Survey
* Your assessment is very important for improving the workof artificial intelligence, which forms the content of this project
* Your assessment is very important for improving the workof artificial intelligence, which forms the content of this project
Structure of the Atom Matter has mass and takes up space. Atoms are basic building blocks of matter, and cannot be chemically subdivided by ordinary means. The word atom is derived from the Greek word atom which means indivisible. The Greeks concluded that matter could be broken down into particles to small to be seen. These particles were called atoms The atom is a basic unit of matter that consists of a dense central nucleus surrounded by a cloud of negatively charged electrons. The atomic nucleus contains a mix of positively charged protons and electrically neutral neutrons (except in the case of hydrogen-1, which is the only stable nuclide with no neutrons). The electrons of an atom are bound to the nucleus by the electromagnetic force. Likewise, a group of atoms can remain bound to each other, forming a molecule. An atom containing an equal number of protons and electrons is electrically neutral, otherwise it has a positive charge if there are fewer electrons (electron deficiency) or negative charge if there are more electrons (electron excess). A positively or negatively charged atom is known as an ion. An atom is classified according to the number of protons and neutrons in its nucleus: the number of protons determines the chemical element, and the number of neutrons determines the isotope of the element. A generic atomic planetary model, or the Rutherford model Atoms are made from smaller subatomic particles. At the centre of an atom is a nucleus containing protons and neutrons. Electrons are arranged around the nucleus in energy levels or shells. Make sure you can label a simple diagram of an atom like this one. Both protons and electrons have an electrical charge. Both have the same size of electrical charge, but the proton is positive and the electron negative. The neutron is neutral. Structure of Atom- Jony Mallik - M.Pharm; MS Page 1 CHARACTERISTICS OF ELECTRON Charge: It is a negatively charged particle. Magnitutide of charge: Charge of electron is -1.6022 x 10-19 coulomb. Mass of electron: Mass of electron is 0.000548597 a.m.u. or 1.1 x 10-31 kg. Symbol of electron: Electron is represented by "e". Location in the atom: Electrons revolve around the nucleus of atom in different circular orbits. CHARACTERISTICS OF PROTON Charge: Proton is a positively charged particle. Magnitude of charge: Charge of proton is 1.6022 x 10-19 coulomb. Mass of proton: Mass of proton is 1.0072766 a.m.u. or 1.6726 x 10-27 kg. Comparative mass: Proton is 1837 times heavier than an electron. Position in atom: Protons are present in the nucleus of atom. For latest information , free computer courses and high impact notes visit : CHARACTERISTICS OF NEUTRON charge: It is a neutral particle because it has no charge. Mass of neutron: . Mass of neutron is 1.0086654 a.m.u. or 1.6749 x 10-27 kg. Compartive mass: Neutron is 1842 times heavier than an electron. Location in the atom: Neutrons are present in the nucleus of an atom. ATOMIC NUMBER Total number of protons present in the nucleus of an atom is called "Atomic number" or "Charge number" Since the total number of protons and the total number of electrons in an atom are equal therefore atomic number may also be defined as: "Total number of electron in an atom is called Atomic number" SYMBOL: It is denoted by "z". Structure of Atom- Jony Mallik - M.Pharm; MS Page 2 MASS NUMBER Total number of protons and neutrons present in the nucleus of an atom is called "Mass number". SYMBOL: It is denoted by "A". A=p+n SYMBOL OF NUCLEUS A nucleus is represented by the symbol : zXA Where X= symbol of element (Cl, Br, H etc.) A= mass number. Z= atomic number. 