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ATOMIC STRUCTUREATOMIC STRUCTURE-:-:made easy:-

by

M.,Anwar SohailM.,Anwar Sohail

Bachelor of Science and Education

Master of Science (Organic Chemistry)CHEMISTRY EDUCATORCHEMISTRY EDUCATOR

Pelham High

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Democritus’ atomDemocritus’ atom[Hypothetical / Not based on [Hypothetical / Not based on

experiments]experiments]

Democritus proposed

that matter is composed

of tiny indivisible

particles called ‘atom’

The word ‘atom’ means

unable to be divided.

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Dalton’s atomic theory Dalton’s atomic theory (1808)(1808)

[Based on experiments][Based on experiments] Every element is made of tiny,

unique particles called atoms that cannot be subdivided.

Atoms of the same element are exactly alike.

Atoms of different elements can join to form molecules.

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Discovery of Discovery of fundamental or fundamental or

subatomic particlessubatomic particles

The electrons, protons and neutrons are called fundamental particles or fundamental subatomic particles.

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Canal Rays and Protons (1886)

Eugene Goldstein noted streams of positively charged particles in cathode rays in 1886.– Particles move in opposite direction of cathode rays. – Called “Canal Rays” because they passed through holes

(channels or canals) drilled through the negative electrode. Canal rays must be positive.

– Goldstein postulated the existence of a positive fundamental particle called the “proton”.

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Discovery of Electrons (1897)

The Discharge Tube

ElectronsElectrons are discoveredby J.J. Thompson when high voltage is applied across a sealedglass tube called the ‘discharge ‘discharge tube’tube’ or CRTCRT at very low pressure.He found that what was called as cathode rays until his time was not “rays” but “particles” travelling from cathode to anode.He called them electrons.electrons.

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Discovery of Neutrons Discovery of Neutrons (1932)(1932)

James Chadwick in 1932 analyzed the results of -particle scattering on thin Be films.

Chadwick recognized existence of massive neutral particles which he called neutrons.– Chadwick discovered the neutron.

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Characteristics of Characteristics of subatomic particles at a subatomic particles at a

glanceglance

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Thomson’s Atomic model(1898)(Also called Plum-pudding model)

Thomson puts together the subatomic particles and comes forward with his atomic model.

In the atom the mass and the positive charge is evenly distributed throughout the atom (like pudding) and the negatively charged electrons are embedded in it like the plum.

He could not experimentally prove his model.

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Rutherford’s Alpha particles Rutherford’s Alpha particles scattering experiment scattering experiment

(1911)(1911) Rutherford

bombarded Alpha particles on a very thin(0.00006cm) gold foil.

Most of the particles passed through, some deflected at large angles and 1 in 20000 deflected back to its own path.

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Inference from Rutherford’s experimentInference from Rutherford’s experiment

Almost all the Alpha particles passed through the gold foil means Most of the atom is empty space.

Some of the + charged alpha particles are deflected at large angles because there is a very tiny dense core of mass and + charge located in the atom. (Called nucleus)

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Rutherford’s Alpha Rays Scattering ExperimentRutherford’s Alpha Rays Scattering Experiment

Gold Foil

Alpha Particles Source

Most of the Alpha Particles passed Through.

Some Deflected at large angles.

One in every 20,000 deflected back on its own path.

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Rutherford’s Atomic model (1920)

Based on his experiment he postulated a

model. The important postulates are:

The atom is mostly hollow. The mass and the positive charge

(protons & neutrons) are located at the center at a very small portion called nucleus.

The electrons revolve around the nucleus like the planets revolve around the sun.

Also called Planetary model

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Atomic (z)# and Mass #Atomic (z)# and Mass # Atomic # is the # of protons present

inside the nucleus of an atom. It’s unique to each element therefore,

its the identity of an element. No two elements can have the same

atomic number. Elements are listed in the periodic

table in the increasing order of their atomic numbers.

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Mass # (A)Mass # (A) Mass # is the sum of the # of protons

and neutrons present inside the nucleus of an atom. Therefore, must be a whole # not decimal.

A = p + nA = p + n OR A = Z + nA = Z + n The periodic table lists the average

atomic mass not the mass #. Atomic mass rounded to nearest

whole number is the Mass #. Mass # has no units.

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Isotopes Atoms of the same element with

different mass #. They have

1. Same atomic #

2. Same symbol

3. Same # of protons and electrons

4. Different # of neutrons & mass #

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Isotopic SymbolIsotopic Symbol

Atomic # (Z)

Mass # (A)Net Charge

Symbol

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The 3 Nuclie of H isotopes

Complete Isotopic SymbolIsotopic Symbol Worksheet

P

N

Z

M

e

P

N

Z

M

e

P

N

Z

M

e

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Average Atomic Mass Weighted average of the atomic

masses of all the naturally occurring isotopes of an element is called average atomic mass.

