HISTORY OF THE ATOM

1910 Ernest Rutherford

He oversaw his students, Geiger and Marsden, carrying

out his famous Gold Foil experiment.

They fired alpha particles (helium nuclei), from the

radioactive decay of radium, at a piece of gold foil which

was only a few atoms thick.

• What did Rutherford expect?

• What did Rutherford find?

• What did Rutherford find?

According to Rutherford:“It was quite the most incredible event that has ever happened to me in my life. It was almost as incredible as if you fired a 15-inch shell at a piece of tissue paper and it came back and hit you. On consideration, I realized that this scattering backward must be the result of a single collision, and when I made calculations I saw that it was impossible to get anything of that order of magnitude unless you took a system in which the greater part of the mass of the atom was concentrated in a minute nucleus. It was then that I had the idea of an atom with a minute massive center, carrying a charge.”

—Ernest Rutherford

Observations They found that:• most of the alpha particles passed

through the gold foil unaffected as they had predicted.

• some were deflected at small angles following a curved path – no real surprise.

• to the astonishment of all - 1 in 8000 -

a few were deflected back at angles greater than 90 degrees.

Observations They found that:• most of the alpha particles passed

through the gold foil unaffected as they had predicted.

• some were deflected at small angles following a curved path – no real surprise.

• to the astonishment of all - 1 in 8000 -

a few were deflected back at angles greater than 90 degrees.

• Rutherford’s new evidence allowed him to

propose a

• more detailed model of the atom with a central

nucleus.

• The nucleus contains most of the mass of the

atom.

• He suggested that the positive charge was all

in a

central nucleus. The moving negative electrons

were

held in place by electrostatic attraction.

• The atom is mainly empty space filled by the

orbiting electrons – like planets around the Sun.However, this was not the end of the story.

Problems with the Rutherford Model• An atomic nucleus composed of

entirely positive charges should fly apart due to electrostatic repulsion.

• The model could not explain the total mass of the atom.

• Nineteenth century physics stated that an electron in motion around a central body must continuously give off radiation. This would mean that atoms should continuously glow.

Problems continued…

• If electrons are releasing radiation then they are losing energy. This loss of energy would cause their orbit radius to decrease and eventually the electron would spiral into the nucleus and matter would collapse.

Along came Neils Bohr…

HISTORY OF THE ATOM

1913 Niels Bohr

Bohr studied under Rutherford at

Victoria University in Manchester.

Bohr proposed a new model of the

hydrogen atom. This model

retained Rutherford’s nucleus but

did not allow the electrons to

move anywhere within the

volume of space around the

nucleus.

HISTORY OF THE ATOMThe Bohr Atomic Model of the Hydrogen Atom• Rutherford’s planetary model is correct. • When an electron is in an "allowed" orbit it does not

radiate. Bohr’s model implied that the classical electromagnetic theory did not apply at the atomic level.

• When an electron absorbs energy from incident electromagnetic radiation, it "quantum jumps" into a higher energy allowed state. This higher energy state corresponds to an allowed orbit with a higher value of the integer n. ( n = 1, 2, 3, 4, …)

• When an electron is in a higher energy state, it can quantum jump into a lower energy state, one with a smaller value of n, emitting all of its energy as a single photon of electromagnetic energy.

HISTORY OF THE ATOM

Emission spectra are produced by thin gases in which the atoms do not experience many collisions (because of the low density). The emission lines correspond to photons of discrete energies that are emitted when excited atomic states in the gas make transitions back to lower-lying levels. A continuum spectrum results when the gas pressures are higher. Generally, solids, liquids, or dense gases emit light at all wavelengths when heated.

An absorption spectrum occurs when light passes through a cold, dilute gas and atoms in the gas absorb at characteristic frequencies; since the re-emitted light is unlikely to be emitted in the same direction as the absorbed photon, this gives rise to dark lines (absence of light) in the spectrum.

HISTORY OF THE ATOM

The Bohr Model Successes• The model successfully predicted the

lines in the visible portion, uv portion , and infrared portion of the spectrum for hydrogen.

The Bohr Model Limitations• The model only explained the spectra

of one-electron systems. It could not explain the emission spectra produced by atoms of two or more electrons.

