Chapter 6: The Periodic Table

JW Dobereiner

Early 1800s

Classified elements into triads based on similar properties

Ca, Sr, Ba

Li, Na, K

Cl, Br, I

J.A.R. Newlands

Law of Octaves – every 8 elements begin to repeat in

patterns

Wasn’t taken seriously because he likened chemistry to music

Dmitri Mendeleev

Father of modern periodic table

Arranged the elements according to increasing atomic

mass

Left blanks for elements that hadn’t been discovered

but that he predicted existed

Made some exceptions (Te and I) because he knew

that iodine was more closely related to Br based on its

properties

Henry Moseley

Assigned atomic numbers to elements (number of

protons)

Arranged the elements in order of increasing atomic

number

No need to make exceptions like Mendeleev had to

Periodic Law

There is a periodic repetition of chemical and physical

properties of elements when they are arranged by

increasing atomic number.

Groups – Vertical families

Group 1 or 1A-Alkali Metals

Group 2 or 2A-Alkaline Earth Metals

Group 13 or 3A-Boron Family

Group 14 or 4A-Carbon Family

Group 15 or 5A-Nitrogen Family

Group 16 or 6A-Oxygen Family

Group 17 or 7A-Halogen Family

Group 18 or 8A-Noble Gas Family

Groups

Note: All of the Group A elements are called the “Representative Elements.”

Group A # = # of valence electrons

The other elements (Group B elements) are known as the transition and inner transition elements.

The Octet Rule- atoms tend to gain, lose, or share electrons to acquire a full set of eight valence electrons.

Some classifications

Metals – grey, typically lustrous or shiny, good

conductors of heat and electricity, and are generally

solids at R.T

Nonmetals – pink, have opposite properties. They can

be solid, liquid or gas at R.T

Metalloids/Semi-metals – blue, along the zigzag

METALS Non-METALS

METALLOIDS

Periodic Trends

Electron Shielding – when the presence of core

electrons diminishes the effect the nuclear charge has

on the valence electrons

Atomic Radii – 1/2 the distance between the nuclei of

two atoms of the same element when the atoms are

joined

Atomic radii: 200 pm

Bond

length

400 pm

Periodic Trends

Atomic Radii

Down a group – atomic radii tends to increase

Why?

Electrons enter higher principle energy levels

Across a period – tends to decrease

Why?

# of protons increase, electrons enter same

principal energy level, no increase in shielding

but increase in nuclear charge

Atomic Radii

Periodic Trends

Ionization energy – the energy required to remove an electron from a gaseous atom

Down a group – tends to decrease

Why?

Electrons are farther from nucleus therefore easier to pull off

Across a period – tends to increase

Why?

More difficult to remove an electron as you get closer to noble gas configuration

There are lots of electrons to remove!

Ionization Energy

First Io

niz

ation

ene

rgy

Atomic number

He

He has a greater IE than H.

same shielding

greater nuclear charge

H

First Io

niz

ation

ene

rgy

Atomic number

H

He

Li has lower IE than

H

more shielding

further away

outweighs greater

nuclear charge

Li

First Io

niz

ation

ene

rgy

Atomic number

H

He

Be has higher IE

than Li

same shielding

greater nuclear

charge

Li

Be

First Io

niz

ation

ene

rgy

Atomic number

H

He B has lower IE than

Be

same shielding

greater nuclear

charge

By removing an

electron we make s

orbital full Li

Be

B

First Io

niz

ation

ene

rgy

Atomic number

H

He

Li

Be

B

C

First Io

niz

ation

ene

rgy

Atomic number

H

He

Li

Be

B

C

N

First Io

niz

ation

ene

rgy

Atomic number

H

He

Li

Be

B

C

N

O

Breaks the pattern because

removing an electron gets to

1/2 filled p orbital

First Io

niz

ation

ene

rgy

Atomic number

H

He

Li

Be

B

C

N

O

F

First Io

niz

ation

ene

rgy

Atomic number

H

He

Li

Be

B

C

N

O

F

Ne Ne has a lower IE than He

Both are full,

Ne has more shielding

Greater distance

First Io

niz

ation

ene

rgy

Atomic number

H

He

Li

Be

B

C

N

O

F

Ne

Na has a lower

IE than Li

Both are s1

Na has more

shielding

Greater distance

Na

First Io

niz

ation

ene

rgy

Atomic number

Periodic Trends

Electronegativity – the relative ability of an atom to attract an electron when the atom is in a compound

Down a group – tends to decrease

Why?

The shielding effect is greater

Across a period – tends to increase (excluding noble gases)

Why?

Greater ‘desire’ to achieve noble gas configuration

Electronegativity

Periodic Trends

Ionic Size

metals - Cations

smaller than their respective atoms

Why?

# of electrons is less # of protons

Example: Na+ is smaller than Na

Periodic Trends

Ionic Size

Non-metals - Anions

larger than their respective atoms

Why?

# of electrons is larger than # of protons

Example: Cl- is bigger than Cl

Comparing ions to ions regarding size

Down a group – ions tend to increase in size

Across a period – tend to decrease

Size of Isoelectronic ions Iso - same

Iso electronic ions have the same # of electrons

Al+3 Mg+2 Na+1 Ne F-1 O-2 and N-3

all have 10 electrons

all have the configuration 1s22s22p6

Size of Isoelectronic ions

Positive ions have more protons so they are smaller.

Al+3

Mg+2

Na+1 Ne F-1 O-2 N-3