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Atomic and Ionic Size Trends. Mr. Shields Regents Chemistry U08 L04. Size Trends. Atomic Radii follows two trends: Radii increases going down a group Radii decreases going across a period. But how do we measure Atomic Radii?. Atomic Radii.

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Atomic and ionic size trends

Atomic and Ionic Size Trends

Mr. Shields Regents Chemistry U08 L04

Size trends
Size Trends

  • Atomic Radii follows two trends:

  • Radii increases going down a group

  • Radii decreases going across a period

But how do we measure Atomic Radii?

Atomic radii
Atomic Radii

Atomic Radius is measured as ½ the distance between

Adjacent nuclei in a molecule.

Why not just measure ½ the diameter of the atom?

Hint: What’s the definition of an atomic orbital?

Atomic and ionic size trends

Atomic Radii

Decreasing radii


Atomic Radii Trends

How can we account for this trend?

Atomic radii1
Atomic Radii

The trend down a group may be easier to explain first

As we move down a group what happens to the principal

Energy level ?

As the principal energy level increases, electrons move

further and further from the nucleus.

Nucleus n=1 2 3 4

Atomic radii in groups
Atomic Radii in Groups

As we move down group 1 each successive S orbital electron

Is further from the nucleus and thus Atomic radii increases

Atomic radii in groups1
Atomic Radii in groups

Since the s orbital is further from the nucleus the radii of the

Atom increases. But …

The nuclear charge is also increasing since Atomic Number is increasing.

Increasing nuclear charge

Diminishes the rate of

Change of increasing

Atomic Radii Down a


Electrons are pulled

Toward the nucleus more


Atomic radii across periods
Atomic Radii across periods

We’ve seen how atomic radii increases going down a group

But what happens when we go across a period?

We’ll, in fact atomic radiidecreases. But why?

We can begin to understand what is happening if we look at

Both the Atomic number and what principle energy level

electrons are being added to.

Atomic radii across periods1
Atomic Radii across periods

As we move across a period Atomic numbers increase

- Pos. Nuclear charge also increases so would expect

the electrons to be pulled closer to the nucleus.

So this could explain decreasing Atomic radius

- BUT … this same thing happens as we move down

a group. And for groups Atomic Radii increases as we

add more electrons? So why does radius increase in

groups but not across a period?

Atomic radii across periods2
Atomic radii across periods

The difference is that when we go across a period electrons do

Not fill higher energy levels. They either occupy lower energy

Levels or the same energy level

Whereas when going down a group electrons occupy

successively Higher principle energy levels. For example …


1 2 3 13

n=3 2-8-1 2-8-2 2-8-3

n=4 2-8-8-1 2-8-8-2 2-8-9-22-8-18-3

n=5 2-8-18-8-1 - - -

But why doesn’t atomic radii remain about the same across

a group?

Atomic radii across periods3





Atomic radii across periods

As atomic number increases across a row additional electrons

are added to the same (or lower) energy level. The effective

Nuclear charge may also be increasing and electrons are pulledin

More strongly towards the larger more positive nucleus.

Unlike when moving down a group, there are no new principal

energy levels being added to counteract the effect of increasing

nuclear charge and increasing effective nuclear charge.



Na: Effective nuclear charge = +1 S: Eff. Nuclear Charge = +6

Atomic and ionic size trends

Nuclear charge (K) = 19

Effective Nuc. Chg. (Grp I) = 1

Nuclear charge (Br) = 35

Effective Nuc. Chg. (Grp VII)= +7

Ionic radii
Ionic Radii

We’ve now seen how atomic radii changes in Periods & Groups

But what happens when Atoms either gain or lose electrons to form ions?

How does Ionic Radii vary down Groups & across Rows?

Remember …

The representative elements of groups 1, 2, 13 and 14 give up electrons to form +1, +2, +3 and +4 ions respectively

Positive ions
Positive Ions

When atoms lose all their valence electrons they lose the

outermost quantum level (n).

Consider Aluminums electron configuration. What is it?

2-8-3 (principle energy levels 1, 2 and 3 are occupied)

What is the electron config after Al loses its 3 valence electrons?

2-8 (only principle energy levels 1 and 2 are occupied)

What is the charge on Aluminum?

Positive ions1
Positive Ions

  • The loss of the outermost valence shell has two effects:

  • The atoms radius shrinks because it loses it’s outermost

  • principle quantum level

  • AND …

  • 2) The Nucleus now has more positive charge than the total

  • negative charge from electrons. The larger effective nuclear

  • charge will now pull electrons in closer to the nucleus

Positive ions2
Positive Ions

Notice that even though the ionic electron config is the same

ionic radius gets progressively smaller moving across the period.

This happens because the positive eff. nuclear charge seen by the same number of electrons increases as we move across a the row

Atomic and ionic size trends

What’s going on


Variation in atomic and Ionic Radii

Negative ions
Negative Ions

Let’s next look at the non-metals, for example Chlorine

Non-metals form ions by gaining electrons

Cl 2-8-7  2-8-8 Cl- (negative ion)

When we add electrons the effective nuclear charge per

electron decreases AND there is increases electron repulsion

So … you would expectthe ionic radius to increase and it does

Cl atomic radius = 99 nm

Cl- ionic radius = 181 nm

Negative ions1
Negative Ions

Moving down groups the principal energy level increases

- This is true for all atoms, Anions (-) & Cations (+)

Li 2-1 Li+ 2 F 2-7 F- 2-8

Na 2-8-1 Na+ 2-8 Cl 2-8-7 Cl- 2-8-8

K 2-8-8-1 K+ 2-8-8 Br 2-8-18-7 Br- 2-8-18-8

So atom & ionic size increases going down Groups

Going across periods Ionic size first decreases then jumps up

When Oxidation states change from positive to negative.

- after the jump up the downward trend in size continues

Atomic and ionic size trends

Mg atom

The ionic compound MgO

Electron Config



O atom

If we were to look at individual atoms Mg would

Actually be larger than Oxygen!