1. 5.1 Development of the Periodic Table
Explain the role of Mendeleev in the development of the periodic table
Explain how the periodic law can be used to predict properties
Describe the modern periodic table.

2. Dmitri Mendeleev
Noticed that when the elements were placed in order of atomic mass
Similar properties appeared at regular intervals
Created a table with columns of elements with similar properties
The First Periodic Table
Many holes were filled in over time

3. Henry Moseley 1887-1915 Recognized a pattern based upon Atomic Number
Periodic Table is ordered by Atomic Number, not Atomic Masses
Ar (At# 18, AtW 39.95) – K (At# 19, AtW 39.01)
Periodic Law – properties of elements are periodic functions of their atomic number.
Cycles are observed as the Atomic Number increases

4. Reading the Periodic Table
Arrangement of elements in order of their atomic number so that elements with similar properties fall into the same column.
Creating Groups or Families.
Periodicity
repeated patterns w/ the increasing of atomic number
Like a spiral growing in diameter outwards
Atomic Number Differences between groups
2,8,8,18,18,32,32

5. 5.2Electron Configuration & The Periodic Table
A fully occupied energy level is stable
Low Energy
Noble Gases
Every element wants to have a fully occupied outer energy level
Np6
Every element wants to have a Noble Gas Config.
Electron configurations can be used to determine the reactivity of elements

6. Electron Configuration7 6 5 4 3 2 1 Lr Lu Ac La
s - block
d – block (s-1)
p - block
f – block (s-2)

7. Noble Gas Configuration using Periodic Table
Move from the left to the right of the period until the element being configured is reached.
Every s and p will have the same energy level, d will be 1 less than the s and f will be 2 less.

8. Writing in Short-hand notation
Al –
Write the short hand notation for the following: V Rb I Hg U W
[Ar] 4s2 3d3
[Kr] 5s1
[Kr] 5s2 4d10 5p5
[Xe] 6s2 4f14 5d10
[Rn] 7s2 5f4
[Xe] 6s1 4f14 5d5

9. Classwork
Use Noble-Gas notation: Es Au Sb Cs Kr

10. Groups, Periods, and Blocks
Groups – Vertical Columns
Periods – Horizontal Rows
Blocks – Sections of P.T. (s,p,d,f)
Main Block Elements – s & p together
Representative group
s-block – Alkali & Alkaline Earth Metals
d-block – Transition Metals
p-block – Metalloids, Halogens, Noble Gases, And other Non-Metals & Metals
f-block – Lanthanide & Actinide

11. The Alkali & Alkaline Metals
Alkaline Earth Metals
Group 2 or 2A
Harder, Denser, & Less reactive than Gr1
Too Reactive to be found as Free Elements
Electron Config. Ns2
Steven Tepsic was here
Alkali Metals
Group 1 or 1A
Li, Na, K, Rb, Cs, Fr
Soft (cut w/ Knife
Extremely Reactive
Not found as Free Elements
Less Dense than H2O
React w/ Non-Metals to form Salts
E-Config. ns1

12. Hydrogen & Helium Both Non-metals Colorless & Odorless Hydrogen Helium
Placed In Group 1
Properties do not resemble elements in Group 1
Helium
Group 18 even though it has a ns2 config.
Placed in Gr. 18 because of its non-reactive nature (similar to the noble gases)

13. The Transition Metals (Groups 3-12)
Compared to s-block Metals
Harder, Denser Hg is an exception (liquid)
Higher Melting Point
Less Reactive
Common Transition metals
Gold, Silver, Copper, Iron, Tungsten, Cobalt, Nickel, Platinum
Electron Config
ns2(n-1)d1-10
Doesn’t always follow a normal pattern
Groups 3-12
Properties
Metallic
Good Conductors
Heat
Electricity

14. Metals vs. Nonmetals Metals Non-metals
Good conductors of heat and electricity.
Luster
Malleable
Non-metals
Poor conductors of heat and electricity.
Dull
Brittle

