Chapter 6:

Periodic Table and Trends

Dmitri Mendeleev• First Periodic Table

• Based on increasing Atomic mass and repeating properties of elements

• Had spaces for “missing” elements that he predicted

Henry G. J. Moseley

• Discovered that elements’ properties were more closely associated with atomic number

• Modern periodic table is based on this discovery

Periodic Law

• When elements are arranged in order of increasing atomic number, their physical and chemical properties show a periodic pattern

Reading the Periodic Table

Alkali MetalsAlkaline

Earth MetalsHalogensNoble GasesTransition

MetalsBoron Family

Carbon Family

Nitrogen FamilyOxygen Family

Groups or FamiliesPeriods

Inner Transition

Metals

Metals

• Left of the stair-step line

• Majority of the elements

• Tend to LOSE electrons

• Most reactive in the s-block

Properties of Metals

• Shiny luster• Good conductors of heat• Good conductors of electricity• Most are solids at room temperature• Malleable• Ductile

goldlead

nickelcopper

Nonmetals

• Right of Stair-Step line• Tend to GAIN electrons• Most reactive group is

halogens• Least reactive group is

Noble Gases

Properties of Nonmetals

• Dull Luster• Poor conductors of heat• Poor conductors of electricity• Brittle• Many are gases at room temperature

CARBON

ARGON

BROMINE

SULFUR

Metalloids

• Along Stair-step line• Have properties of

metals AND nonmetals• Many are used in

transistors, found in electronics

Antimony Silicon

Boron

Alkali Metals

• Group 1 (Except H)• All have only 1 valence electron• Most reactive metals; never found in

elemental form in nature• Soft and shiny• Relatively low melting points

Alkaline Earth Metals

• Group 2• All have 2 valence electrons• Second most reactive metals; never

found in pure state in nature• Harder, denser, and stronger than

alkali metals• Have higher melting points than

alkali metals

Transition Metals• Groups 3-12• All have 1 or 2

valence electrons (in s sublevels)

• Do not fit into any other group or family

• Have many irregularities in their electron configurations

Boron Family

• Group 3A• Have 3 valence electrons• Boron is a metalloid• All others are metals

Carbon Family

• Group 4A• All have 4 valence electrons• Carbon is a nonmetal• Si and Ge are metalloids• Sn and Pb are metals

Nitrogen Family

• Group 5A• All have 5 valence electrons (s and p

sublevels)• N and P are nonmetals• As and Sb are metalloids• Bi is a metal

Oxygen Family

• Group 6A• All have 6 valence electrons• Oxygen, Sulfur, and Selenium are

nonmetals• Tellurium and Polonium are

metalloids

Halogens

• Means “salt former”• Group 7A• All have 7 valence electrons• Most reactive nonmetals• All are nonmetals

Noble Gases

• Group 8A• 8 Valence electrons makes a full

electron shell: s2 p6

• Complete, stable electron configuration (Complete outer energy level)

• Least reactive of all elements

Rare Earth Elements(Inner Transition metals)

• Found in 2 rows at bottom of periodic table• Lanthanide series follows La• Actinide series follows Ac• Little variation in properties• Actinides are radioactive; only first three and

Pu are found in nature

Summary

• Groups: Up and Down• Periods: Across• Main Group Elements are in groups 1-2, 13-

18• Elements along the stair step line are

metalloids• Elements to the left of the stair step line are

metals• Elements to the right of the stair step line are

nonmetals

Octet Rule“Noble Gas Envy”

• Atoms tend to gain, lose, or share electrons in order to acquire a full set of valence electrons (typically 8)

Periodicity

• Properties of the elements change in a predictable way as you move through the periodic table

• These properties include• Atomic Radius• Ionization energy• Electronegativity

Atomic Radius

• Distance from nucleus to outermost valence electrons

Atomic Radius

• Increases down groups• Decreases from left to right

Ionization Energy

• The energy needed to remove 1 of an atom’s electrons

• Decreases as you move down a group• Increases from left to right, across a period• Successive ionization energies increase for

every electron removed

1st ionization energy

Electronegativity

• Reflects an atom’s ability to attract electrons in a chemical bond

• Related to its ionization energy and electron affinity

• Increases from left to right, across a period• Decreases from top to bottom, down a group

Shielding

• Shielding electrons are electrons located between the nucleus and the valence electrons

• For example:• Chlorine has the following electron configuration:

1s2 2s2 2p6 3s2 3p5

• The shielding electrons would be 1s2 2s2 2p6 • The valence electrons would be 3s2 3p5

• Then we have 10 shielding electrons and 7 valence electrons, right?

Shielding

• You try one• Try Sodium (Na)… Remember, first you have to know

the electron configuration

• Did you get:• Electron configuration: 1s2 2s2 2p6 3s1

• Shielding electrons: 1s2 2s2 2p6 • Valence electrons: 3s1

• So we have 10 shielding electrons and 1 valence electron, right?

Zeff

• Effective nuclear charge (Zeff) is the charge felt by the valence electrons after you’ve taken into account the number of shielding electrons that surround the nucleus.

• Huh?• Let’s put it in an equation

• Zeff =# of protons - # of shielding electrons

• So, calculate the effective nuclear charge for the all the elements in period 3

• Now, calculate the effective nuclear charge for all the elements in group 2

• What pattern do you see arising?

• What is the correlation between Zeff and atomic radius? (Remember opposite charges attract)

• The greater the Zeff the smaller the atomic radius • I’m still lost…..• Greater effective nuclear charge means that

the valence electrons are feeling a greater pull toward the nucleus, making the atom smaller in size

In summary…

• Effective nuclear charge can be used to predict trends in atomic radius• Increases from left to right and decreases from top to

bottom• Zeff = Z - σ

• Effective nuclear charge is dependent upon electron shielding

• Electronegativity increases from left to right and decreases from top to bottom