# Important Chemistry Notes for IITJEE/NEET Preparation- Structure of Atom

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Class 11 Chemistry is a broad subject that requires a thorough understanding of the concepts and topics covered. As a result, we have provided** Chemistry Notes PDF for IIT JEE/NEET** to students and NEET aspirants. **Structure of Atom Class 11 Notes PDF** for NEET can be found below. With the help of detailed syllabus, Class 11 students learn what they need to study, how many points are assigned to each unit, and how much time is allotted for each unit. As a result, they can easily plan their study schedule.

Check out the Structure of Atom Class 11 notes PDF for your IIT JEE/NEET Preparation based on the IIT JEE/NEET Chemistry Syllabus. The Structure of Atom notes PDF is designed in such a way that it is very useful for IIT JEE/NEET aspirants.

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**Chemistry Notes PDF for IIT JEE/NEET**to students and NEET aspirants.

**Structure of Atom Class 11 Notes PDF**for NEET can be found below. With the help of detailed syllabus, Class 11 students learn what they need to study, how many points are assigned to each unit, and how much time is allotted for each unit. As a result, they can easily plan their study schedule.

### ATOMIC RADII

### ELECTRON

- It was discovered through the study of Cathode rays (discovered by Zulius Plucker) and the name was proposed by Stoney.
- Charge : It was determined by Mullikan by oil drop method as –1.602 × 10–19 coulombs or 4.803 × 10–10 e.s.u.
- Mass : It was found by J. J. Thomson as

9.11 × 10–28 g. - Specific charge : e/m ratio is called specific charge and was determined by Thomson as 1.76 × 108 coulombs/gm.
- Radius : It is of the order 10–15 cm.
- Density : 2.17 × 1017 g/cc.
- Mass of electron at speed v is m =
- Atomic mass unit : It is 0.0005486 amu.
- Mass of one mole of electron : It is 0.55 mg.

### CATHODE RAYS

#### SOME PROPERTIES OF CATHODE RAYS

- They cast shadow of the object in their path
- Rotate a mica wheel
- Deflected by electric and magnetic fields in a direction showing negative charge.

### PROTON (H+)

- Mass : It was found to be 1.672 × 10–24 g or 1.672 × 10–27 kg or 1.00728 amu. It is about 1837 times heavier than an electron.
- Charge : It carries unit positive charge 1.602 × 10–19 coulombs or 4.803 × 10–10 esu.
- Specific charge : It is 9.58 × 104 coulomb/gm. It varies with nature of gas and is maximum if H2 is used.
- Charge on 1 mole of proton is 96500 coulomb or 1 Faraday.
- Volume : The volume for proton is approximately 1.5 × 10–38 cm3.

### NEUTRON (0N1)

- Mass : Its mass is 1.675 × 10–24 gm or 1.675 × 10–27 kg or 1.00866 amu.
- It is heavier than proton by 0.18%.
- Density : Its density is 1.5 × 1014 g/cm3.
- Specific Charge : It is zero.
- Stability : It is least stable of all elementary particles present in an atom.
- Disintegration : Isolated neutron is unstable and disintegrates into electron, proton and neutrino.
- Among all elementary particles neutron is the heaviest and least stable.

### OTHER SUBATOMIC PARTICLES

- Positron (Positive electron +1e0). Discovered by Dirac (1930) and C. D. Anderson (1932). They are highly unstable and produce -rays on combining with electrons.
- Neutrino and Antineutrino are particles of small mass and no charge as stated by Fermi (1934). Anti-neutrino spin clockwise and neutrino spin anticlockwise.
- Meson : They are unstable particles and include pions (+,– or 0) Kaons (K+, K–, K0, K–0) and eta meson ().

- Anti proton (–1p1) : Negative proton produced by Segre and Weigland (1955) by proton-proton and proton-neutron collisions.
- v-particles : They may be positive, negative or neutral. Discovered by G. D. Rochester and C C. Butler v– and v0 are 2200 times heavier than electron. Heavier disintegrate into pions and lighter into mesons.

### THOMSON'S ATOMIC MODEL

### RUTHERFORD'S NUCLEAR MODEL

### NUCLEUS

- Nucleus : It is small heavy and positively charged material located in the centre of atom and electrons are distributed in extra nuclear part of atom and revolve around the nucleus.
- Radius : It is of the order 1.5 × 10–13 cm to

6.5 × 10–13 cm (1.5 – 6.5 Fermi). In general

- Volume : It is about 10–39 cm3. and that of atom is 10–24 cm3
- Density : It is about 1014 g/cm3.
- Diameter : It is about 10–15 m or 1 fm (1 fm = 10–15 m)
- Nucleus contains neutrons and protons, collectively known as nucleons.

