1. Development of the Periodic Table

2. Development of the periodic table
Dmitri Mendeleev (mid-1800s) – arranged elements by similarities in properties and mass.
Left blank spaces in the table for undiscovered elements.
Henry Moseley arranged the elements by atomic number (number of protons), which is the current form of the table.

3. Modern Periodic Table Horizontal rows are called periods.
Vertical Columns are called groups or families.

4. Modern Periodic Table (cont)
Periodic Law – When the elements are arranged in order of increasing atomic number, there is a periodic repetition of their physical and chemical properties.
As you move across a period, the properties of the elements change.
As you move down a group, the properties are similar.

5. Modern Periodic Table (cont)
Group 1(1A) – Alkali Metals
Group 2 (2A) – Alkaline Earth Metals
Groups 3-12 (3B-8B, 1B, 2B) –Transition and Inner Transition (Lanthanides & Actinides) Metals
Group 17 (7A) – Halogens
Group 18 (0) – Noble (Inert) Gases

6. Modern Periodic Table (cont)
B, Si, Ge, As, Sb, Te, Po are metalloids.
C, N, O, P, S, Se are non-metals.

7. Modern Periodic Table (cont)
Most elements are solids at 25 C.
Exceptions:
Liquids – Mercury (Hg) and Bromine (Br)
Gases – Hydrogen (H), Nitrogen (N), Oxygen (O), Fluorine (F), Chlorine (Cl), all of Group 18 (He, Ne, Ar, Kr, Xe, Rn)
Man-made elements – Technetium (Tc) and all elements 93 and above.

8. Chemical Periodicity

9. Classification of the Elements
Of the major subatomic particles, electrons play the most significant role in determining the properties of an element
Therefore, there is a relationship between electron configurations of elements and their position on the periodic table

10. Classification of the Elements
Elements can be classified into four categories according to their electron configurations:
Noble Gases – outermost s and p sublevels are filled
Representative Elements – outermost s and p sublevels are only partially filled

11. Classification of the Elements
Transition Metals – outermost s and nearby d sublevels contain electrons
Inner Transition Metals – outermost s and nearby f sublevels contain electrons

12. Explaining Vertical Patterns
Shielding - the decreased attraction between an electron and the nucleus of an atom with more than one energy level (shell).
As energy levels increase shielding does too. This causes a decreased attraction between electron and the nucleus

13. Explaining Horizontal Patterns
Effective Nuclear charge (Zeff) - amount of nuclear charge “felt” by a electrons in valence shell
The charge (Z) of the nucleus attracts electrons. The magnitude of this electrostatic attraction is proportional to the charge of the nucleus.

14. Periodic Trends: Atomic Radius
The quantum-mechanical model does not have a defined boundary
Therefore, the radius of an atom cannot be measured directly
For solids, x-ray diffraction can provide an estimate of the distance between nuclei

15. Periodic Trends: Atomic Radius
Atomic radius – size of the atom from the nucleus to the valence shell (orbital)
Group trend – atomic size generally INCREASES as you move DOWN a group
Periodic trend – atomic size generally DECREASES as you move across a period

16. Periodic Trends: Ionic Radius
Ionic Radius – If ions are not being compared to their original atoms and are not isoelectronic, the trend is the same as atomic radius. (Ex: K+ vs O2-, K+ is larger  Shielding)
Isoelectronic Ions: When ions have the same number of electrons, the ion with more protons has a smaller radius (Ex: Ba2+ vs I-, Ba2+ is smaller  Zeff)

17. Periodic Trends: Ionic Radius
Positive ions (cations) are smaller than the original atom
Negative ions (anions) are larger than the original atom

18. Periodic Trends: Ionization Energy
Ionization Energy – the energy required to overcome the attraction of the nuclear charge and remove an electron from the valence shell(orbital)
Removing one electron results in the formation of a positive ion with a 1+ charge
The energy required to remove this electron is called the first ionization energy

19. Periodic Trends: Ionization Energy
Group trends – the first ionization energy generally DECREASES as you move DOWN a group
Periodic trends – the first ionization energy generally INCREASES as you move across a period (for the representative elements)

20. Periodic Trends: Ionization Energy
Atoms of metallic elements have low ionization energies – they form positive ions easily
Atoms of nonmetals easily form negative ions

21. Periodic Trends: Electronegativity
Electronegativity – the tendency for the atoms of an element to attract electrons in a chemical bond
Electronegativity values help predict bonding type
Group trends – the electronegativity generally DECREASES as you move DOWN a group
Periodic Trends – the electronegativity generally INCREASES as you move across a period
*** NOBLE GASES HAVE ZERO EN ***

22. Periodic Trends: Metallic Activity
Metallic Activity: how readily a metal gives up an electron. The more easily an electron is given up the more vigorous the reaction
Using shielding and Effective nuclear charge predict what the trend is.

23. Periodic Trends