Chemistry Chapter (6) Atomic Structure

Points to remember

• The atom was believed to the smallest indivisible particle of matter and that all matter composed of such atoms. Dalton in 1808 put forth his famous atomic theory incorporating the idea of indivisibility of an atom. It was the first scientific approach to the structure of matter.
• Afterwards the experiments conducted by the brilliant research workers like J.J Thomson (1897), Rutherford (1911), Niel Bohr (1912), Chadwick, Mosley and other proved that the atom itself had a complex structure. If shown that an atom consists of several small particles, three important ones of which are electron, proton and neutron.

Atom

Modern Concept: The smallest individual particle of an element that takes part in chemical reaction is called atom.

Composition Of Atom

Fundamental    Particle	Massin Gram(CGS)	Mass in gm(SI)	Charge in esu(CGS)	Charge in coulomb S.I	Charge Unit	Symboll	Discovere
Electron	9.108×10-28 g	9.108×1031kg	4.8×10-0 coulomb	1.602×10-19 coulomb	-1	-1e0	J.J Thomson
Proton	1.6735×10-24    g	1.6735×10-27 kg	-	1.602×10-19 coulomb	+1	+1p0	Goldstein
Neutron	1.675×10-24 g	1.675×10-27 kg	0	0	0	0n1	chadwick

• The absolute mass of electron is called REST MASS which is equal to 9.1096×10^-31 kg.
•  Mass of moving electron = rest mass/ (1-(v/c)²)^1/2
• Where v =velocity of electron, c = velocity of light.
• The e/m value of electron was determined by Thomson and is equal to 5.27×10^-17 esu/g = 1.76×10¹¹ coulomb per kg.
• The value of radius of electron is 2.8×10^-13 cm.
• The radius of proton is about 10^-13 cm.
• The radius of neutron is about 10^-13 cm.

SOME OTHER FUNDAMENTAL PARTICLES

 Positron: 

•      It is represented by +1e^o. It is antiparticle of electron (discovered by Anderson).

  ·       Meson:

•  It may be positive, negative or neutral. It is represented by µ- mesons called ‘muons’ and π-mesons called ‘pions’ (discovered by Yukawa).

  ·       Neutrino: 

•      It has no charge and the mass is less than electron. It is represented by v(_ov^o) (discovered by pauli).

  ·       Nucleus: 

•     The central core of an atom that contains most is called Nucleus. Nucleus is positively charged. The size of nucleus is measure in Fermi.(1 fermi =10^-13 cm).

  ·       Atomic Number: 

•     The  number of proton present in the nucleus of an atom is known as atomic number. It is represented by Z (Zeta).
•         It neutral atom, No. of electrons. Therefore atomic number = No. 01 protons = No. of electrons.

  ·       Mass Number: 

•    The sum of number of protons and neutrons present in the nucleus is called mass number.

It is represented by A. The atom is represent by symbol

                                                                                                                              A----Mass number

                                                                                                                                                    X ------Atom

      Z -----Atomic number 

  ·       Nucleon Number: 

•     The number of nucleon i.e. protons and neutrons present in an atomic nucleus of a particular nuclide is called nucleon number. It is also called mass number.

  ·       Cathode Rays:

•     Cathode rays (electrons) are produced when electricity passes through evacuated tubes.

  ·       X-Rays: 

•     When cathode rays are allowed to strike a metal target placed in their path, radiations are produced. These radiations are called X- rays.

  ·       Radioactivity: 

•      Radioactivity is the emission of radioation by unstable atomic unstable atomic nuclei and the most common types of radiation are alpha (α) rays, beat (β) rays and gamma (γ)
•             Alpha particles are Helium nuclei, beta particles are electrons and gamma rays are high energy electromagnetic radiation.

  ·       Rutherford’s Atomic Model:

•      According to it, atom consists of a very small positively charged nucleus and extra nuclear electrons. The nucleus  consists of protons and neutrons and contains practically all the mass of an atom.

  ·       Drawbacks Of  Rutherford’s Model:

•     It could not explain the stability of an atom .
•     It could not explain the spectral lines the spectrum of hydrogen atom.

