Unit 5 – The Periodic

Table & Periodicity

Development of the Modern

Periodic Table

• Antoine Lavoisier – 1789 – 33 elements (23 actual elements)

Lavoisier's table of simple substances

Further reading:

Gases

New names (French) Old names (English translation)

Lumière Light

Calorique

Heat

Principle of heat

Igneous fluid

Fire

Matter of fire and of heat

Oxygène

Dephlogisticated air

Empyreal air

Vital air

Base of vital air

Azote

Phlogisticated gas

Mephitis

Base of mephitis

Hydrogène Inflammable air or gas

Base of inflammable air

Metals

New names (French) Old names (English translation)

Antimoine Antimony

Argent Silver

Arsenic Arsenic

Bismuth Bismuth

Cobolt Cobalt

Cuivre Copper

Étain Tin

Fer Iron

Manganèse Manganese

Mercure Mercury

Molybdène Molybdena

Nickel Nickel

Or Gold

Platine Platina

Plomb Lead

Tungstène Tungsten

Zinc Zinc

Nonmetals

New names (French) Old names (English translation)

Soufre Sulphur

Phosphore Phosphorus

Carbone Pure charcoal

Radical muriatique Unknown

Radical fluorique Unknown

Radical boracique Unknown

Earths

New names (French) Old names (English translation)

Chaux Chalk, calcareous earth

Magnésie Magnesia, base of Epsom salt

Baryte Barote, or heavy earth

Alumine Clay, earth of alum, base of alum

Silice Siliceous earth, vitrifiable earth

Development of the Modern

Periodic Table

• John Dalton – 1808 – 36 elements

O

Oxyge

n

H

Hydrogen

N

Nitrogen

C

Carbon

S

Sulphur

P

Phosph

orus

Au

Gold

Pt

Platinum

Ag

Silver

Hg

Mercu

ry

Cu

Copper

Fe

Iron

Ni

Nickel

Sn

Tin

Pb

Lead

Zn

Zinc

Bi

Bismuth

Sb

Antimony

As

Arseni

c

Co

Cobalt

Mn

Manganese

U

Uranium

W

Tungsten

Ti

Titaniu

m

Ce

Ceriu

m

K

Potassiu

m

Na

Sodium

Ca

Calciu

m

Mg

Magnesium

Ba

Barium

Sr

Strontiu

m

Al

Aluminium

Si

Silicon

Y

Yttri

um

Be

Berylliu

m

Zr

Zirconiu

m

Development of the Modern

Periodic Table

• Jöns Berzelius -1814 – 47 elements

– Used letters for symbols

Element Berz. present

Aluminium Al

Argentum (Silver) Ag

Arsenic As

Aurum (Gold) Au

Barium Ba

Bismuth Bi

Boron B

Calcium Ca

Carbon C

Cerium Ce

Chromium Ch Cr

Cobalt Co

Columbium Cl (Cb) Nb

Cuprum (Copper) Cu

Ferrum (Iron) Fe

Fluoric Radicle F

Element Berz. present

Glucinum Gl Be

Hydrargyrum (Mercury) Hg (Hy) Hg

Hydrogenium H

Iridium I Ir

Magnesium Ms Mg

Manganese Ma (Mn) Mn

Molybdenum Mo

Muriatic Radicle (Chlorine) M Cl

Nickel Ni

Nitric Radicle N

Osmium Os

Oxygenium O

Palladium Pa Pd

Phosphorus P

Platinum Pt

Plumbum (Lead) Pb (P) Pb

Element Berz. present

Potassium Po K

Rhodium Rh (R) Rh

Silicium Si

Sodium So Na

Stibium (Antimony)* Sb (St) Sb

Strontium Sr

Sulphur S

Tellurium Te

Tin Sn (St) Sn

Titanium Ti

Tungsten Tn (W) W

Uranium U

Yttrium Y

Zinc Zn

Zirconium Zr

Development of the Modern

Periodic Table

• Johann Döbereiner – – Classified the elements

into “triads” or groups of

3 elements with similar

chemical whose physical

properties varied in a

predictable way

according to their atomic

masses

1829

Development of the Modern

Periodic Table

1H 7Li 9Be 11B 12C 14N 16O

19F 23Na 24Mg 27Al 28Si 31P 32S

35Cl 39K 40Ca 52Cr 48Ti 55Mn 56Fe

• John Newlands -1864 – “Law of Octaves”

– Organized by increasing atomic mass

– Repeating properties

Development of the Modern

Periodic Table

• Dmitri Mendeleev – 1869 – 63 elements

– Rows (later columns) of similar

chemical properties

– Increasing atomic mass

– Missing elements

Mendeleev’s Table of 1869

In 1875, a French chemist

discovered Gallium

(eka-aluminum) and its

properties were very close

to what Mendeleev

predicted!