1 23 35 14 27 1H , 11Na , 17Cl , 7N , 13Al Structure of Atom- Jony Mallik - M.Pharm; MS Page 3 Rutherford Atomic model The Rutherford model or planetary model is a model of the atom devised by Ernest Rutherford. Rutherford directed the famous Geiger-Marsden experiment in 1909, which suggested, upon Rutherford's 1911 analysis, that the so-called "plum pudding model" of J. J. Thomson of the atom was incorrect. Rutherford's new model for the atom, based on the experimental results, contained the new features of a relatively high central charge concentrated into a very small volume in comparison to the rest of the atom and with this central volume also containing the bulk of the atomic mass of the atom. This region would be named the "nucleus" of the atom in later years. RUTHERFORD’S ATOMIC MODEL Rutherford's atomic model shows the existence of nucleus in the atom, nature of charge on the nucleus and the magnitude of charge on the nucleus. APPARATUS FOR EXPERIMENT Alpha particles. Gold foil. (0.0004 cm thick) Zinc sulphide screen. Electron Gun. EXPERIMENT In his experiments, Rutherford bombarded alpha particles on very thin metallic foils such as gold foil. In order to record experimental observations, he made use of circular screen coated with zinc sulphide. OBSERVATIONS He observed that most of the alpha particles were pass through the foil undeflected. Very few particles were deflected when passed through the foil. One particle out of 8000 particles was deflected at 90o. Structure of Atom- Jony Mallik - M.Pharm; MS Page 4 Few particles were deflected at different angles. For latest information , free computer courses and high impact notes visit : MAIN POINTS OF RUTHERFORD’S THEORY Major portion of the atom is empty. The whole mass of the atom is concentrated in the center of atom called nucleus. The positively charged particles are present in the nucleus of atom. The charge on the nucleus of an atom is equal to (+z.e) where Z= charge number, e = charge of proton. The electrons revolve around the nucleus in different circular orbits. Size of nucleus is very small as compare to the size of atom. EXPLANATION OF POSTULATES 1. Since most of the alpha particles were passed through the foil undeflected, therefore, it was concluded that most of the atom is empty. 2. Small angles of deflection indicate that positively charged alpha particles were attracted by electrons. 3. Large angles of deflection indicate that there is a massive positively charged body present in the atom and due to repulsion alpha particles were deflected at large angles. DEFECT OF RUTHERFORD’S THEORY There were two fundamental defects in Rutherford's atomic model: According to classical electromagnetic theory, being a charge particle electron when accelerated must emit energy. We know that the motion of electron around the nucleus is an accelerated motion, therefore, it must radiate energy. But in actual practice this does not happen. Suppose if it happens then due to continuous loss of energy orbit of electron must decrease continuously. Consequently electron will fall into the nucleus. But this is against the actual situation and this shows that atom is unstable. If the electrons emit energy continuously, they should form continuous spectrum .But actually line spectrum is obtained ELECTRONEGATIVITY "Relative tendency or relative power of an atom to attract shared pair of electrons towards itself is called ELECTRONEGATIVITY." Structure of Atom- Jony Mallik - M.Pharm; MS Page 5 E.N depends upon the size of atom . Small atoms have large values of E.N. Big atoms have small values of E.N. E.N decreases in a group. E.N increases in a period. Most Electronegatively element is "Flourine". E.N = 4 Bohr’s atomic model Various postulates of Bohr’s atomic model are: 1. In an atom, the electrons revolve around the nucleus in certain definite circular paths called orbits, or shells. 2. Each shell or orbit corresponds to a definite energy. Therefore, these circular orbits are also known as energy levels or energy shells. 3. The orbits or energy levels are characterized by an integer not, where, n can have values 1, 2, 3, 4……. The integer not (= 1, 2, 3…) is called the quantum number of respective orbit. The orbits are numbered as 1, 2, 3, 4………… etc., starting from the nucleus side. Thus, the orbit for which n=1 is the lowest energy level. The orbits corresponding to n = 1,2,3,4…..etc., are also designated as K,L,M,N……….etc., shells. When the electron is in the lowest energy level, it is said to be in the ground state. Since, electronics can be present only in these orbits, hence, these electrons can only have energies corresponding to these energy levels, i.e., electrons in an atom can have only certain permissible energies . 