Electron Energy and Light

Quantum Model of the Atom

Electromagnetic Spectrum

• Broad range of energetic emissions

• Waves tranfer energy

Niels Bohr

• Hydrogen emission spectrum
• Discovered "energy levels" of electrons

Electron Energy Levels

• Electrons can only have certain amounts of energy, so Bohr realized they can move around the nucleus only at certain distances

• Change in energy is like climbing up and down a ladder

Electrons and Light

• Electrons absorb energy and jump to a higher energy level (excited state)

• They give off energy and return to the original energy level (ground state) and release light

Niels Bohr

Bohr Model 
- Nucleus is in the center
- n = the electron orbits 

How to Represent Atoms Using Bohr Models

• We can use Bohr models to show how electrons are arranged in shells around an atom

• There are specific rules for how the electrons fill the shells

• Example: Oxygen

How to Represent Atoms Using Bohr Models

• First, we identify the ATOMIC NUMBER of that atom

• Atomic # = # of protons = # of electrons
  

• You can find the atomic # on the periodic table

Multiple Choice

If Magnesium (Mg) has an atomic # of 12, how many protons does it have?

• 6

• 12

• 2

• 24

Multiple Select

What does the number 26 represent in the image of Iron (Fe)? (select ALL correct answers)

• Atomic #

• # of protons

• # of neutrons

• # of electrons

Rules for placing electrons in a Bohr Model

• Electrons are placed around the nucleus using the following rules:

• 1st  orbital gets 2 electrons

• 2nd orbital gets 8 electrons
  

• 3rd orbital gets 8 electrons

• Start with the 1st orbital. Once full, extra electrons are placed in the next orbital

Filling in the electrons - First 10 elements

Multiple Choice

How many orbitals would Silicon (Si) need if it has an atomic # of 14?

• 3

• 2

• 1

• 14

Silicon has 3 orbitals

Beyond Bohr

How we got to our current model

Heisenberg Uncertainty Principle

• States that the position and the velocity of an object cannot both be measured exactly at the same time

• Electrons are thus not in neat fixed orbits like Bohr thought, but in orbitals

• Orbital: 3-D region around the nucleus that indicates the probable location of an electron

Erwin Schrodinger

• He took the Bohr model one step further

• He used math!! to describe the likelihood of finding an electron in a certain position

• This model is known as the Quantum Mechanical Model of the atom

Schrodinger's Equation

• The solutions of his equation can tell us things about the electron

• 1) The size of the electron cloud

• 2) The energy of the electron cloud

• 3) The shape of the cloud

Electron Cloud Model

Electrons are NOT in neat, planet-like orbits. There are regions where they are LIKELY to be found.

Quantum Mechanical Model or Electron Cloud Model

Orbitals

• Electrons are not in neat fixed orbits like Bohr proposed, but in orbitalsOr

• Orbitals: 3-D region around the nucleus that indicates the probable location (90%) of an electron

S-Orbitals

S orbital has one shape (sphere)

It can hold 2 electrons

P-Orbitals

3 different orientations

Each orientation can hold 2 electrons

Multiple Choice

How many electrons are needed to fill in ALL orientations of p-orbitals

• 2

• 4

• 6

• 12

D-orbitals

5 different orientations of the d-orbitals

Each orientation can hold 2 electrons

F-Orbitals

7 different orientations

2 electrons in each orientation

How to remember them all?

• S = sphere

• P = peanut shape

• D = double peanut + doughnut

• F = flower

• Each orbital # is counting odd #s in order: s=1, p=3, d=5, f=7

Electron Orbitals

These orbitals can overlap, cause interference with each other, and effect the overall energy of each other.

Electron Configurations

• The arrangement of the electrons in an atom

• Identify the probable locations of electrons based on energy level and orbital shape

• Assume all atoms are at the ground state

Open Ended

What does "ground state" mean when we are talking about atoms?