Mendeleev predicted the

existence of unknown

elements like eka-aluminum

Mendeleev’s Revised Table

1871

Development of the Modern

Periodic Table

• Henry Moseley – 1913 – Rearranged the table by increasing atomic

number

Development of the Modern

Periodic Table

• History Review – Lavoisier – list of elements

– Berzelius – symbols as letters

– Mendeleev – table with similar chemical

properties and increasing atomic mass

– Moseley – increasing atomic number

Development of the Modern

Periodic Table

• Modern Periodic Table

– Group (or family) – column

– Period – row

Development of the Modern

Periodic Table

• Periodicity is the tendency to recur

at regular intervals.

• Periodic law – when organized by

increasing atomic number, there is a

periodic repetition of chemical and

physical properties.

Why?

What is similar for all elements

in a group or family?

They have the same number of outer level electrons!!!

These have 1

These have 3

Classification of the Elements

Electrons, particularly

the valence electrons,

control many of the

chemical and physical

properties of atoms!

Valence Electrons

Classification of the Elements

• Elements can be classified into

four categories based upon their

electron configurations 1. Noble gases

• He, Ne, Ar, Kr, Xe, and Rn

• Full outer s and p sublevels

2. Representative elements

• Groups 1A – 8A (this includes Noble Gases)

• Partially filled s and p sublevels

3. Transition metals

• d block

• Outer s and nearby d sublevel contain

electrons

4. Inner transition metals

• f block

• Outer s and nearby f sublevel contain

electrons

Classification of the Elements

Noble Gases

Inner Transition Metals

Transition Metals

Representative Elements

• Classifying by properties 1. Metals – left of stairs

– Lustrous (shiny)

– Malleable (not brittle)

– Ductile (drawn into a wire)

– Good conductor of heat and electricity

2. Nonmetals – right of stairs

– Dull-looking

– Brittle

– Poor conductors

3. Metalloids – on the stairs (minus aluminum)

– Properties of both metals and nonmetals

Classification of the Elements

Blue = metals

Green = metalloids

Yellow = nonmetals

• Families – Noble Gases – group 8A, 0, or 18

– Alkali metals – group 1A (minus hydrogen)

– Alkaline Earth metals – Group 2A

– Halogens – group 7A

– Other families are just referred to as the

element at the top of the column

• Carbon family for group 4A

Classification of the Elements

Periodic Trends

• Arrangement of elements on the periodic table is linked to the electron configuration, many trends can be used to predict chemical and physical behavior.

• To understand and use these trends we must first understand how NUCLEAR CHARGE and SHELLS AND SHIELDING influence electron behavior.

ALL TRENDS ARE EXPLAINED BY NUCLEAR CHARGE AND

SHELLS AND SHIELDING

Periodic Trends

• Nuclear Charge – As you move across a period or down a

group, the atomic number increases.

– This means the number of protons in the

nucleus is increasing.

– With more protons, the positive pulling

strength (nuclear charge) of the nucleus is

increasing

Nuclear Strength

Increases

Increases

Periodic Trends

• Shells – Energy levels

• The higher the level, the farther from the nucleus

• Across – highest energy does not change

• Down – energy levels increase

• Shielding – Inner level electrons interfere or shield the valence

electrons from the nucleus • Across – shielding is constant

• Down – shielding increases

Shells and Shielding

Constant

Increases

Nuclear Charge verses Shells

and Shielding

• Period Trend – Nuclear Strength Wins!!!

• Because Shells and Shielding are constant across a

period they don’t affect period trends

• Therefore, ALL PERIOD TRENDS are caused by

increasing nuclear charge.

Nuclear Charge verses Shells

and Shielding

• Group Trend – Shells and Shielding Win

• Even though nuclear charge is increasing, more Shells (your farther from the nucleus) and Shielding (inner level electron interference) decreases the effective nuclear strength.

• Therefore, ALL GROUP TRENDS are caused by shells and shielding or their effect on the nuclear charge.

