope I) x(Its percentage in nature) + (Atomic mass of isotope II)x(Its percentage in nature) 100 For example take 98% Carbon-12u and 2% Carbon-13u This does not mean that any Carbon atoms exist with the mass number of 12.02u. If you take a certain amount of Carbon, it will contain both isotopes of Carbon, and the average mass is 12.02 u. Isobars Atoms of different elements with different atomic numbers, which have the same mass number, are known as Isobars. For example, Calcium and Argon: both have the same mass number – 40 20Ca40 and 18Ar40. ** Elements are defined by the number of protons they possess.