1. The Periodic Table
A few elements, including copper, silver, and gold, have been known for thousands of years
There were only 13 elements identified by the year 1700.
Chemists suspected that other elements existed.
As chemists began to use scientific methods to search for elements, the rate of discovery increased.

2. The Periodic Table
Early chemists attempted to organize the known elements
Some used the properties of the elements
Dobereiner was a German chemist who published his classification of the elements
He organized the elements into triads

3. The Periodic Table
A triad is a set of three elements with similar properties.
The elements shown here formed one triad. Chlorine, bromine, and iodine may look different, but they have very similar chemical properties.
Dobereiner noted a pattern in his triads. One element in each triad tended to have properties with values that fell midway between those of the other two elements.

4. The Periodic Table
In 1869, a Russian chemist and teacher, Dmitri Mendeleev, published a table of the elements.
The organization he chose was a periodic table.
Elements in a periodic table are arranged into groups based on a set of repeating properties.
This concept is known as periodicity.

5. The Periodic Table Mendeleev left spaces in his table
He predicted that elements would be discovered to fill those spaces, and he predicted what their properties would be based on their location in the table.

6. The Periodic Table
Mendeleev’s table was so successful because chemists were able to make predictions from it.
However his table had one error
It was arranged by increasing atomic mass.

7. The Periodic Table Henry Moseley modified Mendeleev’s table
He arranged the elements by increasing atomic number.
This led to the Periodic Law
The physical and chemical properties of elements are periodic functions of their atomic numbers.

8. The Periodic Table
There are seven horizontal rows or periods in the table
Each period corresponds to the energy level
There are eighteen vertical columns or groups in the table

9. The Periodic Table
Although Mendeleev’s table only had 60 elements, today’s table has 118 elements
Many elements have been discovered since his original work
The noble gases - due to their unreactivity
The lanthanides and actinides series – many are man-made

10. The Periodic Table Periodicity can be observed in the periodic table
Each group has similar properties
The electron configuration tells an element’s position in the periodic table

11. The Periodic Table
The elements can be grouped into three broad classes based on their general properties.
Three classes of elements:
Metals
Nonmetals
Metalloids

12. The Periodic Table The majority of elements are metals
Metals have three key properties
Shiny or luster
Flexible (malleable – hammer into a sheet and ductile – drawn into a wire)
Good conductor of energy (electricity and heat)

13. The Periodic Table
Although there are fewer nonmetals, they are more abundant on Earth
Nonmetals have three key properties
Dull
Brittle
Poor conductor of energy

14. The Periodic Table Metalloids have properties of metals and nonmetals
Their properties can be changed by conditions
They are found along the stair step of the periodic table
Silicon is the most famous metalloid
It is responsible for computer chips

15. The Periodic Table Remember:
s block – groups 1 & 2
p block – groups 13 – 18
d block – groups 3 – 12
f block – bottom two rows
s & p block elements are called main group (representative) elements

16. The Periodic Table s block Group 1 – Alkali metals
Most reactive metals
So reactive, not found in nature as elements
Group 2 – Alkaline Earth metals
Less reactive than group 1 metals

17. The Periodic Table p block Group 13 – 16 Mixed group
Group 17 – Halogens
Most reactive nonmetals
Group 18 – Noble Gases
Completely unreactive nonmetals

18. The Periodic Table d block f block Groups 3 – 12 – Transition metals
Less reactive than groups 1 & 2
f block
Bottom two rows
Known as the lanthanides and actinides (inner transition metals)

19. The Periodic Table There are several trends in the periodic table
Atomic radii
Valence electrons
Ionic radii
Ionization energy
Electron affinity
Electronegativity

20. The Periodic Table
Atomic radii is one half the distance between the nuclei of identical atoms that are bonded together

21. The Periodic Table Atomic Radius Trend:
Down a group – atomic radii increases
This happens because of the increased number of energy levels
The energy levels shield the electrons from the attraction of protons in the nucleus
Across a period – atomic radii decreases
This happens because as more electrons are added to the same energy level, those electrons are pulled closer due to the increased number of protons in the nucleus
Largest atomic radii – francium
Smallest atomic radii – fluorine

22. The Periodic Table
Valence electrons are the electrons found in the outermost energy level
These are the electrons available to be gained, lost, or shared
All atoms want 8 valence electrons or a full outer energy level
Noble gas electron configuration
Valence electrons determine the chemical properties of the atom

23. The Periodic Table
Group Number
# of Valence Electrons 1 2 13 3 14 4 15 5 16 6 17 7 18 8
Valence electrons can be represented using Lewis Dot Diagrams

24. The Periodic Table
Atoms are neutral because there are equal numbers of both protons and electrons
Sometimes atoms can gain or lose electrons to form ions
An ion is an atom or group of atoms that has a positive or negative charge
Losing electrons results in a positive ion called a cation
Gaining electrons results in a negative ion called an anion

25. The Periodic Table Metals (left side of the table) form cations
Cations are smaller than their atom counterparts because they are losing an electron (and sometimes an energy level)
More positive charges have a greater pull on less negative charges

26. The Periodic Table Nonmetals (right side of the table) form anions
Anions are larger than their atom counterparts because they are gaining an electron
Less positive charges cannot pull in the greater number of negative charges

27. The Periodic Table
Ionization energy is the energy required to remove an electron from an atom
a low IE means it is easier to remove the electron
Atoms can lose an electron, to form an ion
They do this to achieve noble gas electron configuration (or 8 valence electrons)
When an atom easily loses electrons, it is said to be active
Metals tend to lose electrons

28. Ionization Energies of Some Common Elements
Symbol
First
Second
Third H
1312
He (noble gas)
2372
5247 Li
520
7297
11,810 Be
899
1757
14,840 C
1086
2352
4619 O
1314
3391
5301 F
1681
3375
6045
Ne (noble gas)
2080
3963
6276 Na
496
4565
6912 Mg
738
1450
7732 S
999
2260
3380
Ar (noble gas
1520
2665
3947 K
419
3096
4600 Ca
590
1146
4941

29. The Periodic Table Ionization Energy Trend:
Down a group – ionization energy decreases
As the valence electrons are farther from the nucleus, the atom gives them up with less energy
Across a period – ionization energy increase
As the number of valence electrons increases in the same energy level, the atom is more resistant to giving up an electron (more energy)
Greatest IE – fluorine
Least IE - francium

30. The Periodic Table
Electron affinity is the energy change required to gain an electron (released energy is a negative value)
When an atom releases a lot of energy, it is said to be active
Nonmetals tend to gain electrons (large energy change)

31. The Periodic Table Electron Affinity Trend:
Down a group – electron affinity decreases (slightly)
Distance from the positive nucleus decreases the pull on the electrons
Across a period – electron affinity increases
As the number of valence electrons added to the same energy level increases, the atom easily accepts another electron (to reach 8)
Greatest EA – fluorine
Least EA – francium

32. The Periodic Table
Electronegativity is the measure of the ability of an atom in a chemical compound to attract electrons
All values are based on fluorine
Fluorine is most electronegative atom with a value of 4.0
The trend decreases in either direction from fluorine