  ·       Plank’s Quantum Theory:

•     The radiation of energy (emitted or absorbed) is not continues but discontinuous in the form of ‘quanta’. The quantum of light energy is called ‘photon’.
•      The amount of energy associated with the quantum of radiation is proportional to the frequency of radiation.

                                              E α v or E = h v

Where h = Planck's constant (h =6.625×10^-37 Js) and v =     frequency.

•      Energy emitted or absorbed by a body is in terms of integral multiples of quantum i.e. E = nhv.

                                       Where n = 1, 2, 3, ………..

  ·       Electromagnetic Radiation:

•      Electromagnetic radiation is the transmission of electric and magnetic fields as a wave motion. The waves are the characterized by the following wave parameters.

i)-Wave Length: 

• It is the distance between two crests and troughs     or adjacent waves. It is represented by λ ( lambda). It is expressed in cm or m or A⁰ or µm or pm.

                1A^o = 10^-1 nm = 10^-8 cm = 10^-10 m.

                1µm = 10^-9 m, 1cm = 10^-2 m = 10⁸ A⁰ = 10⁴ µm =10⁷ nm

                1nm = 10^-9 m, 1 pm = 10^-2 A⁰ = 10^-10 cm.

ii)- Frequency: 

•  It is the number of wave or cycles which pass a given point on the wave per second. It is expressed in cycles (or waves) Per second (cps) or herts (Hz). It is also expressed in units of reciprocal seconds (i.e. Sec^-1).

  ·       Frequency (v) is inversely proportional to wavelength (λ).

                    i.e. v α 1/ λ or v = c/λ.

·       Where c = velocity of light (3× 10¹⁰ cm/sec).

iii)-Velocity: 

• It is the distance travelled by the wave in one second and denoted by c. It is equal to the product of wavelength (λ) and frequency (ν).      C =  λν

iv)- Wave Number: 

• It is the number of wavelength per centimeter and is equal to the reciprocal of wavelength. It is denoted by the ν (nu bar). It is expressed in cm^-1 or m^-1.

v)- Amplitude: 

• The  height of a crest or depth of a trough is called amplitude.

  ·      Spectrum: 

•      An array of the components of radiation separated in the order of increasing or decreasing wavelength is called spectrum. It can be broadly divided into two types.

            i)- Emission Spectrum     ii)- Absorption Spectrum

  ·      Absorption Spectrum: This spectrum obtained when electron passes from a lower energy level to a higher energy level. This type of spectrum contains dark lines on bright background.

  ·       Emission Spectrum: 

•     This  spectrum obtained when electron passes from to lower energy level. This type of spectrum contains bright lines on dark background. Emission spectra are of following two types.

             i)-   Continuous spectrum    ii)- Discontinuous spectrum

  ·    Continuous Spectrum:

•     The spectrum obtained without any discontinuity in wavelength of all color red to violet is called continuous spectrum. Example : sunlight.

   ·       Discontinuous Spectrum: 

•     It may be band spectrum or line spectrum.
•      Atomic Spectrum: this kind of spectrum is the characteristic of the atom of the element and hence used to identify the atom in a material i.e. the spectrum of element is the finger print of that elements. It is also known as line spectrum.

·       Band Spectrum: This type of spectrum obtained from the molecule in gaseous state. It is also known as molecular spectrum or fluten spectrum.

·       Hydrogen Spectrum: The atomic spectrum of hydrogen is characterized by the lines corresponding to the radiation quanta of sharply defined energy is called hydrogen spectrum.

   Electronic Transition And Values Of n₁ And n₂ For Various     Spectral Series Of Hydrogen Spectrum

SERIES  OF                       LINES	n1	n2	SPECTRUM REGION	WAVELENGTH (oA)
Lyman series	1	2,3,4,	Ultraviolet	Less than 4000 Ao
Balmer series	2	3,4,5,	Visible	Between 4000 Ao and 7000 Ao
Paschen series	3	4,5,6,	Near infra red	More than 7000 Ao
Brackett series	4	5,6,7,	Far infra red	More than 7000 Ao
Pfund series	5	6,7	Far infra red	More than 7000 Ao

  ·       Bohr’s Atomic Model: 

•     It is base on Planck's quantum theory. Main facts of the model.
•         Electrons move round the nucleus in certain circular orbit.
•    Only those orbits are permitted in which angular momentum of electron is equal to nh/ 2π where n = 1,2--------.
•    Energy is neither radiated nor absorbed when electrons move in their orbits.
•    When electron jumps from one orbit to another energy is evolved or absorbed E₂ – E₁ = ∆E =hv.

