Chapter Chapter 6.1-6.36.1-6.3

Periodic Table LecturePeriodic Table Lecture

Do members of the same family,

generally behave the same?Yes

The Periodic TableThe Periodic Table

The Alkali MetalsLithium, Sodium, Potassium, Rubidium, Cesium, and francium very reactive 1 valence electron s1 sublevel is filled

Alkali Earth Metals

Beryllium, Magnesium, Calcium, Strontium, Barium, and Radium2 valence electronss2 sublevel is filled

The Transition Metalsmetals with atomic numbers 21-112 highest s & d sublevels have electrons

MetalloidsLike metals & nonmetals

Boron, Silicon, Germanium, Arsenic, Antimony, Tellurium, Polonium

Nonmetals• Consists of Carbon, Nitrogen, Oxygen, Phosphors, Sulfur, Selenium• poor conductors of heat and electricity compared to metals• dull and brittle

Halogens• Consists of Fluorine, Chlorine, Bromine, Iodine, Astatine• nonmetals• have 7 valence electrons• very reactive• want one more electron (octet rule)

Noble Gases• consists of Helium, neon, Argon, Krypton, Xenon, Radon• unreactive stable inert because they already have 8 valence electrons

Inner Transition metals• consists of elements with atomic numbers 58 through 71 and 90 through 103•F sublevels partially filled

• the Lanthanide Series has atomic numbers 58 -71 and the Actinide Series has atomic numbers 90-103

Other Metals

Define the term inert gas?

noble gas –unreactive & stable

Group 1A

Group 2A

Group 3A

Group 4A

Group 5A

Group 6A

Group 7A

Representative Elements #1 – Group IA-VIIA outer s & p orb partially filledAlkali Metals

Alkaline Earth

Nonmetals/Metalloids

Halogens

ns1

ns2

ns2 np1

ns2 np2

ns2 np3

ns2 np4

ns2 np5

Group 0

8 or 18

Noble Gases ns2 np6

Representative Elements #1 Lewis dot structure

1s2 2s2 2p6 1s2Na

Group B Transition Metals

Filling the “d” orbital

Group 58-71 Lanthanides Filling the “4f” orbital

Group 90-103Actinides Filling the “5f” orbital

A. Ionic SizeA. Ionic Sizemetals (group 1A-3A)

lose electrons to become stable cation

non-metal (group 5A-7A) gain electrons to become stable anions.

1A =

2A =

3A =

5A =

6A =

7A =

Loses 1 e-

Loses 2 e-

Loses 3 e-

Gains 3 e-

Gains 2 e-

Gains 1 e-

7PERIODS

! !v

A Family is a Group living between Columns

Periodic Table Song by Tom Lehrer abovePeriodic Table Song by Tom Lehrer above

End of Lecture 6.1End of Lecture 6.1

Next Lecture 6.2Next Lecture 6.2

http://www.privatehand.com/flash/elements.html

Who designed the 1st periodic table in 1869?

Dmitri Mendeleev

grouped w/ similar chemical and physical properties & ordered by atomic mass.Ex:

Co Ni

Ar K

Te I

http://www.youtube.com/watch?v=y7dmRtlXaYQ

http://www.youtube.com/watch?v=zUDDiWtFtEM

Lecture 6.3Lecture 6.3Periodic TrendsPeriodic Trends

I. Periodic Trends - Atomic Size

Atomic Radii:

NucleusDistance between

nuclei

Atomic Radius

Measured as 1/2 distance between nuclei 2 atoms

Atomic Size generally INCREASES as you move down a group on the periodic table.

Why? down a group

increases # of energy levels

Example:Ca atom larger than a Mg atom. Why?

An energy level is added!

Atomic Size generally DECREASES across a row on the periodic table.Why? adding more p+ pulls in extra

electrons

Na < ionization energy than O because less protons pull.

RELATIVE ELECTRONEGATIVITY, IONIZATION RELATIVE ELECTRONEGATIVITY, IONIZATION ENERGY, RADII, SHIELDING ETC…ENERGY, RADII, SHIELDING ETC…

Hydrogen

2.1

Oxygen

3.5

Carbon

2.5

Sodium

0.9

Electro negativities:Hydrogen has the smallest atomic radius

B. Ionization EnergyB. Ionization Energyenergy needed to pull an electron away from an atom.

B. Ionization EnergyB. Ionization Energy

Example : Na Na+1 + e-

Ionization energy decreases as you move down a group.increased distance from protons

reduces attractive force

Why?

Period TrendPeriod Trend: : Ionization energy generally increases as you move across a period.

nuclear charge increases (more protons)

which increases attractive forces

Why?

energy required to remove the energy required to remove the 1st1st outermost electron is outermost electron is 1st ionization 1st ionization energy.energy.

What is the second ionization energy?

Which is harder to remove?

Why?

What happens to the shielding of the nucleus as you move across a period?

•ONLY adding electrons, NOT a new energy level.

Remains constantWhy?

What happens to the shielding of the nucleus as you move down a group?

another energy level that shields those valence electrons.

IncreasesWhy?

CaCa++ions ions –– smallersmaller than the original than the original atomatom

When electrons lost,

a whole energy level lost

decreases radius.

Why?

Negative anions grow Negative anions grow largerlarger

there are more e- than p+

(increased electron repulsion),

Why?

Natom N-3

anion

from group 5A to the right,from group 5A to the right,aannions ions gradually decreasegradually decrease in sizein size

groups 6A &7A only gain 1 or 2 e-

Have Same # of e-, but increased # of p+

Why?

N-3 O-2 F-1

anion anion anion.

B. ElectronegativityB. Electronegativity

Noble gases no electronegative #

Why?inert / don’t form compounds.

Can’t force a noble gas to take an electron – they have s2 p6

3. Period Trends3. Period Trendsleft to right electronegativity increases. Why?High ionization energy = high electronegativity

Resists electron

loss

Attracts electrons

Fluorine is the most electronegative!

4. Group Trends4. Group TrendsElectronegativity decreases down a group.

Why?

Increased energy levels and shielding

Cs has the lowest electronegativity

3 alkalis.MOV