Chapter 4The Periodic

Table

Chemistry I 5.0

4-1 How are the elements organized?

• Late 1700s – Only 30 elements were identified

• Mid 1800s – 65 elements were now identified with the help of spectroscopy.

History of Periodic Table

• J.W. Dobereiner: Organized the elements into groups with similar properties.

• He called these groups triads.• The middle element is often the

average of the other two.

Triads on the Periodic Table

History of the Periodic Table

• J.A.R. Newlands– Law of octaves. He said that properties

repeated every 8th element.

History of the Periodic Table

• Mendeleev: Father of the Periodic Table– Organized the 1st periodic table according to

increasing atomic mass and put elements with similar properties in the same groups.

Mendeleev (cont.)

• He rearranged some elements out of atomic mass in order to keep them together with other elements with similar properties. He also left three blanks in his table and correctly identified the properties of these 3 unidentified elements that were later identified and match his predictions.

Mendeleev’s Work

History of the Periodic Table

• Moseley– Each element has a certain

amount of positive charge in the nucleus which are called protons.

– Moseley reorganized the periodic table by atomic number.

Glenn Seaborg “Seaborgium” Sg #106

• Born in 1912 in Michigan, Seaborg proposed reorganizing the Periodic Table one last time as a young chemist working on the Manhattan Atomic Bomb Project during WWII by pulling the “f-block” elements out to the bottom of the table. He was the principle or co-discoverer of 10 transuranium elements. He was awarded the Noble prize in 1951 and died in 1999.

History of the Periodic Table

• The Periodic Law: When elements are arranged in order of increasing atomic number, their physical and chemical properties show a regular (periodic) pattern.

– Valence electrons: outermost electrons which are responsible for chemical bonding.

Valence Electrons in the Periodic Table

Periodic Law

• Vertical Column – Group– Similar properties– A.K.A. Family

Horizontal Row - Period

Properties of Metals, Nonmetals, and Metalloids

• Metals: luster, malleable, ductile, high density, solids at room temp., good conductors, react with acids to produce hydrogen gas

• Nonmetals: brittle, dull, neither ductile or malleable, nonreactive with acids, nonconductors

• Semimetals (A.K.A. Metalloids): properties of both metals and nonmetals

Parts of the Periodic Table

Aluminum Metal

Sulfur Nonmetal

Silicon Metalloid

1A 1

2A2

3A 4A 5A 6A 7A13 14 15 16 17

8A18

1 = Alkali Metals and Hydrogen Group

13 = Boron Group

18 = Noble Gas Group

17 = Halogen Group

16 = Oxygen or Chalogen Group

15= Nitrogen Group

14 = Carbon Group

2 = Alkaline Earth Metals

Transition Group

Lanthanide SeriesActinide Series

Rare Earth Metals

3B 4B 5B 6B 7B 8B 9B 10B 1B 2B3 4 5 6 7 8 9 10 11 12

Trends in the Periodic Table

• Atomic Radius

– The distance from the nucleus to the outermost electrons.

– Atoms get larger going down a group and from right to left in a period.

Atomic Radii vs Atomic Number

Trends in the Periodic Table

• Ionic Size

– When atoms gain electrons, they become larger.

Trends in the Periodic Table

• Ionic Size– When atoms

lose electrons, they become smaller.

– Ions become larger when you go down a group.

Relative Sizes of Positive &

Negative Ions

The sodium ion lost an electron, and therefore the positive

protons in the nucleus exert a stronger pull on the remaining negative electrons, shrinking

the orbitals. Thus positive ions are smaller than their atoms.

The chloride ion gained an electron, and therefore the fewer positive protons in the nucleus

exert a weaker pull on the extra negative electrons, increasing the size of the orbitals. Thus negative ions are larger than

their atoms.

Trends in the Periodic Table• Ionization Energy

– The energy needed to remove electrons from atoms.

– Elements that do not want to lose their electrons have high I.E.

– I.E. increases going up a group due to electron shielding.

– I.E. increases going from left to right in a period.

Ionization Energy of the 1st 20 Elements

Ionization Energy vs. Atomic Number

D. Successive Ionization Energies:

1. Energy required to remove electrons beyond the 1st electron.

2. Ionization energies will increase for every electron removed.

3. Na [Ne]3s1 Na• 1st = ____ kJ 2nd = ____ kJ4. Mg [Ne]3s2 Mg: 1st = ____ kJ 2nd = ____ kJ 3rd =

____kJ5. Al [Ne]3s23p1 Al: 1st = ____kJ 2nd = ____kJ 3rd = ____kJ 4th

= ___kJ

738

4560

496 145

0773057

71816

2744

11,600

Trends in the Periodic Table

• Electronegativity– Reflects an atom’s ability to attract

electrons in a chemical bond.

– Increases going up a group.

– Increases going from left to right in a period.

Trends in the Periodic Table• Electron Affinity

– Energy change that occurs when an atom gains an electron.

– General rule – Nonmetals have more negative electron affinities than metals (except for the Noble Gases)

- A highly negative electron affinity attracts electrons. (nonmetals)

- A positive electron affinity does not attract electrons. (metals)

Electron Affinity