Today in Chemistry...
Chapter 5- Electrons in Atoms
I.) Light and Quantitized Energy:
- The light emitted from a substance is based on the electron configuration of the substance. (Elements)
- When electrons absorb energy it will jump to the next highest energy level.
- As soon as it jumps up, it will release a photon (light particle) and drop back down to its original level. (its ground state)
- The photon that is released by the electron will contain the exact same amount energy as it absorbed.
- Each element will emit its own color or series of colors when its electrons become exated
- Electromagnetic Radiation- (wave features) all travel in a wave-like manner
-does NOT require matter to travel from one point to another (Ex.) Light, Radio, TV, Microwaves, X-rays...
a.) Standing Wave (Transverse Wave)- Example at Bottom of Page
Frequency- (V-nu) the number of waves that pass through a given point in one second
-measured in Hertz- Hz ( 1 H2=1/s )
*Inverse-When wave length increases the frequency decreases and energy increase (vice versa)
ROY G BIV- white light is made up of all colors of the rainbow
-wavelength
-frequency
C- speed of light
- speed of light in a vaccum 3.0 x 10 to the 8th power m/s (constant)
(Ex.) A wave has a fequency of 4.0 x qo to the 12th power H 2. what is the wavelength?
wavelength= speed of light/ frequency
wavelength= 3.0 x 10 to the 8th power m/s /4.0 x 10 to the 12th 1/s= 7.5 x 10 to the -5th power m
II.) Particle Nature of Light- explains why different substances emit their own unique spectra of colors when heated.
-also explains why when a cetain frequency of light shines on metals, the metal will emit a unique frequency of light.
- Quantum Concept- matter can emit a certain quanta (amount) of enregy
-matter can only absorb a certain quanta of energy
*both are based on the element and its electron configuration.
- Quantum- the minimum amount of energy that a substance can gain or lose.
A.) Planck- studied the light emitted from heated matter
-Equantum =hv ( h= Planck's Constant)
*6.626 x 10 to the -34th power J/s (Joule)
* The energy (radiation) of light increased as its frequency increased.
* The energy can only be absorbed or emitted in small whole numbers.
-----------ROY G BIV---------------> Red-lowest frequency
Violet- higher frequency
(visual spectrum)
B.) Photoelectric Effect- when a certain frequency of light is absorbed by a metal it will emit a certain frequency of light.
-the frequency of light shone on the metal must be equal to or greater than the frequency emitted.
- Albert Einstein- 1st to explain the "wave=particle duality" of light
*Photon- term he used to explain the small packets of energy (light)
Ephoton=hv (Planck's constant reworded)
2. Atomic Emission Spectra- the series of frequencies given off by an element when its electrons are excited.
- Each element has its own unique series of colors (frequencies)
III.) Quantum Theory- (as applied to the atom)
- Bohr- 1st to propose that electrons had specific energy levels
- Ground State- the lowest energy levels that the electrons of an atom would be in
- The smaller the energy level, the lower the energy state of those electrons
- The Closer to the nucleus, the lower the energy
- The larger energy levels (further from the nucleus) would require a greater amount of energy
- Bohr gave each energy level a numerical value 1,2,3...7
2.) Hydrogen line Spectra-
- Hydrogen has only one electron in the lowest energy level
- When energy was added, the electron moved up to the next highest frequency level
- Immediately the elctron gave off a photon corresponding to the difference in energy betweeen the 2 energy levels (1 and 2)
E=Ehighest=Elowest=Ephoton=hv
- Bohr's model of the atom as explained, only works for Hydrogen but all other elements are explained using this model
-Balmer Series- visible light (e- drops to n=2)
-Lyman Series- Ultravilent light (e- drops to n=1)
-Pachen Series- infrared light (e- drops to n=3)
3.) The Quantum Mechanical Model-
- Bohrs model did not explain the electron in terms of a wave (only a particle)
- De Broglie- stated that if Bohrs calculations were true (they were) that the electrons, if in a fixed orbital, would have to have a fixed frequency, wavelength, and energy associated with them.
