Periodic Table and Trends 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 Nitrogen Family Oxygen Family Groups or Families Inner Transition Metals Periods Carbon Family Boron Family Alkali Metals Alkaline Earth Metals Halogens Noble Gases 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

gold lead copper nickel

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 BROMINE SULFUR ARGON

Metalloids Along Stair-step line Have properties of metals AND nonmetals Many are used in transistors, found in electronics

Silicon Antimony 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 Did you get: 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 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