Periodic Table and

Periodicity

BHS Chemistry 2010- 2011

In 1869, Dmitri Mendeleev, a Russian chemist noticed patterns in certain elements.

He discovered a way to arrange the elements so that they were organized by their chemical and physical properties.

Moseley’s Contribution

Henry Moseley is

credited for further

arranging the elements

on the periodic table in

order of the number of

protons they contained.

Circle Periodic Table

Attention:

New Additions to Periodic Table

WOMANIUM (Wo)

Physical properties: Boils at nothing and may freeze any time.

Chemical properties: Very active and highly unstable. Possesses

strong affinity with gold, silver, platinum, and precious stones. Violent

when left alone. Turns slightly green when placed next to a better

specimen.

Usage: An extremely good catalyst for dispersion of wealth.

MANIUM (Xy)

Physical properties: Solid at room temperature but gets bent out of

shape easily.

Chemical properties: Becomes explosive when mixed with Childrium

for prolonged period of time.

Usage: Possibly good methane source.

Caution: In the absence of WO, this element rapidly decomposes and

begins to smell.

Periods

The table is arranged in horizontal

rows (going across) called periods.

There are 7 periods.

The period tells you how many

electron energy levels the atom

has.

1

5 4

3

2

7

6

Group

The table is also arranged in vertical

columns (going down) called groups.

There are 18 groups.

Members of each group have similar

physical and chemical properties.

1

2

3

2

4

2

5 6

2

7 8 9 10 11 12

2

13 14 15 16

18

17

III. Properties of Metals, Nonmetals and Metalloids

A. Properties of Metals

1. They are malleable, and have luster.

2. They reflect heat and light

3. Good conductors of heat and electricity.

4. Typically solids at STP

5. High melting points

6. They lose electrons in chemical reactions to become cations (+)

B. Properties of Nonmetals

1. They are dull and brittle.

2. Don’t conduct heat and electricity well.

3. Low boiling and freezing points.

4. They exist in all three phases at STP, but most are gases.

5. They gain electrons in chemical reactions to become anions (negative ions)

C. Properties of Metalloids

(B, Si, Ge, As, Sb, Te)

1. They possess

intermediate properties

between metals and

nonmetals.

2. They are

semiconductors at

higher than room

temperatures.

3. They are all solids at

STP.

IV. Special Groups

A. Alkali Metals – Group 1 (IA)

on the periodic table.

1. They are soft and easily

cut.

2. They are highly reactive. (especially with H2O)

Highly Reactive Video

3. All have an electron

configuration ending in s1

4. Gives up 1 electron in

bonding (+1)

5. Has 1 valence electron

B. Alkaline Earth Metals – Group 2 (IIA) on the periodic table.

1. They are less reactive than the alkali metals.

2. They have an electron configuration ending in s2

3. Gives up 2 electrons to form a (2+) charge

4. Has 2 valence electrons

C. Halogens – Group 17 (VIIA) on the periodic table.

1. They are highly reactive and react violently with hot metals.

2. They form diatomic molecules

3. They have an electron configuration ending in s2p5

4. Accepts 1 electron to have a (-1) charge

5. Has 7 valence electrons

D. Noble Gases – Group 18 (VIIIA) on the periodic table.

1. Extremely stable and unreactive

2. They exist as single atoms

3. They have an electron configuration ending in s2p6

4. Has 8 valence electrons

E. Transition Metals – Groups 3 - 12 (B groups) on the periodic table.

1. They possess characteristics of active metals to varying degrees.

2. They form compounds that are usually brightly colored.

F. Inner Transition Metals – two rows at the bottom of the periodic table.

1. Many of the inner transition metals are radioactive.

2. Many of the actinides are synthetic.

“The Periodic Table”

The Periodic Law says: When elements are arranged in order

of increasing atomic number, there is a periodic repetition of their physical and chemical properties.

Horizontal rows = periods

Vertical column = group (or family)

• Similar physical & chemical prop.

ALL Periodic Table Trends Influenced by three factors:

1. Energy Level

• Higher energy levels are further

away from the nucleus.

2. Charge on nucleus (# protons)

• More charge pulls electrons in

closer. (+ and – attract each other)

3. Shielding effect

Shielding The electron on the outermost

energy level has to look through all

the other energy levels to see the

nucleus.

This effect decreases

the attraction of the

nucleus for the outer

electrons..

What do they influence?

Energy levels and Shielding have an

effect on the GROUP ( )

Nuclear charge has an effect

on a PERIOD ( )

Atomic Size

Measure the Atomic Radius - this is half the distance between the two nuclei of a diatomic molecule.

} Radius

#1. Atomic Size - Group trends

As we increase

the atomic

number (or go

down a group). . .

each atom has

another energy

level,

so the atoms get

bigger.

H

Li

Na

K

Rb

#1. Atomic Size - Period Trends

Going from left to right across a period,

the size gets smaller.

Electrons are in the same energy level.

But, there is more nuclear charge.

Outermost electrons are pulled closer.

Na Mg Al Si P S Cl Ar

Ions

Some compounds are composed of

particles called “ions”

• An ion is an atom (or group of atoms)

that has a positive or negative charge

Atoms are neutral because the number

of protons equals electrons

• Positive and negative ions are formed

when electrons are transferred (lost or

gained) between atoms

Ions Metals tend to LOSE electrons,

from their outer energy level

Nonmetals tend to GAIN one or

more electrons

Ions

Here is a simple way to remember

which is the cation and which the

anion:

This is a cat-ion. This is Anion.

She’s unhappy and

negative.

+ +

Ionic Radius

The size of an ion

Cations are smaller (lost e-) and

anions are larger than the atoms they

cam from.

Ionic radius Group trends

Each step down a

group is adding an

energy level

Ions therefore get

bigger as you go

down, because of

the additional

energy level.

Li1+

Na1+

K1+

Rb1+

Cs1+

Ionic radius Period Trends Across the period from left to

right, the nuclear charge

increases - so they get smaller.

Notice the energy level changes

between anions and cations.

Li1+

Be2+

B3+

C4+

N3- O2-

F1-

Trends in Ionization Energy

Ionization energy is the amount of energy required to completely remove an electron.

Ionization Energy - Group trends

As you go down a group, the first IE decreases because...

• The electron is further away from the attraction of the nucleus, and

• There is more shielding.

Ionization Energy - Period trends

IE generally increases from left

to right.

Same shielding.

But, increasing nuclear charge

The arrows indicate the trend:

Ionization INCREASE in these

directions

Trends in Electronegativity

Electronegativity is the tendency

for an atom to attract electrons to

itself.

An element with a big

electronegativity means it pulls the

electron towards itself strongly!

Electronegativity Group Trend The further down a group, the farther

the electron is away from the nucleus, plus the more electrons an atom has.

Thus, more willing to share.

Going down a group, EN decreased

Electronegativity Period Trend Metals

• They want to lose electrons

• Low electronegativity

Nonmetals.

• They want more electrons.

• Going across a period, the EN

increases

The arrows indicate the trend:

Electronegativity INCREASE in these

directions

Textbook:

Atomic Radius: pg. 141

Ionization Energy: pg. 143

Ionic Radius: pg. 149

Electronegativity: pg. 151