PERIODIC TABLE TRENDS

If you were to look carefully at many of the properties of the elements, you would notice something besides the similarity of the properties within the groups. You would notice that many of these properties change in a fairly regular fashion that is dependent on the position of the element in the periodic table. And of course, that is what you will do next. As you compare elements from left to right across the periodic table, you will notice a trend or regular change in a number of properties. The same thing happens if you go up and down on the periodic table and compare the properties of the elements.  

Atomic Radius

You know that each atom has a nucleus inside and electrons zooming around outside the nucleus. It should seem reasonable that the size of an atom depends on how far away its outermost (valence) electrons are from the nucleus. If they are very close to the nucleus, the atom will be very small. If they are far away, the atom will be quite a bit larger. So the atomic size is determined by how much space the electrons take up. 


1) As you move down a group, atomic radius increases.

WHY? - The number of energy levels increases as you move down a group as the number of electrons increases.  Each subsequent energy level is further from the nucleus than the last.  Therefore, the atomic radius increases as the group and energy levels increase.

2) As you move across a period, atomic radius decreases.

WHY? - As you go across a period, electrons are added to the same energy level.  At the same time, protons are being added to the nucleus.  The concentration of more protons in the nucleus creates a "higher effective nuclear charge."  In other words, there is a stronger force of attraction pulling the electrons closer to the nucleus resulting in a smaller atomic radius.

First Ionization Energy

Definition:  The energy required to remove the outermost (highest energy) electron from a neutral atom in its ground state.

1) As you move down a group, first ionization energy decreases.

WHY?
Electrons are further from the nucleus and thus easier to remove the outermost one. "SHIELDING" - Inner electrons at lower energy levels essentially block the protons' force of attraction toward the nucleus.  It therefore becomes easier to remove the outer electron

2) As you move across a period, first ionization energy increases.

WHY? - As you move across a period, the atomic radius decreases, that is, the atom is smaller.  The outer electrons are closer to the nucleus and more strongly attracted to the center.  Therefore, it becomes more difficult to remove the outermost electron.

Exceptions to First Ionization Energy Trends

1) Xs2 > Xp1  e.g. 4Be > 5B WHY? - The energy of an electron in an Xp orbital is greater than the energy of an electron in its respective Xs orbital.  Therefore, it requires less energy to remove the first electron in a p orbital than it is to remove one from a filled s orbital.  2) Xp3 > Xp4  e.g.  7N > 8O WHY? - After the separate degenerate orbitals have been filled with single electrons, the fourth electron must be paired.  The electron-electron repulsion makes it easier to remove the outermost, paired electron.

Second and Higher Ionization Energy

Definition:  Second Ionization Energy is the energy required to remove a second outermost electron from a ground state atom.
Subsequent ionization energies increase greatly once an ion has reached the state like that of a noble gas.  In other words, it becomes extremely difficult to remove an electron from an atom once it loses enough electrons to lose an entire energy level so that its valence shell is filled.

Ionization Energies (kJ/mol)

Element Na Mg Al

1st IE 495.8 737.7 577.6

2nd IE 4562.4 1450.6 1816.6

3rd IE 6912 7732.6 2744.7

4th IE 9543 10,540 11,577

Metals, Non-Metals, and Metalloids

Metals

Common characteristics:

Metallic luster (shine)

Generally solids at room temperature

Malleable

Ductile

Conduct heat and electricity

Exist as extended planes of atoms

Combine with other metals to form alloys which have metallic characteristics

Form positive ions, e.g.  Na⁺, Mg²⁺, and Al³

Non-Metals

Common characteristics:

Rarely have metallic luster (shine)

Generally gases at room temperature

Neither malleable nor ductile

Poor conductors of heat and electricity

Usually exist as molecules in thier elemental form

Combine with other nonmetals to form covalent

Generally form negative ions, e.g.  Cl^-, SO₄^2-, and N^3-

The differences in the characteristics of metals and nonmetals can be explained by the following:

Metals have relatively few electrons in their valence shells.

