Download Racheli Taubes Model Lesson on Periodic Table Trends Essential

Racheli Taubes
Model Lesson on Periodic Table Trends
Essential Question:
How does the placement of elements on the periodic table affect and determine trends and
patterns?
Goals:
1. The learner will identify the periodic trends (across periods and down columns) for
ionization energy, electronegativity, atomic radius, and ionic radius
2. The learner will know use Table S in the NYS Regents Reference Book to help her
identify periodic trends.
Materials:
1.
2.
3.
4.
Regents Reference Book Table S
Envelope with strips
Oxygen vs Sulfer chart
Handout booklet with charts
Engage:
Group activity separating characteristics of Oxygen and Sulfur.
Procedure:
What are the properties of Oxygen? What are the properties of Sulfur?
If they are so different, why are they so close to each other on the Periodic Table? Why are they
in the same column?
So they seem to have different physical properties, but if we look on the Periodic Table, they are both
in Group 6A, meaning they have same number of valence electrons, which causes them both to form
similar type of compounds. – form ionic compounds with active metals, form similar compounds like
H2O / S2O and CO2 / CS2
The Periodic Table is very organized and orderly. Elements are placed in the same group and period
because they have similarities. And there similarities cause trends and patterns in the elements.
As we move across a row or down a column of the periodic table see trends within a row or column
and are able to make predictions about the physical and chemical properties of the elements.
Today we will be learning a few trends in the Periodic Table.
1. Ionization Energy
Ionization energy of an atom =amount of energy needed to remove an electron from the atom forming a
cation. The greater the ionization energy, the more difficult it is to remove an electron.
The ease with which electrons can be removed from an atom or ion has a major impact on chemical
behavior.
Na=sodium
ION=atom or molecule in which the protons do not equal the electron, so have charge.
1st Ionization: Na(g) -----> Na+(g) + e2nd Ionization: Na+ (g) --->Na+2(g) + eThe second ionization energy is the energy needed to remove a second electron, etc. I1<I2<I3
WHY is this so? Because as electrons are pulled away the ion becomes more and more positive so the
electrons that are still there need increasingly more energy to be pulled away.
What are the elements with the highest ionization energy? Just eyeball it... NOBEL GASES. Why?
They are so stable and do not give up their electrons easily.
Going DOWN a group:
Atomic #
Symbol
Ionization Energy
3
Li
520
11
Na
496
19
K - potassium
419
37
Rb - rubidium
403
55
Cesium
376
87
Francium
393
Going ACROSS a period:
Atomic #
Symbol
Ionization Energy
3
Li
520
4
Beryllium
900
5
B
801
6
C
1086
7
N
1402
8
O
1314
9
F
1681
10
Ne
2081
11
Na
496
Trends on Periodic Table:
1. Within each column/group, as go down, the ionization energy decreases because the valence
electrons are further away from the positive nucleus and therefore less energy is required to
remove the valence electrons
Smaller atoms vs larger atoms
Smaller atoms have higher first ionization energies. Increasing the effective nuclear charge or
decreasing the distance between the nucleus and electrons increases the attraction between the electrons
and nucleus so hard to remove those electrons. *Get 3 volunteers +1 and act out Nucleus, Electron on
orbital 1 and orbital 2 and see the distance and strength of attraction.
2. Within each row/period, as go from left to right, increase in first ionization energy - - because
closer to “magic number 8” and also because the atomic radii decrease as go across
Irregularities due to stability
*Irregularities between 2A and 3A (there is a decrease). The elements in 2A have a full
s shell and when move across to 3A, the electron is placed on the p shell, which was empty.
* decrease from 5A and 6A (N-O;P-S;As/Arsenic-Se/Selenium) because of Hund's rule that each
electron in the p3 configuration reside in different orbitals which minimizes the electron-electron
repulsion among the 3 electrons. But as soon as add a fourth electron it is paired up and have a
repulsion.
2. Electronegativity
Electronegativity is the attraction of an atom in a molecule for electrons. Simply put, when atoms have
covalent bonds, they share electrons to achieve stable electron configuration, but they don't share the
electrons equally. The greater the electronegativity the greater its ability to attract electrons to itself.
