1. Periodic Trends

2. Organizing the Elements
Chemists used the properties of elements to sort them into groups.
JW Dobereiner ( ) grouped elements into triads according to similar properties such as their willingness to react with metals.
Mendeleev (1869)originally arranged the elements in his periodic table in order of increasing atomic mass.
From his organization of the elements that had been discovered at the time he was able to predict the properties of elements that were yet to be discovered.
When Mendeleev developed his table the structure of an atom was not yet known.
In 1913 Henry Moseley determined the atomic number for each known element.
Organizing the Elements

3. When elements are arranged in order of increasing atomic number, there is a periodic repetition of their physical and chemical properties.
The current periodic table has seven rows or periods which correspond to each energy level. There are more elements in the higher numbered periods because there are more orbitals in the higher energy levels.
The elements within a column, or group, have similar properties
The properties of the elements within a period change as you move across a period from left to right; however the pattern or properties repeats as you move from one period to the next
The Periodic Law

4. Electron Configuration in Groups (Things We Know)
Nobel gases
Are the elements in Group 8A
Have a filled outer shell (meeting the octet rule)
Are used in “short-hand” electron configuration
Representative Elements
Are the elements in Groups 1A-7A
Their group # is equal to the number of electrons in the highest occupied energy level
Transition Elements
Are the elements in Group B
Transition metals have electrons in d orbitals
Inner transition metals have electrons in f orbitals
Electron Configuration in Groups (Things We Know)

5. Periodic Table Showing Electron Configuration
Representative Elements
Noble Gases
Representative Elements
Transition Elements
Inner Transition Elements
Periodic Table Showing Electron Configuration

6. Classification of Elements
Representative Elements are further broken down into:
Alkali Metals-First column on left
Alkaline Earth Metals-Second column on left
Other Metals-between transition metals and stairsteps
Metalloids-stairsteps (Boron, Silicon, Germanium, Arsenic, Antimony, Tellurium, Astatine)
Nonmetals-to the right of stairsteps, before noble gases
Halogens- Group 7A, column right before noble gases
Classification of Elements

7. In general, atomic size increases from top to bottom within a group and decreases from left to right across a period
As the atomic number increases within a group, the charge on the nucleus increases and the number of occupied energy level increases.
The increase in positive charge draws electrons closer to the nucleus.
The increase in the number of occupied energy levels shields electrons in the highest
The atomic radius is one half the distance between the nuclei of two atoms of the same element when the atoms are joined.
The distances are extremely small so they are measured in picometers (pm=1012m)
Atomic Radius
Trends in Atomic Size

8. Trends in Ionization Energy
First ionization energy tends to decrease from top to bottom within a group and increase from left to right across a period
Ionization energy- The energy required to remove an electron from an atom.
This energy is measured when an element is in its gaseous state.
The energy required to remove the first electron from an atom is called the first ionization energy producing a cation with a +1 charge.
The second ionization energy is the energy required to remove an electron from an ion with a +1 charge producing an ion with a +2 charge
The third ionization energy is the energy required to remove an electron from an ion with a +2 charge producing an ion with a +3 charge.
Trends in Ionization Energy

9. Cations are always smaller than the atoms from which they form
Cations are always smaller than the atoms from which they form. Anions are always larger than the atoms from which they form.
For metals, when they lose an electron it increases the attraction between the remaining electrons and the nucleus, drawing the remaining electrons in and making the ion smaller.
The metals in the representative elements also tend to lose their outermost electrons decreasing the number of occupied energy levels.
The trend is the opposite for nonmetals. These ions are larger than the atom. As the number of electrons increases the attraction of the nucleus for a particular electron decreases.
Trends in Ionic Size

10. Trends in Electronegativity
In general, electronegativity values decrease from top to bottom within a group. For representative elements, the values tend to increase from left to right across a period.
Electronegativity- The ability of an atom of an element to attract electrons when the atom is in a compound. Expressed in units called Paulings.
The least electronegative element is cesium. It has a tendency to lose electrons and form positive ions.
The most electronegative element is fluorine. It has a really strong tendency to attract electrons, when it is bonded to any other element it either attracts the shared electrons or forms a negative ion.
Trends in Electronegativity