Why do the atoms of elements get smaller when moving from left to right within a row (period) across the periodic table?

Why do the atoms of elements get smaller when moving from left to right within a row (period) across the periodic table?

Why do Atoms get smaller?

To answer this question let’s look at the 2nd row in the periodic table

Charged particles inside the atom determine its size. A neutral lithium (Li) atom has 3 protons in its nucleus and 3 electrons in 2 energy levels in the space surrounding the nucleus.

protons

electrons

The diameter of an atom is determined by the strength with which the protons in the nucleus pull on the outermost shell of electrons. This electron “sees” 2 electrons and 3 protons when “looking” inward. An effective nuclear charge of 1+. The larger the effective nuclear charge, the stronger the pull, the shorter the distance across the atom.

Now let’s compare a Li atom to a Be atom.

LiBe

What does an electron in the outer energy level of a Be atom “see” when it “looks” inward?4 protons and 2 electrons. Electrons don’t see electrons in the same energy level.

An effective nuclear charge of 2+.

Be

This larger effective nuclear charge pulls the outermost level (valence shell) of electrons closer to the nucleus.

BeThis makes the distance from one side of the atom to the other smaller.

Now let’s compare a Be atom to a B atom.

Be

What does an electron in the outer energy level of a B atom “see” when it “looks” inward?5 protons and 2 electrons.

An effective nuclear charge of 3+.

BAn effective nuclear charge of 2+.

This larger effective nuclear charge pulls the outermost level (valence shell) of electrons closer to the nucleus.

B

BThis makes the distance from one side of the atom to the other smaller.

Now let’s recap what happens to the atomic diameter as one moves across the periodic table from left to right.A Li atom has an effective nuclear charge of1+A Be atom has an effective nuclear charge of2+A B atom has an effective nuclear charge of3+The larger the effective nuclear charge the_____________(bigger, smaller) the atom.Smaller. Let’s look at diagrams to confirm

Li3 p1+, 2e1-

ENC - 1+ Be4 p1+, 2e1-

ENC - 2+

B5 p1+, 2e1-

ENC - 3+

Smaller, Why?Larger effective nuclear charge

Now let’s model the atoms of row 3 elements

11Na row 3

3 energy levelsgroup 1

1 e1- in outer (valence) shell

11 e1- total, 11 p1+

11p1+ 2e1-8e1-1e1-

Outermost electrons “see”11p1+, 10e1-

Effective Nuclear Charge is1+

12Mg row 3



12 e1- total, 12 p1+

12p1+ 2e1-8e1-2e1-



12p1+ 2e1-8e1-2e1-

Larger effective nuclear charge makes a Mg atom smaller than a Na atom

13Al row 3



13 e1- total, 13 p1+

13p1+ 2e1-

8e1-

3e1-



13p1+ 2e1-

8e1-

3e1-

Larger effective nuclear charge makes an Al atom smaller than a Mg atom

14Si row 3



14 e1- total, 14 p1+

14p1+ 2e1-

8e1-

4e1-



13p1+ 2e1-

8e1-

4e1-

Larger effective nuclear charge makes a Si atom smaller than an Al atom

Now let’s look at a graph of Atomic Radius in 10-10 m vs Atomic Number.

Atomic Radius vs Atomic Number

Radius m x 10-10

0.5

1.0

1.5

2.0

2.5

5 10 15 20Atomic Number

Li

Be

B C N

OF Ne

Na

MgAl

Si P S

ClAr

K

Ca

Here is a graph for the 1st 40

Elements

Atomic Radius vs Atomic #

0

0.5

1

1.5

2

2.5

3

0 10 20 30 40 50

Atomic #

Ato

mic

Rad

ius

(m x

10

-10)

In general Atomic Radius tends to get smaller across the periodic table and larger down the periodic table. This can be summarized by

Smaller atomic radius


Most of the trends in the Periodic Table can be explained by comparing atomic radii.For instance the trend in Electronegativity.Electronegativity is the relative attractive force of an atom for the electrons in a bond. Generally the smaller the atom the greater the electronegativity. This is because the positively charged nucleus is closer to the shared electrons in a smaller atom.

+ +

Electronegativity is a measure of the attractive force

+ +

Greater the distance theweaker the force



Increasing ElectronegativityIncreasing Electronegativity

Ionization energy is a measure of the amount of energy required to remove an electron from an atom. In general, the smaller the atom, the harder it is to remove its electrons. As an atom gets smaller the energy required to remove its outermost electron increases.This can be illustrated by the following:

+ 2e1-8e1-1e1-

+ 2e1-1e1-

Stronger attractive force

Weaker attractive force




Increasing 1st ionization energy


Multiple Ionization Energies

12p1+ 2e1-

8e1-

2e1-energy

1st e1- 2nd e1- 3rd e1-

If multiple electrons are removed from a single atom the energy required changes in predictable ways.


energy

1st e1- 2nd e1- 3rd e1-

After the 1st electron is removed the valence shell gets a little closer so it takes more energy to remove the 2nd electron.

12p1+ 2e1-

8e1-

1e1-


energy

1st e1- 2nd e1- 3rd e1-

Since the 2nd shell is much closer than the valence shell it takes a lot more energy to remove an electron from this shell.

12p1+ 2e1-

8e1-


energy

1st e1- 2nd e1- 3rd e1-

12p1+ 2e1-

7e1-


energy

1st e1- 2nd e1- 3rd e1-

12p1+ 2e1-

7e1-

4th e1-


energy

1st e1- 2nd e1- 3rd e1-

12p1+ 2e1-

6e1-

4th e1-

When looking at a graph of multiple ionization energies the number of valence shell electrons can be deduced. If the element’s period is known then the element can sometimes be deduced.

energy

1st e1- 2nd e1- 3rd e1- 4th e1-

If this atom comes from an element in period 4 what element is it?

Ga

Group 3

energy

1st e1- 2nd e1- 3rd e1- 4th e1-

If this atom comes from an element in period 5 what element is it?

Rb Since the 4th ionization energy is not even 2x’s greater than the 3rd, and the 2nd is more than 2x’s bigger than the 1st this atom must only have 1 valence shell electron.

Electron Affinity is the amount of energy released when an atom gains an electron. In general the smaller the atom the greater the quantity of energy released when it gains an electron.This can be summarized as a trend using the Periodic table.






Increasing electron affinity


Metallic properties are due to the ease with which metal atoms lose their electrons. In general the more easily metal atoms lose electrons the more metallic its properties.The smaller a metal’s atoms are, the more tightly held are its electons, the less metallic its properties.








Decreasing metallic properties

Decreasing m

eta llic properties

When metals react they usually do so by losing electrons. The more easily they lose electrons the more reactive they are. The smaller a metal’s atoms the less reactive they are. This results in the following trend for metallic reactivity:








Decreasing metallic properties

Decreasing m

eta llic properties

Decreasing metallic reactivity

Decreasing m

etallic reactivity

Non-metals which don’t have completely filled outer shells can react by gaining electrons. In general the smaller these non-metallic element’s atoms are the stronger their attraction for electrons, the more reactive they are. In general reactive non-metals get more reactive from bottom to top, and more reactive from left to right in the peridic table.

Increasing non-metallic reactivity

Increasing non-metallic reactivity

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Why do the atoms of elements get smaller when moving from left to right within a row (period) across the periodic table?