Sec. 4.3- Electron Structure

The Story of the Electron Hotel A man built a hotel for electrons with a restaurant next door. He was making so much money that he decided to add on with some more rooms and a parking garage. He still had high demand and decided to add on some more rooms and a shopping center. He used the last space he could to put some rooms above the shopping center.

Restaurant Parking Garage Shopping Center

This man had some very strange ideas about how to run his hotel This man had some very strange ideas about how to run his hotel. He insisted four things: The lowest possible level must be used first (actually it was the fire inspector that insisted on this one) There can only be one electron in a room until all rooms at that level contain one electron No more than 2 electrons to a room When two electrons are in a room, they must point in different directions (either up or down)

If he had 8 electrons to put in rooms, where would they go? Restaurant Parking Garage Shopping Center

What about 21 electrons? Restaurant Parking Garage Shopping Center

42 electrons? Restaurant Parking Garage Shopping Center

Where do Electrons Actually Live? They don’t actually live in a hotel- they live around the nucleus inside all atoms What term did we use for the location of electrons?

Electron clouds are composed of energy levels these energy levels are not actually locked orbits like Bohr displays These energy levels are further broken down into “subshells” Subshells are then composed of atomic orbitals What were atomic orbitals? Electrons live in these atomic orbitals- max of 2 electrons per orbital

Each one of these structures can be related to features on the electron hotel Electron cloud= Electron hotel Energy level= section of hotel (Above restaurant, parking garage, shopping center) Subshell= Floor Orbital= Which room

Energy levels- the “rings” in Bohr’s model Not actually rings, but can be useful to picture them that way- just understand that the electron location will not always be locked in that particular orbit Energy levels are given numerical values- 1, 2, 3… Energy level 1= electrons with lowest energy, closest to the nucleus As energy levels increase, the electron energy increases, and moves further away from the nucleus

Subshells- show shape of electron cloud Depending on the energy level an electron is located in, the shape of its orbit/ cloud of probability will change These subshell shapes are given letter designations

Four types of subshells: s, p, d, f Lowest energy= s Highest energy= f Subshells are described by their energy level and type Ex: 2s (energy level 2, s subshell)

The four types of subshells have energy level requirements Subshell s= energy level 1 or higher Subshell p= energy level 2 or higher Subshell d= energy level 3 or higher Subshell f= energy level 4 or higher

Every subshell has a set of orbitals based off of the shape/orientation of the subshell itself Remember- orbitals are areas of high probability of finding an electron- an approximation, not a pinpoint location

For this class, here is what you need to know: Subshell s= 1 orbital Subshell p= 3 orbitals Subshell d= 5 orbitals Subshell f= 7 orbitals Each orbital can hold a max of 2 electrons

Flow Chart- Breakdown of Electron “Addresses”

Types of subshells Subshell Begins in Energy Level Number of orbitals Max # of electrons held s 1 p 2 3 d 5 f 4 7

Electron Configurations Electron configurations show us the grouping/position of electrons in an atom The number and configuration of electrons determines how something glows…so it’s important to know “where the electrons live” for an atom!

Electrons will fill subshells starting with the lowest possible energy first Which energy level did we say was lowest energy? Which subshell was lowest energy? Highest?

So, electrons will fill subshells in this order: 1s, 2s, 2p, 3s, 3p, 4s, 3d, 4p, 5s, 4d, 5p, 6s, 4f, 5d, 6p, 7s, 5f, 6d, 7s, 5f, 6d, 7p Woah…. Lots of subshells- but don’t worry, lucky for us, the periodic table is an extremely useful tool for remembering what order electrons fill subshells

When writing electron configurations out, there are two ways of writing them Orbital diagrams Spectroscopic Notation

We will start with orbital diagrams: a horizontal line is drawn for each orbital in a given subshell. The number of lines drawn depends on the number of orbitals in that subshell. Arrows are drawn on each line to represent the electrons Ex:

Electron Configuration-Rules Aufbau Principle: electrons fill subshells (and orbitals) so that the total energy of atom is the minimum What’s that mean? Electrons fill lowest available subshells before moving on to next higher energy level Why do they fill this way?

How many orbitals did a p subshell have? Hund’s Rule: Place electrons in unoccupied orbitals of same subshell before doubling up How many orbitals did a p subshell have? Orbital diagram of p subshell: If you have 3 electrons in a p subshell, you have to place one electron in each orbital before doubling up

Pauli Exclusion Principle: If two electrons occupy the same orbital, they MUST have different spins Huh? Only spins= up and down So if 2 electrons in same orbital (on same line)- 1 up, 1 down Ex: 4 electrons in a p subshell (you will need to double up)

Orbital Diagram of Cl First- how many electrons does Cl have? Look back at list of lowest to highest subshells- start at the beginning, then once lowest energy subshell fills, continue to next (Aufbau principle) One electron in each orbital in same subshell before doubling up (Hund’s Rule) If electrons in same orbital, opposite spins (Pauli Exclusion) Orbital diagram- label orbitals with their subshells- number of lines for each subshell depends on number of orbitals that subshell has

Spectroscopic Notation Shorthand way of showing electron configurations Write the subshell labels (1s 2s 2p 3s…), with the number of electrons in that subshell as a superscript Ex: Orbital diagram of Cl: Spectroscopic notation:

Writing Spectroscopic Notation 1: Determine number of electrons needing homes 2: Follow Aufbau Principle for filling order 3: Fill in subshells until they reach their max (s=2, p=6, d=10, f= 14) 4: The sum of all the superscripts should equal the total number of electrons (check your work) Spectroscopic for S:

Noble Gas Notation Noble gases? All gases in group 8A- last group on the right side of periodic table. They contain full valence shells Valence shell? Valence Electrons? Noble Gas notation- even more shorthand- Noble gas used to represent notation of inner electrons, and only the valence electrons are written out after that Ex: Spectroscopic for Br: Noble Gas notation:

Which Noble Gas do I choose? Think Price is Right How do you win the initial bid? Choose the Noble gas that is closest to the element without going over! Which Noble gas would you use for Br?

Noble Gas Notation Example 1: Determine number of electrons to place 2: Determine which Noble Gas to use 3: Start where Noble gas left off and write spectroscopic notation for valence electrons Ex: Write the Noble Gas Notation for As