Chapter 4: Arrangement of Electrons in Atoms

Electrons can move as

Electrons can move as

Particles and Waves

Wave Nature

• Proposed by Louis de Broglie

• Electrons could be considered waves confined to the space around the NUCLEUS

• Only certain wavelengths and frequencies could exist

A wave can be describes at

wavelength ( labda)

Frequency v (nv)

Energy: E

Wavelength

• Symbol: Lambda

• The distance between the crests of a wave

• UNITS: meters or nanometers

• 10^9nm = 1m

Frequency

• Symbol: v (nv)

• The number of waves that pass/second

• UNITS: 1/sec = 1 hertz

Speed of Light

• A constant

• Symbol: C

• fastest possible speed

Speed of light equation

3.00 x 10^8 meters/second

lambda and v are ….

inversely related

long wavelength =

low frequency

short wavelength =

long frequency

Electromagnetic Spectrum

350 nm = violet to 750 nm = red

Particle Nature of Light

Some properties of light cannot be explained by wave theory: white hot objects and the photoelectric effect

“White Hot” Objects

• When objects are heated they EMIT LIGHT

• Wave theory predicts that only UV LIGHT would be emitted but white, yellow, and orange light

Photoelectric Effect

• Emission of electrons by certain metals when light shines on them

• The light must be of a certain energy/frequency to “knock” an electron from the surface

  (Albert Einstein observed in 1905)

Quantum Energy

• MAX PLANCK

• White hot objects and the photoelectric effect both emit small, specific amount of energy call QUANTA or PHOTONS

• A specific amount “ bundle” of energy that can be gained or lost by an atom

Energy of a Quantum

• An individual Quantum is known as a PHOTON of light and has evergy

Energy of a Quantum equation

• h = Plack’s constant

• h = 6.626 x 10 ^-34 joules x seconds

• E = hv

Energy + wavelength equation

E = hc/lambda

• inversely related

Wavelength of colors

750-625 nm - Red

625-590nm - Orange

590-565 nm - Yellow

565-520 nm - Green

520-435nm - Blue ( indigo)

435-350 nm - Purple

Wavelengths ( highest to lowest)

Gamma Ray

X-Ray

Ultraviolet light

Visible

Infread

Microwave

Radio

Bohr Model

Why did Bohr study hydrogen?

- After observing the H line emission spectrum

• Simplest atom - only has one electron

• The electrons of the gas are easily excited by a current

1. How are the electrons moving

2. Where are the located

3. How much energy do they have

1. As particles with a definite circular path

2. In rings, or orbits around the nucleus

3. Th energy of an electron can be calculated

1. Electrons closer to the nucleus have …

2. At higher energy levels, the energy is …

1. Lower energy values

2. Higher

• each orbit is an energy lever designated by the variable n

• N =1 is the ground state

• N = infinity is when the electron has been removed from the atom

• Energy is quantized, meaning there are no levels in between the levels designated by n = 1,2,3,4, ect

Spectroscopy

When atoms are excited by a outside energy source ( heat, flame, or electric current) the electrons can be promoted to higher energy states. However, this situation is highly unstable and the electron will eventually return to a lower energy state.

• when the electron returns to a lower energy level, energy is given off in the form of ( quanta), also known as a (photon of light)

Energy input:

From LOWER levels to HIGHER level

Energy output:

From HIGHER levels to LOWER levels

Hyrogen produces lines in the

visible, ultraviolet, and infrared regions of the electromagnetic spectrum

Balmer series

visible

Lyman series

Uv

Paschen series

IR ( infread)

Energy at each level

E = -2.178 x 10^-18 J ( Z^2/ n^2)

Z = number of protons

n = energy level

• Calculation: deltaE = Efinal - Einital

• -2.178 x 10^-18 ( Z^2/n^2 (efinal) -Z^2/n^2 (einitial) )

Bohr’s Model ( hydrogen)

• Cannot successfully predict line spectra for an other element besides hydrogen

• Electrons must not be moving in circular orbits at set distances from the nucleus

Electrons

• Moving as waves

• Not in orbits

• Located in Orbitals

• Areas of high probability based on the wave motion and energy of an electron

• Every orbital has an “address” given by a set of 4 Quantum numbers

• Energy: determined by h (like Bohr)

• Higher n= higher energy level = more energy = further from the nucleus

• Schrodinger

Proposed a complex mathematical relationship to predict where e- is located

DeBroglie

• Proposed the particle to describe electron motion

Heisenberg

• Proposed the Heisenberg uncertainty

• The more you know about e- position, the less you know about it’s velocity (Inversely)

• Pauli

• Every electron in an atom must have a unique set of quantum numbers

• Electrons in the same orbital must have opposite spins

Hund’s

• Electrons will fill an unoccupied within a sublevel before pairing up with another electron in an orbital

• This minimizes repulsion and leads to stability

Principle Quantum Number

• n

• Tells what energy level the electron level is in

• Possible values: 1 to infinity

• Orbital Quantum Number

• l

• Tells about the Shape of the orbital where the electron is housed

• Possible values: 0 to ( n-1)

• Magnetic Quantum Number

• Ml

• Tells about the orientation of the orbital around the nucleus

• Possible values: -l to +l including 0

Spin Quantum Number

• Ms

• Tells the direction of the electron spin within an orbital

• Possible values: +1/2 and -1/2

• +1/2 : Clockwise spin

• -1/2 : Counterclockwise spin

S orbital:

• spherical ( l =0)

• 0

P orbital

• peanut shaped ( l =1)

• -1,0,1

D orbital:

• daisy shaped ( l =2)

• -2,-1,-0,1,2

F orbital

flower shaped ( l =3)

-3,-2,-1,0,1,2,3

Energy level Diagrams

• A way to keep track of electrons

• Includes line boxes that represent Orbitals

• Also shows orbital types and energy levels

• Once filled, it is possible to find the quantum number set for any electron in the atom

Aufbau Principle

• Meaning: Build up

• Electrons must be filled from the lowest energy to the highest energy

• Do not move to the next until the one below has been completely filled

Hund’s Rule

• Electrons will fill an unoccupied within a sublevel before pairing up with another electron in an orbital

• This minimizes repulsion and leads to stability

Octet Rule

• Atoms become stable when their outermost energy level contains 8 electrons

• Noble gases have a full OCTET

• The outermost electrons are called Valance Electrons (except He)

Electron Configuration

• Shows every electron in the atom starting with the 1s orbital

• Valence electrons exist in the highest energy level listed

Noble Gas Configuration

• Uses the preceding noble gas core to shorten the electron configuration

• Allows the valence electrons to be determined easily