Structure and Properties Flashcards
Aufbau Principle
Orbitals are filled in order of increasing energy
Hund’s Rule
Orbitals of a particular sublevel must be half filled with 1e- before the e- are paired
Pauli Exclusion Principle
Each orbital can hold a maximum of 2 e- which must be spinning in the opposite direction
Orbital Filling Chart
Used to create energy level diagrams
How to draw energy level diagrams
Determine the #of e- that need to be placed, follow aufbau principle
Energy level diagrams for Anions
ion formed by the addition of e- to the p sublevel until it is full.
Energy level diagrams for Cations
A positive ion that is formed by removing e- from s, p, or d sublevels. Removed from the highest principle quantum # and the highest energy level first
Electron Configuration
A short hand notation that shows the arrangement of e- in each of the occupied sublevels of an atom or ion. (a simplified way of describing the e-level diagram)
Orbital Filling Chart and the Periodic Table
1s, 2s, 2p, 3s, 3p, 4s, 3d 4p, 5s, 4d, 5p, 6s, 5d, 6p, 7s, 6d, La+Ac - 4f + 5f
Core Electron Comfigurations
Also known as shorthand e- configurations. Places the noble gas in square brackets at the beginning of the e- configuration (show orbitals when bonding)
Anomalous Electron Configurations
For d sublevel only: If possible, e- are promoted from a full s orbital to the d level
Lewis Symbols of Atoms
Visual representations of the valence e- in s and p orbitals. ex: O - 2s^2(2p^4), 2+4 = 6e-
Atomic Spectra Phenomenon
Sir Issac Newton: lines imply that electronic energy is quantized
Bohr and Hydrogen
Succeeded in calculating H+ wavelengths but did not work for larger atoms so original idea of fixed orbit and distance was abandoned
Louis de Broglie
1923 - particles are transported by a wave
Werner Heisenburg
1927 - you cannot know the position and velocity of a particle
Erwin Schrodinger
1926 - developed equations that describe motion of an e-
1st Quantum #: Principle Quantum # (n)
e- are located in energy levels or shells and are identified by a PQ#. n can be any value from 1 to infinity and describes the orbital size and energy. the greater the n, the further an e- from the nucleus
2nd Quantum #: Sublevels (l)
each n contains sublevels of slightly differing energies, identified by a letter -s, p, d or f. (s-sharp, p-principle, d-diffuse, f-fundamental). L = 0-(n-1)
3rd Quantum #: Orbitals [m(l)]
each sublevel contains a certain # of orbitals, each holding max of 2e-.
4th Q#: Spin
each e- has opposite spins in one orbital
What is an Orbital
A 3D region of space aound the nucleus of an atom where the probability of finding an e- is high.
Magnetic Quantum # m (l)
Each sublevel contains a certain # of orbitals with a specific orientation m (l)= -l -> +l
Types of Bonds
Ionic: EN=1.7-3.3
Covalent: EN=0-0.2
Non-polar: EN=0.3-1.6
Metallic: sea of electrons
VSEPR
2-1, 3-1, 2-2, 4-1, 3-2, 5-1, 4-2, 3-3
2-1: Bent. 3-1: trigonal pyramidal. 2-2: Bent. 4-1: seesaw. 3-2: Tshaped. 5-1: square-based pyramidal. 4-2: square planar. 3-3: Tshaped.
Valence Bond Theory
Orbits overlap to form new orbital w/ opposite spin e-.
Sigma Bond
Direct end to end overlap of s or p orbitals. Single covalent bond is always sigma bond.
Pi Bond
Two parallel orbitals share a pair of e-. (Double covalent is shorter and stronger than single covalent)
How many sigma and pi?
Single bond
Double bond
Triple bond
Single: 1sigma
Double: 1sigma,1pi
Triple: 1sigma,2pi
Polar covalent bonds
E- pair is shared unequally, most closer to atom w/ higher EN, formation of dipoles
Free Radicals (Paramagnetic)
contain a single unpaired e- in the structure
Resonance Structures
Two or more Lewis structures that show the same relative position of atoms but different position of double/triple bonds
Coordinate covalent bonds
a bond in which the shared pair of e- is contributed by one atom and an atom/ion that requires 2e- to complete it’s s or p sublevel
Hybrid Orbitals
when an e- from a full orbital is promoted to an empty orbital of higher energy in same PQ# occurs in the bonding process
