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Chapter 7 - Coggle Diagram
Chapter 7
7.3 - The Bohr Model
Elements can be identified by their emission spectra.
Electrons
absorb
energy to rise in energy level and
lose
energy when falling in energy level
n2 -> n1 drop releases energy outside the visible light spectrum.
Electron levels get closer as they get further from the nucleus.
Calculating the energy of a level / level change
E = -2.178 * 10e-18 (z²/n²)
E = -2.178 * 10e-18 ((z²/nf²)-(z²/ni²))
ni = initial energy lvl, nf = final energy lvl. z = # of protons
Only specific numbers are allowed, because an electron can't be between levels.
Energy is quantized
7.4 and 7.5 - Quantum Numbers
Big Ideas
Heiseinberg Uncertainty Principle
- You cannot know both position and momentum of an electron at the same time
Schrodinger treated electrons like waves to predict the location of electrons.
Orbital
- An amount of space dedicated to an electorn. 90% chance of being there
Quantum Numbers
Four quantum numbers
make up a set
Principal quantum #
(n)
Tells us size and energy of orbital was well as row on periodic table.
Ex. 1, 2, 3, 4...
Angular momentum quantum #
(l)
0, 1, 2, 3. s, p, d, f.
s is a circle shape. p looks like an hourglass.
Anything beyond f is theoretical but possible. Next would be g
Maximum angular momentum depends on n. n-1 is the l max.
Ex. n = 1 only has 1s.
n = 2 has 1s 2s 2p
Tells you shape of orbitals.
Magnetic quantum #
(ml)
Tells you the number of orbitals and the orientation or the orbitals.
Can range from -l to l. Count the number of numbers and that's the number of orbitals.
Ex. 2p. p is 1. -1 to 1 has 3 numbers. 3 orbitals.
Spin quantum #
(ms)
Can be +1/2 or -1/2. The direction an electron is spinning in an orbital.
Two electrons per orbital
7.6 - Electron Configurations
Big Ideas
Electron shielding is the reason that the orbitals are filled up in a wonky order
The Aufbau Principle
- Comes from the german word "to build". Electrons are built up in orbitals from low to high energy.
This is where we get the periodic table orbital filling trend
Hund's Rule
- Electrons want their own orbital before they pair up. First fill every hole before you start putting two in each.
Pauli Exclusion Principle
- Two electrons cannot share the same quantum number. AKA. spins must be different
Shorthand Notation
The last noble gas in brackets, and then everything that comes after that
7.1 - EM Radiation
The electromagnetic spectrum.
Waves
Types of Waves
Microwave
Infrared
Visible
ROYGBIV
Ultraviolet
Radio
X-ray
Gamma
Frequency is how often a wave repeats. Hertz
Wavelength is the distance between troughs or peaks of a wave. Meters
Angstrom conversion = 10e-10
Nanometer conversion = 10e9
Travel at speed of light (2.998e8 m/s)
7.2 - Wave VS. Particle
Is light a wave or a particle?
Our modern conclusion: It can be both!
Scientists
Max Planck
Studied energy emitted from solids when heated a LOT. Kept cutting it in half until he found the smallest transferrable amount of E.
Discovered the quantum.
Energy is quantized
Light is a particle!
E = hv
Isaac Newton
Only observations. Source: trust me bro
Light is a particle!
Different colors are different sized particles
Thomas Young
Light is a wave!
The Double Slit Experiment
Interference pattern is wave behavior, therefore light is a wave
Albert Einstein
E = mc²
Light is a particle!
Light is also quantized. One quantum of light =
photon
Louis deBroglie
SCREW YOU MATTER IS A WAVE
Invented deBroglie formula to find the wavelength of ANYTHING.
THE THREE FORMULAS
E = hv
λ = h/mv
Nu (v) = c/λ
7.7 - Periodic Trends
Trends that you can discern from the periodic table
Electronegativity -
How much an element pulls electrons towards itself in a covalent bond
Most Electronegative - Top right (not including noble gases), so Fluorine
Atomic Radius -
Half the distance between two identical nuclei
Goes
up
when going down a group because the number of electron levels goes up.
Goes
down
when going across a period because stronger nuclear charge - The electrons pull the protons in closer
Largest Atomic Radius - Bottom left, so Francium
Ionization Energy
- The energy required to remove an electron
Goes
up
when going across a period because as the radius gets smaller the pull of the electrons on the protons is stronger and its harder to pull them away,
Goes
down
when going down a period because the outermost electron is held further from the protons, and so is easier to remove
Highest ionization energy - Top right, so Helium
Electron Affinity
- Energy given off when a gaseous element gains an electron
Goes
up
across a period because electrons are pulled closer to the nucleus when there are more protons
Goes
down
when going down a period due to the increase in atomic radius
Highest electron affinity - Top right (not including noble gases), so Fluorine
Differences left and right are smaller than up and down