ATOMIC STRUCTURE

Particle Relative Charge (C) Mass Charge unit of electric charge

(m) proton + 1 +1.6 x 10-19 1.00732

neutron no charge 0 1.00871

electron - 1 - 1.6 x 10-19

0.00055

ATOMIC STRUCTURE•  

Mass Number - the total numberof protons and neutrons in an atom

Atomic Number - the number of protons in an atom

By definition, the word atom implies a neutral particle. Therefore, there are an equal number of protons and electrons in any atom.

SUMMARY

1. The Atomic Number of an atom = number of

protons in the nucleus.

2. The Atomic Mass of an atom = number of

Protons + Neutrons in the nucleus.

3. The number of Protons = Number of Electrons.

4. Electrons orbit the nucleus in shells.

5. Each shell can only carry a set number of electrons.

ATOMIC STRUCTURE

Electrons are arranged in energy levels or

shells around the nucleus of an atom.

• first energy level a maximum of 2 electrons

• second energy level a maximum of 8

electrons

• third energy level a maximum of 18 electrons

(when it is the outer most

energy level a max. of 8)

max. # of electrons in an energy level = 2n2

ATOMIC STRUCTURE

There are two ways to represent the atomic

structure of an element or compound;

1. Electronic Configurations: the arrangement of the electrons around the nucleus

2. Energy Level Diagrams

Carbon• Electron Configuration– 1s22s22p2

• Lewis Dot Diagram for Carbon–Why can carbon form 4 bonds?

• Energy level diagram– Ground state/ excited state– http://www.youtube.com/watch?v=Vb6kAxwSWgU– http://www.youtube.com/watch?v=K-jNgq16jEY

Rules for making energy levels diagrams

• Pauli exclusion Principle:  Electrons can not have the same 4 quantum numbers.  Electrons with opposite spins can occupy the same atomic orbital.

• Aufbau principle: Electrons are placed in the orbitals, starting with the lowest energy orbitals first.  (MAX 2 electrons per orbital). A sublevel must be filled before moving onto the next higher sublevel.

• Hund’s Rule: When electrons are placed in a set of orbitals of equal energy, they are spread out as much as possible to give as few paired electrons as possible.

Quantum Numbers (aka the electron’s address)

• The Principal quantum number, n is the main energy of an electron, where n= 1,2,3,4,….   (Energy LEVEL)

• The Secondary quantum number, l represents the shape of the electron orbit.  • This describes the additional electron energy sublevel. • The number of values that l can have equals the principal quantum number. 

The values for l = 0,1,2,3

• The Magnetic Quantum number, ml gives the direction of the electron orbit (or orientation in space) –  orbital• If l = 0, then m  = 0• if l = 1, then m  = -1, 0, 1• If l = 2, then m  = -2,-1,0,1,2   etc….

• The Spin Quantum number, m   gives the electron’s spin• The values for m  = +1/2, -1/2 (clockwise or counter clockwise)

•  

Quantum NumbersnPrincipal Quantum number(major shell or energy level)

lSecondary Quantum Number- shape(subshell or sublevel)   

mMagnetic quantum number(orbitals: volume of space- 2 e /orbital)

SSpin quantum number(spin)

Subshell = quantum #

max # of e  Sub-shell

# of orbitals – quantum number

1 s = 0 2 -     total 2 s 1        0 Clockwise (+1/2)

2 s = 0p = 1

26   -   total 8

sp

1        03     -1,0,+1

or counterclockwise (-1/2)

3 s = 0p = 1d = 2

2610 –  total 18

spd

1         03     -1,0,+15   -2,-1,0,+1,+2

4 s =0p = 1d = 2f =3

261014 –  total 32

spdf

1          03       -1,0,+15 -2,-1,0,+1,+27-3-2-10+1+2+3

Some exceptions to energy level diagrams

EXPECTED• Cr  [Ar]4s23d4 Cu [Ar]4s23d9

ACTUAL • Cr  [Ar]4s13d5 Cu [Ar]4s13d10

What happens? Electrons are ‘borrowed’ or ‘promoted’ from the 4s subshell to give a 3d subshell that is exactly half- filled for Cr or completely filled for Cu.  A half filled or completely full subshell has a special stability as it is lower in energy.

Some exceptions to energy level diagrams

• Ag    [Kr]5s14d10

• Au   [Xe]6s14f145d10

• What happens?  Again the electrons are promoted from the s subshell to the d orbital as it is more stable, of lower energy…  more energetically favourable to do this.

• U   [Rn]7s2 5f3 6d1

• Why is Fe +2 and +3 ?  Sn??• http://www.chemguide.co.uk/inorganic/group4/oxstates.html