The Periodic Table

History of Periodic Table

• Developed by Dmitri Mendeleev 1869

• Arranged 70 known elements by increasing atomic mass

• Noticed a periodic recurrence of physical and chemical properties

Mendeleev’s Table

• Not completely correct– I and Te properties did

not match elements in same row, switched

• Predicted elements that had not been discovered– Ge, predicted mass, phy

and chem properties based on location

Henry Moseley (1914)

• Developed modern arrangement of the periodic table– According to # of protons

• Periodic Law: the physical and chemical properties of the elements are periodic functions of their atomic number

General arrangement

Periods• Horizontal Rows• Elements in the same period have same valence energy level

General arrangement

Groups (families)• Vertical Columns• Elements in the same

group have same # of valence electrons

• Have similar chemical and physical properties

Groups of Elements

Group 1 – Alkali Metals

• Have 1 valence electron• Form ions with a +1

charge• Soft• Extremely reactive, do

not occur as free elements in nature

Alkali Metals (cont’d)

• React with halogens (group 17) to form compounds in a 1:1 ratio

• General FormulaMHwhere, M = metal

H = Halogen

• React with oxygen to form compounds with a 2:1 ratio

• General Formula: M2O

Ex) Lithium oxide

Alkaline Earth Metals

• Have 2 valence electrons

• Form ions with +2 charge

• do not occur as free elements in nature

Alkaline Earth Metals

• React with halogens to form compounds General Formula: MH2

Calcium chloride

• Reacts with oxygen1:1 ratio

Magnesium oxide

Halogens – Group 17

• Have 7 valence electrons

• Most reactive nonmetals

• Fluorine is most reactive

Transition Metals Groups 3 - 12

• Have multiple charges• Lose electrons from

both s and d sublevels• Form colored

compounds

Noble Gases

• Have full valence shells• Extremely nonreactive• Called inert gases

Periodic Trends

• Atomic radius (table S)

• Def: half the distance between the nuclei of 2 like atoms

Atomic size

Atomic RadiiTrend in a Group

• looking from top to bottom in a group, the number of principal energy levels in each element increases

• valence electrons are shielded by the inner electrons the valence electrons are held more loosely to the

atom• As you go from top to bottom down a group, atomic

size INCREASES

Atomic RadiiTrend in a Period

• As you go from left to right across a period atomic size DECREASES.

• as # of protons increases, nuclear charge increases

• the inner electrons (all but valence) are the same

• Valence electrons are added to same energy level, but nuclear charge increases

• valence electrons are more closely attracted to the nucleus

Ionic Radii

• as atoms become ions, the size of atom changes due to the change in # of electrons

Anion Formation

• when an atom gains electrons, ion has neg charge

• radius of the atom increases• nuclear charge is spread across more

electrons, less attraction to nucleus

Cation Formation

• when an atom loses electrons, ions have a + charge

• radius of the atom decreases• atoms lose valence electrons, loss of outer

shell

Ions with the same Electron Configurations (Isoelectronic)

• When comparing ions that are isoelectronic, the radius of the ion with the greatest number of protons will be the smallest.

Ex) What is the Bohr configuration of the following ions? Which is largest ion?

Na+ Mg+2 O-2 F-1

Ionization Energy

• Def: the amount of energy needed to remove the outermost electron from a neutral atom

• Atoms may have more than one ionization energy

• Second Ionization energy is the amt of energy needed to remove the 2nd outermost electron from an atom.

• 1ST ionization energies of elements listed on Table S

Ionization EnergyTrend in a Group

• the lower the IE the more loosely bound the electron

• the larger the atom the more loosely bound the electron is to the atom

• the ionization energy of atoms in a group will decrease looking from top to bottom

Ionization EnergyTrend in a Period

• looking from left to right in a period, atoms get smaller in size

• Ionization energy increases

Homework Answers

• Page 406#9 Which element in each pair has the larger

ionization energy?sodium, potassium

Magnesium , phosphorous

Page 409

#19 Distinguish between the 1st and 2nd ionization energy of an atom.

Page 409

# 20 Indicate which element in each pair has the greater 1st ionization energy.a. Lithium, boronb. Magnesium, strontiumc. Cesium, aluminum

Page 409

#21 Would you expect metals or nonmetals to have higher ionization energies?

Page 409

#22 Arrange the following elements in order of increasing ionization energy.a. Be, Mg, Sr

b. Bi, Cs, Ba

c. Na, Al, S

Page 409

#23 Why is there a large increase between the first and second ionization energies of the alkali metals?

Electronegativity

• Chemical compounds are formed because atoms will lose, gain, or share valence electrons

• Electronegativity – the ability of an atom to attract a pair of electrons when bonded to another atom

Electronegativity (cont’d)

• The closer the pair of electrons is to the nucleus, the greater the attraction

• as size of the atom decreases, its electronegativity increases

• Electronegativity increases across a period, and decreases down a group

Metallic Character

• Def: the measure of an atom’s ability to lose electrons and form positive ions

• The stronger the metallic character, the easier it will lose electrons

Properties of elements with greatest metallic character:

• Low ionization energies

• low electronegativity

Trend in Metallic CharacterLeft to right in a Period:

Metallic character decreases

Top to bottom in a Group:Metallic character increases