Chem 6-10

wavelength

wavelength

the distance from one point to a similar point on a wave

amplitude

the distance from the origin to the crest or origin to the trough

frequency

the amount of waves per second

Electromagnetic spectrum

a range of all types of electromagnetic radiation

electromagnetic waves

radio waves, microwaves, visible light, ultraviolet radiation, x-rays, gamma rays

photon

a packet of light

what is h

planck’s constant

What number is planck’s constant

6.6262 X 10^-34

what is v

frequency (Hz)

what is c

speed of light (m/s)

what number is the speed of light

3.00 X 10^8

what is lambda

wavelength (m)

Energy equation

E= hv

Speed of light equation

c= lambda(v)

Excited atoms

absorbed energy and electrons are bumped up to upper energy levels and are unstable

Ground State

When electrons are at their lowest energy configuration

Heisenberg uncertainty principle

It is impossible to determine accurately the momentum and position of a particle (You can never know where electrons are)

Orbital

an area of high probability of finding an electron

n

energy levels

l

sublevel (s=0,p=1,d=2)

m

orbital (-l → +l)

s

spin ( +/- 1/2)

Electron Configuration

Organizing electrons in an atom with the lowest energy level to the orbital with the highest energy level

S orbital

O- shape

P orbital

infinity- shape

Aufbau principle

electrons will fill the lowest energy level first (arrow thing)

Hund

electrons will fill one orbital at a time (-1, 0 , 1)

Pauli

No two electrons can have the same quantum numbers (opposite spin)

Valence electrons

electrons in the outermost energy level

core electrons

electrons that are not in the outer energy level

Periodicity

periodic trend

Atomic radius

1/2 the distance between 2 nuclei of 2 like atoms in a solid crystal

electronegativity

the attraction that an atom has between a shared pair of electrons and its positive nucleus

e- affinity

energy change for adding an electron to an atom in a gaseous state

ionization energy

energy change for the removal of 1 mole of electrons from 1 mole in a gaseous state

Periodic trends

know about AR, IE, E- affinity, electronegativity

ionic bonds

an electrostatic attraction due to opposite charges

octet rule

atoms will lose, gain, or share e- to achieve the same electron configuration to the nearest Noble gas.

Covalent Bond

Sharing electrons

Resonance structures

2 or more valid electron dot formulas for a single molecule (can write the Lewis dot structure multiple different ways)

Sigma Bond

Bonding orbitals that overlap head to head (each bond has 1 )

Pi bond

Bondin orbitals that overlap side by side (2 in triple bonds, 1 in double bond)

VSPER Theory

electron pairs want to be as far away from each other as possible. (Why we get the shapes and angles)

Polar

unequal sharing of electrons

Non-Polar

equal sharing of electrons

Unshared pair (lone pair)

the 2 dots on dot structure

unpaired electron

alone dot on lewis dot structure

Hybridization

When S and P mold to make SP

Polar molecule

one end of the molecule is slightly negative

Non-polar molecule

symmetric polarity within the molecule

Atomic Radius Periodic trends

Side to side makes more protons and more electrons in the same energy level = less nuclear pull so decrease. Smaller . Down is more energy levels Bigger so increase (Fr is highest)

Ionization Energy Periodic Trend

side to side there is more valence electrons so increases. Up is Smaller the atom the more nuclear pull-on valence electron so smaller (He is the largest)

Electron affinity Periodic Trend

Side to side is more protons so increse. Down is more shielding. Bigger the atom the less nuclear pull so decrease (F highest)

Electronegativity Periodic trends

Side to side is more protons so increase. Down is bigger so less nuclear pull so decrease. (F highest)

shielding

inner electrons block outer electrons. Larger the atom this is increased

Orbitals of equal energy

degenerate

Isoelectronic

Species atoms/ and or ions that have the same number of electrons (same electron configuraation

intermolecular forces

weak forces of attraction between two or more molecules

Dispersion forces

the positive end of a momentary dipole is attracted to the negative end of another momentary dipole

dipole- dipole forces

the occurrence of an electrostatic attraction between the positive end of a dipole to the negative end of another

Momentary dipoles

asymmetric distribution of electrons in an atom at a gien time

Hydrogen Bonding

a dipole-dipole force of attraction between two or more molecules but is due the hydrogen bond within a molecule, specifically between H-N, H-O, and H-F

Mega

Know types of bonds chart

Know types of Bonds charts

Metallic bonds

an attraction of positive ions to a “sea” of electrons