ATOMIC STRUCTURE

ATOM - smallest part of an element that can exist by itself.Cannot be created nor destroyed

The ATOM

400 B.C.- Democritus- invisible atom • Earth• Water• Air• Fire

The ATOMMiddle Ages- The ALCHEMISTS

Tried to change materials into Gold.

Failed, but developed lab techniques and discovered elements.

Atomic Theory

1803 John Dalton's Atomic Theory

Dalton's Atomic Theory

1. All matter composed of atoms.2. Atoms of elements are identical.3. Atoms cannot be created nor destroyed.4. Atoms are combined in specific ratio's.

Cathode Ray Tubes-

Electric currents pass through them. 

Cathode Ray Tubes-1897 Sir Thomson-

Suggested that cathode ray tubes consists of (-) particles.These particles are deflected by both magnetic and electric fields.

Plum Pudding ModelDevised by Thomson.Negative “plums” surrounded by positive “pudding”.

ELECTRONS-

sub atomic particle.negatively charged1/1837th the mass  of the proton.

ELECTRONS

Robert Millikan-1909 Oil drop experiment.Found the mass of an electron.

ELECTRONS-

Atoms are neutral, so there must be something positive (+).Electrons  are small, so there must be something larger that accounts for the mass.

Lord Ernest Rutherford-

1908 Known for two major

experiments.

GOLD FOIL EXPERIMENT

Rutherford’s ExperimentBombarded metal foil with fast moving positive particles.

1 out of 8,000 bounced straight back.

Some powerful force repelled them. 

Rutherford ModelThe atom is mostly empty space.The center must have a positive charge.This center must occupy a small amount of space.It is very dense.

NUCLEUS

Positive charged. Dense in center. Contains most mass little volume. Atoms mostly empty space,

electrons surround.

NUCLEUS

Contains  nucleons which include  neutrons   and protons

PROTONS

Positive charged.Located in the nucleus.

NEUTRONS

Electrically neutral- no charge.Equal in mass to protons.Discovered by James Chadwick in 1932.

MASS DISCOVERED -28

Electron - 9.1 X 10 1/1837 Thomson 1897 -24

Proton + 1.67 X 10 1 Early 1900’s -24

Neutron 1.68 X 10 1 Chadwick 1932

 

Atomic Number

Number of protons. Protons determine the atom.Each element has its own unique number of protons Atomic number stays constant

Atomic Number

Mass Number  (atomic mass)-

Total # of protons + neutrons

U-235  92 = Atomic NumberMass Number = 235 - 92 = neutrons ---------------------------------------

23 11 Na=  11 protons                           12 neutron                           11 electrons

Isotopes

Same element with different # of neutrons.Number of protons do not change.Nuclides - name given to any isotope

IONS

Charged atoms.Positive have lost electrons Negative have gained.

 

Atomic Mass Unit-AMU

Atomic Mass Unit  (u) =  1/12 the mass of C-12 1.66 X 10 -24  g

Average Atomic Mass

Average Atomic Mass

Average Atomic Mass- weighted average of all Natural atomic masses.All isotopes put together by %

Cl=34.969u     47.3% Cl=36.966 u    52.7%

Average Atomic Mass

Average Atomic Mass = % (mass) + % (mass)

Average Atomic Mass

Find the Average Atomic Mass of element X. 40.7% weighs 13.78 grams, 30.3% weighs 14.12 grams, and 29.0% weighs 14.99 grams…

= .407 (13.78) + .303 (14.12) + .290

(14.99)

Electromagnetic Spectrum

Travel at speed of light  186,000 m/s

Can travel through a vacuum

Electromagnetic Spectrum

Radiowaves ------longest wavelength Microwaves   infra- red   ROY-G-BIV    Ultra -violet   X-ray    Gamma------shortest wavelegth

The longer the wavelength– the slower the frequency

Quantum Theory

1900 Max Planck- German Physicist- light has particle like propertiesa Quanta is a finite quantity of energy.

Photoelectric Effect

Photoelectric Effect- The emission of electrons by

metals when light shines on it. Energy given off.

Quantum Theory

1905 Al Einstein-

Spectroscopy Spectroscopy- Study of the emission spectra of certain substances.  Each atom gives off Specific lines.  Continuous spectrum- separates into a band of colors that blend.Bright line spectrum - separate into specific lines of colors.

Spectroscopy

Each element has it's own specific  bright-line spectrum.  In this way, the composition of various substances may be found.

Bohr

Neils Bohr- (1913) 

Bohr felt that electrons could only be in certain areas around the nucl eus.

Bohr Theory

Electrons revolve around nucleus in

specific energy levels.

Have specific orbits - K, L, M, N

Principle Energy Level, electron shells, + energy levels all mean the same.

Bohr Theory

Shell       # of electrons  (2n 2 )1 22 83 184 325 50

Bohr TheoryGround state- lowest energy state of atom.Excited state-higher energy state than ground. Electrons are further away.

