Atomic theory

orbital shapes

s - spherical

p - perpendicular dumbbells

d - 4 leaf clover

number of electrons in each orbital

s - 1 orbital, 2 electrons
p - 3 orbitals, 6 electrons
d - 5 orbitals, 10 electrons
f - 7 orbitals, 14 electrons

aufbau principle

• fill the lower energy first, the ones closest to the nucleus

• energy sublevel must be filled before moving on to the next higher level

london dispersion forces

• weakest intermolecular force

• is between all types of molecules

• temporary dipoles

dipole-dipole forces

• polar molecules only

• electrostatic attraction

ion-dipole forces

• attraction between an ion and a polar molecule (water)

• ex. NaCl breaks up because the ion-dipole attraction is stronger than the attraction of Na+ to Cl-

• forms an electrolyte solution

hydrogen bonding

• between hydrogen and oxygen, nitrogen, or fluoride (highly electronegative)

• hydrogen bonding is the strongest form of intermolecular bonding however it is not as strong as normal covalent bonds

• more lone pairs = more stronger

boiling point - how does it work ?

• in order for a liquid to turn into a gas, the intermolecular forces must be broken

• the stronger these forces are, the higher the boiling point is

ionic solid properties

• high melting/boiling points: strong electrostatic forces

• hard but brittle: rigid lattice structure

• conduct electricity

• solubility: Usually dissolve in polar solvents like water

Bohr’s Postulates

1. Electron moves in a circular orbit around the nucleus

2. Electron has a constant energy

3. Electrons move from higher-lower energy levels by absorbing or emitting energy as photons

Ground state

Electrons in their original energy level

Quantum Model of Atom

• Electrons have wave-like properties (move as circular standing waves)

• Location of electron is based on probability and uncertainty

• Electrons have sublevels and energy levels depending on energy

• Electrons move by absorbing or emitting energy as photons

Principal Quantum Number (n)

Describes the number of energy levels / energy level we’re in

• n = {1,2,3…}

Secondary Quantum Number (l)

Describes the number of sublevels inside main energy level

• {0,1,2,3…n-1} (always one less than principal number)

• Value also corresponds to subshell type {0=s,1=p,2=d,3=f)

Magnetic Quantum Number (m_l)

Describes number of orbitals in sublevels

• {-l…+l}

• e.g: n = 2, l = 1, m_l = -1,0,1, #orbitals:3

• Mas number of e^-: multiply number or orbitals by 2 (recall: max is 2n²)

Spin Quantum Number (m_s)

Describes the orientation of the electron in said orbital

• Two spin states represented as +1/2 (up) and -1/2 (down)

Hund’s Rule

Orbitals in the same energy level/sublevel have the lowest amount of energy configuration for an atom in the one with the maximum number of unpaired electrons allowed by the Pauli Exclusion principle

Pauli Exclusion Principle

States that no two electrons in an atom can have the same set of four quantum numbers

Electron Configuration

Describes the location and number of electrons in the energy levels of atoms

Spherical (s) Block on periodic table

Groups 1&2 Including Helium

Dumb-bell (p) Blocks on Periodic Table

Groups 13-18

Diffuse (d) Block on Periodic Table

Groups 3-12 (transition Metals)

Fundamental (f) Blocks on Periodic Table

Radioactive Elements

Metallic Bonds

Electrostatic attraction between closely packed cations

• Greater delocalized electrons: stronger bond and higher melting point

Polar Molecule

• Net dipole/molecule dipole (movement of electrons in one direction)

Net Dipole Moment

Movement of electrons in one direction

Non-polar molecules

• Non polar bonds

• dipole sum of zero (all dipoles cancel out)

Sigma Bond

• Covalent bond

• end to end overlap

• strong relative strength

• no freedom to rotate

Pi Bond

• Covalent bond

• Side -to - side overlap

• weak relative strength

• ability to rotate freely

Valence Bond Theory

Explains how atoms form bonds by overlapping their outermost electron orbitals:

• Orbital Overlap: Atoms form bonds when their atomic orbitals overlap.

• Electron Pairing: A bond forms when two electrons with opposite spins occupy the overlapped orbitals.