15. Halogens, Metalloids, & Noble Gases (p-block elements)
semiconductors
mostly brittle solids
metal & non-metal properties
Noble Gases (Group 18 or 8A)
Inert (non-reactive)
Stable (low energy)
Monatomic Gases
Elements Reactivity
P- Block Elements
Electron Config. np(1-6)
Halogens (Group 17 or 7A)
np5 Configuration
Most reactive of non-metals
React w/ metals to form salts

16. Lanthanides & Actinides ‘Inner Transition Metals’
4f (1-14) Configuration
“Rare-Earth Metals”
Which are not actually rare
Shiny Luster
Reactive
Metals
All have practical uses
Phosphors in TV’s contain many Lanthanide metals.
Actinides
5f (1-14) Configuration
Unstable & Radioactive
1st four are found in nature
The rest are synthetic
Not very Practical
Uranium is an exception
Nuclear Power
Elements made in the Stars

17. 5.3 Periodic Trends
Elements are grouped according to their Physical and Chemical Properties.
Period Trends – how a characteristic increases/decreases across a period
Group Trends – how a characteristic increases/decreases up & down a group

18. Atomic & Ionic Radii pg151&159
2 Types of Ions
Cations ( + Charge)
Loss of Electrons
Anions ( - Charge)
Gaining of Electrons
Cations – smaller than neutral counterpart
Less electrons
Anions – larger than neutral counterpart
Gained Electrons
Atomic Radius – ½ the distance between the nuclei of identical atoms joined in a molecule
Period Trend
Decreases L to R
Greater # of p+ pulls electrons closer
Group Trend
Increases T to B
Increase in energy level – greater size

20. Valence Electrons Valence Electrons
Valence Electrons – electrons available to be lost, gained, or shared in the formation of compounds.
Representative Group
Electrons in the outer s & p sublevels
Valence Electrons
Group Number 1 2 13 14 15 16 17 18
Valence Electrons 3 4 5 6 7 8

21. Ions –atoms that have gained or lost a # of electrons.
Cations – Atoms that need to lose electrons to reach a complete valence shell.
Na will lose 1e- to become Na 1+
Na = 1s2 2s2 2p6 3s1
Na 1+ = 1s2 2s2 2p6
Anions – Atoms that need to gain electrons to reach a complete valence shell.
Cl will gain 1e- to become Cl 1-
Cl = 1s2 2s2 2p6 3s2 3p5
Cl 1- = 1s2 2s2 2p6 3s2 3p6

22. Noble Gas Configuration
When atoms form ions their electron configurations resemble that of a noble gas element.
Pseuodo noble gas-
Ions that resemble noble gases, but contain a d or f block configuration.

23. Ionization Energy
Ionization – any process that results in the formation of an ion.
Ionization energy – the energy required to remove an electron from an atom
A + energy  A+ + e-
Measured w/ atoms in gas phase to eliminate influence from nearby atoms.
Period Trend – Increase from L to R
Group Trend – Decreases from T to B

25. Multiple Ionization Energies
More electrons more energy needed.
Usually there is a large jump in ionization energy when an atom loses an entire energy level.
Example: Sodium has 1 electron in valence shell.
1st = 496
2nd = 4562
3rd = 6912
2nd ionization energy,
Energy needed to remove the second electron from an atom’s valence shell.
3rd ionization energy,
Energy needed to remove the third electron from an atom’s valence shell.

26. Electron Affinity
Electron Affinity – is the energy change that occurs when an electron is acquired by a neutral atom.
Some atoms want electrons
A + e -  A- + Energy ( Exothermic)
Some atoms do not want electrons
A + e - + Energy  A- (Endothermic)
Will be unstable & lose electron spontaneously
Period Trend – More easily acquires electrons as you move across the period. ( More Negative/ More Exothermic)
Group Trend – Harder to acquire electrons as you move down the table. (Less Negative in general)

27. Electronegativity
Electronegativity – the ability of an atom in a compound to attract electrons
Fluorine – has the strongest electronegativity
arbitrarily assigned a value of 4
all other atoms were then compared to F
Period Trends – increases from L to R
Group Trend – decreases from T to B

29. Trend summary Electronegativity Increases Ionization Increases
Electron Affinity Increases
Atomic Radii / Ionic Radii Decreases
Electronegativity Increases
Ionization Increases
Electron Affinity Increases
Atomic Radii / Ionic Radii Decreases