Transcript Chapter 14 Lesson 2

Chapter 6
Lesson 3
“Periodic Trends”
The Big Idea…
Nuclear
Charge – the
effect protons of an atom
have on its size and shape.
I. Trends in Atomic Size
A. Atomic Radius
1. Def – one half of the distance between the
nuclei of two like atoms in a diatomic molecule.
-radius = from center to
outer edge
2. Measured in picometers (pm)
-1 x 10-12 m
3. Measured in angstroms (Å)
-1 x 10-10 m
3. The distance between the nuclei of two
covalently bonded flourine atoms is 128
pm. What is the atomic radius of 1
fluorine atom? 64 pm
4. The distance between the nuclei of two
covalently bonded nitrogen atoms is 142
pm. What is the atomic radius of 1
nitrogen atom in nanometers? .0710 nm
5. The distance between the nuclei of two
covalently bonded iodine atoms is 276
pm. What is the diameter of 1 iodine
atom in pm? 276 pm
B. Group/Family Trends
1. Size increases as you move down a column.
-Why?
electrons are added to energy levels
farther away from the nucleus.
-How can electrons get into energy levels
away from the nucleus?
electrons are shielded from the positive
nuclear charge by 1s electrons…2s electrons…
2p electrons… and so forth.
2. Shielding demo
In Class Assignment
1. Place the following atoms in order from
smallest to largest in terms of their atomic
radius.
Br, I, Be, He, Rn
He < Be < Br < I < Rn
2. Why is a sodium atom smaller than a
potassium atom?
Because potassium has more electrons and these electrons fill
energy levels that are farther away from the nucleus of a
sodium atom.
C. Periodic Trends
1. Atomic radii decreases as you move
from left to right.
Why?
electrons are added to the same
principle energy level. Nuclear charge
pulls electrons in closer to the nucleus.
Group #1
Group #2
Period #2
Period #3
In Class Assignment
1. How would you describe the atomic
radius of a period 2 alkaline earth metal
compared to a period 4 alkaline earth
metal? The atomic radius of the period 2 a.e.m.
would be smaller than the period 4 a.e.m.
2. How would you describe the atomic
radius of a period 3 alkali metal and a
period 3 halogen?
The atomic radius of a period 3 a.m. would be
larger than a period 3 halogen due to nuclear
charge.
II. Trends in Ionization Energy
A. Def – the amount of energy required to
overcome the attraction of the nuclear
charge and remove an electron from an
atom.
Energy
e-
e-
e-
ee-
Na
e-
ee-
e-
e-
e-
1+
e-
e-
e-
ee-
Na
e-
ee-
e-
e-
1. 1st Ionization Energy
-removing the 1st electron
2. 2nd Ionization Energy
-removing an electron from a 1+ ion.
3. 3rd…4th…5th… and so on
**Note table 6.1
B. Group Trends
1. Ionization energy decreases as you move down
a group.
Why?
The farther electrons are from the nuclear charge,
the easier they are removed.
C. Period Trends
1. Ionization energy increases as you move from
left to right.
Why?
The nuclear charge increases, but the electron
shield does not.
**Remember noble gases do not want to give up
electrons**
III. Trends in Ionic Size
A. Ion Review
1. metals  low ionization energy
2. non-metals  high ionization energy
**How does losing/gaining an electron affect the
size of an ion???**
B. Cations and Anions
1. Cations = always smaller than their neutral
atoms
Why?
loss of an electron causes the nuclear charge
to increase. Thus the remaining electrons are
pulled in farther.
ee-
e-
2. Anions = always larger than their
neutral atom.
Why?
The increase of another electron
causes the nuclear charge to decrease.
Thus the size of the ion increases.
Increases
C. Group Trends
1. Ionic size increases as you move down a
Anions decrease
group on the periodic table.
Cations decrease
D. Periodic Trends
1. Ionic size decreases from left to right for both cation and anions
1. How does the ionic radius of sodium compare to that of cesium?
Cesium is larger
2. How does the ionic radius of boron compare to that of fluorine?
Fluorine is larger
IV. Electronegatvity
A. Def – the tendency for an atom to attract
electrons from another atom.
1. similar to magnetism
2. Electronegativity is a relative value based on
ionization energy.
3. Noble Gases are not included…Why???
-they don’t form compounds
B. Electronegativity Trends
1. E-negativity decreases as you move down a
column.
2. E-negativity increases as you move across a
period.
-Fluorine = most electronegative element
-Cesium = least electronegative element
3. Metals = low e-negativity
Non-metals = high e-negativity
E-negativity = tug of war
Fluorine or Cesium? Fluorine
Calcium or Sulfur? Sulfur
Oxygen
Oxygen or Magnesium?
Oxygen or Nitrogen? Oxygen
4. E-negativity values help predict types of
bonds.