By Aaradhya

Learning Standards

• Atomic Structure Broad Concept: Atomic models are used to explain atoms and help us understand the interaction of elements and compounds observed on a macroscopic scale.,

–Recognize discoveries from Dalton (atomic Recognize discoveries from Dalton (atomic theory), Thomson (the electron), Rutherford (the theory), Thomson (the electron), Rutherford (the nucleus), and Bohr (planetary model of atom) and nucleus), and Bohr (planetary model of atom) and understand how these discoveries lead to the understand how these discoveries lead to the modern theory.modern theory.

–Describe Rutherford’s “gold foil” experiment that Describe Rutherford’s “gold foil” experiment that led to the discovery of the nuclear atom. Identify the led to the discovery of the nuclear atom. Identify the major components (protons, neutrons, and electrons) major components (protons, neutrons, and electrons) of the nuclear atom and explain how they interact.of the nuclear atom and explain how they interact.

–Write Write thethe electron configurations for the first electron configurations for the first twenty elements of the periodic table.twenty elements of the periodic table.

All matter is composed of atomsAtoms cannot be made or destroyedAll atoms of the same element are identicalDifferent elements have different types of atomsChemical reactions occur when atoms are rearrangedCompounds are formed from atoms of the constituent elements

Atomic theory proposed by John Dalton

Atom Definition

Atom

What is inside the atom?What is inside the atom?

Discovery of Protons

• 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”.

� Passing an electric current makes a beam Passing an electric current makes a beam appear to move from the negative to the appear to move from the negative to the positive end.positive end.

Thomson’s Experiment AndThomson’s Experiment AndDiscovery of ElectronsDiscovery of Electrons

Voltage source +-

Voltage source

Thomson’s ExperimentThomson’s Experiment

� By adding an electric field he found By adding an electric field he found that the moving pieces were negative.that the moving pieces were negative.

+

-

The electron was discovered in The electron was discovered in 1897 by Thomson. He imagined 1897 by Thomson. He imagined the atom as a “raisin pudding” with the atom as a “raisin pudding” with electrons stuck in a cake of electrons stuck in a cake of positive charge.positive charge.

J.J. Thomson’s Model of Atom

• Plum Pudding Model, 1896

• Thought an atom was like plum pudding– Dough was cloud– Raisins were electrons

– Didn’t know about neutrons at this time

Rutherford’s experiment anddiscovery of nucleus

• English physicist Ernest Rutherford (1911)

• Shot alpha particles at fluorescent screen.

• When an alpha particle hits a fluorescent screen, it glows.

Lead block

Uranium

Gold Foil

Fluorescent Screen

What he expected

He ExpectedHe Expected

The alpha particles to pass through without The alpha particles to pass through without changing direction very much.changing direction very much.

What he got

He thought the mass was evenly distributed in the atom

The Nuclear Atom

Since some particles were deflected at large angles, Thomson’s model could not be correct.

How Rutherford explained results…..

• Atom is mostly empty space.

• Small dense, positive piece at center is (NUCLEUS)

Rutherford’s Findings

a) The nucleus is smallb) The nucleus is densec) The nucleus is positively

charged

“Like howitzer shells bouncing off of tissue paper!”

Conclusions:

Most of the particles passed right through A few particles were deflected VERY FEW were greatly deflected

The model created by Rutherford had still some serious discordance. According to the classic science, electron moving around the nucleus should emit an electromagnetic wave. Electron should than move not by the circle but helical and finally collide with the nucleus. But atom is stable.

Rutherford also realized that the nucleus must contain both neutral and positively charged particles. The neutron was then discovered in 1932 by Chadwick.

Atomic number and Mass number :-

• a) Atomic number (Z) :-• The atomic number of an element is the number of protons present in the • nucleus of the atom of the element.• All the atoms of an element have the same atomic number.• Eg :- Hydrogen – Atomic number = 1 (1 proton)• Helium - Atomic number = 2 (2 protons)• Lithium - Atomic number = 3 (3 protons)• b) Mass number (A) :-• The mass number of an element is the sum of the number of protons and • neutrons (nucleons) present in the nucleus of an atom of the element.• The mass of an atom is mainly the mass of the protons and neutrons in the nucleus of the atom.• Eg :- Carbon – Mass number = 12 (6 protons + 6 neutrons) Mass = 12u• Aluminium – Mass number = 27 (13 protons + 14 neutrons) Mass = 27u• Sulphur – Mass number = 32 (16 protons + 16 neutrons) Mass = 32u

• In the notation of an atom the Mass number• atomic number and mass number E g :- N• are written as :- Atomic number

Symbol of element

14

7

Isotopes

• Isotopes are atoms of the same element having the same atomic numbers but different mass numbers.

