1. Chapter 5 Periodic Law
Chapter 4 Periodicity

2. Development of the Periodic Table
About 70 elements had been discovered by the mid-1800’s, but no one had found a way to relate the elements in a systematic, logical way.

3. Dimitri Mendeleev
He develop the 1st
periodic table of the
elements.
Arranged elements in order of increasing
atomic mass and created columns with elements having similar properties.

4. Mendeleev left blank spaces in the table because there were no known elements with the appropriate properties and masses.

5. Mendeleev and others were able to predict the physical and chemical properties of the missing elements. Eventually these elements were discovered and were found to have properties similar to those predicted. There were many exceptions in his table, however.

6. Henry Moseley (1887 – 1915)
In 1913, arranged elements in order of increasing atomic number thus reversing the order of the elements and correcting the drawbacks found in Mendeleev’s table.

8. The Elements Song
Neat-o Animation

9. Periodic Law
Periodic law states the physical and chemical properties of the elements are periodic functions of their atomic number. In other words, when the elements are arranged in order of atomic number, elements with similar properties appear at regular intervals.

10. A Group or Family is a column on the periodic table
A Group or Family is a column on the periodic table. Elements in the same column have similar chemical properties.

11. 2 conventions for numbering:
1-18
A/B elements

13. Blocks Main Group Elements Transition Metals Inner Transition Metals
C. Johannesson

14. Group A Elements
Group A elements all have electrons in the outer s, or s and p orbitals. These are known as representative elements. The group number indicates the number of valence, or outer shell, electrons except with helium which has 2.
Examples:
IIA - Ca (20) 1s22s22p63s23p64s2
VIA – S (16) 1s22s22p63s23p4

15. Group 18 (VIIIA)
Group 18 (VIIIA) elements are the noble gases with 8 valence electrons, except helium which has 2. Noble gases are inert (nonreactive) in nature. They do not form ions.
Group 18 Noble Gases
Unreactive
Also called inert gases
No stable compounds on 3 of the noble gases, He, Ne, Ar, have ever been prepared
Few compounds of the others, but very difficult to do
Stability of element increases compared to the stability of the element without that electron configuration
Uses- Ne, Ar signs; He blimps, some hot air balloons

16. Group 18 (VIIIA)
Full s & p orbitals in the highest principal energy level
Electron configuration very stable, making them inert
When other atoms of other elements gain or lose electrons in reactions, they achieve electron configuration of noble gases

17. Group B Elements
Group B elements or transition elements (d block) have electrons in their outer d orbitals. The have varying number of valence electrons.
Example: Zn (30) 1s22s22p63s23p64s23d10

18. Lanthanides and Actinides Series
Lanthanides and Actinides elements (f-block) have electrons in their outer f orbitals. These elements have varying numbers of valence electrons.
Example: Nd (60)
1s22s22p63s23p64s23d104p65s2 4d105p66s25d14f3

20. Hydrogen
No group number
Only element in family
Most common element in the universe
Very reactive
Compounds of H very common- H2O
Found in proteins, carbs, and fats with C and O

21. Periodic Law
When elements are arranged in order of increasing atomic #, elements with similar properties appear at regular intervals.
C. Johannesson

22. What trends are found in the periodic table?
Trends on table occur vertically and horizontally
Group 1
Li increasing reactivity with H2O Na K
Knowing the trends enables you to predict chemical behavior.

23. You know from the quantum mechanical model that an atom does not have a sharply defined boundary that sets the limit of its size.
Therefore, the radius of an atom cannot be measured directly. There are, however, several ways to estimate the relative size of atoms.
The atomic radius is one-half of the distance from center to center of 2 like atoms.
Atomic radius- ½ the distance

24. Group Trends
Atomic size generally increases as you move down a group of the periodic table.
Why?
Li 2s1 adding principal energy levels
Na 3s1
K 4s1
Atoms getting larger with more energy levels
Electrons getting further away from + charged nucleus

25. Periodic Trends
Atomic size generally decreases as you move from left to right across a period.
Why?
As you go across a period, the principal energy level remains the same. Each element has one more proton and one more electron than the preceding element.
The electrons are added to the same energy level, causing the increasing positive charge to pull them in closer.
Realize at some point this effect is less pronounced more electrons, more reaction between them repulsion and that force is greater than positive attraction of nucleus

26. shielding effect- the reduction of the attractive forces between a nucleus and its outer electrons due to the blocking effect of inner electrons
The shielding of the nucleus by electrons also increases with the additional occupied orbitals between the outermost orbital and the nucleus.

28. Trends in Ionization Energy
When an atom gains or loses an electron, it becomes an ion.
The energy required to overcome the attraction of the nuclear charge and remove an electron from a atom is called the ionization energy.
Removing one electron results in the formation of a positive ion with a 1+ charge.
Na(g) Na+(g) + e-

29. The energy required to remove this first outermost electron is called the first ionization energy.
To remove the outermost electron from the 1+ ion requires an amount of energy called the second ionization energy, and so forth.
Nifty swell animation

30. Group Trends
Ionization energy generally decreases as you move down a group of the periodic table.
This is because the size of the atoms increases as you descend, so the outermost electron is farther from the nucleus.
The outermost electron should be more easily removed, and the element should have a lower ionization energy.

31. Periodic Trends
For the representative elements, ionization energy generally increases as you move from left to right across a period.
The atomic number and therefore positive charge increases and the shielding effect is constant as you move across. A greater attraction of the nucleus for the electron leads to the increase in ionization energy.
Also, electron configuration/ noble gas configuration harder to remove electron to get to a more stable electron configuration- easier to gain an electron
Ionization energy generally decreases as you move down a group of elements but increases across a period.
within a group, as you travel down a family there is more space between a positive nucleus and outer electrons
easier to take electrons from the atom
easier = less energy
shielding effect also plays role
within a period, atomic # increases, protons increase
positive pull is harder, making it more difficult to remove electron Cl

33. Trends in Electronegativity
The electronegativity of an element is the tendency for the atoms of the element to attract electrons when they are chemically combined with atoms of another element.
Electronegativity generally decreases as you move down a group.
As you go across a period from left to right, the electronegativity of the representative elements increases.

34. The metallic elements at the far left of the periodic table have low electronegativities. By contrast, the nonmetallic elements at the far right (excluding the noble gases), have high electronegativities.
The electronegativity of cesium, a metal, the least electronegative element, is 0.7; the electronegativity of fluorine, a nonmetal, the most electronegative element, is 4.0.
Because fluorine has such a strong tendency to attract electrons, when it is chemically combined to any other element it either attracts the shared electrons or forms a negative ion.
In contrast, cesium has the least tendency to attract electrons.
Groovy animation