1. Chapter 5
The Periodic Law

2. Vocabulary
Section 1 Periodic Law Periodic Table Lanthanide Actinide Section 2 Alkali Metals Alkaline-Earth Metals Transition Metals Main-Group Elements Halogens Section 3 Atomic Radius Ion Ionization Ionization Energy Electron Affinity Cation Anion Valence electrons Electronegativity

3. History of the Periodic Table
Mendeleev—arranged the elements in order of increasing atomic mass
Created a table in which elements with similar properties were grouped together
In this way, properties of undiscovered elements could be predicted
Moseley—arranged atoms according to increasing atomic number
Gave us the periodic law—the physical and chemical properties of the elements are periodic functions of their atomic numbers
Produced a table in which elements with similar properties fell into the same column, or group.

4. Electron Configuration and the Periodic Table
Section 2
Electron Configuration and the Periodic Table

5. S-Block Elements Group 1—alkali metals Group 2—alkaline earth metals
Silvery appearance; Soft enough to cut with a knife
Extremely reactive, found as compounds in nature
Contains 1 valence electron (outermost energy level)
Group 2—alkaline earth metals
Harder than alkali metals
Although less abundant, still found as compounds in the Earth’s crust
Contains 2 valence electrons

6. P-Block Elements 6 Groups Boron Carbon Group 13
Always found combined with other elements in nature
Contains 3 valence electrons
Carbon
Group 14
Group varies in many properties because?
Contains 4 valence electrons
Allotropes – forms of an element in the same state of matter, but have different physical and chemical properties
Example: diamond and graphite are both made of carbon and are both solids, they differ in their properties because they have different structures – they are isotopes.

7. Allotropes of Carbon
Example: diamond and graphite are both made of carbon and are both solids, they differ in their properties because they have different structures – they are isotopes.

8. P-Block Elements Continued
Nitrogen
Group 15
Properties vary as in the carbon group
Contains 5 valence electrons
Oxygen
Group 16
Act as nonmetals, meaning they tend to gain electrons
Contain 6 valence electrons

9. P-Block Elements Continued
Halogens
Group 17
Form compounds with metals (salts)
A salt is a combination of a metal and a halogen
Reactive nonmetals that tend to gain 1 electron
Contains 7 valence electrons
Noble Gases
Group 18
Colorless
Unreactive
Contains 8 valence electrons – except helium which has 2 (filled orbitals)

10. d-Block Elements Known as Transition metals
Share properties such as conductivity, luster, malleability, atomic size, electronegativity, ionization energy, melting points, and boiling points
May have multiple oxidation states
Tend to lose electrons to become positively charged ions

11. f-Block Elements: Inner Transition Metals
Lanthanide series (4f)
Silvery metals
Similar reactivity like Group 2 (alkaline-earth metal) elements
Actinide series (5f)
Radioactive elements
Only four exist in nature; the rest are synthetic elements that are created in particle accelerators

12. Bell Ringer
For the following elements, name the block and group in which these elements are located on the periodic table; and whether high or low reactivity
[He] 2s2 2p3
[Xe] 4f14 5d10 6s1
[Rn] 7s1
[Ar] 3d10 4s2 4p5

13. Section 3
Periodic Trends

14. Atomic Radii What is it? Period Trend Group Trend
One-half the distance between the nuclei of identical atoms that are bonded together
Period Trend
Gradual decrease in atomic radii from left to right
Electrons pulled closer by the increasing positive charge of the nucleus
Example:
Lithium is larger than Fluorine
Group Trend
Gradual increase down a group
As electrons occupy sublevels in higher energy levels farther from the nucleus, the sizes of the atoms increase
Positive charge of the nucleus is overridden by sheer number of electrons
Iodine is larger than fluorine

16. Ionization Energy (IE)
What is it?
The energy required to remove one electron from the outer shell of an atom
A + energy  A+ + e-
Ion—an atom or group of elements that has a positive or negative charge
Periodic Trend
Ionization Energy increase from left to right across a period
Caused by increasing positive charge of nucleus
Example: It takes more energy to remove an electron from fluorine than from lithium
Group Trend
Decreases from top to bottom down groups
Increase in distance from the nucleus = decrease in attraction between the nucleus and electrons
Example: It takes less energy (easier) to remove an outer electron from Rubidium than from lithium

17. Removing electrons from positive ions
Second ionization energy—energy required to remove additional electrons from positive ions.
Requires much more energy than first ionization energy.
With each additional electron removed, the effect of nuclear charge increases on the electrons, making them increasingly more difficult to remove from the outer shell.

18. Recall Rank in order of increasing atomic radius:
Mn, Ca, Cu, Br
Li, Rb, H, K
Rank in order of increasing ionization energy:
Al, P, S, Si
Sb, P, As, N

19. Electron Affinity What is it? Period Trend Group Trend
The difference in energy when an electron is added to a neutral atom.
Most atoms release energy (F and Cl). A + e-  A- + energy
Represented by a negative number.
Some require energy to gain an electron (N). (must be forced) A + e- + energy  A-
Represented by a positive number.
Period Trend
Increases from left to right (halogens gain electrons most readily)
Caused by increase in effective nuclear charge; want to be noble gases.
Group Trend
Generally decreases down the group.
Caused by dominant increasing atomic radius which decreases effective nuclear charge.

21. Ionic Radii What is it?? Periodic Trend Group Trend
The radii of the most common ions of the elements.
Positive ions—cations (loss of electrons) A+
Negative ions—anions (gain of electrons) A-
Periodic Trend
Cation radii—decrease from left to right
Electron cloud shrinks due to increasing nuclear charge and the removal of higher- energy level electrons.
Anion radii—increase from left to right
Electron cloud increases due to a decrease in nuclear charge as electrons are added, and an increase in repulsion due to increased number of negatively-charged electrons.
Group Trend
Gradual increase in ionic radii as you go down a group.

22. Valence Electrons
Chemical compounds are form because electrons are lost, gained or shared
Only the outer electrons are involved in forming compounds
Valence electrons – electrons available to be lost, gained or shared in the formation of chemical compounds
Ex: an electron lost from the 3s sublevel of Na would form Na+

23. Fluorine is the most electronegative atom
Electronegativity
What is it??
The ability of an atom in a compound to attract electrons from another atom in the compound.
Periodic Trend
Increases from left to right
Due to increasing effective nuclear charge
Group Trend
Decreases from top to bottom
Fluorine is the most electronegative atom