### ATOMIC NUMBER/MOSELEY'S POSTULATES

### MASS NUMBER

### AVERAGE ATOMIC MASS

### ISOTOPES

### ISOBARS

### ISOTONES

### ISODIAPHERS

### ISOELECTRONIC SPECIES

Nuclear charge) 1.1 1.2 1.3 1.4

### FAILURE OF RUTHERFORD'S MODEL

### PLANCK QUANTUM THEORY

E = h =h

(or 6.6726 × 10–27 ergs in c.g.s. units).

### INTENSITY OF LIGHT

### ELECTROMAGNETIC RADIATION

#### WAVELENGTH

#### FREQUENCY

#### VELOCITY

m sec–1 and denoted by c.

#### WAVE NUMBER

#### AMPLITUDE

### ELECTROMAGNETIC SPECTRUM

#### ATOMIC SPECTRUM

#### DISPERSION

#### CONTINUOUS SPECTRUM

#### LINE SPECTRUM

#### ABSORPTION SPECTRUM

#### EMISSION SPECTRUM

#### HYDROGEN SPECTRUM

n1, n2 = electronic levels involved in transition, = Wave number

- For calculation of longest wavelength line use n2 nearest and for shortest wavelength line use n2 infinity e.g. value of longest wavelength in Balmer Series of hydrogen spectrum use n1 = 2 and n2 = 3.
- Last line of spectrum is called Series limit. Last line is the line of shortest wavelength and high energy when n2 = we get last wavelength

- Number of Lines in a Transition : Mathematical formula for number of lines is follows as

### BLACK BODY RADIATION

### PHOTOELECTRIC EFFECT

- Threshold frequency (v0) : The minimum frequency of incident radiation to cause the photoelectric effect is called threshold frequency.
- Work function : A part of the photons energy that is absorbed by the metal surface to release the electron is known as work function of the surface denoted by . The remaining part of the energy of photons goes into the Kinetic energy of the electron emitted.

- K.E. is independent of the intensity of light.
- Number of photoelectrons Intensity of light
- K.E. is directly proportional to frequency of incident light.
- is known as Einstein's photoelectric equation.
- Energy required to stop the ejection of electrons is given by eV0 where e is the electric charge and V0 is stopping potential.

### BOHR’S MODEL OF ATOM

- Electrons revolve around nucleus only in certain selected circular orbits. These orbits are associated with definite energies and are called energy shells or levels.
- Electrons can move only in those circular orbits where angular momentum is a whole number and multiple of h/2. i.e. mvr =. or simply an integral number of wavelengths should fit in given electron orbit of radius r i.e. n=2r.
- Electrons energy in a particular orbit is constant.
- Lowest energy state is called ground state and when electron absorbs energy and jumps to higher state, it is called excited state
- Electronic energy is negative because at infinite distance there is no interaction between electron and nucleus thus energy is zero. While when close to nucleus, attraction takes place, energy is released and it becomes negative as it was already zero. The energy of electron increases with the value of n, but the difference of energy between two successive orbits decreases. Thus

- Energy of electron in nth orbit

- Radius of nth orbit

- Velocity of electron in nth orbit,

- Kinetic energy of electron in nth orbit,

- Potential energy of electron in nth orbit,

- Total energy of electron in nth orbit,

- Number of revolutions per second in nth orbit,

- Angular velocity
- Angular momentum = mvr
- Number of spectral lines when electron jumps from the nth to ground level =
- The electrons energy is generally expressed in kcal or kJ mol–1 or in electron volts eV.

- Some important values :

#### LIMITATIONS OF BOHR’S MODEL

- Explains the spectrum of elements having only one electron
- Does not explain splitting of spectral lines under magnetic field (Zeeman effect) and electric field (stark effect)
- Does not explain quantisation of angular momentum.
- It goes against the Heisenberg’s uncertainity principle.

### SOMMERFIELD MODEL

- Motion of electrons is in closed elliptical paths of definite energy levels having nucleus on either of the focii.
- Angular momentum is quantized
- where k = 1, 2 ---------n.
- It does not explain distribution of electrons in extranuclear part of atom and also does not explain for de Broglie concept.

### QUANTUM MECHANICS

- Photoelectric effect
- Black-body radiation
- Line spectra of H-atom
- Variation of heat capacity of solids with temperature.

