Electron Configurations

Parts of the atom Protons identify the element. Neutrons add mass and help glue the nucleus together ( all those protons are NOT happy being stuck next to each other).

Parts of the atom Electrons and their ground state energy addresses tell how an element reacts. Since chemistry is about reactions we care a lot about those addresses

Quantum Theorists The quantum theorists found out that electrons do not move in circular orbits but in complex three-dimensional shapes called orbitals.

What options do the electrons have? As there are more electrons, there are more possible ground state energy addresses These energy addresses are called quantum numbers

What options do the electrons have? The locations an electron can fill are called orbitals, only two electrons can fit in any orbital. When the possible orbitals are graphed they have shapes that determine how molecules will form

Quantum numbers Each electron has 4 quantum numbers Principal Quantum Number - Main energy level Orbital- Shape of the orbitals (also called sublevels)

Quantum numbers Each electron has 4 quantum numbers Magnetic- Which part of the orbital group the electron is in Spin- Which way the electron is pointing its magnetic field (North/South or Up/Down)

Rules for Electron Energy Addresses 1. Aufbau (Building Up) Principle Electrons fill orbitals with lowest energy first.

Rules for Electron Energy Addresses 1. Aufbau (Building Up) Principle Electrons fill orbitals with lowest energy first 2. Pauli Exclusion Principle Each orbital can only hold two electrons and they must have opposite spins 3. Hund s Rule Electrons will fill the orbitals at an energy level one at a time until each has one electron. They will all have the same spin. Only after all the orbitals have one electron will electrons of opposite spins start to double up.

Practice Use the three rules to put the electrons in their addresses

Writing Electron Configurations Write a number for the energy level Write a lower case letter for the sublevel Write a superscript for how many electrons are in all the orbitals of the sublevel Example Helium (2 electrons, both in 1s) 1s 2 Example Fluorine (9 electrons) (2 in 1s, 2 in 2s, 5 in 2p) 1s 2 2s 2 2p 5

Representing Electron Location Lets look at how we represent Lithium s electrons Li - 1s 2 2s 1 Lithium has 3 electrons 2 in energy level 1 in the s orbital 1 in energy level 2 in the s orbital

The periodic table is a tool Electron location can be determined using the periodic table

Vertical columns are called groups.

1 2 3 4 5 6 7 Each horizontal row is called a period

Energy Levels Each period represents an energy level Example: Period 1 consists of Hydrogen and Helium Hydrogen has 1 electron Helium has 2 electrons

1 2 3 4 5 6 7 H He

Orbitals Four types of orbitals s, p, d, f The orbitals fill up in a regular pattern. This pattern is dependent upon the number of electrons an element has in its electron cloud

What s Orbitals Look Like

What p Orbitals Look Like

What d Orbitals Look Like

What f Orbitals Look Like

Electron Configurations and the Periodic Table

s 1 s 2 S- block In this section of the periodic table, the highest energy level electrons are filling s orbitals.

The P-block p 1 p 2 p 3 p 4 p 5 p 6 In this section of the periodic table, the highest energy level electrons are filling d orbitals.

d block In this section of the periodic table, the highest energy level electrons are filling d orbitals. d 1 d 2 d 3 d 5 d 5 d 6 d 7 d 8 d 10 d 10

D orbitals fill up after previous energy level so first d is 3d even though it s in row 4. 1 2 3 4 5 6 7 3d orbitals

F - block inner transition elements f 1 f 2 f 3 f 4 f 5 f 6 f 7 f 8 f 9 f 10 f 11 f 12 f 13 f 14

1 2 3 4 5 6 7 f orbitals start filling at 4f 4f 5f

The Last Electron in the Ground State The last electron entered to the ground state configuration has the highest energy, and can be used to determine the chemical properties of that element.

