1. The Periodic Table

2. Periodic Trends Atomic Radius (radius of an atom)
atomic radius is measured (just like any other radius)
the atomic radius of an atom increases each period and decreases from group to group
The shielding effect is how the atomic radius of an atom decreases from left to right on the periodic table
Ionic Radius (radius of an ion)
ions have a different radius than their non-ionic counterparts.
Example: H+ has a different radius than H because it has no electrons
protons pull electrons towards the center, preventing them from escaping orbit
electrons pull away from the center because they are in constant motion
cations are smaller than their neutral form because they have more protons than electrons (more protons means more force keeping electrons in)
anions are larger than their neutral form (more electrons means they can orbit further away)

3. Ionic Radius

4. Periodic Trends Ionization Energy and Electronegativity
Ionization Energy is the amount of energy required to lose an electron from a regular atom (how easily it will give up an electron)
Electronegativity/Electron Affinity is an atom’s ability to draw electrons toward it when it has a chemical bond
ionization energy and electronegativity share the same patter on the periodic table
They both increase as the group number increases, but decreases as the period number increases

5. Periodic Trends

6. Valence Electrons
Valence electrons are the electrons that occupy the highest energy level (1-7) of the electron cloud.
The highest energy level is equal to the period
Example:
Carbon has the electron configuration 1s2 2s2 2p2.
Carbon belongs to period 2, so its highest occupied energy level is 2.
Carbon’s 2nd energy level has 2 electrons in sublevel s and 2 electrons in sublevel p, therefore carbon has 4 valence electrons

7. Valence Electrons Group 1 has 1 valence electron
Group 2 has 2 valence electrons
Groups 3-12 (transition metals) vary
Group 13 has 3 valence electrons
Group 14 has 4 valence electrons
Group 15 has 5 valence electrons
Group 16 has 6 valence electrons
Group 17 has 7 valence electrons
Group 18 has 8 valence electrons

8. Formation of Ions Ionization – The formation of an ion
Groups 1-14 elements make cations
These ions ALWAYS have 0 valence electrons
Groups elements make anions
These ions ALWAYS have 8 valence electrons
Group 18 elements (noble gases) don’t make ions because they already have 8 valence electrons
The charge of an ion can be determined from the number of valence electrons
1 valence e- makes a +1 charge because it had to lose 1 electron to equal zero
6 valence e- makes a -2 charge because it had to gain 2 electrons to equal 8
4 valence e- makes either a +4 or a -4 charge, depending on the element

9. Formation of Ions

10. Properties of Groups
Group 1 Alkali Metals: Most metallic and reactive group
form +1 cations
Group 2 Alkaline Earth metals: Very reactive group
form +2 cations
Block D Transition Metals: Semi-reactive
form cations with different charges
Block F Lanthanides and Actinides: Radioactive and extremely heavy metals
Group 16 Chalcogens: Very reactive and electronegative
form -2 anions
Group 17 Halogens: Extremely reactive and electronegative group
form -1 anions
Group 18 Noble Gases: Completely unreactive because of complete valence shell

11. Types of Elements
There are 3 types of elements: metals, metalloids and nonmetals
1) S, D and F block Metals: Malleable (can change shape without breaking), shiny, conducts electricity
2) P Block Metalloids: Show both metallic and non-metallic character. Metalloids touch the staircase of the periodic table
3) P Block Nonmetals: dull, brittle, and poor at conducting electricity because they hog electrons
Note: metallic character trend follows the same pattern as atomic radius

12. History
Mendeleev was the scientist responsible for seeing the pattern of properties of the elements in the periodic table
he was the first to organize those elements into periods and groups.
Elements within the same group have the same physical and chemical properties.
When he organized his version of the table, there were holes in it because there were only 60 out of 110 known elements at the time. After he died, the holes were later filled by newly discovered elements, verifying what he knew to be correct:
you can predict how an atom will behave by knowing where its atomic number falls on the periodic table.
Periodic law – Both physical and chemical properties of the elements are functions of their atomic numbers.