DEVELOPMENT OF A NEW ATOMIC MODEL

I. Behavior of electrons related to absorption and emission of light
A. ____________________: taking in energy
B. ____________________: giving off/releasing energy

II. Properties of light
A.  Light as a wave
1. Light = form of electromagnetic (EM) radiation
2. All types of EM radiation together make up the EM ___________________
a. all types move at the speed of light (c = 3.00 x 10⁸ m/s)
b. all have ____________________ (distance between same pt. on different waves -
measured in ____ - abbreviated ____)
c. all have ____________________ (# of waves that pass a certain point in 1 sec. -
measured in ____ - abbreviated ____)
d. c = l ^. n
e. frequency and wavelength are ___________________ proportional

III. Photoelectric effect
A. Emission of e- from a metal when light shines on the metal
B. Only light of certain frequency could cause the photoelectric effect
1. Wave theory alone could not explain this
2. Light as particles
a. Object emits energy in small, specific amounts (called _________)
b. Suggested by ____________________
c. E = h ^. n
1) E =
2) h =
3) n =
d. energy and frequency are ____________________ proportional
C. Einstein suggested that light had properties of both waves and particles
Referred to as the ________________________________________ of light
D. _________________: particle of EM radiation carrying a quantum of energy

IV. Hydrogen atom line emission spectrum
A. ____________________: lowest energy state of an atom
B. ____________________: atom has higher potential energy than in ground state
C. Light (of a certain amount of energy) shone through a prism, separated into
specific frequencies of light (corresponds to different colors)
D. Quantum theory
1. When atom falls from excited state to ground state, __________________
____________________________________________________________
2. Energy of photon = difference ___________________________________
____________________________________________________________
3. Energy states of atoms are fixed

V. Bohr model of the hydrogen atom
A. said that e- circled the nucleus in fixed paths
B. when in path, has fixed amount of energy
C. e- cannot exist in space between paths
D. drawback of Bohr’s model =

EMISSION & ABSORPTION SPECTRA

According to the Bohr atomic model, electrons orbit the nucleus within specific energy levels.  These levels are defined by unique amounts of energy.  Electrons possessing the lowest energy are found in the levels closest to the nucleus.  Electrons with higher energy are located in progressively more distant energy levels.

If an electron absorbs sufficient energy to bridge the "gap" between energy levels, the electron may jump to a higher level.  Since this change results in a vacant lower orbital, this configuration is unstable.  The "excited" electron releases its newly acquired energy as it falls back to its initial or ground state.  Often, the excited electrons acquire sufficient energy to make several energy level transitions.  When these electrons return to the ground state, several distinct energy emissions occur.  The energy that the electrons absorb is often of a thermal or electrical nature, and the energy that an electron emits when returning to the ground state is called electromagnetic radiation.

In 1900, Max Planck studied visible emissions from hot glowing solids.  He proposed that light was emitted in "packets" of energy called quanta, and that the energy of each packet was proportional to the frequency of the light wave.  According to Einstein and Planck, the energy of the packet could be expressed as the product of the frequency (n) of emitted light and Planck's constant (h).  The equation is written as        E = hn

If white light passes through a prism or diffraction grating, its component wavelengths are bent at different angles.  This process produces a rainbow of distinct colors known as a continuous spectrum.  If, however, the light emitted from hot gases or energized ions is viewed in a similar manner, isolated bands of color are observed.  These bands form characteristic patterns - unique to each element.  They are known as bright line spectra or emission spectra.

By analyzing the emission spectrum of hydrogen gas, Bohr was able to calculate the energy content of the major electron levels.  Although the electron structure as suggested by his planetary model has been modified according to modern quantum theory, his description and analysis of spectral emission lines are still valid.

In addition to the fundamental role of spectroscopy played in the development of today's atomic model, this technique can also be used in the identification of elements.  Since the atoms of each element contain unique arrangements of electrons, emission lines can be used as spectral fingerprints.  Even without a spectroscope, this type of identification is possible since the major spectral lines will alter the color.

ELECTRON ARRANGEMENT

GENERAL LOCATION -------------------- ____________________

|

V

____________________

|

V

____________________

|

V

SPECIFIC LOCATION -------------------- ____________________

ENERGY LEVELS -

- numbered from closest to farthest away from nucleus

nucleus .

E 1 E 2 E 3 E 4

Electrons will occupy the location with the ________________________ amount of energy.

SUBLEVELS -

- designated by letters

- number of sublevels in an energy level =

nucleus .

E 1 E 2 E 3 E 4

ORBITALS -

- number of orbitals in an energy level =

- ___ sublevel has 1 orbital; ___ sublevel has 3 orbitals;
___ sublevel has 5 orbitals; ___ sublevel has 7 orbitals

HOW MANY ELECTRONS CAN AN ORBITAL HOLD?

