Chapter 7
7.1 - EM Radiation
7.2 - Wave VS. Particle
7.3 - The Bohr Model
7.4 and 7.5 - Quantum Numbers
7.6 - Electron Configurations
7.7 - Periodic Trends
The electromagnetic spectrum.
Waves
Types of Waves
Frequency is how often a wave repeats. Hertz
Wavelength is the distance between troughs or peaks of a wave. Meters
Microwave
Infrared
Visible
Ultraviolet
Radio
X-ray
ROYGBIV
Gamma
Is light a wave or a particle?
Our modern conclusion: It can be both!
Scientists
Max Planck
Isaac Newton
Thomas Young
Albert Einstein
Louis deBroglie
Travel at speed of light (2.998e8 m/s)
Only observations. Source: trust me bro
Light is a particle! Different colors are different sized particles
Light is a wave!
The Double Slit Experiment
Interference pattern is wave behavior, therefore light is a wave
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
E = mc²
Light is a particle!
Light is also quantized. One quantum of light = photon
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
Elements can be identified by their emission spectra.
Electrons absorb energy to rise in energy level and lose energy when falling in energy level
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
n2 -> n1 drop releases energy outside the visible light spectrum.
Only specific numbers are allowed, because an electron can't be between levels. Energy is quantized
Angstrom conversion = 10e-10
Nanometer conversion = 10e9
Big Ideas
Quantum Numbers
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
Four quantum numbers make up a set
Principal quantum # (n)
Angular momentum quantum # (l)
Magnetic quantum # (ml)
Spin quantum # (ms)
Tells us size and energy of orbital was well as row on periodic table.
Ex. 1, 2, 3, 4...
0, 1, 2, 3. s, p, d, f.
Maximum angular momentum depends on n. n-1 is the l max.
s is a circle shape. p looks like an hourglass.
Ex. n = 1 only has 1s.
n = 2 has 1s 2s 2p
Tells you shape of orbitals.
Tells you the number of orbitals and the orientation or the orbitals.
Anything beyond f is theoretical but possible. Next would be g
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.
Can be +1/2 or -1/2. The direction an electron is spinning in an orbital.
Two electrons per orbital
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
Trends that you can discern from the periodic table
Electronegativity - How much an element pulls electrons towards itself in a covalent bond
Atomic Radius - Half the distance between two identical nuclei
Ionization Energy - The energy required to remove an electron
Electron Affinity - Energy given off when a gaseous element gains an electron
Most Electronegative - Top right (not including noble gases), so Fluorine
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
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
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