The Periodic Table• Dmitri Mendeleev

– designed periodic table in which the elements were arranged in order of increasing atomic mass

• Henry Moseley– designed periodic table in

which the elements were arranged in order of increasing atomic number

• Periodic law states that when the elements are arranged by atomic number, their physical and chemical properties vary periodically.

• We will look in more detail at three periodic properties: atomic radius, ionization energy, and electron affinity.

The properties of the elements exhibit trends. These trends can be predicted using the periodic table and can be explained and understood by analyzing the electron configurations of the elements.

There are two important trends. First, electrons are added one at a time moving from left to right across a period. As this happens, the electrons of the outermost shell experience increasingly strong nuclear attraction, so the electrons become closer to the nucleus and more tightly bound to it.

Second, moving down a column in the periodic table, the outermost electrons become less tightly bound to the nucleus. This happens because the number of filled principal energy levels (which shield the outermost electrons from attraction to the nucleus) increases downward within each group.

These 2 trends explain the periodicity observed in the elemental properties of atomic radius, ionization energy, electron affinity, and electronegativity.

Periodicity of atomic radii

• Atomic radii increase down a group

Li Cs; 2s 7s

• Atomic radii decrease going across a period – the

effective nuclear charge , Zeff increases

- a proton is added to the nucleus and

shielding remains constant

Zeff = Zactual – electron shielding

• The nuclear charge felt by an electron in an outer

shell is called the effective nuclear charge

Atomic radii of main-group and transition elements

Opposing forces: Changes in n andchanges in Zeff

Overall Trends

(A) n dominates within a group; atomicradius generally increases in a group

from top to bottom

(B) Zeff dominates within a period; atomicradius generally decreases in a period

from left to right

Periodicity of atomic radius

Large size shifts whenmoving from oneperiod to the next

Ranking Elements by Atomic Size

PLAN:

SOLUTION:

PROBLEM: Using only the periodic table, rank each set of main group elements in order of decreasing atomic size.

(a) Ca, Mg, Sr (b) K, Ga, Ca (c) Br, Rb, Kr (d) Sr, Ca, Rb

Size increases down a group; size decreases across a period.

(a) Sr > Ca > Mg These elements are in Group 2A.

(b) K > Ca > Ga These elements are in Period 4.

(c) Rb > Br > Kr Rb has a higher energy level and is far to the left. Br is to the left of Kr.

(d) Rb > Sr > Ca Ca is one energy level smaller than Rb and Sr. Rb is to the left of Sr.

IONIZATION ENERGY• Cations are formed when an atom loses one or more

electrons

Na Na+ + e-

Mg Mg2+ + 2e-

• Energy input is required for this process

• The first ionization energy, Ei1, is the minimum

amount of energy required to remove the outermost

electron from an isolated gaseous atom

H + 1312 kJ H+ + e-

• In some cases a second and even a third electron

may be removed

Ca + 590 kJ Ca+ + e- Ei1

Ca+ + 1145 kJ Ca2+ + e- Ei2

Al Al3+ Ei1, Ei2, Ei3

• For a given element Ei1 < Ei2 < Ei3

- because it is much harder to remove an

electron from a positively charged ion than from

the corresponding neutral atom

• Compare Ei2 vs. Ei3 for Mg

Mg+ Mg2+ + e- 1451 kJ

Mg2+ Mg3+ + e- 7733 kJ

• Ionization energy shows a clear periodic trend

Ei decreases as a group is descended

e.g. Ei for Li > Na > K > Rb > Cs

- the electron is lost from successively higher

energy levels which are further away from the

nucleus

• There is a gradual increase in Ei as a period is

traversed

Na < Si < Cl

• Part of the reason: atomic radii decrease making the

outermost electrons closer to the nucleus and thus

harder to remove

• The increase across the period is not smooth –

breaks occur at Be/B and N/O

Ei for B < Be

Be: 1s2 2s2 Be+: 1s2 2s1

Be: 1s2 2s2 Be+: 1s2 2s1

B: 1s2 2s2 2p1 B+: 1s2 2s2

• In Be, to form Be+ a filled shell is being broken –

this is very energy expensive

• On the other hand, B+ has a filled shell; in

addition it is easy to remove the single 2p electron

•In the next case:

