• The structure of the periodic table corresponds directly to the fundamental structure of the atoms of the elements that comprise it. Specifically, it is related to the way that electrons orbit around the nuclei of these atoms.
• Electrons are arranged into shells around the atom and each shell is subdivided into orbitals. Each orbital holds two electrons.
• The most basic, the "s" orbital, is a simple sphere. There is only one type of s orbital and only one per electron shell. It is the only orbital in the first shell.
• The second type of orbital, the "p" orbital, is found in shells two and up. There are three types of p orbital, and these orbitals are only filled after the s orbital in the same shell has been filled.
• The third type, the "d" orbital is found in the third shell and up, but only begins to be filled in each shell when the next shell's s orbital has already been filled. There are five d orbitals.
• The fourth type of orbital is the "f" orbital. It only is found in shells four and higher. It is only filled after the s orbital of the shell two ahead of it is full. There are seven f orbitals.
• The first two elements, Hydrogen and Helium, are the physical manifestation of the filling of the first electron shell. This shell contains only the s orbital, which contains two electrons like all orbitals. Coincidentally, this first period contains TWO elements.
• Each period contains the same number of elements as its corresponding shell contains electrons. It is not as simple as that, of course. The elements that correspond to the d3 orbitals (the d orbitals in shell three) is included in period four because the spots in the d3 orbital are filled only after the s4 orbital has been filled. This same "one behind" pattern applies to all d orbitals, and a similar "two behind" pattern applies to the f orbitals.
• It is this organization of electrons into orbitals that gives elements their chemical properties. It is the structure behind everything from electronegativity to reactivity to oxidation states. Each of these things is the manifestation of elements trying to fill their outermost electron shell completely.

-the s orbital is the first one to be filled and contains two electrons

-the p orbitals are the second and contain six electrons

-the d orbitals are the third and contain ten electrons

-the f orbitals are the fourth and contain fourteen electrons

QUANTUM NUMBERS
n=orbital size and energy
l=orbital shape
m(l)=which orbital
m(s)=electron spin direction

Z*, Effective nuclear charge

• Z* is the nuclear charge experienced by the outermost electrons
• Explains why E(2s) < E(2p)
• Z* increases across a period owing to incomplete shielding by inner electrons.

Ionization Energy

• energy required to remove an electron from an atom in the gas phase

Flame Test
The following metals give emit certain colors of light when their atoms are excited.

Metals	Color
Lithium (Li)	Pink/Red
Potassium (K)	Purple
Copper (Cu)	Green
Barium (Ba)	Yellow/Orange
Calcium (Ca)	Pink
Strontium (Sr)	(Charmander) Red
Sodium (Na)	Yellow

Heisenberg Uncertainty Principle

• It is impossible to simultaneously define the position and momentum of an electron.

Superposition Principle

• It is possible for something to be in multiple states at once
• *You don't know whether the cat is alive or dead, so it's both.
  

Pauli Exclusion Principle

• No two electrons can have the same set of quantum numbers

-That is, each electron has a unique address.
Hund's Rule

• Electrons are most stable when they are in their own orbital, unpaired.
  

EINSTEIN AND THE PHOTOELECTRIC EFFECT:

• Definition: The ejection of electrons when light strikes the surface of a metal
• How it works:

1. Electrons will be ejected only if the frequency of the light is high enough
2. If the frequency of the light is not high enough, then increasing the intensity of the light will not have any effect
3. However, if the frequency is high enough and you increase the intensity of the light, then more electrons will be ejected
4. If you increase the frequency of the light, the velocity at which the electron ejects will increase.

ATOMIC SPECTRA (for the hydrogen atom):

• Lyman Series: arises from electrons moving from states with n>1 to the n=1
• Balmer Series (visible light): arises from electrons moving from states with n>2 to n=2

EQUATIONS:

electromagnetic radiation: c= λf, where c is speed of light (3.0 x 10^8 m/s) , λ is wavelength and f is frequency

Planck's Equation: E=hf , where E is energy, h is Planck's constant (6.626 x 10^-34 J s) and f is frequency

Rydberg's Equation: 1/λ = R(1/nf2 - 1/ni2), where λ is wavelength, R is Rydberg's constant (1.0974 x 10^7 m^-1) and n is the electron orbital

Potential energy of electron in the nth level= En= -Rhc/(n^2) : R is Rydberg constant, H is Planck's constant, and c is the speed of light.

de Broglie's Equation: λ = h/ mv, where λ is wavelength, h is Plank's constant, m is mass and v is velocity

CHAPTER 8

CHEMICAL PERIODICITY:

• Atomic Size: Generally decreases across a period and increases down a group.

-This is because Z* (effective nuclear charge) increases across a period so proton-electron attraction increases across a period, thus decreasing the size.
- Down a group, electron shielding increases and there is more electron-electron repulsion so the electrons in the outer shell are further away from the nucleus so atomic size increases.

• Ionization Energy: Generally increases across a period and decreases down a group.
• Electron Affinity: Generally increases across a period and decreases down a group.
• Electronegativity: Generally increases across a period and decreases down a group.
• Melting & Boiling Points: Generally transition metals have high melting points and noble gases have extremely low melting points; generally melting point increases down a group except for alkali metals, alkali-earth metals, and elements in group 12 (the last group of the transition metals).