1. TABLE
PERIODIC

2. Development of the Periodic Table
Newlands – law of octaves
Mendeleev – the periodic table

3. Periodic Law Elements in order of increasing atomic number Fe Co Ni Cu
Atomic no. above
Atomic mass below

4. Periodic Law
Repeating pattern of characteristics (periodicity)

6. Lewis dot diagrams Valence shell electrons
Atomic Structure and Periodicity
Lewis dot diagrams Valence shell electrons

7. Large Groups Metals, Nonmetals, Metalloids Transition Metals
Rare Earth Elements

8. READING THE TABLE Groups (Families) Cu, Ag, Au
List 5 characteristics that Cu, Ag, and Au have in common

9. READING THE TABLE Periods – 1, 2, 3…
What characteristics do the elements in period 1 all have in common?

10. Metals Left side of periodic table, below staircase Luster - shine
Good conductors of heat and electricity
Solids (room temp), exception-mercury
Malleable-can be hammered into thin sheets
Ductile-can be drawn into fine wires
Lose electrons in bonds
Make positive ions

11. Nonmetals Far right side of periodic table above staircase
Poor conductors of heat and electricity
Neither malleable nor ductile
Gases or solid (room temp), 1 liquid (bromine)
Gain electrons in bonds
Make negative ions

12. Semimetals Metalloids ex. Si, As, Ge, Sb, Te
Both properties of metals and nonmetals
Form + ions
Semiconductors

13. Alkali Metal Li, Na, K …
Shiny solids, malleable, ductile, good conductors
Low densities, low melting point
Very reactive (Na most active)
Soft enough to be cut by a knife
Intense reactions with water and air
One valence electron
Make 1+ ions
Na, K most abundant of group, always found in nature in compounds (due to reactivity)
Na: baking soda, table salt, bleach

14. Alkaline Earth Metals Be, Mg, Ca, Sr, Ba, Ra
Densities and melting points > group 1A
2 Valence electrons
2+ charge on ions
Not found in nature as elements
Mg and Ca most abundant in group
Mg wheels, tools
CaCO3 shells, chalk

15. Halogens F, Cl, Br, I “Salt Former”
All exist in elemental forms as Diatomic Molecules
F2 Cl2 Br2 I2
High reactivity (React with most metals and nonmetals)
7 valence electrons
Make 1- ions
F corrosive gas, most reactive element
CFC refrigerators, air conditioners
Cl pool cleaner, PVC plastic, bleach
Br pesticides, photo film
I iodized salt, alcohol solution for cuts

16. Noble Gases He, Ne, Ar, Kr, Xe, Rn Least reactive of all elements
8 valence electrons
Do not make ions
Ar light bulbs
Ne signs
He balloons
Rn is a radioactive gas

17. Transition Metals Iron and Titanium, abundant; Platinum, rare
High densities, high melting points
Make alloys
Relatively non-reactive
Make colored compounds
Luster
Good conductors of electricity and heat
Make variable ions

18. Inner Transition Metals
Lanthanides (begins with La)
-soft, silvery metals, tarnish in air
-making steel alloys
Actinides (begins with Ac)
-all isotopes are radioactive
-nuclear fuels

19. Hydrogen Nonmetal, colorless, odorless 1 valence electron 1+ charge
Elemental H2
Most is combined with oxygen as water
Most abundant element in universe
Major use--making ammonia for fertilizer

20. PERIODIC TRENDS Atomic Radius Ionization Energy Electronegativity
Ion Size
Electron Affinity

21. Atomic Radius (pm) Plotted against Atomic Number

22. Atomic Radius
Distance from the center of the nucleus to the outside shell of electrons

23. Atomic Radius
Why does atomic radius increase as you go down a column or family?
Why does atomic radius decrease as you go across a row or period?

24. Practice Problems
K Al P Cl
As S F He
Ra Mg Be Ca
F B O N
Li Al Sb At

25. Atomic radius increases as you go down a group because you are adding more electron shells/ adding more energy levels.
Atomic radius decreases across a period because you are adding more protons, increasing the effective nuclear charge and the electrons are being added to the same energy level.
The nucleus pulls that electron shell in tighter

26. Ionization Energy

27. Ionization Energy Energy needed to remove an e- in gas state
Tells you how strongly an atom holds onto its valence electron
Unit-joules (J or kJ)

28. Ionization Energy

29. Ionization energy increases across a period because the valence electrons are closer and closer to 8.
High IE means they don’t want to give up electrons.
Low IE means they don’t mind giving up electrons.
Ionization energy decreases down a group because of “shielding” – they have more electron shells, further from the nucleus (+), so they are easier to pull away

30. Successive Ionization Energies
Mg  Mg+ + e- 735 kJ/mol Mg+  Mg2+ + e kJ/mol Mg2+ Mg3+ + e kJ/mol

31. Successive Ionization Energies
Al  Al+ + e- 580 kJ/mol Al+  Al2+ + e- 1,815 kJ/mol Al2+ Al3+ + e- 2,740 kJ/mol Al3+ Al4+ + e- 11,600 kJ/mol

32. Practice Problems
K Al P Cl
As S F He
Ra Mg Be Ca
F B O N
Li Al Sb At

33. Electronegativity

34. Electronegativity
The attraction that an atom has for electrons in a chemical bond
F is 4 (highest)
Fr is lowest

35. Practice Problems
K Al P Cl
As S F He
Ra Mg Be Ca
F B O N
Li Al Sb At

36. Ionic Size
Cations (Positive Ions) - smaller than atom A more positive nucleus pulls electron shells in tighter Anions (Negative Ions) - larger than atom More electrons= more resistance to positive nucleus

37. Ionic Radii

38. Atomic Radius Ionic Radii
Electronegativity
Whichever element is closest to have 8 electrons in its valence shell has the highest electronegativity. The closer an element has to 8 electrons in its valence shell, the it wants to attract electrons.
Noble gases have a full valence shell and DO NOT want to attract electrons. They have the lowest electronegativity
For elements in the same group, the element furthest down has the most electron shells, so the valence shell is less affected by the attraction of the nucleus (lowest electronegativity)
Ionization Energy
Whichever element is closest to have 8 electrons in its valence shell has the highest ionization energy. The closer an element has to 8 electrons in its valence shell, the more difficult it is to pull an electron away.
For elements in the same group, the element furthest down has the most electron shells, so the valence shell is further from the nucleus, making it easier to pull an electron away (lowest ionization energy)
Atomic Radius
Whichever element is furthest down has the most energy levels/electron shells (has the largest radius)
For elements in the same period, the element furthest to the right has the highest effective nuclear charge, which pulls the electrons in tighter (has the smallest radius)
Ionic Radii
Whichever ion is the most negative has the largest atomic radius because it is less affected by the positive nucleus.
Whichever ion is the most positive has the smallest radius because it has a higher effective nuclear charge.