Mrs. Ermann

Regents Chemistry

Basic Periodic Table

And

The Quantum Mechanical Model of the Atom

Regents Chemistry

Periodic Table/Electron Configuration Notes

Periodic Table Basics:

Elements in a column are called a “group.” The number on top of the column is the “group number.”

Some groups have specific names. The ones you have to know are:

Group 1 – Alkali Metals – very reactive metals

Group 2 – Alkaline Earth Metals

Groups 3 through 12 are called the Transition Metals

Group 17 – Halogens – very reactive non-metals

Group 18 – Noble Gases – these elements are NOT REACTIVE – they will not bind to another element to make a compound.

Each row on the table is known as a “period”

Period numbers are listed on the left hand side of the table.

The periodic table is divided into three types of elements:

1. Metals

2. Non-metals

3. Metalloids (also called semimetals)

They are located as shown:

The Lanthanides and Actinides are the two rows on the bottom.

Elemental Properties

  • Properties of metals
  • Have luster (are shiny)
  • Malleable (can be hammered into shapes)
  • Ductile (can be pulled into wires)
  • Good conductors of heat and electricity due to low specific heat
  • Form cations (lose electrons)
  • High melting points
  • Properties of nonmetals
  • Dull
  • Can be brittle, hard, or soft
  • Poor conductors of heat and electricity
  • Form anions (gain electrons)
  • Low melting points (some are liquids or gases at room temp)
  • Properties of metalloids
  • Have properties of both metals and non-metals

Reactivity and Metallic Character

  • Metalslose electrons and form cations when they react
  • Nonmetalsgain electrons and form anions when they react
  • From left to right on the table, metallic character DECREASES (electrons are not easily removed, nonmetals are on the right side)
  • From top to bottom on the left side of the table, metallic character INCREASES (electrons are more easily lost)

Periodic Table Development

  • Dimitri Mendeleev, a Russian Chemist, was the first to create begin to create the modern periodic table.
  • He organized the elements by their chemical and physical properties and atomic mass. He left blank spaces for undiscovered elements where he predicted their properties. However, not all the elements fit in the table correctly based on the chemical and physical properties. Why this happened was a mystery until protons were discovered in the early 1900’s.
  • Henry Moselyreorganized the Periodic Table to list elements in order by atomic number (this is the modern periodic table you see today) and all elements fit in perfectly. This is the Periodic Law:
  • Periodic Law = when elements are arranged in order of increasing atomic number, there is a periodic repeating pattern to their properties

The Bohr Model of the Atom

There are many models of the atom that have developed over time. In 1915, the physicist, Neil Bohr presented his model of the atom based on his research. You are probably familiar with it from middle school. It looks like this:

His model has the following:

  1. A very small nucleus in the middle of the atom which contains a positive charge (Bohr did not know about neutrons – they had not been discovered yet.)
  2. The electrons are located outside the nucleusin shells (or orbits) moving around the positively-charged nucleus similar to the planets orbiting the Sun.
  3. The shells (or orbits) have a specific amount of energy associated with them and they are called ENERGY LEVELS.

Main points of the Bohr Model – VERY IMPORTANT – KNOW THESE!

i)Electrons move around the nucleus in specific shells (or orbits) just like the planets orbit around the sun. The shells that have a set size and energy.

ii)The lowest energy shell (or orbit) is located closest to the nucleus.

iii)Energy is absorbed or given off when the electron moves from one orbit to another.

iv)THE ELECTRONS CANNOT BE LOCATED BETWEEN SHELLS! THEY MUST BE IN A SHELL!

Some of the types of energy that might enter the atom:

  1. Light
  2. Heat
  3. Sound
  4. Electrical

NOTE: THE BOHR MODEL HAS BEEN DISPROVEN, BUT IT IS A USEFUL MODEL TO UNDERSTAND THE STRUCTURE OF THE ATOM!

The Quantum Mechanical Model of the Atom

Bohr’s model is important, but it wasn’t correct.

So… a new model was needed… the quantum mechanical model which is based on many complicated math equations. Scientists would put in different variables and plot each answer. When they were finished, the following were the shapes of the orbitals:

Atomic orbital or electron cloud = region of high probability (likelihood) of finding an electron with a particular energy (these have the above shapes).

Orbital Diagrams and Electron Configurations

Electron configurations (how electrons are arranged around the nucleus) give us insight into an element’s properties and chemical behavior.