It is measured in amu (atomic mass unit)

1 amu is 1/12th of the mass of C-12 atom.

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Calculating Average Atomic Mass

(Mass of A*%) + (Mass of B*%)

100

Complete Average Atomic Mass – 1 Average Atomic Mass – 1 Worksheet

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Isotones Isotones are atoms of different

elements with same # of neutrons. Examples: S – 32 and P – 31 Ca – 40 and K – 39

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Problems with Rutherford’s modelProblems with Rutherford’s model

As per the classical laws of Physics: if a particle (electron) is revolving around oppositely charged particle (positive nucleus), the revolving particle loses its energy continuously and finally falls in to the central particle. Therefore the atom should collapse.But this is not happening in nature.

If the negatively charged electron is revolving around positively charged nucleus, the atomic spectra should be a band spectrum but in nature the atomic spectrum is line spectrum.

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Band spectrumBand spectrum When white light is passed through a

prism, it splits in to 7 different colors and they appear as bands of 7 colors on a film or screen. (Example in nature: Rainbow)

This is called a band spectrum. It is not a characteristic of an atomic spectrum

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Line Spectrum or Atomic emission Line Spectrum or Atomic emission spectrumspectrum

When electricity is passed through a tube filled with a gas (Ex.CRT), light will be emitted. If the emitted light is passed through the prism and its image is recorded on a film it appears as ‘sharp lines on black background’. This is called line spectrum or “atomic emission spectrum” .Every element has a characteristic emission spectrum of its own.

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Absorption spectrumAbsorption spectrum An absorption spectrum is formed by

shining a beam of white light through a sample of gas.– Absorption spectra indicate the wavelengths of

light that have been absorbed by the gas.– It appears as dark lines on bright background.

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Characteristics of LightCharacteristics of Light

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Characteristics of LightCharacteristics of Light Velocity (cc): Distance traveled by light in 1 second.

It’s a constant 3.00 x 10 3.00 x 10 88 m/s m/sc = c =

Wave length (): Distance between any two similar points on successive waves. Measured in mm or nmnm (nano meters) 1nm = 10 – 9 m = 10 – 7 cm = c / = c /

Frequency ( ): # of waves that cross a given point in 1 second. Measured in Hertz (HzHz) or cycles per second (cpscps) = c / = c /

Amplitude: Height of a crest or depth of a trough. Refers to the intensity of light.

Energy (E): Energy contained in a wave. Measured in Joules (J) E = h

Where h is Planck’s constant (6.626 x 10-34 J)

Complete Characteristics of lightCharacteristics of light Worksheet

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Frequency, Wavelength & Energy relationships

When frequency increases: Energy increases Wavelength decreases

When Wavelength increases: Energy decreases Frequency decreases

When amplitude decreases: intensity (brightness of light) decreases

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Bohr’s atomic modelNeils Bohr presented his atomic

model retaining the basic idea of Rutherford’s model. The important postulates are:

1. Electrons revolve around the nucleus in definite, closed, circular paths called orbits.

2. Each orbit is associated with a definite amount of energy therefore also called as energy level.

3. These orbits or energy levels are numbered 1,2,3,4….. or K,L,M,N…. from inside onwards. Bigger the orbit, greater is the energy associated with it.

They are also called principal quantum levels, represented by ‘n’.

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Bohr’s model Continued:-

4. More than one energy levels are possible for an electron. However, as long as an electron is in an energy level its energy remains constant.

5. When an electron gains energy it jumps from lower energy level to higher.

6. When it jumps back from higher energy level to lower, it loses energy in the form of light.

7. The energy released ( E )can be calculated by:

Where ‘h’ is Planck’s constant, ‘v’ is the frequency of light emitted.

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Modern ModelModern Model

oror

Wave Mechanical ModelWave Mechanical Model

or or

Quantum Mechanical Model Quantum Mechanical Model

of the atomof the atom

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Particle nature of light(1901)Max Planck's Quantum TheoryMax Planck's Quantum Theory

Max Planck studied the radiation emitted by various objects at high temperatures and came to a conclusion that:

Light is absorbed or emitted by matter in the form of discrete packets of energy. Each energy packet is called Photon and the energy it holds is termed Quantum.

The energy contained in each PHOTON of light is directly proportional to its frequency and can be calculated by the equation: E=hv

Where h is Planck’s constant (6.626 x 10-34 J) Planck’s quantum theory helped understanding the

phenomenon of Photoelectric effect (ejection of electrons from the surface of metal when light of a certain frequency

falls on it.)

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DeBroglie’s DeBroglie’s Dual nature of electronDual nature of electron

Based on Planck’s quantum theory and Bohr’s quantized orbits, De Broglie suggested that:

every moving particle exhibits a wave nature so also the electrons.

electrons behave more like waves on a vibrating string than like particles.