  ·       Defects Of Bohr’s Theory: 

•    It fails to explain the spectra of multi-electron atoms.
•     It considers electron only as a particle. The wave nature of electron is not explained.
•    It is unable to explain the splitting of spectral lines into groups of fine lines under the influence of magnetic field (Zeeman effect) and electric field (Stark effect).

  ·       Quantum Numbers: 

•     These are the identification numbers for an individual electron in an atom. These number are constant which are used to solve Schrodinger wave equation.
•      The quantum numbers explain the orbital concept of atomic model i.e. size, shape, orientation and spin of electron cloud.

  ·        Principal Quantum Number(n): 

•    It describes the energy and position of an electron in a shell.
•     Maximum number of electrons that can be present in a principal quantum number = 2n².

  ·       Azimuthal Quantum Number (I): 

•    It describes the shape of electron cloud. It also tells that energy level consists of sub- levels.
•    Number of sub- shells in a shell = n.
•    Different values of ɭ’ from 0 to (n – 1).
•    l = 0 corresponds to s orbital. l = 1,2,3 corresponds to p,d and f orbitals respectively.
•  Maximum number electrons in sub – shell = 4f +2.
•  Increasing order of energy’ S< P <d <f.

  ·       Magnetic Quantum Number (m):

•    It tells as about the orbitals of which a given a sub- shell is composed.
•    Value of m’ from -l through zero to +l.
•     Maximum number of orbitals in a shell = n².

       ·   Spin Quantum Number (S): 

•      It explains of spinning of electron around its own axis. An electron can be spin either clockwish. Thus, ‘s’ can have two possible values, + 1/2 or – 1/2.
•     When two electrons are present in an orbital, they have opposite spins(↑↓).

  ·       Orbit: 

•    It is well-defined circular path around the nucleus in which electron revolves. It represents that an electron moves around nucleus in one plane.

  ·       Orbital: 

•    It is the region in three – dimensional space around nucleus where probability of finding electron is maximum.

  ·       Nodal Plane : 

•     The region where the probability of finding the electron is almost zero is called node or nodal plane. The region of  high electron  density is called lobe.

  ·       Pauli’s Exclusion Principle:

•     No two electron in the same atom can be have the same values of all the four quantum numbers.

Rules For Filling Of Orbitals in an atom

•        The filling of orbitals by electrons in an atom is governed by the

i).Aufbau Principle: 

• The orbitals are successively filled in order of  their increasing energy , the lowest – energy orbitals available being filled up first.
• The relative energies of orbitals are in the following order:

 Is <2s <2p <3s <3p <4s <3d <4p <5s <4d <5p <6s <4f <5d <6p <7s  and so on.

ii). (n +l) Rule: 

• An energy level for which (n+l) is low, has a low energy. If (n+l) is the same for two sub- levels, then the one having low value of principle quantum number ‘n’ has the lower energy. The energy level having lower energy is filled first.

iii). Hund’s Rule: 

• Pairing of electrons does not take place in any p,d or f orbital until the orbitals of same sub- level contain one electron each.

iv). Stability Of Orbitals: 

• Orbitals in the same sub-level tend to become full or exactly half-full of electrons. This configuration is very stable.

v). Wave Nature Of Electron: 

• According to wave mechanical theory (given by de Broglie), the electrons, protons and even atoms when in motion, possess wave properties.

vi). Heisenberg’s Uncertainly Principle: 

• It is impossible to determine simultaneously the exact velocity (or momentum) and the exact position of a material particle.

Sher Afzal Ranais

Author & Editor

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