Principles that help us understand electron configurations:

  • Aufbau Principle: Electrons occupy orbitals of lowest energy first
  • Pauli Exclusion Principle: No more than 2 electrons can occupy any orbital; if 2 electrons are to be in one orbital, they must have opposite spins
  • Hund’s Rule: When filling p, d or f orbitals, they will half fill first (receiving 1 electron), before any will double up

Principle energy levels actually are made up of sublevels containing orbitals with different shapes and energies. Each orbital can hold 2 electrons (as long as the electrons have opposite spins)

There are 4 types of sublevels (s, p, d, and f), each made up of orbitals with their own shapes

  • s sublevel has 1 orbital
  • p sublevel has 3 orbitals
  • d sublevel has 5 orbitals
  • f sublevel has 7 orbitals

Electron configuration = system of showing the arrangement of electrons in an atom using letters and numbers to represent electrons in the orbitals

Writing Electron Configurations and Drawing Orbital Diagrams

  • The key is to apply Aufbau, Pauli, and Hund
  • Electron configuration
  • Large number on left = principle energy level
  • Letter = sublevel (orbital)
  • Superscript number = number of electrons in that sublevel
Principal Energy Level / Number of Sublevels / Type of Sublevels / Maximum Number of Electrons
1 / 1 / 1s / 2
2 / 2 / 2s, 2p / 8
3 / 3 / 3s, 3p, 3d / 18
4 / 4 / 4s, 4p, 4d, 4f / 32

Directions:

  1. Find your atom's atomic number on the periodic table.
  2. Memorize the orbitals that are on the periodic table:

  1. Starting with the 1s block, and moving from left to right, write down all the orbitals and their respective electrons until you get to your element.

Example:

Write the electron configuration for Magnesium using the periodic table.

Magnesium is located in Period 3, Group 2. Therefore, the electronic configuration would be:

Write the electron configuration using the periodic table for Nickel.

Electron Configuration Homework

Write the electron configuration for the following atoms:

  1. Oxygen: Example: 1s2 2s2 2p4
  1. Argon
  1. Bromine
  1. Aluminum
  1. Cobalt
  1. Calcium
  1. Potassium
  1. Nickel
  1. Silicon
  1. Beryllium
  1. Titanium
  1. Chromium
  1. Sodium
  1. Phosphorous
  1. Copper

Orbital Notation Homework

Write the orbital notation for the following atoms:

  1. Oxygen: example:

1s / 2s / 2px / 2py / 2pz
↑↓ / ↑↓ / ↑↓ / ↑ / ↑
  1. Argon
  1. Fluorine
  1. Aluminum
  1. Magnesium
  1. Calcium
  1. Potassium
  1. Carbon
  1. Silicon
  1. Beryllium
  1. Sodium

Noble Gas Core Configuration (Noble Gas Abbreviation)

This is becoming a lot of work writing down these ever-expanding electron configurations!The more electrons we have in the atom, the longer the electron configuration.We have a method we can use to simplify writing these configuration as we become more familiar with them.If we look closely, we notice that the electron configuration for rubidium is the same as the previous element, argon, with a single 5s electron added on.For the electron configuration of argon, let us simply write [Kr].The simplified electron configuration for rubidium then becomes:

Rubidium / [Kr]5s1

We can do this for any element, BUT, we must useonly noble gasesin the brackets.This is called thenoble gas core configuration.In this method of writing electron configurations, the last noble gas before we get to the element of interest is the noble gas we put into the brackets.For instance, for the element aluminum we write

Sulfur / [Ne]3s23p4
Chlorine / [Ne]3s23p5
Argon / [Ne]3s23p6
Potassium / [Ar]4s1
Calcium / [Ar]4s2

We may NOT use any element in the brackets, only noble gases.

Noble Gas Notation Homework

Write the Noble Gas notation for the following atoms:

  1. Chlorine: Example: [Ne] 3s2 3p5
  1. Cadmium
  1. Francium
  1. Arsenic
  1. Bismuth
  1. Bromine
  1. Chromium
  1. Nitrogen
  1. Phosphorous

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Electron Configuration

  1. Write the electron configuration for the following atoms:

B. Write the noble gas configuration for the following elements.

C. Write the orbital notation for the following elements.

  1. Nitrogen

A. ______

B. ______

C.

  1. Chlorine

A. ______

B. ______

C.

  1. Titanium

A. ______

B. ______

C.

  1. Beryllium

A. ______

B. ______

C.

  1. Cobalt

A. ______

B. ______

C.

  1. Calcium

A. ______

B. ______

C.

Mixed Electron Configuration

  1. Electron configuration for Fluorine
  1. Orbital notation for Boron
  1. Electron configuration for Helium
  1. Noble gas notation for Strontium
  1. Orbital notation for Lithium
  1. Noble gas notation for Xenon
  1. Electron configuration for Neon
  1. Orbital notation for Iron
    Electron Configurations for Ions

Ions are atoms that have a charge due to the loss or gain of electrons.

Metals lose electrons and become positively charged.

Non-metals gain electrons and become negatively charged.

Since the number of electrons for ions are different than for neutral atoms, the electronic configuration and orbital diagrams will be different as well.

EXAMPLE:

1.Write the electron configuration and orbital diagram for Na+1.

Since sodium LOST one electron, we will indicate one LESS electron:

2.Write the electron configuration and orbital diagram for F-1.

Since fluorine GAINED one electron, we will indicate one MORE electron:

1.Write the electron configuration and orbital diagram for the following ions and indicate the noble gas with which they are isoelectronic:

  1. Calcium with a +2 charge
  1. Lithium with a +1 charge
  1. Phosphorous with a -3 charge
  1. Sulfur with a -2 charge
  1. Chlorine with a -1 charge
  1. Aluminum with a +3 charge
  1. Strontium with a +2 charge
  1. Oxygen with a -2 charge

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