Metals have lower ionization energies than nonmetals.

Metals have smaller electron affinities than nonmetals.

Metals have larger atoms than nonmetals.

1) As you move across a period, metallic character decreases and nonmetallic character increases.

2) As you move down a group, metallic character increases and nonmetallic character decreases.

 Metalloids

A class of 8 elements that have properties of both metals and nonmetals.

B

Si

Ge

As

Sb

Te

Po

At

Common characteristics:

Generally look metallic but are brittle (not malleable or ductile) Neither good conductors or insulators; instead they are semiconductors.

 Electronegativity

• Electronegativity - The tendency of an atom to draw electrons in a bond toward itself.
• There are two periodic trends concerning electronegativity.
• As you move down a group, electronegativity decreases.
• As you move across a period, electronegativity increases.

H 2.20

He  *

Li 0.98

Be 1.57

B 2.04

C 2.55

N 3.04

O 3.44

F 3.98

Ne *

Na 0.93

Mg 1.31

Al 1.61

Si 1.90

P 2.19

S 2.58

Cl 3.16

Ar *

K 0.82

Ca 1.00

Sc 1.36

Ti 1.54

V 1.63

Cr 1.66

Mn 1.55

Fe 1.83

Co 1.88

Ni 1.91

Cu 1.90

Zn 1.65

Ga 1.81

Ge 2.01

As 2.18

Se 2.55

Br 2.96

Kr *

Rb 0.82

Sr 1.95

Y 1.22

Zr 1.33

Nb 1.6

Mo 2.16

Tc 1.9

Ru 2.2

Rh 2.28

Pd 2.20

Ag 1.93

Cd 1.69

In 1.78

Sn 1.96

Sb 2.05

Te 2.1

I 2.66

Xe *

Cs 0.79

Ba 0.89

La 1.1

Hf 1.3

Ta 1.3

W 2.36

Re 1.9

Os 2.2

Ir 2.20

Pt 2.28

Au 2.54

Hg 2.00

Tl 2.04

Pb 2.33

Bi 2.02

Po 2.0

At 2.2

Rn *

Fr 0.7

Ra 0.9

Ac 1.1

Unq

Unp

Unh

Uns

Uno

Une

Bonds can be classified according to the difference in electronegativities of the atoms (EN).  Bonds are:

• ionic if EN > 1.8.
• polar covalent if 1.8 EN  0.4
• nonpolar covalent if EN < 0.4

The larger the difference in the electronegativities of the atoms in a bond, the stronger the strength of the bond.  As the bond becomes stronger, melting and boiling points generally increase as well. When comparing bonds with the same atoms but with different oxidation states, we have to consider another factor besides the difference in the electronegativities of the atoms since this will be the same if the atoms are the same.  As the oxidation number on an atom increase, its ability to draw electrons in a bond toward itself increases as well.  For example, the electronegativity of an Mn atom in Mn₂O₇ (oxidation number of +7) is much greater than the electronegativity of an Mn atom in MnO (oxidation number of +2). 

Melting and Boiling Points

Melting Point is when the element turns from solid to liquid, whereas the boiling point is when the element turns from liquid to gas. The graph pattern is fairly similar between the two, so we will discuss the melting point.

For each period the melting point rises from Group 1 to Group 14, then falls to the lowest value at Group 18. If the d-block elements are also included periodicity can be seen between rows of these elements, but as periodicity becomes less apparent with increasing atomic number this is less obvious than for the s- and p-block elements. 

In other words, the melting and boiling points increase as we move accros a period, and increase as we move down a group. 

Density

Density can be simply defined as the mass of the atom divided by its volume.

1) It increases as you move accross a period, untill it reaches non-metals and metalloids where it starts decreasing again.

2) It increases as you move down a group.

The majority of metals have higher densities than the majority of nonmetals. The high density of most metals is due to the tightly packed crystal lattice of the metallic structure. The strength of metallic bonds for different metals reaches a maximum around the center of the transition metal series, as those elements have large amounts of delocalized electrons in tight binding type metallic bonds.