1. Unit 3 – The Periodic Table

2. Mendeleev (1869) arranged elements in order of atomic mass
He then noticed certain properties
Ex.) unreactive gas
Ex.) reactive metal that combines
1:1 with chlorine
.....re-appeared periodically
Each element was placed below the
preceding one with similar properties
noble gases

3. Elements are now arranged by atomic number
Vertical columns are called groups
--contain elements with similar
chemical properties
because they have similar electron arrangements
same group = same number of valence electrons
Horizontal rows are called periods
--properties of the elements change
predictably
same period = same number of principal energy levels
(period no. = no. of PELs or shells)

4. Periodic Properties
-- change in a predictable way as you move across or down
the table
1. Atomic Radius:
half the distance between 2 nuclei
in the solid state
-- measured in picometers (pm) = meter
Related to:
a. number of e- energy levels:
radii: Li
less
small Fr
more
large
b. strength of attraction between nucleus and outer e-
number of protons
(nuclear charge):
Radii: Li
Less
large Ne
More
small

6. Ionic Radius – radius after electron gain or loss
metals lose electrons
+ charge attracts electrons to
nucleus, radius gets smaller
non-metals gain electrons
– charge repels electrons,
radius gets larger

7. 2. Ionization Energy
Periodic Properties, continued
--the amount of energy required to remove the most loosely held
electron from an atom
-- measured in kilojoules per mole (kJ/mol) [see ref table S]
-- depends on radius
a. large radius = weak attraction for outer e-
= low ionization energy
b. small radius = strong attraction for outer e-
= high ionization energy
3. Electronegativity
-- the tendency to gain electrons in a bond
-- arbitrary scale (no units) [see ref table S]
-- depends on radius
large radius = low electronegativity
small radius = high electronegativity

8. 4. Metallic Character
Periodic Properties, continued
-- ease of electron loss
-- depends on radius
large radius = high metallic character
small radius = low metallic character

9. Trends in Periodic Properties
Electronegativity
Electronegativity

10. Types of Elements 1. Metals -- luster -- malleable -- good conductors
Metals have a weak hold on their electrons
-- low ionization energy and electronegativity
-- react with non-metals by losing electrons, forming + ions
-- metallic solids (except Hg)
2. Non-Metals
no luster, brittle, poor conductors
Non-metals have a strong hold on their electrons
-- high ionization energy and electronegativity
-- react with metals by gaining electrons, forming - ions
-- react with other non-metals by sharing electrons to form
covalent bonds

11. Non-metals can be:
a. molecular solids -- C P S Se I2 At
b. liquids -- Br2
c. gases -- N2 O2 F2 Cl2 H2
3. Metalloids
-- properties midway between the above
-- used as semiconductors
-- all solids: B Si Ge As Sb Te
4. Noble Gases
-- do not easily gain, lose, or share electrons
-- all gases: He Ne Ar Kr Xe Rn

12. The Elements by Group 1 and 2 -- very reactive metals (except H)
-- not found pure in nature
-- 1 more reactive than 2
-- within each group, reactivity increases as
you move down the group
-- atoms lose e- easily: easiest for larger elements
( group 1, lower elements)
-- usually form ionic compounds
--Transition elements
-- less reactive metals
-- many found pure
-- have incomplete inner e- shell:
: similar properties
: ions can have several
different charges

13. 13 – less reactive metals (except B, metalloid)
Al – self protective metal
14 – increasing metallic character as you move down the group
C Si Ge Sn Pb
non-metal metalloids metals
C – basis of life; organic chemistry
-- 3 allotropes (different forms of the same element)

14. Allotropes have different properties due to their different structures:
Si + Ge – used in computer circuits
Pb – easily worked metal
subtly toxic; formerly used in pipes, paint, gasoline, solder
15 – Also increasing metallic character as you move down the group
N2 – very unreactive; found pure
N N strong triple bond
must react (“fix”) nitrogen before it can be used
found in protein; used in fertilizers and explosives

15. O2 -- reactive gas; found pure due to photosynthesis
16 mostly nonmetals
O2 -- reactive gas; found pure due to
photosynthesis
-- allotropes O2 and O3
(ozone)
ozone in the atmosphere absorbs harmful UV energy
light
O2  O• + O•
O2 + O•  O3
chlorofluorocarbons like Freon (CF2Cl2) reach
the ozone layer and destroy it
light
CF2Cl2  CF2Cl + Cl•
Cl• + O3  ClO• + O2
map showing ozone depletion
over Earth’s southern hemishpere

16. S – less reactive than O2
-- an impurity in fossil fuels and metal ores
S + O2  SO2
SO2 is an irritating pollutant that combines with moisture in
the air to produce acid rain:
SO2 + H2O  H2SO3

17. 17 – The Halogens
-- all are diatomic, reactive, not found pure
-- reactivity decreases as you the group
(easier for small atoms to gain e-)
-- Fluorine is the most electronegative
-- melting and boiling points increase
as you the group
Fluorine
gas
Chlorine
gas
Bromine
liquid
Iodine
solid

18. 18 – The Noble Gases
--very stable; full outer electron levels
He Ne Ar -- no compounds
Kr Xe Rn -- can react with F2 or O2
-- used when reactions are undesirable
Ar – to fill light bulbs
He – non-flammable; for balloons and blimps
helium
hydrogen