Chem 101 Test 2

Heisenberg Uncertainty Principle

It’s impossible to know the exact position and momentum(mass x velocity) of a particle, at the same time

formal charge

(# valence e-) - (# lone pair e-) - 0.5(bonding e-)

bond order

(# of bonds around central atom) / (# bonding sites)

Quantum mechanics

Field of study that includes quantization of energy, wave-particle duality, and the Heisenberg uncertainty principle to describe matter.

Shortest to largest bond

triple, double, single

bond order increases —> bond length decreases —>

bond strength increases

Energy levels with same n→

degenerate energy levels

Degeneracy

the energies of all the orbitals with the same n are the same

Pauli Exclusion Principle

no two e- can have the same set of quantum numbers(no more than 2 e- can occupy the same orbital

-larger n → larger orbital size

-larger orbital size→ more distance→ weaker e- attraction to nucleus→ higher orbital energy

• atomic number - # of core electrons

noble gases

• High IE and low EA

Nonmetals

High EA and gain e- easily

Metals

have low IE and lose e- easily

ionic compounds

• higher melting and boiling points; good conductors of electricity when dissolved in water; has higher stability than constituent atoms

Aufabu Principle

e- will fill orbitals in order of increasing energy (think about the electron configuration things like 1s^2, 2s^2, etc.)

Hund’s Rule

 e- fill lowest e- orbitals first (thing about the arrow drawings for e- config.)

play most important role in chemical reactions

valence e-

Cation

increased Zeff

Anion

Decreased Zeff

increase down a group

atomic size and covalent radius

decrease down a group

Ionization Energy, electron affinity, and electronegativity

decrease left—> right

atomic size and covalent radius

increase left—> right

I.E., electron affinity, and electronegativity

more arrows in orbitals

greater first I.E.

negative Electron affinity

exothermic

electron affinity

change in energy for the process of adding an electron to a gaseous atom to form an anion

greater Zeff

easier to add e-

P.E. of two atoms decreases

the closer they get

nonpolar covalent

equal e- sharing

pure covalent

-atoms are identical (ex: diatomic ones)

more electronegativity difference

greater bond polarity and bond dipole

same atoms

nonpolar bond

nonpolar bond

change in electronegativity is 0mor

e polar

higher electronegativity difference

Most stable structure has

most negative formal charge on most electronegative atom

greater bond (ex: single vs. double bond)

smaller bond length and stronger bond

expanded octet(can have more than 8 e- when drawing the bond out)

possible for elements in 3rd and higher periods

Axial

there are two axial positions that are 180° apart

Equatorial

there are three equatorial positions that are 120° apart

take up more space

lone pairs

Smaller dipole moment

smaller electronegativity difference between the two atoms

Dipole moment = 0

nonpolar

Dipole moment > 0

polar

Bond dipole depends on

direction and magnitude

If bond dipoles are not equal length in opposite directions

they don’t cancel out and become nonpolar

The number of hybrid orbitals in a set is equal to

the number of atomic orbitals that were combined to produce the set

A covalent bond forms when

 an orbital from one atom overlaps with an orbital from another atom

Single bonds

sigma bonds

double bond

1 sigma, 1 pi

triple bond

1 sigma, 2 pi

Delocalized bonding

is resonance

E(-) diff < 0.4

nonpolar

E(-) diff: 0.4 to 1.9

polar

E(-) diff > 1.9

ionic