History of the atom

• Dalton• J.J. Thompson• Rutherford• Bohr

Dalton’s Atomic Theory• All elements are composed of

atoms which are indivisible• Atoms of the same element are

identical• Atoms of different elements

can mix together in simple whole number ratios to form compounds

• Chemical reactions occur when atoms are separated, joined or rearranged. Atoms of one element can’t be changed into atoms of another element

Law of Definite Proportions and Law of Multiple Proportions

Law of Conservation of Matter

Early Atomic Models

• J.J. Thompson’s plum-pudding model

• Rutherford’s Model of the atom

Problems with Rutherford Model

• Couldn’t account for the chemical properties of the elements (why do elements react in the way that they do?)

The Wave Nature of Light

• Electromagnetic RadiationForms of energy that exhibit wave like characteristics

• Wavelength – How far apart the waves are

• Frequency– The # of waves to pass a

point in a certain timeframe• Wavelength and Frequency

are Inversely Proportional

Wavelength and Frequency• Not Working very hard

• Long Wavelength• Low Frequency

Wavelength and Frequency• Working VERY hard

• Short Wavelength• High Frequency

What does this tell you about the relationship between wavelength, frequency and energy?

The Photo Electric Effect

• In the early 1900s an experiment was done that COULD NOT be explained by light being a wave.

• Different colors were shined on to a metal plate

• Electrons would come off the metal plate for ONLY CERTAIN FREQUENCIES (colors) OF LIGHT

If light only behaved like a “wave” ANY frequency of light would cause the electrons to be released.

Brick Wall Analogy

• Different balls thrown at wall

• Each one with different mass

• All at same speed of 90 mile/hour

• Each ball is like a color or light - each has its own energy

Ping Pong Ball

No effect

Softball

No effect

Super Dense Steel Ball

Dislodges Brick and send it flying

JB

Einstein and the Photo Electric Effect

• This observation led Einstein to believe that light acted like a particle and a wave

• This is called the “dual nature” of light

• Light carried packets of energy called quanta, or photons

Neils Bohr and the New Model of the Atom

• Bohr hypothesized that electrons could only be at certain energy “levels” around the nucleus

• Electrons could “jump” from lower to higher energy states by absorbing a quantum of energy

• When an electron releases the energy it gained, specific colors or wavelengths of light are emitted

Mercury Line Spectra

Electrons and the Atom

Nucleus

Electron at “Ground State”

e-

Add Energy

e-

electron at “Excited State”

Electron at “Ground State”

e-

Just like a jumper has potential energy at the top of the jump, the electron has stored potential energy in the higher orbit.

Quantum Leap

Electrons disappear from one orbit and reappear at another without visiting the space in between!

Release Energy in the form of colored light

JB

Energy and the Electron

• Ground State – lowest energy state for an electron

• Excited state – high energy state for e-

• Quantum – exact amount of energy to move an electron from one energy level to another

Heisenberg Uncertainty Principle• It is impossible to

pinpoint the exact location and velocity of an electron at any point in time

• You can estimate where an electron will be 90% of the time

• An electron cloud shows where an electron spends most of its time

Quantum Mechanical Atomic Model

• Similar to Bohr model except that e- cannot be found in distinct orbits around the nucleus

• Determines how likely it is for an electron to be found in various regions around the nucleus.

Atomic Orbitals

• Region around the nucleus where an electron of a given energy is likely to be found

• Each orbital has a characteristic size, shape, and energy

• There are four different orbitals: s, p, d, f

Different types of atomic orbitals

Principle Energy Levels

• Symbolized by n = 1,2,3,4 etc.• Each energy level

contains sublevels denoted by a number and a letter (ex. 1s)

• Each sublevel contains a certain # of orbitals

• Every orbital can hold 2 e-

Energy Level # of Sublevels Type of sublevel and orbitals

n=1 1 1s (1orbital)

n=2 2 2s (1 orbital)2p (3 orbitals)

n=3 3 3s (1 orbital)3p (3 orbitals)3d (5 orbitals)

n=4 4 4s (1 orbital)4p (3 orbitals)4d (5 orbitals)4f (7 orbitals)

See table 5.1 on page 131

Electron Configurations• Aufbau principle: e- occupy

the lowest energy sublevels 1st (see pg. 133 figure 5.7)

• Pauli Exclusion Principle: every orbital can hold a maximum of 2 e- (paired e- spin in the opposite direction)

• Hund’s Rule: e- fill all empty orbitals in a sublevel BEFORE they pair up

Aufbau Diagram

Valence Electrons• Valence electrons –

electrons found in the highest energy level of an atom

• Valence e- determine the chemical properties of the element

• A filled outer energy level (stable octet) of 8 e- makes an atom stable (or for He 2e- fills its outer energy level)

Group Valence e- (and the orbitals they occupy)

1A 1 (s1)

2A 2 (s2)

3A 3 (s2p1)

4A 4 (s2p2)

5A 5 (s2p3)

6A 6 (s2p4)

7A 7 (s2p5)

8A 8 (s2p6)

Electron Dot Diagrams

• Shows only number of valence e-

• Electrons shown as dots

Carbon has 4 valence e-