Nuclear Charge verses Shells

and Shielding

Nuclear Strength Increases S/S

Increases lo

werin

g

effective N

uclear S

trength

Atomic Radius

• Atomic radius – estimated as ½ the distance

between the nuclei of 2 like atoms in a

diatomic molecule.

• As atomic radius increases, the element

increases in size.

Atomic Size

Atomic Size

• Group Trend – Increases down a group

• Caused by an increase in Shells and Shielding

• Period Trend – Decreases across a period

• As nuclear strength increases the nucleus

pulls the outer electrons closer

Atomic Size

Size Decreases Size In

creases

Ionization Energy

• Ionization energy – energy needed to remove

an electron from an atom

Na(g) Na+

(g) + e-

Ionization Energy

• To remove an electron you have to overcome the nucleus’ hold (nuclear charge) on the electron. – 1st Ionization Energy – Energy needed to remove

first electron.

– 2nd Ionization Energy – Energy needed to remove a second electron.

• This is always higher than the 1st ionization energy.

• When an electron is removed, the nucleus has a stronger hold on the remaining electrons.

• When you have a noble gas electron configuration it becomes very difficult to remove an electron.

Ionization Energy

Ionization Energy

• Group Trend – Decreases as you go down a group

1. Shells – the farther the outer electrons are farther

from the nucleus, the weaker the pull.

2. Shielding – the inner level electrons block the

nucleus’ ability to attract the valence electrons.

• Period Trend – Increase as you go across a period

– Greater nuclear strength makes it harder to

remove electrons

Ionization Energy

1st Ionization Energy increases

1st Io

nizatio

n E

nerg

y d

ecreases

Ion Formation

• Octet Rule – atoms will gain or lose electrons

(sometimes even sharing like in molecules) to

acquire a full set of 8 valence electrons.

– Metals will lose electrons

– Nonmetals gain electrons

To hydrogen atoms are walking

down the street.

Hey, I think I

just lost an

electron! Are you

sure!

Yeah!

I’m

POSITIVE

!

Ion Formation

• Cation – formed when electrons are removed from a neutral atom – Nucleus has a stronger pull on the remaining electrons

decreasing the size

– Usually involves a decrease in number shells

The Cation is

smaller than the

neutral atom!

Ion Formation

• Anion – formed when electrons are added to a neutral atom – Nucleus has a weaker hold on the

increased number of valence electrons and the ion size increases

The Anion is

larger than the

neutral atom!

Ion Size

• Period Trend – Cations and Anions both decrease in size across a

period

• Increased nuclear charge pulls in the valence electrons

• Group Trend – Cations and Anions both increase in size down a group

• Increased shells and shielding mean greater ion size

Ionic Size

Anions decrease Cations decrease

Both

increase

Electronegativity

• Electronegativity – attraction of

one atoms nucleus to another

atoms electrons when they are

chemically bonded

O H H

Water Molecule

Measure of how

strong the oxygen

nucleus attracts the

hydrogen’s electron

Electronegativity

H

2.20

Li

0.98

Be

1.57

B

2.04

C

2.55

N

3.04

O

3.44

F

3.98

Na

0.93

Mg

1.31

Al

1.61

Si

1.90

P

2.19

S

2.58

Cl

3.16

K

0.82

Ca

1.00

Ga

1.81

Ge

2.01

As

2.18

Se

2.55

Br

2.96

Rb

0.82

Sr

0.95

In

1.78

Sn

1.96

Sb

2.05

Te

2.1

I

2.66

Cs

0.79

Ba

0.89

Tl

1.8

Pb

1.8

Bi

1.9

Po

2.0

At

2.2

Electronegativity

• Period Trend – Increases across a period

• Nuclear strength increases

• Greater hold on electrons

• Group Trend – Decreases down a group

• Shells and shielding mean the effective nuclear strength decreases down the group

• Weaker hold on the electrons

Electronegativity

Electronegativity increases

Electro

neg

ativity

decreases

Chapter 7

• Allotrope - 2 or more different molecular forms of

the same element in the same state.

– Oxygen

• O2 – air

• O3 – ozone

– Carbon

• 8 allotropes so far

• amorphous carbon allotrope, carbon nanofoam,

carbon nanotube, the diamond allotrope, fullerene

allotrope, graphite, lonsdaleite, and ceraphite

allotrope.

Graphite Diamond

Three

allotropes of

carbon

http://www.creative-chemistry.org.uk/molecules/carbon.htm

Coal

Other allotropes of carbon