Electron Configuration

There are rules to follow when filling in orbitals with electrons

Aufbau Principle

Atoms fill the orbitals from lowest energy level to the highest

(start at 1s then move on up)

Hund's Rule

Electrons occupy equal-energy orbitals to have a maximum number of unpaired electrons

(each orbital of one type gets 1 e- before doubling up)

Pauli Exclusion Principle

• Each orbital holds a maximum of 2 electrons

• Electrons in the same orbital must spin in opposite directions

• (one arrow up, one arrow down)

Types of Electron Configurations

We have 3

Orbital Notation

Represent electrons in orbitals using arrows

Electron Configuration Notation

We use a standard notation to describe where the electrons are

Noble Gas Notation

Abbreviates part of the electron configuration by using the noble gas symbol just prior to the element. Then add the rest of the configuration.

Electron Configurations from the Periodic Table

• Energy Level is the ROW

• Orbital shape is the BLOCK

• Number of e- is the COLUMN

• Example: H = 1s¹

• Row 1, s-block, 1st column

Multiple Choice

What is the electron configuration for Lithium?

• 1s²2s¹

• 1s²

• 1s²2s²2p¹

Introduction to the Periodic Table

The Evolution of the Periodic Table

• By 1860, chemists discovered 60 elements and determined their atomic masse

• Elements with similar properties were grouped

• Ex: copper, silver, and gold were the coinage metals

• GOAL: To organize elements in a system that shows similarities while acknowledging difference

Dobereiner's Triads

• In 1829, J.W. Dobereiner classified elements in groups of 3, called triads

• Properties of the middle element were often approximate averages of the 1st and 3rd

Multiple Choice

Given the melting point of Calcium is 842^oC and the melting point of Barium is 727^oC, estimate the melting point of Strontium.

• 1569^oC

• 115^oC

• 792^oC

• 784.5^oC

The First Periodic Table

• Published by Dmitri Mendeleev in 1869

• Organized by increasing atomic mass and similar properties

• Patterns of repeating properties appeared

Henry Mosely

• In 1914, Henry Mosely determined the atomic number of each of the known elements

• He realized that if the element were arranged in the order of increasing atomic # instead of atomic mass, they had a better fit within the table

The Periodic Law

• The physical and chemical properties of the elements repeat in a regular pattern when arranged by increasing atomic number

• Atomic Number

• 1) Identifies the element

• 2) Number of protons and electrons (assuming the atom is neutral)

Determining Atomic Structure

• Periods (Rows 1-7)

• Indicate the number of energy levels

• Groups/Families (Columns 1-18)

• Indicate number of valence electrons

Valence Electrons

• Electrons in outermost energy level

• Indicated by Group # on Periodic Table

• Interact with other atoms to determine many physical and chemical properties of an element

• Electrons that are lost, gained, or shared in a chemical reaction

• Each energy level hold a max amount of electrons

• 1st hold 2 e-, 2nd hold 8 e-, 3rd holds 18 e-

Multiple Choice

How many valence electrons does Aluminum have?

• 3

• 13

• 2

• 4

Structure Determines Properties

• Differences in properties are due to differences in electron arrangement

• Metals: valence electrons are loosely bound to nucleus and easily lost

• Freedom of motion accounts for conductivity

• Nonmetals: valence electrons are tightly held and not easily lost

Modern Periodic Table

• Horizontal Rows = Periods

• Vertical Columns = Groups or Families

• Group 1: Alkali Metals (highly reactive)

• Group 2: Alkaline Earth Metals

• Group 17: Halogens (highly reactive)

• Group 18: Noble Gases (unreactive or inert)

Metal Groups (no, not Metallica)

• Transition Metals (Groups 3-12)

• Unpredictable behavior and properties due to complicated atomic structure

• Lanthanides (Element #58-71)

• Sometimes called "rare earth metals"

• Actinides (Element #90-103)

• All are radioactive, most are unnatural