Post on 11-Jan-2016
Chapter 6 The Periodic Table and
Periodic LawThe elements, which make up all
living and non-living matter, fit into a orderly table. When interpreted
properly, the table describes much of the elements physical and chemical
properties.
What is the Periodic Law and how was it formulated?
• Demitri Mendeleev is known as the father of the periodic table
• He arranged the elements in families (groups) and periods (rows,series) according to atomic mass and properties
• Mendeleev noted that the chemical and physical values for elemental properties would either be high or low depending upon the group under observation.
• He proposed the first Periodic Law "The properties of the elements are a periodic function of their atomic masses"
• Left blanks in his table for undiscovered elements
Moseley’s Modern-Periodic Law
• There was some inconsistencies with Mendeleev’s table
• In the early 1900’s Moseley was able to experimentally determine the atomic number of all known elements
• Moseley then proceeded to rearrange the elements according to increasing atomic numbers.
• New/Modern Periodic law states that the properties of elements are a periodic function of their atomic number
The Modern Periodic Table
• Glenn T. Seaborg won the Nobel Prize for his work in nuclear chemistry
• In 1944, formulated the “actinide concept” of heavy element electronic structure. This concept predicted that the fourteen actinides
Some Characteristics of Groups
Group 1 (IA) - Alkali Metals
( metal characteristics - shiny, malleable, ductile, good conductors)
• Very active metals - activity increases as you go down group
• All have one valence electron - form +1 cations by losing an electron
• react violently with water
Group 2 (IIA) - Alkaline Earth Metals
• activity increases as you move down the column not as reactive as alkali metals
• Ca, Sr, and Ba react violently when they come into contact with water
• All have two valence electrons
• Form +2 cations by losing 2 electrons
Group 17 (VIIA) - Halogens
• All gain one electron to form anions with a charge of -1.
• All are nonmetals except for At which is a semimetal
• All are diatomic in their elemental form
Group 18 (VIIIA) - Noble (Rare) Gases
• Mistakenly labeled as "inert gases" until about 30 years ago because it was thought that these gases did not react with anything.
• Noble gases have filled valence (outermost) shells.
Periodic Trends
• Atomic Radii
1) As you move down a group, atomic radius increases.
2) WHY? - The number of energy levels increases as you move down a group as the number of electrons increases.
• As you move across a period, atomic radius decreases.
• WHY? - As you go across a period, electrons are added to the same energy level. At the same time, protons are being added to the nucleus. The concentration of more protons in the nucleus creates a "higher effective nuclear charge."
First Ionization Energy
• Definition: The energy required to remove the outermost (highest energy) electron from a neutral atom in its ground state.
1) As you move down a group, first ionization energy decreases.
2) WHY? Electrons are further from the nucleus and thus easier to remove the outermost one - + shielding effect of other electrons
3) As you move across a period, first ionization energy increases.
4) WHY? - As you move across a period, the atomic radius decreases, the attraction from the positive nucleus gets larger
Electronegativity
• Definition: The ability of an element to attract electrons in a chemical bond
1) As you move down a group,
electronegativaty decreases
2) As you move across a period, it
increases.