Chemistry I Ch. 11 Modern Atomic Theory Notes Ch. 11

advertisement

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.

II.

a.

Examples – light, heat, x-rays, microwaves

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.

1

Measures waves in one second of time – waves/sec OR 𝑠𝑒𝑐

OR s

-1 ii.

1 wave/sec = 1 hertz (Hz) c.

Speed = wavelengths x frequency OR 𝑚 𝑠

=

𝑚

1

×

1 𝑠

III.

d.

Ex: 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)

Diagram of the electromagnetic spectrum

Highest energy Lowest energy

Highest frequency Lowest frequency

Shortest wavelength Longest wavelength

IV.

Dual Nature of Light – is it a wave or a particle? 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

I.

The Hydrogen Atom

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.

III.

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)

Model of Bohr Hydrogen http://www.visionlearning.com/librry/flash_viewer.php?oid=1347&mid=51

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.

IV.

Pauli Exclusion principle – a.

Each orbital can hold ONLY 2 electrons. b.

The electrons have opposite spin

Hund’s Rule – each orbital of equal energy is occupied with one electron until all are filled, then 2 nd

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

II.

nucleus of an atom.

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.

Download