The Periodic Table of the Elements

Unit III

Figure 1: The Periodic Table of the ElementsFrom: http://www.vcs.ethz.ch/chemglobe/ptoe/periodic.gif

I. The Periodic Table: Historical Development

a) Dimitri Mendeleev: Russian chemist; developed early periodic law; organized elements of the periodic table based upon atomic mass.

• Problem: gaps existed between elements and same elements were found to have different atomic masses (isotopes).

b) Henry Moseley: British physicist; revised Mendeleev’s periodic table to produce today’s version; organized elements based upon atomic numbers.

II. Structure of the Periodic Table

• The periodic table is organized into groups and periods.

a) Groups: Characteristics

• Groups are vertical families of elements.

• The elements within groups share similar chemical properties.

• Elements within groups have the same number of valence electrons.

• The groups of the periodic table range from 1-18.

Na. Li.

Group 1 Elements

Be: Mg:Group 2 Elements

b) Periods: Characteristics

• Periods are horizontal families of elements.

• The elements within periods have electrons that occupy the same energy levels.

• Elements within periods do not have the same number of valence electrons.

Groups (1-18)

Periods 1-7

1. Group 1: Alkali Metals

• The alkali metals are the most active metals.

• They each contain one valence electron.

• The alkali metals have oxidation numbers of +1.

• They will form positively charged ions (cations).

• They will have a tendency to lose electrons.

• Francium and Cesium are most active.

2. Group 2: Alkaline Earth Metals

• The Alkaline Earth Metals are active metals.

• They contain two valence electrons.

• The Alkaline Earth Metals have oxidation numbers of +2.

• They will form positively charged ions (cations).

• The Alkaline Earth Metals will have a tendency to lose electrons.

• Radium and Barium are most active.

3. Groups 3-12: Transition Elements

a) Characteristics

• multiple positive oxidation numbers.

• form colored aqueous solutions.

• Transition elements are all metals.

• have incomplete inner d-sublevels.

4. Group 17: Halogens

Characteristics:

• contain four diatoms (F2 (g), Cl2 (g), Br2 (l), I2 (s)).

• most active nonmetals.

• have elements that exist in all three states of matter.• all have primary oxidation numbers of -1.

• contain seven valence electrons.

• will have a tendency to gain electrons.

• will have a tendency to form negatively charged ions (anions).

• Fluorine is the most active.

5. Group 18: The Noble (Inert) Gases – Monoatomic Molecules

Characteristics

• most stable elements.

• all have complete valence shells.

• are extraordinarily inactive.

• are all gases at STP.

• are also known as the monatomic molecules.

NOTE: Bromine (Br) is the only nonmetal that is a liquid at STP.

Mercury (Hg) is the only metal that is a liquid at STP.

III. Electronegativity, Ionization Energy, Electron Affinity, and Atomic Radii

a) Electronegativity:

• The force of attraction that an atom’s nucleus has on its own valence electrons and those of other elements.

• Electronegativity is a measure of the attraction that an atom has for electrons in a covalent bond.

• Two scales of electronegativity are in common use: the Pauling scale (proposed in 1932) and the Mulliken scale (proposed in 1934).

• The difference in electronegativity values between two elements enables chemists to predict the type of bonds in chemical compounds.

1) The Pauling scale

• was devised in 1932.

2) Mulliken Scale

• The most electronegative element (fluorine) is given an electronegativity value of 4.0

• The least electronegative element (francium) has a value of 0.7, and the remaining elements have values in between.

• Electronegativity values for all elements are found on Table S of The Physical Setting Reference Tables.

• On the Mulliken scale, numbers are obtained by averaging ionization potential and electron affinity.

• Consequently, the Mulliken electronegativities are expressed directly in energy units, usually electron volts (eV).

b) Ionization Energy

• The amount of energy needed to remove the most loosely bonded electron from an atom.

• Metals will generally have small ionization energies. This is due to the fact that metals will easily loose electrons due to their low electronegativity values.

• Nonmetals will generally have large ionization energies. This is due to the fact that nonmetals will NOT easily loose electrons due to their high electronegativity values.

c) Atomic Radii

IV. Classes of Elements: Metals, Nonmetals, and Metalloids

a) Metals:

• account for 2/3 of the elements on the periodic table.

• are found on the left side of the periodic table (left of the steps).

• have a tendency to lose electrons.

• form positively charges ions (cations).

• are excellent conductors of heat and electricity.

• exist as solids at STP (except Hg).

• are malleable (can be flattened) and ductile (can be stretched into a wire) .

• are lustrous (shiny).

• have low electronegativities and ionization energies.

• have high atomic radii.

b) Nonmetals:

• account for 1/3 the elements on the periodic table.

• have a tendency to gain electrons.

• are found on the right side of the periodic table (right of the steps).

• have a tendency to form negatively charged ions (anions).

• are poor conductors of heat and electricity (good insulators).

• exist in all three states of matter.

• are brittle as solids.

• have high electronegativities and ionization energies.

• are dull as solids.

• have low atomic radii.

c) Metalloids:

• are found along side of the steps.

• have properties intermediate of metals and nonmetals.

• Examples include boron (B), arsenic (As), silicon (Si), and tellurium (Te).

NOTE: Hydrogen is the only nonmetal that is found on the left side of the periodic table.

V. Trends of the Periodic Table

a) Electronegativity: the force of attraction that an atom’s nucleus has on its own valence electrons.

b) Ionization Energy: the amount of energy needed to remove the most loosely bound valence electron from a given atom.

• Across a period: Increases due to increased nuclear charge (number of positively charged protons).

• Down a group: Decreases due to a greater distance between the positively charged nucleus and outer electrons and shielding.

• Across a period: Increases due to the stronger force of attraction that atom’s positively charged nucleus has on the negatively charged electrons. Meaning that it take more energy to remove an electron from that atom.

• Down a group: Decreases due to the weaker force of attraction that atom’s positively charged nucleus has on the negatively charged electrons. Nucleus and electrons are further apart.

c) Atomic Radius:

• ½ of the distance between the nuclei of two adjacently bonded atoms.

• Across a period: Decreases due to the greater electronegativity. The nuclei of atoms are closer together.

• Down a group: Increases due to the lower electronegativity. The nuclei of atoms are further apart.