the periodic table dimitri mendeleev russian scientist developed the first published table in 1869....

The Periodic Table

Dimitri Mendeleev

• Russian scientist developed the first published table in 1869.

• Arranged elements in order of atomic mass

• Elements with similar properties were placed in columns

• left spaces for undiscovered elements

Periodic Law the pattern of the table

• Mendeleev’s table was called a periodic table because……..

• If the elements were arranged in order of increasing atomic mass then elements with similar properties would show up PERIODICALLY.

Henri Moseley (1914)

• Used X-rays to reveal atomic numbers of several elements.

• Suggested that the elements should be arranged in order of atomic number instead of atomic mass. – Today that is how it is arranged

Reading the Periodic Table

• Families or Groups– Up and down

• Numbered 1 to 18• Elements in groups (families) have

similar characteristics

• Periods–Across the table 1-7

Classifications of Elements

MetalsNon-metals Metalloids (semi-metals)

Metals

• Elements to the left of the steps• except H, Ge & Sb

MetalsGe

Sb

Metal characteristics

• Solids at STP (except mercury)

• good conductors of electricity and heat– Mobile electrons

• shiny luster• ductile - drawn (made) into wire

• malleable - hammered into sheets

• when combined in compound have a positive oxidation state (+, cation)

• Metallic character increases R to L–Most reactive metals found in

Group 1

Non-metals

H

18

B

Si

As

Te

Non-metal characteristics

• Solid, liquid or gas at STP• used as insulators• solids are dull• brittle• when combined in a compound

have negative oxidation state (-, anion)

• Metallic character decreases L to R

Metalloids

• a.k.a. Semimetals

• both metal and non-metal properties

• “sit on the steps”

• 2 hide underneath

BSi

AsTeSb

Ge

Elements are (s), (l) & (g)

• Solid – particles have vibratory motion and are

tightly packed– Almost ALL of the elements

• Liquid – particles can move throughout substance – particles are farther apart than in solids– conforms to shape of container

– ONLY Hg and Br

• Gas– molecules in constant, random,

straight line motion– molecules fill container– large space between “volumeless”

molecules– conforms to the shape of the

container– H, N, O, F, Cl and Group 18

Allotropes

• Different forms of an element found naturally in the same state– Different molecular structures

Carbondiamond, coal, graphite

OxygenO2 & O3

N O F

Cl

Br

I

Diatomic elements

• Elements found bonded to itself “uncombined”

• N2 O2 F2 Cl2 Br2 I2 and H2

H

Monoatomic elements

18

Ne

Ar

Kr

Xe

Rn

HeValence shellfilled with 2e-

Valence shellfilled with8 e-

• Inert• Don’t react

or form compounds

Valence electrons

• The number of valence electrons correspond to group number

• 1, 2 , 13 - 18

Last digit• The Energy Level (shell) they are in

is the same as the period #

Group / Family Characteristics

Each group has characteristic properties that are directly related to electron configuration & especially the number of valence electrons

Group 1

• Alkali Metals• all elements except hydrogen• most reactive metal group• SO reactive they are never found

alone in nature– always bonded to another element

• form +1 ions

Group 2

• Alkaline Earth Metals• second most reactive metals group• SO reactive they don’t occur alone

in nature– always bonded to another element

• form +2 ions

d- block elements

• Transition Metals

COLORful

ions and solutions

• Transition metals not as reactive as other metals

• some found in free state– Au, Ag, Pt

• multiple positive oxidation states– Iron (IV) oxide FeO2

– Iron (II) oxide FeO

Group 17

• Halogens• Most reactive nonmetal group• SO reactive that they don’t occur

by themselves in nature– at least bonded to themselves

(diatomic)• form -1 ions

Group 18

• Noble gases• “Inert” gases• don’t like to combine with other

elements– Valence shells filled

Ionization Energy

Ionization Energy (first)

• Definition: – energy needed to remove the most

loosely held e-

Valence electronhigh E electronOutermost electron

What’s holding the e- in place?

