1. Periodic Properties of Elements
Tadas Rimkus
Krishna Trehan
Rachel Won
Soo Jeon

2. What is the SWBAT???? Objective
Know a brief history of the periodic table
Know the effective nuclear charge
Examine periodic tends in the atomic size, ionization energy, and electron affinity
Examine the sizes of ions and their electron configurations
Explore some of the difference in the physical and chemical properties of metals, nonmetals, and metalloids.
Discuss some periodic trends.

3. 7.1 The Periodic Table
Mendeleev, Meyer, Moseley strived to investigate the possibilities of classifying elements in useful ways.
Mendeleev and Meyer - developed the periodic table on the basis of the similarity in chemical and physical properties. He predicted Ga, Ge, and Sc.
Moseley- established that each element has a unique atom number.
- found out that each element produces X rays of a unique frequency; the frequency increases as the atomic mass increases
Rutherford- proposed the nuclear model of the atom.
Element in the same column of the periodic table have the same number of electrons

5. Periodic Table

6. 7.2 Sizes of Atoms and Ions Electron shells in atoms
- As we move down a column of the P.T, we change the principal quantum number, n, of the valence orbital of the atoms.
-All the orbitals with the same value of n
= shell
Before Bohr, Gilbert N. Lewis had suggested that electrons in atom arranged in spherical shells around the nucleus.

7. Distance from nuclueus
The quantity plotted on the vertical axis is called radial electron density.
Red line = Ne
Blue line = He
The reason that 1s shell in Ne so much closer to the nucleus than the shell in He is..
nuclear charge of Ne is +10, but He is +2.
Because the 1s electrons are the innermost electrons of the atom, they are not shielded from the nucleus very effectively by other electrons.
Probability of finding the electron at a particular distance from the nucleus.
As the nuclear charge increases, the 1s electrons are pulled closer and closer to the nucleus.

8. Atomic sizes
Ionic size – the size of an ion plays an important role in determining the structure and stability of ionic solids.
Periodic trend – metals tend to lose electrons. The size of the atom becomes smaller with the loss of each.
Bonding atomic radius – is based on the distances separating atoms when they are chemically bonded to one another. The radius of the atom is determined by the sizes of the orbitals occupied by the outermost electrons.

9. Sample Problem
Arrange these atoms and ions in order of decreasing size:
Si, Si+2 and C+2.
Solution- Cations are smaller than their parent atoms. So,
The Si+2 ion is smaller than the Si atom. Because Si is below C in group of 4A of the periodic table, Si is larger than C+2. ( Si > Si+2 > C )

10. The distinction between non-bonding and bonding atomic radius
Non-bonding atomic radius
Bonding atomic radius

11. 7.3 Ionization Energy
the amount of energy need to remove an electron from a specific atom or ion in its ground state
1st Ionization Energy
energy needed to remove the 1st electron from an atom

12. Ionization Energy… CONTINUE
ACROSS- increase
As you go across the periodic table, the electrons are closer to the nucleus increasing the energy necessary to remove an electron
DOWN- decrease
As you go down the periodic table, the electrons are farther from the nucleus

13. The ionization energies of the transition elements increase slowly as we proceed from left to right in a period.
Alkali metals have the lowest ionization energy in each row.
Noble gases are the highest.

14. Sample problem
of the following, which has the highest and lowest first
ionization energy?
C and Al
Answer: Highest- C
Lowest- AL

15. 7.4 Electron Affinities
Ionization Energy: energy required to remove an valence electron
Cl(g)  Cl+(g) + e- ΔE = 1251 kJ/mol
See how E is positive?
This is the energy which you must add, in order to remove an electron from Chlorine

16. Electron Affinity
Did you know that atoms can gain electrons to gain a negative charge?!
This energy change that occurs is called ELECTRON AFFINITY because it measures the ATTRACTION between the newly added electron and the atom
It shows how much the atom WANTS the electron!

17. Electron Affinity
So……. FLUORINE wants an electron MORE than LITHIIUM, meaning it would have a LARGER electron affinity!

18. Continue?
The elements on the right, are one electron away from GETTING A FULL OCTET!
They will have a high affinity, so they can gain the last electron.
(The more negative the electron affinity, the greater the attraction for an electron)

19. Continue! Nuclear Charge
the change in the nucleus or the number of protons
ACROSS- increase
DOWN- increase
Atomic Radius
one half the distance from center to center of like atom down-increases
As you go down the periodic table, a new energy level is added increasing the size of the atoms
Ex) of the following, which has the highest and lowest first ionization energy?
Answer: Highest- C
Lowest- AL

20. Periodic Trends Of Electron Affinity
Electron Affinity tends to decrease as we go from left to right!
BUT WHY?!?!?!?!?!?!?!?!?!?!?!
This is because the elements on LEFT side of the Periodic Table of LESS valence electrons.

21. 7.5 metals, nonmetals, and metalloids
Metals are the shiny things we have learned to love and treasure
They are ductile, conduct heat and electricity, and malleable.
They are all solid (except Mercury!) at room temperature.
They can have high melting points, for example Chromium has a melting point at 1900 C.

22. Metals
Metals have low ionization energies (energy required to remove an valence electron).
WHICH MEANS, they tend to form positive ions EASILY!!!!!!!

23. Continue…
Periodic Trend - Nonmetals tend to gain electrons. The size of the atom becomes larger with the addition of each consecutive electron. The nuclear charge has less of a pull on the electrons with the gain of each electron due to greater electron repulsion!