4. The electrons present in an atom can move from a lower energy level (Elower) to a level of higher energy (Ehigher) by absorbing the appropriate energy. Similarly, an electron can jump from a higher energy level (Ehigher) to a lower energy level (Elower) by losing the appropriate energy. The energy absorbed or lost is equal to the difference between the energies of the two energy levels, i.e., ΔE= Ehigher - Elower Structure of Atom- Jony Mallik - M.Pharm; MS Page 6 AMOUNT OF ENERGY Energy released or absorbed by an electron is equal to the difference of energy of two energy levels. Let an electron jumps from a higher energy level E2 to a lower energy level E1.The energy is emitted in the form of light . Amount of energy released is given by: E = E 2 - E1 E2 - E1= h Where h = Planck's consrant ( 6.6256 x 10-34 j.s) = Frequency of radiant light ANGULAR MOMENTUM OF ELECTRON Angular momentum of an electron in an energy level is given by: m v r = nh /2 Where n =1, 2, 3, ……….. m = mass of electron V = velocity of electron r = radius of orbit OR Only those energy levels or orbits are possible for which angular momentum of electron is an integral multiple of h /2 . There are particular types of Nuclide. They are: Isotopes Isobars Isotones Structure of Atom- Jony Mallik - M.Pharm; MS Page 7 Here, A= Mass Number. Z=Atomic Number n=Number of charge (+ or-ve) Isotopes: The atoms having same atomic number but different atomic mass number are called Isotope. Isobars: Nuclides having the same mass number but having the different Proton/Atomic number are called Isobar. Isotones: Atoms of different elements having different mass number and different atomic number but same neutron number are called Isotones. Hydrogen Hydrogen atom (Z = 1) has no neutrons. Number of protons = 1Number of electrons = 1Number of neutrons = 0It has been reported that the hydrogen element has atoms with mass number 2 and 3 also i.e., Atoms of elements having the same atomic number with different mass numbers are called isotopes. Structure of Atom- Jony Mallik - M.Pharm; MS Page 8 Nuclear composition of isotopes of chlorine: Nuclear composition of isotopes of carbon: Structure of Atom- Jony Mallik - M.Pharm; MS Page 9 Characteristics of Isotopes All isotopes of an element have the same number of valence electrons thus have identical chemical properties. The physical properties of the isotopes are different due to the difference in the number of neutrons in their nuclei. The densities, melting points and boiling points etc., are slightly different. Reason for Fractional Atomic Masses of Elements Atomic masses of many elements are in fractions not in whole numbers. Example: Cl - 35.5Cu - 63.5The fractional atomic masses of elements are due to the existence of isotopes having different masses. Example:1 Natural chlorine consists of two isotopes: atomic mass of chlorine. Structure of Atom- Jony Mallik - M.Pharm; MS Calculate the average Page 10 Example: 2 A naturally occurring sample of Lithium contains 7.42% of 6Li and 92.58% of 7Li. The relative mass of6Li is 6.015 and that of 7Li is 7.016. Calculate the atomic mass of a naturally occurring sample of lithium. Solution: Example: 3 Which of the following two nuclei are isotopes of each other? Solution: The two isotopes are: Orbit As postulated by Bohr, an orbit is a definite circular path at a definite distance from the nucleus in which the electron revolve round the nucleus. Orbits are designated by the capital letters K, L, M, N…. etc. Orbits are circular in shape. It represents the planer motion of the electron. An orbit indicates an exact position of an electron in an atom. The maximum number of electro in an orbit is equal to 2n^2, where n is the number of the orbit. Structure of Atom- Jony Mallik - M.Pharm; MS Orbital As postulate by wave nature of an electron, an orbital is a threedimensional region around the nucleus within which the probability of finding an electron is maximum., Orbital’s are designated by s, p, d… etc. Orbital’s have different shapes. It represents the three-dimensional motion of the electro round the nucleus. An orbital does not specify the exact position of an electron in an atom. An orbital cannot accommodate more than two electrons. Page 11 Structure of Atom- Jony Mallik - M.Pharm; MS Page 12 QUANTUM NUMBER Different numbers of electron can be present at the different energy shell or orbit that surrounds the central dense nucleus. From the concept of the spectral analysis of an atom, the shell can be spherical or semispherical. A very important concept regards to the electron is that, the electron not only surrounds the nucleus it also surrounds its own during the time of surrounding the nucleus. Which electron present at which energy shell? The energy shell is spherical or not? What is the axis of rotation of electron, clockwise or anti-clockwise? All the things can be described by some special numbers called Quantum numbers. Electrons can be labelled using the subshell and orbital or by using the four quantum