The Life of Atom

http://www.youtube.com/watch?v=hhbqIJZ8wCM

Birth

• In 1809 Dalton came up with the first picture of the atom. • Tiny indestructible sphere

Childhood

• In 1897 Thomson discovered very light weight negatively charged particles (electrons)

• Chemists determined that the negative charge must be balanced by a positive charge: the raisin bun model

Adolescence

• In 1911 Rutherford published the results from the gold-foil experiment

The Results…

• Until now atoms were thought to be solid • Most of the alpha particles went right through

the foil• Some alpha particles curved when they went

through• Only a few alpha particles deflected back

(which is what was expected)

Gold-Foil Conclusions

• The atom is made up of mostly empty space• Alpha particles are positive and curved when

they got too close to the small nucleus• Only alpha particles that hit the nucleus got

deflected back (rarely happened so must be small)

Teen Years

• Entirely positive nucleus would explode (+ charges repel)

• The total mass number of the atom could not be accounted for

• In 1932 Atom gets a girlfriend…the neutron is discovered!

Rutherford’s Model of Atom

• The nucleus is small and made up of protons and neutrons

• The electrons circle around the nucleus

Trouble in paradise?

Rutherford’s model doesn’t quite work• Electrons should lose energy and crash into

the nucleus but this doesn’t happen• 19th century physics dictates that a body in

motion must continuously give off energy- seen as a continuous spectrum through a spectroscope- but we see a line spectrum

Bohr’s Addition to Atom

• In 1913 Bohr explains why a line spectrum is seen instead of a continuous spectrum

• Electrons are only giving off certain frequencies of light

• Electrons travel in defined spaces called orbitals, which have defined energy

How does a line spectrum say all that?

• When an electron is excited it jumps from one orbital to a higher one

• The electron does not stay excited and eventually goes back to its ground state (original orbital)

• A wave of light is emitted (photon) from this process which can be seen as a line on a line spectrum

Problems with Bohr’s Theory

• Bohr couldn’t explain why lines appeared in ones, threes, fives and sevens

• Physicist Max Planck supported Bohr’s idea that atoms can absorb or emit only discrete quantities of energy called quantums

• Einstein called these packets of energy photons

Adulthood

• In 1926 Schrodinger derived the quantum mechanical model of the atom

• Described atoms as having wave-like properties which came from de Broglie’s hypothesis

• Mathematically determined the shape of the orbitals and the probably of an electron being in a certain place at a certain time- orbitals are not just spheres anymore

• In 1927 Heisenburg stated: Although the shape of the orbital is predictable, the exact location of an electron cannot be determined.

This is the Heisenburg uncertainty principle

Quantum Theory

• Quantum is the new and improved Bohr-Rutherford model

• Model shows electron placement and helps to determine valence electrons and stability of the atom

• Each orbital can hold a maximum of 2 electrons

Orbital Shapes and Orientation

S is a sphere shape– 1 orientation = 1 orbital = 2 electrons

P is a figure 8 or dumbbell shape– 3 orientations = 3 orbitals = 6 electrons

D orbitals have a clover shape– 5 orientations = 5 orbitals = 10 electrons

F orbitals have many shapes– 7 orientations = 7 orbitals = 14 electrons

Rules for Quantum

1. Aufbau Principle - each electron is added into the subshell with the lowest energy orbital available

2. Hund’s Rule - Each orbital subshell gets a single electron first and then electron can pair. All electron are ‘up’ when single

3. Pauli Exclusion Principle - no electron can have the same 4 quantum #s in an atom - electron sharing an orbital have opposite spins

Quantum Number

Principle quantum number ‘n’• Describes the orbital’s energy level and relative sizeOrbital shape quantum number ‘l’• Describes the orbital’s shape, energy of subshellsMagnetic quantum number ‘ml’• Describes the orientation in spaceSpin quantum number ‘ms’• Describes the behaviour of a specific electron in an

orbital

Summary of Quantum numbers of electrons in atoms

Name Symbol Allowed Values Property

Principal (Shell) n Positive integers (1, 2, 3, etc.)

Orbital size and energy level

Orbital shape (subshell)

l Integers from 0 to (n-1)

Orbital shape (l values 0, 1,2 and 3 correspond to s, p, d and f orbitals)

Magnetic ml Integers from –l to +l

Orbital orientation

Spin ms +1/2 or -1/2 Spin orientation