Please read the information below the periodic table!!!

Section 6.1 and 6.2

Mrs. Johnson's suggestions:

1. Complete a well-labeled periodic table on one of the blank periodic tables in your yellow pages, showing where the representative elements, metals, non-metals, metalloids, groups, periods, diatomic elements, ionic charges, ionic exceptions, akali metals, alkaline earth metals, transition metals, inner transition metals, halogens, and noble gases are.

2. Two things you may want to know are who Mendeleev is, and what the Periodic Law is. (Extra credit questions!)

Mendeleev: A Russian scientist credited with creating the foundation of the modern periodic table. See an article about him here.

Periodic Law: When elements are arranged in order of increasing atomic number, there is a periodic repetition of their physical and chemical properties.

Section 6.3 Periodic Trends - for Unit 2, you need to know Atomic and Ionic size (radius).

This is often a section that students either find confusing, or they simply memorize the trend but have no idea why it happens. You must understand why the trends occur if you wish to do well on the test!!

Understanding this diagram will help you to understand the trends. I will list them below with a short explanation.

The two major reasons why the trends occur are:

1) the shielding effect - as more and more energy levels are added around a nucleus, the nucleus is less and less able to 'grab' the outside (valence) electrons tightly. We say that the inner 'rings' or energy levels of electrons shield the outside ring from the inward pull of the nucleus. As a consequence, the valence is held more loosely, causing a larger atom. Shielding effect increases down a group only.

2) nuclear charge - as more and more protons are added across the periods of the table, the nucleus becomes more and more positively charged. As nuclear charge increases, the nucleus of the atom is able to pull it's electrons more tightly inward, creating a smaller atom. Nuclear charge increases from left to right across periods only.

To summarize these reasons: 

Down a group: shielding increases, nuclear charge also increases, but the shielding effect is stronger than the nuclear charge effect. Net result: larger atoms.

Across a period: shielding is held constant (across any period all atoms have the same number of energy levels) but nuclear charge increases. Net result: smaller atoms.

Periodic Trends FAQ

Interesting fact (and you will NEED this fact for the lab!) The acidity of every element increases going from the left to the right across a period. The acidity of the elements decreases as you go down a group.

Atomic radius (size) increases down a group. Why?

1. Because each element in the group adds an energy level as we go down the group. Each additional 'ring' of electrons makes an element larger.

2. Because the shielding effect increases. The nucleus has more difficulty attracting/holding the outside (valence) electrons because there are so many inner layers of electrons "getting in the way". 

Ionic radius (size) increases down a group. Why?

This happens for the same reasons that atomic radius increases down a group. 

Ionization energy (the energy needed to take an electron away from an atom) decreases down a group. Why?

As energy levels are added down a group, shielding increases. This makes it harder for the nucleus of the atom to hang on to outer electrons. Less energy is needed to remove an electron from an atom with a lot of shielding energy levels. When an electron is removed, the atom becomes an ion. This is why it is called "ionization energy". It is the energy needed to make an atom into an ion.

Electronegativity (the ability of an atom to attract electrons when it is part of a compound, like HCl for example) decreases down a group. Why?

This is based on several factors, including an element's ionization energy. The most electronegative element on the table is fluorine - with 9 protons and only two energy levels, it does not experience much shielding, but does experience a fairly strong nuclear force for it's size. The least electronegative element on the table is Cesium - with 55 protons and six energy levels, it experiences quite a lot of shielding, which is stronger than its nuclear force.

Why do atoms get smaller as we go across a period?

Remember that across a period shielding never changes. As a constant, we can ignore it. What does change is the nuclear force. As we go across a period each element has one more proton than the last, so each new element experiences a stronger nuclear force, being pulled inward/becoming smaller.

Why do ionization energy and electronegativity increase across a period?

Again, shielding is a constant, so we can ignore it. But as the nuclear force gets stronger moving across toward the halogens, the ability to attract electrons - whether you are alone or in a compound - also gets stronger. These concepts (ionization energy and electronegativity) are strongly related to each other.

For more on trends: http://www.scribd.com/doc/496624/Periodic-Trends