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The Periodic Table

History

1829 German J. W. Dobereiner

Grouped elements into triads

One of these triads included chlorine, bromine, and

iodine; another consisted of calcium, strontium, and

barium. In each of these triads, the atomic weight of

the intermediate element is approximately the

average of the atomic weights of the other two

elements. The density of that element is

approximately the average of the densities of the

other two elements.

History

The problem with this arrangement was that

Dobereiner’s model became outdated as new

elements were identified.

A good model is able to incorporate newly

understood information.

Dobereiner’s Triad Model was not useful,

since several newly discovered elements did

not “fit” into it.

Not all elements had triads

History

Russian scientist Dmitri Mendeleev

taught chemistry in terms of

properties

Mid 1800 – atomic masses of

elements were known

Wrote down the elements in order of

increasing atomic mass

Found a pattern of repeating

properties

Mendeleev’s TableGrouped elements in columns by similar

properties in order of increasing atomic

mass

Found some inconsistencies - felt that

the properties were more important than

the mass, so switched order.

• Found some gaps

• Must be undiscovered elements

• Predicted their properties before they

were found

The Modern TableElements are still grouped by properties

• Similar properties are in the same

column

Order is in increasing atomic number

Added a column of elements Mendeleev

didn’t know about.

• The noble gases weren’t found because

they didn’t react with anything.

Horizontal rows are called periods

There are 7 periods

Vertical columns are called groups.

Elements are placed in columns by

similar properties.

Also called families

1A

2A 3A 4A 5A 6A7A

8A

0

The elements in the A groups

are called the representative

elements

Metals

Metals Luster – shiny.

Ductile – drawn into wires.

Malleable – hammered into sheets.

Conductors of heat and electricity.

Transition metals The Group B

elements

Non-metals Dull

Brittle

Nonconductors

- insulators

Metalloids or Semimetals Properties of both

Semiconductors

These are called the inner

transition elements and they

belong here

Group 1A are the alkali metals

Group 2A are the alkaline earth metals

Group 7A is called the Halogens

Group 8A are the noble gases

Group Characteristics Alkali metals

• Group 1: very reactive metals which do not occur freely in nature.

1 electron in outer shell

Alkaline Earth Metals

• Group 2: next reactive metals, found in earths crust but not in elemental

form.

2 electrons in outer shell

Transition Elements

• Group 3-12: metals with varying reactivities. Greater density than Group 1

or 2 elements.

1-2 electrons in outer shell

Lanthanides and Actinides

• These elements are also transition elements but have been taken out to

prevent the periodic table being so wide.

Boron Group

• Group 13: reactive, contains metal and metalloid.

3 electrons in outer shell

Carbon Group

• Group 14: contains metalloids, metals and non metals.

4 electrons in outer shell

Nitrogen Group

• Group 15: contains metalloids, metals and non metals.

5 electrons in outer shell

Oxygen Group

• Group 16: contains contains metalloids, metals and non metals. Reactive

6 electrons in outer shell

Halogens

• Group 17: non-metals, very reactive.

7 electrons in outer shell

Nobel gas

• Group 18: non-metals, non reactive.

8 electrons in outer shell

Periodic Table Part 2

Periodic trends

Identifying the patterns

What we will look for

Periodic trends-

• How properties vary as you go across

a period

Group trends

• How properties vary as you go down

a group

Why?

• Explain why they vary

What we will investigateAtomic size

• how big the atoms are

Ionization energy

• How much energy to remove an electron

Electronegativity

• The attraction for the electron in a compound

Ionic size

• How big ions are

Ionization Energy

The amount of energy required to

completely remove an electron from

a gaseous atom.

• Removing one electron makes a +1

ion

The energy required is called the first

ionization energy

What determines IE

The greater the nuclear charge the

greater IE.

Increased atomic radius decreases IE

Group trendsAs you go down a group first IE

decreases because of

• Bigger atoms…so outer electron

less attracted even though there

are more +charges in the nucleus

Periodic trends

All the atoms in the same period

• Same approximate size

• Increasing nuclear charge

So IE generally increases from left to

right.

Discussion Question

What is Atomic Radius?

• The radius of an atom

• The distance an atom travels

• The width of an atom

• The distance around an atom

Atomic Size

First problem where do you start

measuring

The electron cloud doesn’t have a

definite edge.

They get around this by measuring

more than 1 atom at a time

Atomic Size

Atomic Radius = half the distance

between two nuclei of molecule}

Radius

Trends in Atomic Size

Influenced by two factors

• Energy Level (Electron Shell)– Higher energy level (shell) is further away

• Charge on nucleus– More charge pulls electrons in closer

Group trendsAs we go down a

group each atom

has another

energy level

(electron shell)

So the atoms get

bigger

H

Li

Na

K

Rb

Periodic TrendsAs you go across a period the radius

gets smaller.

Same energy level (electron shell)

More nuclear charge

Pulls outermost electrons closer

Na Mg Al Si P S Cl Ar

What is Ionic Radius?

• the radius of an atom

• the radius of the most common ion of

an atom

• the size of an atom

• the width of an ion

Ionic Size

Cations are positive ions

• Cations form by losing electrons

• Metals form cations

Cations are smaller than the atom

they come from

Ionic size

Anions are negative ions

• Anions form by gaining electrons

• Nonmetals form anions

Anions are bigger than the atom they

come from

Group trends

Adding energy level

(electron shell)

Ions get bigger as

you go down

Li1+

Na1+

K1+

Rb1+

Cs1+

H1+

Periodic Trends

Across the period nuclear charge

increases so they get smaller.

Energy level (electron shell) changes

between anions and cations

Li1+

Be2+

B3+

C4+

N3-

O2- F1-

Electronegativity

The tendency for an atom to attract

electrons to itself when it is

chemically combined with another

element.

• How “greedy”

Big electronegativity means it pulls

the electron toward it.

Group Trend

The further down a group

• Bigger atoms (outer electrons further

from nucleus)

• More electrons per atom

Less attraction for electrons

Lower electronegativity.

Periodic Trend

Metals - left end

• Low nuclear charge

• Low attraction for extra electrons

Low electronegativity

Right end - nonmetals

• High nuclear charge

• Large attraction for extra electrons

High electronegativity

Not noble gases- no compounds

Ionization energy, electronegativity

INCREASE

Atomic size increases,

Ionic size increases

Nuclear Charge

Energ

y L

evels

& S

hie

ldin

g

How to answer why questions

Trend

• Periodic and Group

Reason

• Nuclear charge

• Energy level and distance from

nucleus

Result

• What happens to which electron