Chapter 4: Arrangement of Electrons in Atoms

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