### DE- BROGLIE PRINCIPLE (1924)

- Proposes that just as radiations have particle nature, the material particles are also associated with wave nature.
- de Broglie wavelength is h = Planck’s constant m = mass of object ; v= velocity and this equation is called the de Broglie equation.

### DAVISSON AND GERMER’S EXPERIMENT

### SCINTILLATION METHOD AND PHOTOELECTRIC EFFECT

### HEISENBERG’S UNCERTAINITY PRINCIPLE

- Mathematically

- As the mass of particle increases, the uncertainity decreases

### QUANTUM MECHANICAL MODEL OF ATOM

- Based on de Broglie's and Heisenberg’s principle.
- Put forward by Schrodinger (1920). Behaviour of electron was described in terms of equation known as Schrodinger wave equation

- Many solutions for this equation are possible for hydrogen but only certain solutions are permissible and are called eigen values
- The solution must be single valued, should satisfy the relation and must be finite and continuous.
- has no physical significance but gives intensity of electrons and thus gives probability of electron in a particular region.

### ORBITALS

- Probability does not become zero even at infinity and is given by .
- Electron orbitals in atoms are called atomic orbitals while those in molecules are called molecular orbitals.
- Orbitals have definite energy and momentum and are quantized. i.e, En = –E1/n2 thus Bohr’s concept of well defined orbits is ruled out.

### QUANTUM NUMBERS

- Four quantum numbers (n, l, m, s) help in providing complete information about an electron in an atom.
- Principal quantum number (n) determines the energy and average distance of electron. It has whole number values also denoted as K, L, M, N. etc. As n increases, distance of electron from nucleus increases and energy increases.
- Azimuthal quantum number (l) determines angular momentum of the electron. It also determines the shape of orbitals and it may have all possible whole number values from 0 to n–1 for each principal energy level. The sublevels are:

mvr

- Magnetic quantum number (m) defines the orientation of electrons cloud in a particular sub shell. Values of m are the number of orbitals associated with a particular sub shell in main shell. Values of m lie from 0 to l. Total values of ‘m’ for a given n is n2. Total values of ‘m’ for a given l is 2l +1. The table shows a clear relation between quantum numbers.

- Spin quantum number (s) tells the spin of the electron. It can have two value (clockwise) and (anticlockwise). Mathematically where s is amplitude of spin quantum angular momentum.

### SHAPE OF ORBITALS

- s orbitals are spherically symmetrical.
- p orbitals are dumbbell shaped.
- d orbitals have five different orientation. Three of them dxy, dyz, dxz are identical in shape but have different orientation.
- The plane passing through nucleus where probability of finding the electron is zero is called a nodal plane. Number of nodal planes in an orbital = l. Number of nodal planes increases with increasing value of n. e.g. 1s has no nodal plane. 2s has one nodal plane. For e.g. : s orbitals (l=0) have no nodal plane, p orbital (l=1) have one nodal plane, d orbitals (l=2) have two nodal planes.

- Orbitals of a sub shell having same energy are called degenerated orbitals.
- Spherical surface within an orbital where probability of finding an electron is zero is called spherical or radial node. Number of spherical nodes = (n–l–1). Angular or non spherical nodes = (l). Thus total nodes = (n – 1).

### PAULI’S EXCLUSION PRINCIPLE

- It is not possible to accomodate more than two electrons in an orbital. In other words. s sub shell can have maximum of 2 electrons p sub shell can have maximum of 6 electrons. Thus max. no. of electrons in a shell can be 2n2.
- Maximum number of electrons in a sub shell can be 2, 6, 10, 14 in s, p, d, f respectively and max. electrons in an atomic orbital can be 2.

### AUFBAU’S RULE

- The order of energies can be calculated by (n + l) rule. i.e. orbitals are filled in order of increasing (n+l) values the one with lower n value is filled first.
- The energy of atomic orbitals for H-atom depends on the value of n only.

### HUND’S RULE OF MAXIMUM MULTIPLICITY

- This arrangement leads to lower energy level.
- Singly occupied orbitals should have same spins giving rise to lower energies.

### RADIAL PROBABILITY DISTRIBUTION CURVES

### ANGULAR PROBABILITY DISTRIBUTION CURVES

The length of line decreases with increasing in the value of angle . Hence there are more chances for finding the electrons along the axes for p orbitals.

### RITZ. COMBINATION PRINCIPLE

### COMPTON EFFECT

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