• A measure of the size of an atom – how close it lies to its neighbor.
• Half the distance between 2 adjacent, identical atoms (crystal or molecule).
• The radius for any atom can vary depending upon the substance it’s in.

• Radii increase down a family – adding a principal energy level which is larger and more electrons internally to shield the valence electrons.
• Radii decrease across a period – same principal energy level with a larger positive charge in the nucleus.

Which has the largest atomic radius: carbon, fluorine, beryllium or lithium?
•Lithium
•has the same outer energy level as the others
•but its nucleus has the smallest positive charge

• Form when atoms gain or lose electrons.
• If an atom gains electrons, it becomes a negative ion.
• If an atom loses electrons, it becomes a positive ion.
• Metals are more likely to lose electrons while nonmetals are more likely to gain electrons.

• When an atom forms an ion, its radius will always change size.
• Negative ions become larger due to a greater electrostatic repulsion among the valence electrons.
• Positive ions become smaller due to a smaller electrostatic repulsion among valence electrons and the fact the outer orbital often becomes empty in the process.

• As you move across a period, the positive ions (on the left) and the negative ions (on the right) decrease in size just like the atomic radii do.
• But negative ions (on the right) will always be bigger than the positive ions (on the left).
• As you move down a group, ionic radii increase – both positive & negative.

1. For each of the following pairs, predict which atom is larger.
a. Mg, Sr d. Ge, Br b. Sr, Sn e. Cr, W c. Ge, Sn
2. For each of the following pairs, predict which atom or ion is larger.
a. Mg, Mg 2+ c. Ca 2 + , Ba 2 + e. Na +, Al 3 + b. S, S 2 d. Cl , I -

1. For each of the following pairs, predict which atom is larger.
a. Mg, Sr d. Ge , Br b. Sr , Sn e. Cr, W c. Ge, Sn
2. For each of the following pairs, predict which atom or ion is larger.
a. Mg , Mg 2+ b. S, S 2 c. Ca 2 + , Ba e. Na + , Al 3 +
2 + d. Cl , I -

• The energy required to remove an electron from a gaseous atom.
• The energy required to remove the first electron is called the first ionization energy and so forth.
• Group 1A has lowest ion. energy w/ group
8A having the highest as you move across a period.

• Successive ionization energies ALWAYS require more energy because there are fewer electrons left to repel one another.
• At some point this will take a huge jump because you’ve reached inner level electrons.
• First ionization energies increase as you move across a period because the nuclear charge increases so the nucleus holds the electrons more tightly.
• However they decrease as you move down a family due to greater atomic size, so electrons are further away from the nucleus.

For each of the following pairs, predict which atom has the higher first ionization energy.
a. Mg, Na d. Cl, I b. S, O e. Na, Al c. Ca, Ba f. Se, Br a. Mg , Na d. Cl , I b. S, O e. Na, Al c. Ca , Ba f. Se, Br

• Atoms tend to gain, lose or share in order to acquire a full set of 8 valence electrons.
• “Eight is great, but two will do.”
• For whatever reason, filled s and p orbitals at any energy level is extremely stable.

• Indicates the relative ability of an atom to attract or gain electrons in a chemical bond.
• Calculated using several factors & always have a value of 3.98 Paulings or less.
• Greatest electronegativity = fluorine (3.98)
• Least electronegativity = franicum (.70)
• It has the same trend of ionization energy which is losing electrons.
• Decreases as you move down a group and increases as you move left to right across the periodic table.