1. Periodic Trends 6.3

2. Atomic Radius
The distance between an atom’s nucleus and its valence electrons.
How closely an atom lies to a neighboring atom
Size of the atom varies from substance to substance

3. Trends within Periods
Atomic Radii DECREASE as you move left to right across a period
Because of Increased nuclear charge(total charge of all protons in the nucleus)
Increased nuclear charge pulls the outermost electrons closer to the nucleus & decreases the atomic radius.

4. Trends within Groups Atomic Radii INCREASE as you move down a group.
Nuclear charge increases and electrons are added to higher energy levels.
Moving down a group:
The Outermost orbital increases in size
Increasing principal energy level

5. Examples
Which element has the smallest atomic radius? Largest atomic radius?
Iodine (I)
Bromine (Br)
Fluorine (F)
Chlorine (Cl)
Fluorine
Iodine

6. Ionic Radius Ion- an atom that gains or loses electrons
When atoms lose electrons they form positive ions and become smaller
The electron lost will always be a valence electron
Loss of valence electrons may leave an empty outer orbital which results in a smaller radius
Repulsion between fewer electrons decreases allowing them to be pulled closer to the nucleus

7. When atoms GAIN electrons they form negative ions and they always become larger.
Increase of an electron to an atom increases the repulsion between the valence electrons forcing them to move farther apart.
This equals larger radius

8. Ionic Radius within Periods
Size of the positive ions gradually decreases
Beginning in group 5A or 6A the size of the much larger negative ions also gradually decreases.
Ionic Radii DECREASE across periods

9. Ionic Radius within Groups
As you move down a group an ion’s outer electrons are in higher principal energy levels resulting in a gradual increase in ionic size
Ionic Radii increase as you move down a group

10. Ionization Energy
The energy required to remove an electron from an atom
“How strongly an atom’s nucleus holds on to its valence electrons”
High IE indicates atom has a strong hold on its electrons
Low IE indicates an atom loses its outer electron easily

11. Ionization Energy
Energy required to remove the 1st electron is the first ionization energy.
Energy required to remove the 2nd electron is the second ionization energy.

12. Ionization Energy within Periods
INCREASE as you move from left to right across a period
The increased nuclear charge of each successive element produces an increased hold on the valence electrons.

13. Ionization Energy within Groups
DECREASE as you move down a group
Occurs because atomic size increases as you move down a group
With the valence electrons farther from the nucleus less energy is required to remove them

14. Octet Rule
Atoms tend to gain, lose, or share electrons in order to acquire a full set of eight valence electrons.
Elements on the right side of the periodic table tend to gain electrons in order to acquire the 8 valence electrons. (Form negative ions)
Elements on the left side of the periodic table tend to lose electrons and form positive ions.

15. Electronegativity of an Element
Indicates the relative ability of its atoms to attract electrons in a chemical bond
Noble Gases are not assigned values
Fluorine is the most electronegative
Fr & Cs are the least
In a chemical bond the atom with the greater electronegativity more strongly attracts the electrons

16. Electronegativity Trends within Periods & Groups
INCREASES as you move across a period
DECREASES as you move down a group
The lowest electronegativities are found at the lower left side
Highest are found at the upper right side

17. Example Problems
Which element has the highest electronegativity? Lowest?
N- Nitrogen
P- Phosphorus
As-Arsenic
Sb-Antimony
Bi- bismuth
N=highest
Bi= Lowest

18. Homework Problems
Pg. 175
#56, 57, 59, 62, & 63