• Nuclear charge– There is a force of attraction between

protons (+) in the nucleus and electrons (-) in the orbitals

– “Opposites attract”

Ionization energy

• Hint: energy needed to make an ion by losing electrons

MOST elements want 8 e- (octet rule)

• Elements with only a few valence electrons will tend to have lower ionization E

• In other words…– It doesn’t take a lot of E to remove

their e-

Metals

MOST elements want 8 e- (octet rule)

• Elements with almost 8 valence electrons won’t give them up so easy– It takes a lot of E to remove their e-

• Non-metals have high Ionization E

Who has the highest IEs?

Graph Ionization Energies

• Label X axis with Atomic #1-20• Label y axis with 1st IE (KJ/mol

atoms)

• AFTER you plot the points– Label data points with element

symbols – Connect data points of same period

only– Separate your graph into Periods

Period 2 Period 3Period 1 Period 4

Fir

st I

oniz

atio

n E

nerg

y (K

J / m

ol)

H

He

Li

Ne

Na

Ar

Ca

What is the periodic trend for ionization energy?

INCREASES

Ionization E trend?

• As move across a period…IE increases. WHY?

3 p+ 9 p+

• As you go across a period–a large proton and a tiny electron

are being added. – more p+ hold the e- tighter

• As move UP a group … IE increases. WHY?Period 1Period 2Period 3Period 4

Let’s figure it out . . .

Potassium

19 p+

+

Sodium11 p+

+

Lithium

3 p+

+

• As you go up a group – the valence shell gets closer to

the nucleus and can hold on to the e- tighter.

because . . .INCREASES

SHIELDING

• Reason why IE decreases as you go down a group

• the distance increases between the p+ and valence e-–AND

• e- in outer shells repel each other

–WHICH IS WHY low IE are so reactive!

Electronegativity

• Ability of an atom to attract electrons of other atoms.

• In other words. . . – Atoms with high e-neg are bullies that

steal electrons

Fluorine

Electronegativity Graph

Electronegativity

• Determine the periodic trend

INCREASES

EXCEPT for Noble gases. WHY?

F

Atoms with high electronegativities also havehigh ionization E

Atomic Radius

• Radius – distance from the center of a circle

(Nucleus) to the outermost edge (valence shell)

R

Atomic Radius Periodicity

DECREASES

DECREASES

Why?????

• Why does atomic radius DECREASE as you move up a group?

• Why does atomic radius DECREASE as you move across a period?

• Losing layers of e-

• Increasing the # of p+ holds the e- in tighter

• Increasing NUCLEAR CHARGE

• What happens to atomic radius as you create a + ion?

• radius decreases

• What happens to atomic radius as you create a - ion?

• radius increases

• Compare the ionic radus of Mg2+ and the atomic radius of Ne.

The radius of the Mg2+ ion is smaller than the atom of Ne, because the Mg2+ ion has more p+ (12) than the Ne atom (10).

Compare . . .

• Fluoride ion & Fluorine atom

• Sodium ion & Sodium atom

• Fluoride ion & Neon atom

General Formulas of compounds

• You can look at the groups on the periodic table and determine how they will combine with elements of different groups.

Write these formulas

• lithium fluoride• sodium fluoride• potassium fluoride• lithium chloride• sodium chloride• potassium chloride

• What do you notice?

• Write the general formula for Group 1 and Group 17 elementsAB

Write these formulas

• beryllium fluoride• magnesium fluoride• calcium fluoride• beryllium chloride• magnesium chloride• calcium chloride

• What do you notice?

• Write the general formula for Group 2 and Group 17 elements

AB2

Write these formulas

• lithium oxide• sodium oxide• potassium oxide• lithium sulfide• sodium sulfide• potassium sulfide

• What do you notice?

• Write the general formula for Group 1 and Group 16 elements

A2B

Write these formulas

• beryllium oxide• magnesium oxide• calcium oxide• beryllium sulfide• magnesium sulfide• calcium sulfide

• What do you notice?

• Write the general formula for Group 2 and Group 16 elements

AB

Formula Writing & Naming Review

• Lead IV oxide• Phosphorus pentoxide• SO3

• Oxygen• Argon• Aluminum oxide• NiO

Different Forms of the Periodic Table

• Changes were made to Mendeleev’s table to look like the modern Periodic table we use today.

• This is not the only form of the periodic table that exists, however it is the most widely accepted.

Stowe’s Physicists p.t.

Benfey p.t.

Zmaczynski p.t.

Alexander Arrangement p.t.

立体周期表の組み立て方 p.t.

THE END