1. Bellringer
No bags in class. If you need to go to your locker do so now, before the bell rings.
What elements are in the same period as Carbon? As Argon?
What elements are in the same family as Carbon? As Argon?

2. The Periodic Table
Chapter 6

3. Organizing the Elements
Chapter 6 Section 1

4. Searching for an Organizing Principle
Prior to 1700 only 13 elements had been discovered and isolated.
The use of scientific methods in science led to the discovery and isolation of many more elements by 1780.
Chemists used the properties of the elements to sort them into groups.

5. Mendeleev’s Periodic Table
Dmitri Mendeleev
Russian chemist and teacher
1869 developed chart (periodic table) for his students to show relationships among elements
Periodic table is arranged into groups based on a set of repeating properties.
Arranged the elements in order of increasing atomic mass
Predicted discovery of more elements

6. Revision of Mendeleev’s Table
Henry Moseley
Discovered number of proton in nucleus (atomic number)
Proposed arranging elements in order of increasing atomic number
Solved problems encountered by Mendeleev
Work resulted in a clear periodic pattern of properties
In the modern periodic table, elements are arranged in order of increasing atomic number.

7. Periodic Law
The chemical and physical properties of the elements are periodic functions of their atomic numbers.

8. Classifying the Elements

9. Classifying the Elements
Metals
Left side of periodic table
Shiny
Good conductors
Ductile and malleable

10. Metals Group 1 = Alkali Metals (IA)
Group 2 = Alkaline Earth Metals (IIA
Both of these groups very chemically active; alkali metals more active
Transition Metals = Groups 3-12 (B)
Large group of metals with characteristic metallic properties
Inner Transition Metals = Lanthanides and Actinides (at bottom of periodic table)
Little known but widely used metals
Actinides are all radioactive

11. Nonmetals Upper right hand side of periodic table
Wider variation in properties
Generally poor conductors
Brittle as solids; many occur as gases at room temperature

12. Nonmetals Group 16 = Chalcogens (chalk formers) (IVA)
Group 17 = Halogens (salt formers) (IIVA)
Group 18 = Noble Gases (chemically inert)

13. Metalloids
8 elements bordering “stair-step” line that have properties of both metals and nonmetals
B, Si, Ge, As, Sb, Te, Po, At
Silicon and germanium used extensively in computer chips and solar cells

14. Classifying the Elements
Chapter 6 Section 2

15. Squares in the Periodic Table
The periodic table displays the symbols and names of the elements, along with information about the structure of their atoms.
Every periodic table must, at the least, display the symbol, atomic number, and atomic mass of each element.
Many periodic tables are color coded to illustrate groups of elements.

16. Electron Configuration in Groups
Elements can be sorted into noble gases, representative elements, transition, or inner transition elements based on their electron configuration.
Electrons and electron configuration determine many of the properties of elements.

17. Electron Configuration in Groups
Noble Gases
Completely full outer s and p orbitals
Octet (8) electrons in outer shell
Most stable electron configuration
Representative Elements (“A” groups)
Elements in the s and p blocks (excluding noble gases)
Wide variety of properties representative of all types of elements
s and p orbitals being filled
For elements in the s and p blocks, the number of valence electrons equals the last digit of the group number (in new system) or the group “A” number (older system).
Transition Metals (“B” groups)
Characterized by filling of d orbitals
Connection between 2 sets of representative elements
Inner Transition Metals
Characterized by filling of f orbitals

18. Chapter 6 Section 3
Periodic Trends

19. Atomic Radius
Generally measured as ½ the distance between the nuclei of identically bonded atoms
For main group elements the trends are:
Increases down a group – addition of energy levels; outer electrons SHIELDED from pull of nucleus by inner electrons
Decreases across a period – electrons added in same energy level; increased positive charge of nucleus pulls electrons closer

20. Atomic Radius Atoms get larger as you move down a group Atoms get
smaller as you
go across a
period

21. SAMPLE PROBLEM 1
Ranking Elements by Atomic Size
PROBLEM:
Using only the periodic table (not Figure 8.15)m rank each set of main group elements in order of decreasing atomic size:
(a) Ca, Mg, Sr
(b) K, Ga, Ca
(c) Br, Rb, Kr
(d) Sr, Ca, Rb
PLAN:
Elements in the same group increase in size and you go down; elements decrease in size as you go across a period.
SOLUTION:
(a) Sr > Ca > Mg
These elements are in Group 2A(2).
(b) K > Ca > Ga
These elements are in Period 4.
(c) Rb > Br > Kr
Rb has a higher energy level and is far to the left. Br is to the left of Kr.
(d) Rb > Sr > Ca
Ca is one energy level smaller than Rb and Sr. Rb is to the left of Sr.

22. Ionic Radius
Positive ions (cations) are always smaller than the atoms from which they formed. Cations form when atoms lose electrons.
Negative ions (anions) are always larger than the atoms from which they formed. Anions form when atoms gain electrons.
Trends
Increases down a group – atoms are larger therefore ions are larger
Across a period – decreases due to smaller size of atoms from which they form (for both positive and negative ions)

23. SAMPLE PROBLEM 2
Ranking Ions by Size
PROBLEM:
Rank each set of ions in order of decreasing size, and explain your ranking:
(a) Ca2+, Sr2+, Mg2+
(b) K+, S2-, Cl -
(c) Au+, Au3+
PLAN:
Compare positions in the periodic table, formation of positive and negative ions and changes in size due to gain or loss of electrons.
SOLUTION:
(a) Sr2+ > Ca2+ > Mg2+
These are members of the same Group (2A/2) and therefore decrease in size going up the group.
The ions are isoelectronic; S2- has the smallest Zeff and therefore is the largest while K+ is a cation with a large Zeff and is the smallest.
(b) S2- > Cl - > K+
(c) Au+ > Au3+
The higher the + charge, the smaller the ion.

24. Ionization Energy
Amount of energy required to remove an outer electron
Removing one or more electrons creates a positively charged CATION
Trends
Decreases down a group – outer electrons are shielded from nucleus; requires less energy to remove
Increases across a period – increased nuclear charge holds outer electrons more tightly

25. SAMPLE PROBLEM 3
Ranking Elements by First Ionization Energy
PROBLEM:
Using the periodic table only, rank the elements in each of the following sets in order of decreasing IE1:
(a) Kr, He, Ar
(b) Sb, Te, Sn
(c) K, Ca, Rb
(d) I, Xe, Cs
PLAN:
IE decreases as you proceed down in a group; IE increases as you go across a period.
SOLUTION:
(a) He > Ar > Kr
Group 8A(18) - IE decreases down a group.
(b) Te > Sb > Sn
Period 5 elements - IE increases across a period.
(c) Ca > K > Rb
Ca is to the right of K; Rb is below K.
(d) Xe > I > Cs
I is to the left of Xe; Cs is furtther to the left and down one period.

26. Octet Rule
Atoms tend to gain, lose, or share electrons in order to acquire a full set of eight valence electrons which is the most stable electron configuration.

27. Electronegativity
Tendency of an atom to attract electrons to itself when bonding with other atoms
Atoms with higher EN have greater attraction of electron
Allows prediction of bond type – ionic or covalent
Trends
Increases across a period – increased nuclear charge, atomic radius, and shielding create greater attraction for electrons
Decreases down a group – bonding electrons in outer shell; farther from nucleus and shielding creates less pull for electrons

28. Periodic Table of Electronegativities

29. Summary of Periodic Trends