N: 1s2 2s2 2p3 N+: 1s2 2s2 2p2

O: 1s2 2s2 2p4 O+: 1s2 2s2 2p3

The first three ionization energies of beryllium (in MJ/mol)

First ionization energies of the

main-group elements

Increase within aperiod and decrease

within a group

Ranking Elements by First Ionization Energy

PLAN:

SOLUTION:

PROBLEM: Using the periodic table, rank the elements in each of the following sets in order of decreasing IE1:

(a) Kr, He, Ar (b) Sb, Te, Sn (c) K, Ca, Rb (d) I, Xe, Cs

IE decreases down in a group; IE increases across a period.

(a) He > Ar > Kr

(b) Te > Sb > Sn

(c) Ca > K > Rb

(d) Xe > I > Cs

Group 8A elements- IE decreases down a group.

Period 5 elements - IE increases across a period.

Ca is to the right of K; Rb is below K.

I is to the left of Xe; Cs is further to the left and down one period.

Identifying an Element from Successive Ionization Energies

PLAN:

SOLUTION:

PROBLEM: Name the Period 3 element with the following ionization energies (in kJ/mol) and write its electron configuration:

IE1 IE2 IE3 IE4 IE5 IE6

1012 1903 2910 4956 6278 22,230

Look for a large increase in energy that indicates that all of the valence electrons have been removed.

The largest increase occurs at IE6, that is, after the 5th valence electron has been removed. The element must have five valence electrons with a valence configuration of 3s23p3, The element must be phosphorus. P (Z = 15).

The complete electronic configuration is: 1s22s22p63s23p3.

ELECTRON AFFINITY• Anions are formed by an atom accepting electron(s)

• This process is also accompanied by an energy

change

• The electron affinity, Eea, is the energy change that

occurs when an electron is added to an isolated

gaseous atom

• The energy change is usually negative – the more

negative the Eea, the greater the tendency to form

anions

Be + e- Be- Eea = 241 kJ mol-1

Cl + e- Cl- Eea = -348 kJ mol-

Cl form anions easier than Be

• The periodic trend is Eea becomes more negative

across a period – trend is not regular

• Again there are breaks at Groups 2A and 5A

• It is very difficult to add an electron to a 2A metalbecause its outer 2s orbital is filled• Values for 5A elements are less negative than expected because they apply to addition of an electron to a relatively stable half-filled

Electron affinities of the main-group elements

Negative values =energy is released whenthe ion forms

Positive values =energy is absorbedto form the anion

Lets tie it all together:

Depicting ionic radii

Periodicity of ionic radii

• For cations: ionic radii is always less than atomic

radii- so Li+ < Li 1s2 1s2 2s1

For Li+; 3 p 2 e- greater attraction here

Li; 3 p 3 e-

• For anions: anions are bigger than their parent

atoms

Cl: 1s2 2s2 2p6 3s2 3p5 Cl-: 1s2 2s2 2p6 3s2 3p6

Ionic vs atomic radius

Ionic size increasesdown a group

Trends in periodsare complex

For atoms that form morethan one cation: thegreater the ionic charge,the smaller the ionic radius

Ranking Ions by Size

PLAN:

SOLUTION:

PROBLEM: Rank each set of ions in order of decreasing size, and explain your ranking:

(a) Ca2+, Sr2+, Mg2+(b) K+, S2-, Cl- (c) Au+, Au3+

Compare positions in the periodic table, formation of positive and negative ions and changes in size due to gain or loss of electrons.

(a) Sr2+ > Ca2+ > Mg2+

(b) S2- > Cl- > K+

These are members of the same Group (2A) and therefore decrease in size going up the group.

These ions are isoelectronic; S2- has the smallest Zeff and

therefore is the largest while K+ is a cation with a large Zeff and is the smallest.

(c) Au+ > Au3+ The higher the positive charge, the smaller the ion.