numbers: n : principal quantum number l : azimuthal quantum number ml : magnetic quantum number ms : spin quantum number Principal Quantum Number, n The principal quantum number, n, is always a positive integer and tells us the energy level or shell that the electron is found in. The maximum number of subshells permitted for a particular shell is equal to n2. The maximum number of electrons permitted in a particular shell is equal to 2 x n2. n Energy Level 1 1st energy level 2 2nd energy level 3 3rd energy level 4 4th energy level Shell K L M N No. Subshells = n2 1 4 9 16 No. electrons = 2n2 2 8 18 32 Azimuthal Quantum Number, l The azimuthal quantum number tells us which subshell the electron is found in, and therefore it tells us the shape of the orbital. l can have values ranging from 0 to n-1. Structure of Atom- Jony Mallik - M.Pharm; MS Page 13 The number of orbitals permitted for a particular subshell is equal to 2l + 1. value of n l=n-1 subshell (orbital shape) No. orbitals = 2l + 1 1 0 s subshell 1 (1 x s orbitals) 2 1 p subshell 3 (3 x p orbitals) 3 2 d subshell 5 (5 x d orbitals) 4 3 f subshell 7 (7 x f orbitals) Magnetic Quantum Number, ml The magnetic quantum number, ml, tells us the orientation of an orbital in space. ml can have values ranging from -l to +l. It is not always possible to associate a value of ml with a particular orbital. value of l subshell values of ml possible orbitals 0 s 0 s 1 p -1, 0, 1 px, py, pz 2 d -2, -1, 0, 1, 2 dxy, dxz, dyz, dx2-y2, dz2 3 f -3, -2, -1, 0, 1, 2, 3 Spin Quantum Number, ms The spin quantum number, ms, tells us the spin of the electron. ms can have a value of +½ or -½. Structure of Atom- Jony Mallik - M.Pharm; MS Page 14 1st - Principle QN 2nd - Orbital QN 3rd - Magnetic QN n l ml 4th - Spin QN ms l goes from 0 to n1 within an energy level n = 1,2,3...7 l values = 0 (for s), 1(for p), 2 (for d), 3 (for f) sublevels Values of ml go from +l to - l , which gives 2l + 1 number of values has 2 values: +1/2 (spin up) and -1/2(spin down) 1. measures the average distance of the e- from the nucleus 1. indicates the shape of the orbital ( set of probable locations of the e- ) 1. identifies the direction the e- orbital has around the nucleus 1. identifies the "spin" or rotation of the e- about its own axis 2. different values of n mean different energy levels 2. diff. values of l mean diff sublevels. In a sublevel all the ehave nearly the same energy. 2. specifies the eorbital in which the eis located within a sublevel. 2. shows that each orbital can contain only 2 e- 3. different values of n mean relatively large differences in the energies of the e-s 3. different sublevels within the same level may have moderately large differences in energy. 3. different values of ml mean little difference in energies of the e- 3. the direction of spin is either in one direction or the other 4. the smallest avgerage distance and the lowest energy occurs when n = 1; each increase in n increases those quantities. 4. within any level, the lowest energy sublevel is s, then p, then d, then f. 4. the number of possible values of ml within a sublevel idenities how many epairs that the sublevel can hold 4. when 2 e- (in an atom) have the same set of QN except for ms, then these e- are called an e- pair 5. the number of epossible in a level is 2n2 5. the number of possible values of l for a level is equal to the value of n Structure of Atom- Jony Mallik - M.Pharm; MS 5.these e- within an e- pair have essentially the same energy Page 15 Electronic Configuration of elements Occupation of Orbitals The first thing to keep in mind is that electrons fill orbitals in a way to minimize the energy of the atom. This would mean that the electrons in an atom would fill the principal energy levels in order of increasing energy (the electrons are getting farther from the nucleus). The order of levels filled would look like this: 1s, 2s, 2p, 3s, 3p, 4s, 3d, 4p, 5s, 4d, 5p, 6s, 4f, 5d, 6p, 7s, 5f, 6d, and 7p One way to remember this pattern, probably the easiest, is to refer to the periodic table and remember where each orbital block falls to logically deduce this pattern. Another way is to make a table like the one below and use vertical lines to determine which subshells correspond to each other. Structure of Atom- Jony Mallik - M.Pharm; MS Page 16 Pauli Exclusion Principle The second major fact to keep in mind is the Pauli Exclusion Principle which states that no two electrons can have the same four quantum numbers. The first three (n,l, and ml) may be similar but the fourth