Chemistry I
Ch. 11 Modern Atomic Theory Notes
Ch. 11-1 Atoms and Energy
I.
Electromagnetic radiation – energy that is transmitted through empty space as a wave.
a. Examples – light, heat, x-rays, microwaves
II.
Characterized by wavelength, frequency and speed
a. Wavelength – (λ called “lambda”) - distance between two peaks or troughs
i. Since it is a distance, it is measured in meters (m)
b. Frequency – (ν called “nu”) – tells the number of peaks/troughs that pass in a
given time
i. Measures waves in one second of time – waves/sec OR
1
𝑠𝑒𝑐
OR s-1
ii. 1 wave/sec = 1 hertz (Hz)
c. Speed = wavelengths x frequency OR
d. Ex:
𝑚
𝑠
=
𝑚
1
×
1
𝑠
6 m/s = 6 m x 1 Hz
6 m/s = __3________ m
x
__2________ Hz
e. Inverse relationship - When wavelength decreases, frequency increases (as long
as speed is constant)
III.
Diagram of the electromagnetic spectrum
Highest energy
Lowest energy
Highest frequency
Lowest frequency
Shortest wavelength
Longest wavelength
Dual Nature of Light – is it a wave or a particle?
IV.
a. Has both characteristics
b. Photon – bundle of light
c. Photoelectric effect - http://phet.colorado.edu/en/simulation/photoelectric
d. First 11 elements
http://web.visionlearning.com/custom/chemistry/animations/CHE1.3-ananimations.shtml
11-2
The Hydrogen Atom
I.
We see light emitted when atoms receive energy, and electrons move to an outer
shell (excited) then fall to their original position (ground state)
a. Emission spectrum – bright line spectrum – energy released by objects
i. Unique to each element
ii. Absorption spectrum – energy absorbed by objects
b. Quantized energy – only certain wavelengths of light are found to be emitted
i. Shows that electrons rise and fall from specific area (energy levels)
II.
Model of Bohr Hydrogen
http://www.visionlearning.com/librry/flash_viewer.php?oid=1347&mid=51
III.
Wave Mechanical Model of the aAtom
a. DeBroglie – devised a way to calculate the energy of a photon, and therefore the
energy of the electron that emitted it.
b. Schrodinger – came up with mathematical equations to describe the space around
a nucleus that electrons would occupy.
i. Orbitals – space around a nucleus most likely to contain an electron
(probablility map)
11 – 3
Atomic Orbitals
I.
Hydrogen Energy Levels
a. Principle Energy Levels (n=1, n=2, n=3…n=7)
b. Sublevels (found within each principle energy level)
i. 4 types - each contains certain orbitals
1. s
2. p
3. d
4. f
c. s sublevel – 1 orbital that is spherical
d. p sublevel – 3 orbitals that are dumbbell shaped in each of x,y,z axes
e. d sublevel – 5 orbitals
f. f sublevel – 7 orbitals
Example of the Aufbau Principle – note this pattern continues for energy levels 6 and 7
II.
Aufbau Principle - Electrons occupy orbitals that require the least amount of energy
(like water takes the path of least resistance)
III.
Pauli Exclusion principle –
a. Each orbital can hold ONLY 2 electrons.
b. The electrons have opposite spin
IV.
Hund’s Rule – each orbital of equal energy is occupied with one electron until all are
filled, then 2nd electron (with opposite spin) may enter the orbital
11-3
I.
Electron Configuration and Atomic Properties
Electron Configuration – description of the arrangement of electrons around the
nucleus of an atom.
II.
Orbital Diagram – box/lines that shows each pair of electrons in their corresponding
orbitals
III.
Electron Configurations correspond to the Periodic Table
a. Valence electron – final electron to be placed in around an atom
b. Period – horizontal row – corresponds to the principle energy level of the valence
electron
c. Group – vertical column (Family)
i. s block – valence electrons placed in the s orbital (2 boxes wide)
1. Groups 1A and 2A
ii. p block – valence electrons placed in the p orbital (6 boxes wide)
1. Groups 3A – 8A
iii. d block – final electrons placed in the d orbital (10 boxes wide)
1. Group B (3-12) elements (sunken down)
iv. f block – final electrons placed in the f orbital (14 boxes wide)
1. (below the main block)
IV.
Atomic Properties and the Periodic Table
a. Position of the valence electron(s) determines the chemical behavior of the
atom/element
b. Atomic theory is an attempt to understand (and therefore predict) why
elements/atoms behave the way they do.
c. Stairstep – separates metals and nonmetals
i. Metals – Left side – tend to lose electrons (form cations +)
ii. Nonmetals – Right side – tend to gain electrons (form anions -- )
iii. Metalloids – found on each side of the stairstep
d. Atomic size – DECREASES as one moves LEFT to RIGHT
i. Number of protons (+ charge) increases
ii. Number of energy levels does not change
iii. Electrons are pulled in tighter
e. Atomic size – INCREASES as one move DOWN a group
i. Number of energy levels increases
ii. Shielding increases (inner electrons prevent nucleus from pulling on outer
electrons)
f. Ionization energy – energy required to REMOVE an electron
i. Trends are the opposite as for Atomic size – for the same reasons
g. Electronegativity – tendency for an atom to GAIN an electron
i. Trends are the same as for ionization energy
ii. Fluorine is the most electronegative element on the periodic table
iii. Francium is the least electronegative element on the periodic table.