HOW MANY ELECTRONS CAN EACH SUBLEVEL HOLD?

HOW MANY ELECTRONS CAN AN ENERGY LEVEL HOLD?

AUFBAU PRINCIPLE –

What is the order in which the sublevels fill with electrons? Use the DIAGONAL RULE.

1s

2s 2p

3s 3p 3d

4s 4p 4d 4f

5s 5p 5d 5f

6s 6p 6d 6f

7s 7p

More Electron Arrangement!

Examples:

Se 1s² 2s² 2p⁶ 3s² 3p⁶ 4s² 3d¹⁰ 4p⁴

Sn 1s² 2s² 2p⁶ 3s² 3p⁶ 4s² 3d¹⁰ 4p⁶ 5s² 4d¹⁰ 5p²

Hg 1s² 2s² 2p⁶ 3s² 3p⁶ 4s² 3d¹⁰ 4p⁶ 5s² 4d¹⁰ 5p⁶ 6s² 4f¹⁴ 5d¹⁰

HOEL (Highest Occupied Energy Level): energy level furthest from the nucleus that contains at least one electron

How to determine this using electron configuration?

~ largest non-exponent number

Se 1s² 2s² 2p⁶ 3s² 3p⁶ 4s² 3d¹⁰ 4p⁴ HOEL = 4

Sn 1s² 2s² 2p⁶ 3s² 3p⁶ 4s² 3d¹⁰ 4p⁶ 5s² 4d¹⁰ 5p² HOEL = 5

Hg 1s² 2s² 2p⁶ 3s² 3p⁶ 4s² 3d¹⁰ 4p⁶ 5s² 4d¹⁰ 5p⁶ 6s² 4f¹⁴ 5d¹⁰ HOEL = 6

Valence Electrons: electrons in the HOEL

How to determine this using electron configuration?

~ add up exponents of terms in HOEL

Se 1s² 2s² 2p⁶ 3s² 3p⁶ 4s² 3d¹⁰ 4p⁴
HOEL = 4 Valence electrons = 2 + 4 = 6

Sn 1s² 2s² 2p⁶ 3s² 3p⁶ 4s² 3d¹⁰ 4p⁶ 5s² 4d¹⁰ 5p²
HOEL = 5 Valence electrons = 2 + 2 = 4

Hg 1s² 2s² 2p⁶ 3s² 3p⁶ 4s² 3d¹⁰ 4p⁶ 5s² 4d¹⁰ 5p⁶ 6s² 4f¹⁴ 5d¹⁰
HOEL = 6 Valence electrons = 2

Noble Gas Configuration: shortcut for electron configuration

How is it written?

~ [ symbol for noble has closest to element with lower atomic # ]

~ [after brackets] next number is the period that the element is located in

~ after that number, write “s”

~ continue electron configuration in diagonal rule order until appropriate # of
electrons is reached

*NOTE: ending of electron configuration and noble gas configuration should be the same*

Se 1s² 2s² 2p⁶ 3s² 3p⁶ 4s² 3d¹⁰ 4p⁴
[Ar] 4s² 3d¹⁰ 4p⁴

^{18 20 30 34}

Sn 1s² 2s² 2p⁶ 3s² 3p⁶ 4s² 3d¹⁰ 4p⁶ 5s² 4d¹⁰ 5p²

[Kr] 5s² 4d¹⁰ 5p²

^{36 38 48 50}

Hg 1s² 2s² 2p⁶ 3s² 3p⁶ 4s² 3d¹⁰ 4p⁶ 5s² 4d¹⁰ 5p⁶ 6s² 4f¹⁴ 5d¹⁰
[Xe] 6s² 4f¹⁴ 5d¹⁰

^{54 56 70 80}

Orbital Notation: drawing of how electrons are arranged in orbitals; will only need to do this for the HOEL

*NOTE: ___ = orbital h or i = electrons

Se 1s² 2s² 2p⁶ 3s² 3p⁶ 4s² 3d¹⁰ 4p⁴ hi hi h h
4s 4p

Sn 1s² 2s² 2p⁶ 3s² 3p⁶ 4s² 3d¹⁰ 4p⁶ 5s² 4d¹⁰ 5p² hi h h __
5s 5p

Hg 1s² 2s² 2p⁶ 3s² 3p⁶ 4s² 3d¹⁰ 4p⁶ 5s² 4d¹⁰ 5p⁶ 6s² 4f¹⁴ 5d¹⁰ hi
6s

Dot Diagrams: symbol represents nucleus and non-valence (“inner-shell”) electrons; dots around symbol represent valence electrons

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