Said electrons werent in fixed circular orbitals (frequency, wavelength and energy can all vary)
4.) Heisenburg Uncertainty Principle-
- You can NOT know the precise location or velocity of an electron at the same time
- You must use photons to detect an electron
- When an electron is struck by a photon the frequency, velocity, and energy it has changes
*De Broglie wave equation:
wavelength= h
mv <----velocity NOT frequency
5.) Schrodinger-
- wave equation explains that electrons have the ability to be in more than one location at any given time
- places electrons around the nucleus based on probability
- Quantum Mechanical Model- does NOT place electrons in fixed orbitals like Bohr had proposed
IV.) Modern Concept of the Atom-
- Electrons are positioned in orbitals based on the 90% probability of their location
- Electron Cloud Model- used this 90% probability to place electrons around the nucleus
1. Principle Energy Levels- 7 total levels (n= 1,2,3,7)
Hydrogen with 1 electron has a principle energy level of n=1
2. Energy Sublevels- 4 types of energy sublevels (s, p, d, f)
- The sublevels have specific shapes
S-spherical p-dumbbell d & f (have multi shapes)
V.) Electron Configurations-
- Arranges electrons around the nucleus based on the atoms ground state
- The ground state is the atom's most stable state.
1. Aufbau Principle- each element is placed in its ground state position
Rules:
- All sublevels within a primary level are of equal energy
- Sublevels increase in order of s, p, d, f
- In a multi-electron atom, the sublevels of a primary level do NOT equal energy values
2.) Hunds Rule- electrons pair up in their specific orbitals
Each electron has the opposite spin of the other in its sublevel
Each sublevel will fill with "up" spin electrons 1st and then fill with "down" spin electrons
^ ("up" spin) v ("down" spin)
- C (carbon) ^ v ^ v ^ ^ __
- 1s 2s 2p
- Ne (neon) ^ v ^ v ^ v ^ v ^ v (Noble gas)
- 1s 2s 2p
3.) Pauli Exclusion Principle- the suborbitals of an atom can hold 2 electrons, these electrons will always have opposite spins
4.) Electron orbital Configuration-
- 2 2 2
- C 1s 2s 2p (carbon)
- 2 2 6 2 3
- P 1s 2s 2p 3s 3p
Diagrams- all will display the atom in its ground state
1.) Orbital Diagram- ^ ("up") v ("down")
- O (Oxygen) ^ v ^ v ^v ^ ^
- 1s 2s 2p
-
- Si (Silicon) ^ v ^ v ^ v ^ v ^ v ^ v ^ ^ __
- 1s 2s 2p 3s 3p
2.) Electron Configuration Notation-
- 2 2 6 2 1
- Al (Aluminum) 1s 2s 2p 3s 3p
- 2 2 6 2 6 2 10 5
- Br (Bromine) 1s 2s 2p 3s 3p 4s 3d 4p
3.) Noble Gas Notation-
- 1
- Na (sodium) [Ne] 3s
- 2 10 5
- Br (Bromine) [Ar] 4s 3d 4p
- 1 5
- Cr (Chromium) [Ar] 4s 3d d & f suborbitals want to
- be filled or 1/2 filled if
- 1 10 possible (move only 1 e-)
- Cu (Copper) [Ar] 4s 3d
- 1 5
- Mo (Molybdenum) [Kr] 5s 4d
***Valence Electrons-
- the electrons in the outter most orbital of the atom
- these are the highest energy
- these give the atom its chemical properties
- these electrons do the chemical bonding between atoms
Lewis Dot Diagrams-
- use the chemical symbol and dots to represent the atom and its electrons
Octet Rule:
- all elements (atoms) want 8 valence electrons
- Except H (Hydrogen) and He (Helium), they want 2
- This will allow them to have a Noble gas electron configuration
- This makes the atom stable
(Examples:)
- .
- (Sodium) Na . (Boron) :B
Bohr Diagram:
- _______
- / \ \ \ \
- / p+11 \ \ \ \
- ! Na ! 2e- ! 8e- ! 1e- !
- ! n 12 ! ! ! !
- \ / / / /
- \ ________/ / / /