• Eg :- Hydrogen has three isotopes. They are Protium, Deuterium (D) and Tritium (T).• H H H• Protium Deuterium Tritium• Carbon has two isotopes. They are :-• C C • Chlorine has two isotopes They are :-• Cl Cl•

1

1

11

2 3

6 6

12 14

1717

35 37

Isobars

• Isobars are atoms of different elements having different atomic numbers but same mass numbers.

• These pairs of elements have the same number of nucleons.

• Eg :- Calcium (Ca) – atomic number - 20 and Argon (Ar) – atomic number 18 have different atomic numbers but have the same mass numbers – 40.

• Iron (Fe) and Nickel (Ni) have different atomic numbers but have the same atomic mass numbers – 58.

• •

Ca Ar40 40

20 18

Fe Ni58 58

27 28

Bohr’s Model of the Atom• Similar to

Rutherford’s model• Thought atom was

mostly empty space• Neils Bohr, 1913

– Nucleus in center is dense, positively charge

– Electrons revolve around the nucleus.

Following Rutherford’s planetary model of the atom, it was realized that the attraction between the electrons and the protons should make the atom unstableBohr proposed a model in which the electrons would stably occupy fixed orbits, as long as these orbits had special quantized locations

Parts of an AtomEach element has a different number of protons in its nucleus

Protons have positive chargeChange the number of protons change elements This is called nuclear physics

The element also has the same number of electrons

Electrons have negative chargeChange the number of electrons ionize the element This is called chemistry

Some elements also have neutrons Neutrons have no chargeThey are in the nuclei of atoms

p

n

e

Subatomic particles

Electron

Proton

Neutron

Name Symbol Charge

Actual mass (g)

e-

p+

n0

9.11 x 10-28

1.67 x 10-24

1.67 x 10-24

Bohr’s model• Electrons move around the nucleus at

stable orbits without emitting radiation.

• Electron in one of these stable orbit has a definite energy.

• Energy is radiated only when electrons make transitions from high energy orbit to a low energy orbit.

In the Bohr model, the electron can change orbits, accompanied by the absorption or emission of a photon of a specific color of light.

Wave Nature of Electromagnetic Radiation

Waves have 3 primary characteristics:1. Wavelength (λ): distance between two consecutive peaks in a wave.

2. Frequency (ν): number of waves (cycles) per second that pass a given point in space.

3. Speed: speed of light is 2.9979 * 108 m/s. We will use 3.00 x108 m/s.

Wavelength and frequency can be interconvert and they have an inverse relationshipv = c/λ

v = frequency (s1) λ = wavelength (m)c = speed of light (m s1)

Wavelength is also given in nm (1 nm = 10-9 m) and Angstroms (Å) (1 Å = 10-10 m).The frequency value of s1 or 1/s is also called “hertz (Hz)” like KHz on the radio.

The Particle Nature of Light

•Blackbody radiation•The photoelectric effect

Blackbody Radiation and the Quantization of Energy

A. The interior of a cold ceramic-firing kiln approximates a blackbody, an object that absorbs all radiation falling on it and appears black. A hot kiln emits light characteristic of blackbody radiation. B Planck’s formula generates a curve that fits perfectly the changes in energy and intensity of light emitted by blackbody at different wavelength for a given temperature

A B

Planck’s Formula

• To find a physical explanation of blackbody Planck made a radical assumption that the hot, glowing object could emit (or absorb) only certain quantities of energy:

• E = nhν• Where E is the energy of the radiation, ν is its

frequency, n is a positive integer (1, 2, 3 and so on) called a quantum number and h is a proportionality constant now called Planck’s constant and has value = 6,626x10-34 J.s

The Photoelectric Effect and The Photon Theory of Light• Current flow when monochromatic light of

sufficient energy shines on a metal plate• The photoelectric effect had certain features: the

presence of a threshold frequency and the absence of a time lag

• Carrying Planck’s idea of packeted energy, Einstein proposed that light itself is particulate, occurring as quanta of electromagnetic energy, called photon

• In terms of Planck’s work we can say that each atom changes its energy whenever it absorbs or emits one photon, one “particle” of light, whose energy is fixed by its frequency

• Ephoton = hν = ∆Eatom

Bohr's model for hydrogen atom

• Niels Bohr adopted Planck’s assumption and explained these phenomena in this way:

1.Electrons in an atom can only occupy certain orbits (corresponding to certain energies).

