Chapter 7: the Electronic Structure of Atoms

quantum theory

idea that atoms/molecules emit energy in discrete quantities called quanta rather than continuously

by Planck

wave

a vibrating disturbance by which energy is transmitted

speed depends on type and nature of medium its travelling through (ex: air, water, vacuum)

wavelength (lambda)

the distance between identical points on successive waves

frequency (v: nu)

the number of waves that pass through a particular point in one second

unit: hertz (Hz) = 1 cycle/s

amplitude

the vertical distance from the midline of a wave to the peak or trough

speed (u)

equals wavelength times frequency

= (lambda) (nu)

electromagnetic wave

has an electric field component and a magnetic field component (both components have same wavelength, frequency, and speed, but travel in mutually perpendicular planes)

Maxwell's theory

provides a mathematical description of the general behavior of light

electromagnetic radiation

the emission and transmission of energy in the form of electromagnetic waves

speed of light (c)

speed of electromagnetic wave in a vacuum

3 x 10^8 m/s, 186,000 miles per sec

v = c / lambda

types of radiation

gamma rays, X rays, UV, visible light (violet to red), infrared, microwave, radio waves (FM to AM)

quantum

the smallest quantity of energy that can be emitted (or absorbed) in the form of electromagnetic radiation

E = hv = h (c / lambda)

h is Planck's constant, v is frequency of radiation

Planck's constant

6.63 x 10^-34 J s

photoelectric effect

a phenomenon in which electrons are ejected from the surface of certain metals exposed to light of at least a certain minimum frequency (threshold frequency: all or none)

number of ejected electrons is proportional to intensity

photons

by Einstein

particles of light

each must possess energy E

KE = hv -W

emission spectra

either continuous or line spectra of radiation emitted by substances

unique for each element, like a fingerprint

line spectra

the light emission only at specific wavelengths

bright lines in different parts of the visible spectrum instead of a continuous spread of wavelengths from red to violet

Rhydberg constant

for hydrogen: 2.18 x 10^-18 J

En = -Rh ( 1 / n^2)

ground state (level)

the lowest energy state of a system

excited state (level)

higher in energy than the ground state

node

points on the string where amplitude of the wave is 0

located at each end and sometimes between ends

greater frequency means more

De Broglie

2pir = n lambda

r is radius of orbit, lambda is wavelength, n = 1, 2, 3

lambda = h / mu

Heisenberg uncertainty principle

it's impossible to know simultaneously both the momentum (mass times velocity) and the position of a particle with certainty

can't know location and momentum at the same time

quantum (wave) mechanics

Schrodinger

electron density

gives the probability that an electron will be found in a particular region of an atom

atomic orbital

the wave function of an electron in an atom

many-electron atoms

atoms containing 2+ electrons

quantum numbers

required to describe the distribution of electrons in hydrogen and other atoms

principle (n), angular (l), magnetic (ml), electron spin (ms)

principle quantum number (n)

integral values: 1, 2, 3, etc.

larger means farther from nucleus

angular momentum quantum number (l)

tells shape of the orbital

dependent on n (can be from 0 to n-1)

designated by s(0), p(1), d(2), f(3), g(4), h (5)

magnetic quantum number (ml)

describes orientation of orbital in space

depends on l: -l, (-l + 1), ... 0 ... (+l - 1), +l

electron spin quantum number (ms)

can be +1/2 or -1/2 depending on spinning motion of electron

boundary surface diagram

encloses about 90% of total electron density in an orbital

s orbital

spherical

size is proportional to n^2

electron configuration

how the electrons are distributed among the various atomic orbitals

electron configuration

ex: 1s^1 (read 1 s 1)

1: principle quantum number n

s: angular momentum number l

^1: number of electrons in the orbital/subshell

orbital diagram

shows the spin of the electron

upward vs. downward arrow

pauli exclusion principle

no 2 electrons in an atom can have the same 4 quantum numbers (change ms if n, l, and ml are the same)

only 2 electrons can occupy the same orbital: they must have opposite spins

paramagnetic substances

contain net unpaired spins and are attracted by a magnet (up up or down down): parallel spins

diamagnetic substances

don't contain net unpaired spins, are slightly repelled by a magnet (up down or down up): antiparallel spins

hund's rule

the most stable arrangement of electrons in subshells is the one with the greatest number of parallel spins

general rules

1. each shell or principle level of quantum number n has n subshells

2. each subshell of quantum number l has (2l + 1) orbitals

3. no more than 2 electrons in each orbital: max number of electrons is 2x number of orbitals used: must have opposite spins

4. determine mx number of electrons in a principle level using 2n^2

1. no 2 electrons can have same 4 quantum numbers (pauli exclusion)

2. most stable arrangement of electrons is one with greatest number of parallel spins (hund's rule)

3. unpaired electrons: paramagnetic, atoms with paired electrons are diamagnetic

4. in hydrogen, energy of electron only depends on n, in many-electron atoms, the energy depends on n and l

5. in many-electron atoms, subshells are filled in order (figure 7.21)

6. s > p > d > f (more energy needed to separate 3s electron than a 3p electron)

aufbau principle

as protons are added one by one to the nucleus o build up the elements, electrons are similarly added to the atomic orbitals

noble gas core

shows in brackets the noble gas element that most nearly precedes the element being considered

transition metals

either have incompletely filled d subshells or readilygive rise to cations that have incompletely filled d subshells

ground-state configurations of elements

see pg. 242 + 243

lanthanides (rare earth series)

have incompletely filled 4f subshells or readily give rise to cations with incompletely filled 4f subshells

actinide series

last row of elements

most aren't found in nature (synthesized)