Obviously it is related to ionization energy. HOW? Ionization energy is about the amount of energy
needed to remove an electron from oneself and electronegativity is about the amount of energy needed
to attract an electron. So if an atom likes electrons it will have high ionization energy since it doesn't
want to let go of its electrons and will have high electronegativity because it has a strong ability to
attract the electrons of other atoms to itself.
Down a group
Atomic #
Symbol
Electronegativity
3
Li
1
11
Na
0.9
19
K - potassium
0.8
37
Rb - rubidium
0.8
55
Cesium
0.8
87
Francium
0.7
Across a period
Atomic #
Symbol
Electronegativity
3
Li
1
4
Be
1.6
5
B
2
6
C
2.6
7
N
3
8
O
3.4
9
F
4
10
Ne
--
11
Na
.9
TREND is the SAME as ionization energy: increase in electronegativity from left to right and down
up.
3. Atomic Radius (SIZE)
Atomic radius is the distance between the nucleus and the outer valence electrons.
Down a group
Atomic #
Symbol
Atomic Radius
3
Li
130
11
Na
160
19
K - potassium
200
37
Rb - rubidium
215
55
Cesium
238
87
Francium
242
Across a period
Atomic #
Symbol
Atomic Radius
3
Li
130
4
Be
99
5
B
84
6
C
75
7
N
71
8
O
64
9
F
60
10
Ne
62
11
Na
160
As go down a column, the element has another shell and the valence electrons are further away from
the nucleus and so the size of the atom is larger and the radius is therefore larger.
TREND: As you go across a period from left to right the radius decreases. Why? There is an increase in
effective nuclear charge, which pulls the electrons in and makes the radius smaller.
Electrons within a shell cannot shield each other from the attraction to protons. Since the
number of protons is also increasing, the effective nuclear charge increases across a period. This
causes the atomic radius to decrease.
THIS IS THE OPPOSITE OF IONIZATION ENERGY
4. Ionic Radius
Ionic radius is the distance between ions in an ionic compound. The size of the ion depends on the
nuclear charge, number of electrons, number of orbitals.
ION=atom or molecule in which the protons do not equal the electron, so have charge.
Cations are smaller than parent atoms; anions are larger than their parent atoms. **handout
Cations lose an electron and so an occupied orbital is now empty and also the electron-electron
repulsion is lessened. Simply think about it as getting smaller because it lost something. Metals become
cations because they will lose the electron to be closer to nobel gas. Nonmetals will become anions
because gain the electrons.
Anions receive an electron and so there are more electron-electron repulsions which causes the
electrons to spread out in more space. Simply think about getting bigger because it gained something.
For ions of the same charge, size increases going down a column of the periodic table.
Again, because there is another layer of orbitals, and so it is larger.
When go across the Periodic Table you are going up in atomic number, meaning have more protons to
pull the electrons in closer to the nucleus.
EX: Arrange these atoms and ions in order of decreasing size: Mg2+, Ca2+, Ca
Well, we know that cations are smaller than their parent atoms so the Ca2+ is smaller than the Ca
atom. And because Ca is below Mg2+ in group 2A, Ca2+ is larger than Mg2+.
Ca, Ca2+, Mg2+
Isoelectric series=series of ions that have the same number of electrons
The size decreases with increasing nuclear charge as the electrons are attracted more strongly to
nucleus.
EX: Isoelectric series with 10 electrons: O2-, F-, Na+, Mg2+, Al3+
Put it in order of increasing atomic number, where then you can see that the nuclear
charge (positive) increases as you move through the series. And because they all have
the same number of electrons, the radius of the ion decreases with the increasing nuclear
charge, as the electrons are more strongly attracted to the nucleus.
Al has the highest atomic number (13).
Assessment:
As a summary of Periodic Table Trends, student will do “thumb up/thumb down” activity.
Thumbs up – increase
Thumbs down – decrease
Teacher will ask: moving Left to Right the Atomic Radius _____ (Decreases)
the Ionization energy _____ (Increases)
the Electronegativity _____ (Increases)
the atomic radius ______ (Decreases)
moving from Top to Bottom the Atomic Radius _____ (Increases)
the Ionization energy _____ (Decreases)
the Electronegativity _____ (Decreases)
the atomic radius _______ (Increases)