• Niels Bohr adopted Planck’s assumption and explained these phenomena in this way:2. Electrons in permitted orbits have specific,

“allowed” energies; these energies will not be radiated from the atom.• Niels Bohr adopted Planck’s assumption and

explained these phenomena in this way: 3.Energy is only absorbed or emitted in such\a way

as to move an electron from one “allowed” energy state to another; the energy is defined by

E = hν

Bohr's model for hydrogen atom

• Lyman ser ies The atom will remain in the excited state for a

short time before emitting a photon and returning to a lower stationary state. All hydrogen atoms exist in n = 1 (invisible).

• Balmer ser ies When sunlight passes through the atmosphere,

hydrogen atoms in water vapor absorb the wavelengths (visible). H atoms exist in n=2.

Similarly it will fill in:• Paschen From n=4,5……… till n=3• Brackutt from n=5,6……… till n=4 • Pfund from n=6,7……… till n=5

Bohr's model for hydrogen atom

The energy absorbed or emitted from the process of electron promotion or demotion can be calculated by the equation:

where RH is the Rydberg constant, 2.18 × 10−18 J, and ni and nf are the initial and final energy levels of the electron.

∆E = −RH ( )1nf2

1ni2

-

Limitation of Bohr’s Model

• Bohr’s model only works for hydrogen atom and other one-electron (hydrogen-like) ionic species, such as He+, Li2+, etc.• For H-atom, electronic energy: En = -2.178 x 10-18 J(1/n2)• For other one-electron particle: En = -2.178 x 10-18 J(Z2/n2)

– (Z = atomic number)• Bohr’s model cannot explain atomic spectra of atoms having more

than one electron;• Bohr’s model also cannot explain why each line in the hydrogen

spectrum appears as double-lines if the discharge tube is placed in magnetic field.

• Perhaps his treatment of electron as having only particulate properties is insufficient.

Aufbau Principle

• As protons are added one by one to the nucleus to build up the elements, electrons are similarly added to these hydrogen-like orbital.

• H : 1s1, He : 1s2, Li : 1s2 2s1, Be : 1s2 2s2

• B : 1s2 2s2 2p1, C : 1s2 2s2 2p2.

Hund’s Rule

• The lowest energy configuration for an atom is the one having the maximum number of unpaired electrons allowed by the Pauli principle in a particular set of degenerate orbitals.

• N : 1s2 2s2 2p3, O : 1s2 2s2 2p4,

• F : 1s2 2s2 2p5, Ne : 1s2 2s2 2p6,

• Na : 1s2 2s2 2p63s1 OR [Ne] 3s1

Heisenberg Uncertainty Principle

x = positionmv = momentumh = Planck’s constantThe more accurately we know a particle’s position, the less accurately we can know its momentum. Both the position and momentum of a particle can not be determined precisely at a given time. The uncertainty principle implies that we cannot know the exact motion of the electron as it moves around the nucleus.

( )∆ ∆x m vh

⋅ ≥ 4 π

Quantum Numbers (QN)

• The principal quantum number (n) is a positive integer (1, 2, 3 and so forth). It indicates the relative size of the orbital and therefore the relative distance from the nucleus of the peak in the radial probability distribution plot

Principal QN (n = 1, 2, 3, . . .)

• The angular momentum number (l) is an integer from 0 to n-1. it is related to the shape of the orbital and is sometimes called orbital-shape quantum number

Angular Momentum QN (l = 0 to n 1)

• The magnetic quantum number (ml) is an integer from –l through 0 to +l. it prescribes the orientation of the orbital in the space around the nucleus and is sometimes called the orbital-orientation quantum number

Magnetic QN (ml = l to l including 0)

Quantum Numbers (QN)

• Spin Quantum Number, (s)This led to a fourth quantum number, the spin quantum number, ms.

Electron Spin QN has only 2 allowed values: +1/2 and −1/2.

• No two electrons in the same atom can have exactly the same energy.

• No two electrons in the same atom can have identical sets of quantum numbers.

(n, l, ml, ms).

• Therefore, an orbital can hold only two electrons, and they must have opposite spins.

Pauli Exclusion Principle