Electron configurations The Amazing Sinking Orbitals

Periodic Table The rows in the periodic table act as a guide to the order in which the sublevels fill

1s 2s 2p 3s 3p 4s 3d 4p 5s 4d 5p 6s 5d 6p 7s 6d 7p

Empty, Half-full, Full Sublevels are particularly stable when they are empty, full or half-full.

Empty, Half-full, Full Sub levels are particularly stable when they are empty, full or half-full. Because of this, if the atom can shuffle electrons to get this conformation, it will

Changing Energy while Filling Orbitals in Sublevels A sublevel with electrons in it has a lower energy than an empty sublevel

Changing Energy while Filling Orbitals in Sublevels A sublevel with electrons in it has a lower energy than an empty orbital You can imagine sublevels sinking, just a tiny bit as electrons are added to them

7 s 6 s 5 s 4s 6 p 5 p 4 d 4p 3p 5 d 4f 3d 2s 3s 2p EMPTY 1s L=0 L=1 l-=2 l-=3

7s 6s 5s 4s 3s 2s 6p 5p 4p 3p 2p 5d 4d 3d 4f 1s L=0 L=1 l-=2 With Electrons l-=3

Half-full orbitals Half full sublevels are stable (low energy) When electrons are filling the orbitals for Chromium, there are 2 choices 4s 2 3d 4 or 4s 1 3d 5 Since 4s 2 3d 4 ALMOST has half filled sublevels and 4s 1 3d 5 has two halffilled sub levels it is lower energy and is more stable

Full and Half-full orbitals Full sublevels are stable (low energy) When electrons are filling the sublevels for Copper, there are 2 choices 4s 2 3d 9 or 4s 1 3d 10 Since 4s 2 3d 9 ALMOST has a full sublevel and 4s 1 3d 10 has one full and one half-filled sublevel it is lower energy and is more stable

NOBLE GASES In their ground state Noble Gases have all their main and sublevels and orbitals FULL SO STABLE!!!!!

Ions An ion is an atom or group of atoms that have gained or lost one or more electrons.

Ions Atoms will generally gain or lose electrons to reach a stable conformation. LIKE THE NOBLE GASES Atoms will gain or lose as few electrons as possible to reach noble gas configuration

Ions Electrons are negative, so ions that are positive (cations) have lost electrons Ions that are negative (anions) have gained electrons

Ions To make positive ions, remove electrons Ca -2e - = Ca +2 To make negative ions, add electrons Br + 1e - = Br -

Ions with d electrons Because the energy of d orbitals drops as they fill, they eventually become lower in energy than the s orbitals that they are next to on the periodic table. Because of this, elements that have d orbitals and s orbitals on the same row of the periodic table lose their s orbital electrons first

Configurations for ions Are done the same way as configurations for neutral atoms Just assign electrons to the right energy level, sublevel (and orbital if you are doing orbital configurations)

Isoelectronic Configurations Ground state atoms have unique configurations But when you start adding or subtracting electrons, you can have species with the same configuration (Like O -2 and Ne). These are called isoelectronic

Isoelectronic Configurations Most species that are isoelectronic are sharing configurations with noble gas elements because they are SO STABLE!

Writing Electron configurations the easy way Yes there is a shorthand

The Shorthand Write the symbol of the noble gas before the element. Then the rest of the electrons. Aluminum - full configuration. 1s 2 2s 2 2p 6 3s 2 3p 1 Ne is 1s 2 2s 2 2p 6 so Al is [Ne] 3s 2 3p 1

More examples Ge = 1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 3d 10 4p 2 Ge = [Ar] 4s 2 3d 10 4p 2 Hf=1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 3d 10 4p 6 5s 2 4d 10 5p 6 6s 2 4f 14 5d 2 Hf=[Xe]6s 2 4f 14 5d 2

The Shorthand Again Sn- 50 electrons The noble gas before it is Kr Takes care of 36 Next 5s 2 Then 4d 10 Finally 5p 2 [ Kr ] 5s 2 4d 10 5p 2