Atomic Theory Development

What is Today’s Model?

Dense, Positively Charged Nucleus Mostly Empty Space Negatively Charged Electron Cloud Most Probable Location of the

Electrons Composed of Protons,

Democritus proposes the 1st _atomic

theory

460 – 370 BC

History of the Atom - Timeline

Antoine Lavoisier makes a substantial number of contributions

to the field of Chemistry

1766 – 1844

John Dalton proposes his atomic theory in

1803

1743 – 1794

0

1856 – 1940

J.J. Thomson discovers the electron and proposes the Plum Pudding Model in 1897_{1871 – 1937}

Ernest Rutherford performs the Gold Foil

Experiment in 1909

1885 – 1962

Niels Bohr proposes the Bohr Model in

1913

1887 – 1961

Erwin Schrodinger

describes the electron cloud in 1926

1891 – 1974

James Chadwick discovered the neutron

in in 1932

1700s 1800s 1900s

Early Greeks

Matter is made of indestructible

particles called “atomos”

Summary for Dalton’s

Atomic Theory

(Father of the Modern Atomic Theory)

All atoms of a single element have the same mass

Atoms of different elements are different.

Atoms can’t be divided, created or destroyed.

Atoms of different elements combine in simple

Discovery of the Electron

In 1897, J.J. Thomson used a cathode ray tube to deduce the presence of a negatively charged particle.

Cathode ray tubes pass electricity through a gas that is contained at a very low pressure.

J.J. Thomson

• ALL atoms must contain

these negative particles.

• atoms did not have a net

negative charge

• there must be something

Ernest Rutherford’s (1871-1937)

 electrons embedded in a positive pudding.

Where exactly are those electrons?

Thomson’s Theory: “Plum Pudding”

 Shoot something at them to see where they are.

Rutherford’s Conclusion (1911)…

Small, dense, positive nucleus.

Equal amounts of (-) electrons at large distances

Neils Bohr’s Atomic model (1913)

Small, dense, positive nucleus.

Equal amounts of (-) electrons at specific orbits

around the nucleus.

This incorrect version of the atom is often used to represented atoms

Bohr’s Calculations

• Bohr calculated the

energies that an e-would have in the

allowed energy levels for the hydrogen

Chadwick

** James Chadwick

discovered neutrons in 1932.

---- _n0 _{have no charge}

and are hard to detect

V.Montgomery & R.Smith 14

V.Montgomery & R.Smith 15

Where are the e- in the atom?

• e- have a dual wave-particle nature (Recall deBroglie)

V.Montgomery & R.Smith 16

Heisenberg’s Idea

• e- are detected by their interactions with

photons

• Photons have about the same energy as

e-• Any attempt to locate a specific e- with a

photon knocks the e- off its course

• ALWAYS a basic uncertainty in trying to locate

e-V.Montgomery & R.Smith 17

Heisenberg’s Uncertainty Principle

• Impossible to determine both the position and

ErwinS Schrodinger

(1887-1961)

• 1926

• the exact location of an electron

cannot be stated

• more accurate to view the

electrons in regions called

electron clouds

• electron clouds are places where

the electrons are likely to be found

• Did extensive work on the Wave

formula  Schrodinger equation • Won a Nobel Prize

Image taken from: nobelprize.org/.../1933/schrodinger

V.Montgomery & R.Smith 19

Schr

Ö

dinger’s Wave Equation

• An equation that treated electrons in atoms as

waves

• Only waves of specific energies, and therefore

frequencies, provided solutions to the equation

• Quantization of e- energies was a natural

Quantum Model

Major points

1. Electrons do not follow fixed paths

2. They move randomly in areas of probability (orbitals)

3. There are specific energies associated with each orbital

New Model =

According to quantum mechanics, each electron is described by four quantum numbers:

•1. Principal quantum number (n)

•2. Angular momentum quantum number (l)

•3. Magnetic quantum number (m_l)

•4. Spin quantum number (m_s)

•The first three define the wave function for a

The Shapes of Atomic Orbitals

• the l quantum number primarily determines

the shape of the orbital

• l can have integer values from 0 to (n – 1)

• each value of l is called by a particular letter

that designates the shape of the orbital

– s orbitals are spherical

– p orbitals are like two balloons tied at the knots

l

= 0, the

s

orbital

• each principal energy

state has 1 s orbital

• lowest energy orbital in a

principal energy state

• spherical

25

2

s

and 3

s

2s

n = 2, l = 0

3s n = 3,

l

= 1

, p

orbitals

• each principal energy state above n = 1 has 3 p orbitals

– m_l = -1, 0, +1

• each of the 3 orbitals point along a different axis

– p_x, p_y, p_z

l

= 2

, d

orbitals

• each principal energy state above n = 2 has 5 d orbitals

Atomic Models-

Simulation

• The question now is what is a fundamental

particle?

• Radioactive decay shows that protons and

neutrons are not fundamental