Chemistry - Chapter 4

electromagnetic radiation

a form of energy that exhibits wavelike behavior as it travels through space

speed of light

3.00 x 10^8 m/s (C)

wavelength

λ (the distance between corresponding points on adjacent waves)

frequency

ν (number of waves to pass a given point in 1 second)

c =

vλ

photoelectric effect

refers to the emission of electrons from a metal when light shines on the metal, cannot be explained by wave nature of light

quantum

the minimum amount of energy that can be gained or lost by an atom, one required to move between energy levels

E =

hv

Planck's constant

h, 6.626 x 10^-34 j/s (m^2 * kg / s)

E =

hc/λ

photon

a particle of electromagnetic radiation with no mass that carries a quantum of energy

ground state

the lowest energy state of an atom

excited state

the atom has a higher potential energy than it has in its ground state

hydrogen's emission-line spectrum

when a narrow beam of the emitted light was shined through a prism, it was separated into four specific colors of the visible spectrum

Niels Bohr

1913- discovered that electrons move around the nucleus in orbits called electron shells.

continuous spectrum

the emission of a continuous range of frequencies of electromagnetic radiation (each color blends into the next in order)

line spectrum

gives off energy--a spectrum showing only certain discrete wavelengths

when atoms absorb energy...

their electrons jump to higher energy levels, when they come back to ground state they emit energy in the form of light

lyman series

all electron transitions that end on the first energy level (n=1), UV range of light

balmer series

all electron transitions that end on the second energy level, only 5 possible lines

paschen series

all electron transitions that end on the third energy level, all involve small energy changes so they are all in the infra red region of the spectrum

major shortcoming of Bohr's model

only worked on hydrogen, electrons do not move in circular orbits

electron dual properties

both properties of particles (jump and fall between high and low energy orbitals) and properties of waves (e- have tiny wavelengths)

principal quantum number

symbolized by n, indicates the main energy level occupied by the e-, n= 1,2,3,4...

azimuthal quantum number

symbolized by l, indicates the shape of the orbital, values 0 to (n-1)

s-orbitals

s-subshell, spherical in shape, can hold maximum of two electrons

p-orbitals

p-subshell, dumbbell shaped, three p-orbitals, hold maximum of 6 electrons

d-orbitals

d-subshell, dumbbell-shaped, 5 d-orbitals, hold maximum of 10 electrons

f-orbitals

f-subshell, max. of 14 electrons, highest in energy

magnetic quantum number

relates to the oriention of the orbital in space relative to the nucleus, -l to 0 to +l

electron spin quantum number

relates to the way the e- spins, +1/2 for clockwise and -1/2 for anticlockwise, produce oppositely directly magnetic fields

aufbau principle

an electron occupies the lowest-energy orbital that can receive it

pauli exclusion principle

no more than two electrons can occupy a space orbital, and these two electrons must have opposite spins

hund's rule

when electrons occupy orbitals of equal energy, one electron enters each orbital until all the orbitals contain one electron with parallel spins

orbital

the area where probability of finding an electrion is high

valence electrons

electrons on the outermost energy level of an atom

core electrons

inner electrons

Heisenberg uncertainty principle

it is impossible to know exactly both the velocity and the position of a particle at the same time, because location of e- can only be determined if it is struck by another particle but after that happens the velocity changes