1. The Periodic Table
Introduction

2. Mendeleev’s Periodic Table
Dmitri Mendeleev

3. Mendeleev’s Table
Grouped elements in columns by similar properties in order of increasing atomic mass.
Found some inconsistencies - felt that the properties were more important than the mass, so switched order.
Found some gaps.
Must be undiscovered elements.
Predicted their properties before they were found.

4. Modern Russian Table

5. Chinese Periodic Table

6. A Spiral Periodic Table

7. “Mayan” Periodic Table

8. The modern table Elements are still grouped by properties.
Similar properties are in the same column.
Order is in increasing atomic number.
Added a column of elements Mendeleev didn’t know about.
The noble gases weren’t found because they didn’t react with anything.

9. Orbital filling table

10. Group 1: Alkali Metals
Most active metals, only found in compounds in nature
React violently with water to form hydrogen gas and a strong base: 2 Na (s) + H2O (l)  2 NaOH (aq) + H2 (g)
1 valence electron
Form +1 ion by losing that valence electron
Form oxides like Na2O, Li2O, K2O

11. Group 2: Alkaline Earth Metals
Very active metals, only found in compounds in nature
React strongly with water to form hydrogen gas and a base:
Ca (s) + 2 H2O (l)  Ca(OH)2 (aq) + H2 (g)
2 valence electrons
Form +2 ion by losing those valence electrons
Form oxides like CaO, MgO, BaO

12. Groups 3-11: Transition Metals
Many can form different possible charges of ions
If there is more than one ion listed, give the charge as a Roman numeral after the name
Cu+1 = copper (I) Cu+2 = copper (II)
Compounds and solutions containing these metals can be colored.

13. Group 17: Halogens Most reactive nonmetals
React violently with metal atoms to form halide compounds: 2 Na + Cl2  2 NaCl
Only found in compounds in nature
Have 7 valence electrons
Gain 1 valence electron from a metal to form
-1 ions
Share 1 valence electron with another nonmetal atom to form one covalent bond.

14. Group 18: Noble Gases
Are completely nonreactive since they have eight valence electrons, making a stable octet.
Kr and Xe can be forced, in the laboratory, to give up some valence electrons to react with fluorine.
Since noble gases do not naturally bond to any other elements, one atom of noble gas is considered to be a molecule of noble gas. This is called a monatomic molecule. Ne represents an atom of Ne and a molecule of Ne.

15. Properties of Metals
Metals are good conductors of heat and electricity
Metals are malleable
Metals are ductile
Metals have high tensile strength
Metals have luster

16. Properties of Metals Lose electrons easily
Low ionization energy and electronegativity.
Form positive ions when combining with other atoms by losing electrons.
Elements with the most metallic properties are in the lower left of the periodic table.

17. Examples of Metals
Potassium, K reacts with water and must be stored in kerosene
Copper, Cu, is a relatively soft metal, and a very good electrical conductor.
Zinc, Zn, is more stable than potassium
Mercury, Hg, is the only metal that exists as a liquid at room temperature

18. Properties of Nonmetals
Carbon, the graphite in “pencil lead” is a great example of a nonmetallic element.
Nonmetals are poor conductors of heat and
electricity
Nonmetals tend to be brittle
Many nonmetals are gases at room temperature

19. Nonmetals Gain electrons easily.
High ionization energies and electronegativities.
Form negative ions when combining with metal atoms by gaining electrons.
Produce covalent bonds by sharing electrons with other nonmetals.
Exist as gases, molecular solids, or network solids at room temp. (Bromine exception).

20. Nonmetals Con’t Do not/are not have luster, good conductors, ductile.
Many are diatomic molecules (Br2, I2, N2, Cl2, H2, O2, F2).
Elements with nonmetallic properties appear in the upper right of the periodic table.

21. Examples of Nonmetals
Microspheres of phosphorus, P, a reactive nonmetal
Sulfur, S, was once known as “brimstone”
Graphite is not the only pure form of carbon, C. Diamond is also carbon; the color comes from impurities caught within the crystal structure

22. Metals & Nonmetals

23. Properties of Metalloids
KNOW THESE!
Metalloids straddle the border between metals and nonmetals on the periodic table.
They have properties of both metals and nonmetals.
Metalloids are more brittle than metals, less brittle than most nonmetallic solids
Metalloids are semiconductors of electricity
Some metalloids possess metallic luster

24. Silicon, Si – A Metalloid
Silicon has metallic luster
Silicon is brittle like a nonmetal
Silicon is a semiconductor of electricity
Other metalloids include:
Boron, B
Germanium, Ge
Arsenic, As
Antimony, Sb
Tellurium, Te

25. Metalloids

26. Determination of Atomic Radius: Periodic Trends in Atomic Radius
Half of the distance between nucli in
covalently bonded diatomic molecule
"covalent atomic radii"
Periodic Trends in Atomic Radius
  Radius decreases across a period
Increased nuclear charge, more protons; electrons
do not get further from nucleus.
  Radius increases down a group
Addition of principal energy levels

27. Table of Atomic Radii

28. Ionic Radii Cations Anions Positively charged ions
  Smaller than the corresponding
atom
Anions
  Negatively charged ions
  Larger than the corresponding
atom

29. Table of Ion Sizes

30. Ionization Energy
The energy required to remove the most loosely held valence electron from an atom in the gas phase.
High electronegativity means high ionization energy because if an atom is more attracted to electrons, it will take more energy to remove those electrons.
Metals have low ionization energy. They lose electrons easily to form (+) charged ions.
Nonmetals have high ionization energy but high electronegativity. They gain electrons easily to form (-) charged ions when reacted with metals, or share unpaired valence electrons with other nonmetal atoms.

31. Another Way to Look at Ionization Energy

32. Electron Affinity - the energy change associated with the addition of an electron
  Affinity tends to increase across a period
  Affinity tends to decrease as you go down
in a period
Electrons farther from the nucleus
experience less nuclear attraction
Some irregularities due to repulsive
forces in the relatively small p orbitals

33. Electronegativity A measure of the ability of an atom in a chemical
compound to attract electrons
  Electronegativities tend to increase across
a period
Why:
  Electronegativities tend to decrease down a
group or remain the same
Why: Electrons farther from the
nucleus experience less nuclear attraction

34. Periodic Table of Electronegativities

35. Summation of Periodic Trends