Bohr Theory

Electrons normally in ground state.When heated they absorb energy and they jump to a higher excited state.

Bohr Theory

Instantaneously the electrons return to the ground state.They do this by releasing (emitting) photons.This can be seen by the various colors produced.

Bohr Theory

Normally in the ground state (closest).

Addition of energy makes the electrons jump to excited levels.

Energy emitted when they go from higher back to closer level.

Energy is released as photons (spectral lines)

Electron Configuration(regents level)

A way to designate electrons and the shells that they are in.Some examples include:

C 2-4 H 1 Na 2-8-1 Cl 2-8-7 K 2-8-8-1****all of these are in the ground

state

Electron Configuration

Some more examples:

O 1-7 Na 2-7-2 Al 2- 6- 5

*** these are all in the excited state

Electron Configuration IONS

+1 +3 -2 -1

K Al O Cl

Valence ElectronsValence Electrons are located in the outermost shell or Principle Energy Level.

OK TIME TO GET TOUGH!

Electron Cloud Theory-

1926 Erwin Schrodinger.

Electrons located in 3-D  region called  Orbitals.

Heisenberg Uncertainty Principle

 

Not possible to know exactly where the electron is.

We can only state the probability of finding a particle at a particular time and place.

 

Werner Heisenberg (1927)

Wave-Mechanical Model

DeBroglie

Wave-Mechanical Model

Electron Dot Diagrams

1916 G.N Lewis

Shows only valence electrons.

QUANTUM THEORY

Sub- Levels

These are the various shapes where electrons can be located.

s-p-d-f

s Orbitals

p Orbitals

d Orbitals

f Orbitals

QUANTUM NUMBERS

Principle quantum # 1, 2, 3 energy shells

Orbital quantum # s, p, d tells the shape 

Magnetic quantum # shows the orientation P has three orbitals Px Py Pz 

Spin quantum # (+) or (-) direction of electron’s spin

Quantum numbers describe the state of a particular electron.

Quantum Numbers- tells the properties of the orbitals and their electrons

  Principle Orbital # of Orbitals Total e-

  2 2

n 1 n 2n 

1 s 1 2------------------------------------------------------------------------------------------------------------------------  

2 s 1 2 p 3 6------------------------------------------------------------------------------------------------------------------------  

3 s 1 2 p 3 6 d 5 10 ------------------------------------------------------------------------------------------------------------------------  

4 s 1 2 p 3 6 d 5 10 f 7 14  tells energy shell tells the shape

(subshells)   

Stupid People Die First

(n)   1    s 

2 s p 3 s p d  4 s p d f orbital 1 3 5 7  electrons 2 6 10 14 

Electrons will fill:

 Ground State- Fills with the electrons as close to the nucleus as possible. Each electron will fill the area with the least amount of energy.The order of increasing energy:

1s 2s 2p 3s 3p 4s 3d 4p 5s 4d 5p 6s 4f 5d 6p 7s 

Electron Configuration

Tells the arrangement of the electrons  2 2 6

1s 2s 2p

1,2 = shell

s,p,d = sub-levels

exponants = # of electronss has 1, p has 3, d has 5 orbitalsonly two electrons per orbital

THE APARTMENT!

*** Only two people per room1st floor: 1 Apartment 1 room, 2 people2nd floor: 2 Apartments- Apartment 2S has 1 room, 2 people Apartment 2P has 3 rooms, 6 people3rd floor: 3 Apartments Apartment 3S has 1 room, 2 people Apartment 3P has 3 rooms, 6 people Apartment 3d has 5 rooms, 10 people

2 2 6 2 6 2 10 6 2

1s 2s 2p 3s 3p 4s 3d 4p 5s

2 2 5

ground state 1s 2s 2p

  2 1 6

excited state 1s 2s 2p 2 6 2

(or) 1s 2p 3s

STOP HERE

RULES

Orbital Notation

LITHIUM 1s 2s 2p

NITROGEN 1s 2s 2p 2p 2p

Aufbau Principle-Aufbau Principle-

Electron occupies the lowest energy orbital that can hold it.

Hund’s Rule-Hund’s Rule-

No orbital in a sublevel may contain 2 electrons until all contain one.

Pauli Exclusion Principle-

Pauli Exclusion Principle- No two electrons will have the same

four quantum #’s

TRY THESE NOW!!!!

Give the following configurations for each: 1) Regents level 2) Advanced level3) Orbital Notation Sulfur Magnesium Iron Phosphorus -2 Rubidium

Cu & Cr

Special cases

Cu---- Its 4s and 3d are very close.Its 4s moves to 3d to become completely filled.

2 9 1 10

4s 3d 4s 3d

4s 3d

Special cases

Cr ----- Its 4s and 3d are very close. One of the 4s moves to 3d, since half-

filled gives more stability.

2 4 1 5

4s 3d 4s 3d

4s 3d