Unit 3: Electron configuration and periodicity

BOHR MODELSGroup 1

Group 2 Group 13 Group 14 Group 15 Group 16 Group 17

Group 18

H

Li Be

He

B C N O F Ne

Na Mg Al Si P S Cl Ar

K Ca

His theory couldn’t explain why metals give off characteristic colors when heated

Or why when iron is heated it first glows dull red, then yellow, then white at higher temperatures

There needed to be a model that better described the behavior or electrons….

His experiments worked with the H atom but not those that have more that one electron.

Bohr did propose that the electrons are in orbits around the nucleus.

Orbits- definite circular path where

an electron is supposed to revolve

around the nucleus

Think of the energy levels as rungs on ladder:

Lowest rung= lowest energy level

Like you can climb up the ladder an electron can move from one energy level to another

Like you can’t stand between rungs an electron can’t exist between energy levels

Like you must move just the right distance to move from one rung to another an electron must gain or lose just the right amount of energy- a quantum of energy is the amount of energy needed to move from one energy level to another

Electrons are on “energy levels”

Energy levels farther from nuclear are higher in energy.

Electrons can exist ON levels, not between levels.

Quantum mechanical model

Determines the allowed energies an electron can have and how likely it is to find the electron in various locations in the electron cloud

Can you predict the exact location of the ferris wheel cars at a given instant?

Electrons are found in atomic orbitals

A 3D region around the nucleus in which there is a high probabliilty of finding an electron

Different from orbits which are a definite circular path in which the electron is supposed to revolve around the nucleus

ELECTRON CONFIGURATION

To describe atomic orbitals and the properties of electrons in those orbitals we use quantum numbers.

They describe the

1. Energy level of electrons (n): These energy levels are assigned numbers (1, 2, 3, 4, etc.).

Within energy levels there are orbitals with different shapes….

2. Shapes of orbital (s,p,d,f):

And each shape of orbital can hold a max numbers of

electrons:

Aufbau principle electrons occupy orbitals of low energy first

Lower energy level orbitals are closer to the nucleus

Pauli exclusion principle

no two electrons can be in the same orbital moving the same way. (opposite spins)

Hund’s rule

When electrons are filling up orbitals of equal energy (say for instance 3 orbitals, which is 6 electrons), one electron enters each orbital until they’re half filled, then they fill with the opposite-spin electrons

Arrangement of electrons in an atom

Shows increasing energy levels and the orbitals on those levels.

“1”, “2”, “3” etc – these represent energy levels

S, p, and d represent orbital shape

1 box – represents one orbital (holds two electrons)

Show how electrons fill orbitals in an oxygen atom.

ORBITAL NOTATION

Na ___ ___ ___ ___ ___ ___

1s 2s 2p 2p 2p 3s

Letters/numbers represent energy level (remember start at bottom of chart)

___ lines represent 1 orbital each.

Show orbital notation for Cl and As

Cl

As

Write coefficient & letter for each energy level.

Superscript (number on top) shows # of electrons at that level.

This method simply takes less space.

Na 1s22s22p63s1

C 1s22s22p2

It can also be expressed in noble gas notation (shorthand notation):

[Ne] 3s1

[He] 2s2 2p2

Be

O

Mg

The “d” orbitals fill up in levels 1 less than the period number, so the first d is 3d even though it’s in row 4.

1

2

3

4

5

6

7

3d

4d

5d

p1 p2 p3 p4 p5 p6

S1 S2

Zn

Ag

Eu (careful- F block)

Instead of writing the whole electron configuration notation, the symbol for the last element in the previous row (noble gas) is written in brackets to represent all the electrons up until that principle level or row.

Examples:

Na becomes [Ne]3s1

Practice:

Ni

Y

Se

Lewis Dot diagrams show only the electrons available for bonding, known as the valence electrons. These are the electrons in the last principle quantum level (highest n value)

Si I H Mg

Practice:

O Cl Ar Ca

Electric currents passing through the gas in each glass tube makes the gas glow it’s own characteristic color

WAVE-PARTICLE DUALITY

A theory that attempts to explain how electrons behave in two different ways

Waves (like light)

Particles (like a ball)

The electromagnetic spectrum is a series of waves

that have different wavelengths.

All electromagnetic radiation (light) consists of electromagnetic waves that travel at 3.00 x 108 m/s

That’s 670 616 629 miles per hour!

Amplitude – height from crest to origin.

Wavelength – distance between crests or troughs of a wave (measured in meters)

*distance between corresponding points on adjacent waves*

Origin

Frequency – number of cycles that pass a given point in a given amount of time.

Measured in Hertz (Hz)

1 Hz = 1 wave passes per second

400 Hz = 400 waves pass per second

½ Hz = ½ wave passes per second

*ALL EM WAVES TRAVEL AT THE SPEED OF LIGHT*

c = fλc = speed of light = 3.00 x 108 m/s

f = frequency (Hz)

λ = lambda = wavelength (m)

As wavelength increases,

frequency __________.decreases

If an EM wave has a frequency of 548 Hz, what is its wavelength?

c = 3.00 x 108 m/s

f = 548 Hz

λ = ?

c = f λ3.00 x 108 m/s = (548 Hz) λ

λ = 547445 m

If an EM wave has a wavelength of 630 nanometers, what is its frequency? (1 nm = 10-9 m)

c = 3.00 x 108 m/s

f = ?

λ = 630 x 10-9 m

c = fλ

3.00 x 108 m/s = (f) (630 x 10-9 m)

f =4.76 x 1014 Hz

The visible spectrum is continuous (there are no breaks and the colors blend together).

White light is a combination of ALL colors of light. A prism breaks up white light into the separate colors so we can see them.

Each color has a definite frequency and wavelength.

*The speed of these colors never changes (always speed of light (c))

Low energy colors have a long wavelength and low frequency.

High energy colors have a short wavelength and high frequency.

Remember that electrons occupy energy levels - the region surrounding the nucleus where an electron is likely to be found (think of rungs on a ladder, fixed levels with space in between)

When electrons are in their lowest energy level, they are said to be in their “ground state”

It is possible for electrons to jump from ground state to a higher energy level (called excited state) by absorbing energy.

When electrons gain energy somehow, they can then occupy a HIGHER ENERGY LEVEL. They jump up to the next level and are said to be in their “excited state”.

When electrons lose energy they will fall back down to their GROUND STATE and release energy, and some of it is released as waves we can see – LIGHT!

Atomic Emission Spectrum

With many electrons jumping to energy levels and falling back, many different shades of light are released and blended.

We can use a prism to separate the light to see the individual shades

Atomic Emission Spectrum

Continuous Spectrum

(no breaks)

Spectroscopy – the science of producing atomic spectra and studying them.

PARTICLE MODEL

The idea that light can act as a particle.

These particles are called photons, or quanta, and can move other matter (knock loose electrons from metal- photoelectric effect).

This powers solar powered calculators.

EINSTEIN

Einstein proposed in 1905 that light can behave as

both a wave and a particle.

Each particle of light is a photon of electromagnetic radiation with no mass and carrying a quantum of energy.

The energy contained in a photon (a quantum) depends on its frequency.

E = hv

E = energy (units are in Joules)

h = Planck’s constant = 6.63 x 10-34 J.s

v = frequency

*Quantum= minimum amount of energy that can be lost or gained by an atom*

If a wave has a frequency of 230 Hz, what is the amount of energy of one photon (quantum) of this wave?

If a wave has a wavelength of 400 nm, what is the amount of energy of one photon (quantum) of this wave?