Ch. 5 - The Periodic Table · Ch. 5 - The Periodic Table 0 50 100 150 200 250 0 5 10 15 20 Atomic...

I II III

III. Periodic

Trends

(p. 140 - 154)

Ch. 5 - The Periodic Table

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A. Periodic Law

When elements are arranged in order of

increasing atomic #, elements with similar

properties appear at regular intervals.

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B. Reaction Patterns

Valence electrons: main group (s & p)

electrons in the outermost energy level

O: 2s2 2p4 = 6 valence electrons

S: 3s2 3p4 = 6 valence electrons

Se: 4s2 3d10 4p4 = 6 valence electrons!!

Similar valence e- within

a GROUP result in

similar chemical

properties in the group.

B. Reaction Patterns

Only main group electrons are in the

valence

This means that 8 is the largest number of

valence electrons possible. (2 + 6 = 8)

This is where the term “full octet” comes

from. A full octet is the most stable

type of electron configuration.

The noble gases are the only

elements that start with full

octets.

1 2 3 4 5 6 7 8

# valence electrons increases, left to right:

B. Reaction Patterns

Valence #s go

by column,

which is why

properties go

by column.

1+ 2+

Ions: charged atoms created when

electrons are gained or lost to create a full

octet (noble) configuration.

B. Reaction Patterns

3+ 4± 3- 2-1- 0

These

“common”

ion charges

are based on

valence

numbers.

Periodic Trends

Periodic trends are due to:

The organization of

electrons

(in energy levels, sublevels and

orbitals due to the laws of

quantum mechanics)

The fact that opposites attract:

The nucleus (positive charge) attracts the

electron cloud (negative charge)

+-

Periodic Trends

Attraction between charged particles

can be mathematically calculated with

Coulomb’s Law:

F𝑒𝑙𝑒𝑐 =kq

1q2

r2

F𝑒𝑙𝑒𝑐 = attraction (force)

k = a constant

r = radius (distance from nucleus to electrons)

q1 = positive charge

(protons)

q2 = negative charge

(electrons)

We’re not going to solve

this equation, but it is the

basis of most periodic

trends, so let’s take a look:

Periodic Trends

Attraction between charged particles

can be mathematically calculated with

Coulomb’s Law:

F𝑒𝑙𝑒𝑐 =kq

1q2

r2

The equation is saying two things:

More protons & electrons = more attraction

Smaller atom = more attraction

Periodic Trends

Vertical patterns:

Going down a column,

energy levels are being added.

For a higher energy level, the

valence electrons are at greater

distance (r) from the nucleus

(Higher energy levels = larger orbitals)

The added distance decreases

attractive force.

1

2

3

4

5

6

7

Periodic Trends

Vertical patterns:

When comparing elements in the

same GROUP (column/family)

Elements at the top of a group

(lower atomic number) hold onto

their valence electrons tightly.

Elements at the bottom of a group

(greater atomic number) hold onto

their valence electrons loosely.

1

2

3

4

5

6

7

Let me explain you a thing…

This next section,

your job is to listen

and UNDERSTAND,

not to write down.

Periodic Trends

It’s all about the Valence Electrons

Only outermost (valence) electrons are

involved in reactions.

Reactions happen because an atom

doesn’t have a stable (full octet)

configuration.

But actually the inner electrons are

always stable already.

Only the valence electrons need work.

Mg: [Ne-10] 3s2

Periodic Trends

Inner electrons are already stable, they don’t do

anything, let’s ignore them.

Valence electrons’ behavior (and element

properties) are based on their attraction to

the protons in the nucleus.

Valence electrons don’t feel the attraction

of all the positive charge in the nucleus.

The valence electrons are

“shielded” from the nucleus

by the inner electrons. Mg: [Ne-10] 3s2

Periodic Trends

The ten inner

electrons cancel

the charge of ten

protons.

Only two protons

are “effective”

Zeff = +2

Mg: [Ne-10] 3s2

Result: Zeff = valence

Zeff = effective nuclear charge

(That’s important.)

Mg: [1s22s22p6] 3s2

10p+

1st energy level shield: 1s2

2p+

2nd E.L. shield: 2s22p6

Periodic Trends

Conclusion:

The shielding effect causes the values

of q1 and q2 to count from 1 to 8 (based

on valence number) for every period.

Which is why patterns restart every

period even though proton number

doesn’t “restart.”

1 2 3 4 5 6 7 8

Periodic Trends

Horizontal patterns:

Going across a period, valence

electrons and Zeff increase. (q1 and q2)

Greater charge (q1 & q2) increasesattractive force.

In a period, higher atomic number =

greater attraction between nucleus and

electron cloud.

1 2 3 4 5 6 7 8

F𝑒𝑙𝑒𝑐 =kq1q2r2

Periodic Trends

Horizontal patterns:

When comparing elements in the

same PERIOD (row)

Elements at the beginning of a

period (lower atomic number) hold

onto their valence electrons loosely.

Elements at the end of a period

(greater atomic number) hold onto

their valence electrons tightly.

1 2 3 4 5 6 7 8

Periodic Trends

Rule of thumb:

Up and to the right, electrons are held

more tightly.

Lower and to the left, electrons are held

more loosely.

C. Johannesson

Atomic Radius

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D. Atomic Radius

Li

ArNe

K

Na

Atomic Radius: SIZE

the distance from nucleus to edge of electron cloud

Or: half the distance between two adjacent nuclei

D. Atomic Radius

Increases to the LEFT in a period

Increases DOWN in a group

D. Atomic Radius

Why larger going

down?

More energy levels

Why smaller to the

right?

Increased nuclear

charge (without

additional shielding)

pulls e- in tighter

D. Atomic Radius

Ionization Energy: Energy required to remove one e- from a neutral atom.

The more stable an atom is, the more energy is required to ionize it.

i.e., the closer to a full octet, the higher the ionization energy.

D. Ionization Energy

First Ionization Energy

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E. Ionization Energy

KNaLi

Ar

NeHe

(First) Ionization Energy

Increases to the right in a period

Increases to the top in a group

E. Ionization Energy

Why is this opposite of atomic radius?

In small atoms, e- are close to the nucleus

where the attraction is stronger

Stronger attraction means more difficult to

remove electrons from.

Why small jumps within each group?

Stable e- configurations (full and half-full

sublevels) don’t want to lose electrons

E. Ionization Energy

First Ionization Energy

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E. Ionization Energy

Full

sublevel

Half-full

sublevel

Successive Ionization Energies

1st I.E. 736 kJ

2nd I.E. 1,445 kJ

Core e- 3rd I.E. 7,730 kJ

Large jump in I.E. occurs when a CORE

(non-valence) e- is removed.

Mg = [Ne-10] 3s2

E. Ionization Energy

Al 1st I.E. 577 kJ

2nd I.E. 1,815 kJ

3rd I.E. 2,740 kJ

Core e- 4th I.E. 11,600 kJ

Successive Ionization Energies

Where is the “jump” for aluminum?

Al = [Ne-10] 3s2 3p1

3 v.e. means it should jump between 3rd & 4th ionizations.

E. Ionization Energy

Ionic RadiusCations (+)

Created by losing e-

smaller than parent

atom

Anions (–)

Created by

gaining e-

larger than parent

atom

G. Ionic Radius

Which atom has the larger radius?

Be or Ba

Ca or Br

Ba

Ca

Examples

Which atom has the higher 1st I.E.?

N or Bi

Ba or Ne

N

Ne

Examples

Which particle has the larger radius?

S or S2-

Al or Al3+

S2-

Al

Examples