quantum number must be different. We are aware that in one orbital a maximum of two electrons can be found and the two electrons must have opposing spins. That means one would spin up ( +1/2) and the other would spin down (-1/2). This tells us that each subshell has double the electrons per orbital. The s subshell has 1 orbital that can hold to 2 electrons, the p subsheel has 3 orbitals that can hold up to 6 electrons, the d subshell has 5 oribtals that hold up to 10 electrons, and the f subshell has 7 oribtals with 14 electrons. Example Structure of Atom- Jony Mallik - M.Pharm; MS Page 17 We have the first three quantum numbers n=1, l=0, ml=0. Only two electrons can correspond to these, which would be either ms = -1/2 or ms = +1/2. As we already know from our studies of quantum numbers and electron orbitals, we can conclude that these four quantum numbers refer to 1s subshell. If only one of the ms values are given then we would have 1s1 (denoting Hydrogen) if both are given we would have 1s2 (denoting Helium). Visually this would be represented as: As you can see, the 1s subshell can hold only two electrons and when filled the electrons have opposite spins. Hund's Rule When assigning electrons in orbitals, each electron will first fill all the orbitals with similar energy (also referred to as degenerate) before pairing with another electron in a half-filled orbital. Atoms at ground states tend to have as many unpaired electrons as possible. When visualizing this processes, think about how electrons are exhibiting the same behavior as the same poles on a magnet would if they came into contact; as the negatively charged electrons fill orbitals they first try to get as far as possible from each other before having to pair up. Example Structure of Atom- Jony Mallik - M.Pharm; MS Page 18 If we look at the correct electron configuration of Nitrogen (Z = 7), a very important element in the biology of plants: 1s2 2s2 2p3 We can clearly see that p orbitals are half filled as there are three electrons and three p orbitals. This is because Hund's Rule states that the three electrons in the 2p subshell will fill all the empty orbitals first before filling orbitals with electrons in them. If we look at the element after Nitrogen in the same period, Oxygen (Z = 8) its electron configuration is: 1s2 2s2 2p4 Oxygen has one more electron than Nitrogen and as the orbitals are all half filled the electron must pair up. The Aufbau Principle Aufbau comes from the German word "Aufbauen" which means "to build". When writing electron configurations, we are building up electron orbitals as we proceed from atom to atom. As we write the electron configuration for an atom, we will fill the orbitals in order of increasing atomic number. However, there are some exceptions to this rule. Structure of Atom- Jony Mallik - M.Pharm; MS Page 19 Example If we follow the pattern across a period from B (Z=5) to Ne (Z=10) the number of electrons increase and the subshells are filled. Here we are focusing on the p subshell in which as we move towards Ne, the p subshell becomes filled. B (Z=5) configuration: 1s2 2s2 2p1 C (Z=6) configuration:1s2 2s2 2p2 N (Z=7) configuration:1s2 2s2 2p3 O (Z=8) configuration:1s2 2s2 2p4 F (Z=9) configuration:1s2 2s2 2p5 Ne (Z=10) configuration:1s2 2s2 2p6 Structure of Atom- Jony Mallik - M.Pharm; MS Page 20 DALTON ‘S ATOMIC THEORY Main postulates of Dalton atomic theory are as follows: 1. Matter is composed of very tiny or microscopic particles called "Atom". 2. Atom is an indivisible particle. 3. Atom can neither be created nor it is destroyed. 4. Atoms of an element are identical in size, shape, mass and in other properties. 5. Atoms of different elements are different in their properties. 6. Atoms combine with each other in small whole numbers. 7. All chemical reactions are due to combination or separation of atoms. DEFECTS IN DALTON’S THEORY: Postulate number 2, 3, 4 and 6 are not correct as described below: DEFECT NO: 1 Atom can be divided into a number of sub-atomic particles such as electron, proton and neutron etc. DEFECT NO: 2 Atoms of an element may be different in their masses. For example: 1H 1 , 1H2, 1H3 17Cl 35 , 17Cl 37 DEFECT NO: 3 All compounds do not have small number of atoms. For example: Decane C10H22. Sugar C12H22O11. DEFECT NO: 4 Atom can be destroyed by fission process in Atom bomb. Nuclear reactor. On the basis of above defects, Dalton's atomic theory has failed now. Structure of Atom- Jony Mallik - M.Pharm; MS Page 21 Structure of Atom- Jony Mallik - M.Pharm; MS Page 22