The wave length of any particle wave can be calculated by the equation: ==h/mvh/mv (De Broglie’s equation)

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Heisenberg’s uncertainty principleHeisenberg’s uncertainty principle

Its impossible to find out both the position and the speed of an electron accurately at the same time.

It is because to locate an electron, light, having wave length shorter than the size of an electron should fall on it and be reflected. Light with such a short wave length will have very high energy, which will energize the electron. Therefore, its velocity around the nucleus will increase.

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Splitting of Bohr’s Spectral linesSplitting of Bohr’s Spectral linesIntroduction toIntroduction to

Orbitals/Quantum #sOrbitals/Quantum #s

Under Magnetic Field

Under High resolution spectroscope

Under Electric Field

With a regular spectroscopeWith a regular spectroscope

sharp principal diffused fundamentalPrincipal Quantum # =(n)Principal Quantum # =(n)

Azimathul Quantum # =(l)Azimathul Quantum # =(l)

Magnetic Quantum # =(m)Magnetic Quantum # =(m)

Spin Quantum # =(s)Spin Quantum # =(s)

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Principle Quantum #

Bohr’s spectral lines. = Energy Level or Bohr’s atomic orbits. Values: any # 1,2,3,….so on from inside

outwards. Value can’t be zero The total # of electrons that can be

accommodated in an energy level is given by 2n2 where n is the Principle Quantum # or energy level.

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Azimathul Quantum # l Splitting of Bohr’s spectral lines under high resolution

spectroscope. = sub energy Levels or orbitals. Values: 0 to n – 1 Total # of “l ” values will be equal to n Ex: For n=4 l values will be: 0,1,2,3 Figure out the l values for 1st, 2nd & 3rd energy levels l = 0 : s orbital l = 1 : p orbital l = 2 : d orbital l = 3 : p orbital

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Magnetic Quantum # (m) Splitting of high resolution lines in

magnetic field. Also called angular momentum Q #. # of m values for each l value = 2l +1

(How many?) Value ranges from – l to 0 to + l (What are they?)

Practice:

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Figure out the m values1. How many m values are there for s orbital?

2. What are they?

3. How many m values are there for p orbital?

4. What are they?

5. How many m values are there for d orbital?

6. What are they?

7. How many m values are there for f orbital?

8. What are they?

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Spin Quantum # (s) Indicates electron spin in the orbital or

electron cloud (either clockwise or counterclockwise)

Values: for each m value there are 2 s values; they are +1/2 and -1/2

This indicates that in each m there are 2 electrons one spinning clockwise and the other counterclockwise.

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Atomic OrbitalsAtomic Orbitals As it is impossible to locate an electron’s exact

position at a given time, therefore: The area around the nucleus where the probability

of finding an electron is maximum is called an orbital.

There are 4 atomic orbitals discovered so far. They are s,p,d,f

The s orbital is spherical shaped electron cloud, the p orbital is a dumbbell shaped electron cloud and the d orbital is a double dumbbell and the f orbital is an 8 lobbed dumbbell.

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Orbitals and ElectronsEnergy level

Principal Quantum # (n)

Types ofOrbitals

Azimuthal Quantum # (l)

# of OrbitalsMagnetic Quantum

# (m)

ElectronsSpin Quantum # (s)

Total # of electrons in the energy level

11 ss 11 22 22

22 ss 11 22 88

pp 33 66

33 ss 11 22 1818

pp 33 66

dd 55 1010

44 ss 11 22 3232

pp 33 66

dd 55 1010

ff 77 1414

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s Orbitalss OrbitalsThere is one ss Orbitals in each energy level

Each one can hold 2 electrons.

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p Orbitalsp OrbitalsThere are 3 p Orbitals in each energy level

(from 2nd energy level on wards)Each one can hold 2 electrons therefore 6 electrons in each p sublevel

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d Orbitalsd OrbitalsThere are 5 p Orbitals in each energy level

(from 3rd energy level on wards)

Each one can hold 2 electrons therefore 10 electrons in each d sublevel

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What does the modern What does the modern atom look like?atom look like?

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Electron Configuration Arrangement of electrons in the various orbitals

of the atom of an element is called electron configuration.

It is governed by 3 laws: Aufbau Principle: Electrons occupy the lowest

energy orbital available. Pauli’s exclusion principle: No more than two

electrons in each orbital. Hunds rule: When degenerate orbitals are

available, Pairing of electrons takes place after half filling.

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Atomic Orbital Energy Diagram

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End of End of PresentationPresentation

Remember the Atomic Remember the Atomic Structure.Structure.

It is the key to learn ChemistryIt is the key to learn Chemistry