Chapter 7 Periodic Relationships Among Elements
Development of the Periodic Table
 Mendeleev (1869) grouped elements with similar properties together by mass. There were
some problems however => the potassium argon - mass problem (also Cu, Ni).
 Mosely (early 20th century)  X-ray experiments led to the discovery of the atomic number
= number of protons. Basing the periodic table on the atomic number rather than the atomic
mass solved the problem of incorrect placement of elements in the initial versions of the
periodic table!
Classification of Elements
 Main group elements are the elements of groups 1A  7A (these elements are also referred to as
the representative elements). All these groups have incompletely filled S or p orbitals.
 Noble gases  completely filled subshells (ex. He  s).
 transition metals  groups 1B  8B. These elements possess incompletely filled d
subshells.
 f- block transition lanthanide and actinide series; their chemistry is mainly due to
incompletely filled f orbitals
 lets have a look at a specific group
Group 1A
Group 2A
ns1 configuration
ns2 configuration
 The outer electrons of an atom which are involved in chemical bonding are known as the
valence electrons. It is these electrons which are primarily responsible for the chemical
behaviour of the element.
 Another example  noble gases ns2np6 (ex. he) almost all are chemically inert.
Electron configurations of cations and anions
 We note that many elements that make up ionic compounds are in the form of anions and
cations. How do we obtain the electron configurations of anions and cations?
Example
 Na (element) ls2 2s2 2p6 3s1 or [Ne]3s1
Na+ number of protons = 11 but the number of e- = 10, since the reason number of protons number of electrons has to add up to the charge of the ion! Na+ ls2 2s2 2p6 or [Ne].
 Note  it’s easy to remove 1 electron from Na, but hard to remove a second electron!
 Similarly Mg
[Ne]3s2
Mg2+ => [Ne]. Note that Mg2+ and Na+ both have the
electron configuration of [Ne]. They are said to be isoelectronic, i.e., atoms and ions that
have the same ground state electron configuration.
Example
F-, Ne, Na+ are all isoelectronic (all have [Ne] for an election configuration).
 Transition metals  these are harder to generalize for the simple reason that they form more
2
than one type of cation (generally) and their cations usually don’t have the noble gas ground
state electron configuration.
Example.
 Determine the ground state electron configuration of Mn2+ from Mn metal.
 Mn metal  [Ar] 4s23d5 whereas the Mn2+ ion has the electron configuration [Ar]3d5! Why?
The energy of the electron is directly proportional to the principal quantum number, i.e., a
higher value of n means higher energy. Therefore, when forming a cation of a transition
metal, the eletrons are first removed from the ns orbital and then the required number of
electrons are removed from the (n-1) d orbital.
Periodic Variation in Chemical Properties
 Effective nuclear charge. To discuss the idea of the effective nuclear charge, we have to explore
the concept of electron screening (shielding) and the notions that electrons in penetrating orbitals
usually shield outer electrons well. Also, we have to discuss the effects of electron repulsion
within the same orbital (e.g., in a case like He 1s2).
 Define Z eff = effective nuclear charge = Zactual (at. number) -  (screening constant)
 Screening Effects (Shielding)
Penetrating vs. non-penetrating orbitals
S orbitals  the e- spends a fair bit of time close to the nucleus  a penetrating orbital.
p  orbitals  less penetrating.
d, f - orbitals  negligible penetration of e-‘s. Therefore, electrons are rarely close to the
nucleus.
 Let’s look at the Li atom, Z =3, electron configuration  ls2 2s1. Therefore, since the nucleus
has a charge of +3, we can (naively) subtract the charge of the two electrons in the 1s subshell
(i.e., ls2 orbital  2 paired e- for a total charge of -2!). Should the effective nuclear charge
felt by the Li 2s1 electron be 3-2 = +1?
 This is not the case, however! The electron is located in a 2s orbital (i.e., a penetrating orbital).
In a penetrating orbital, the electron does spend some time near the nucleus, i.e., inside the
90% probability sphere of the ls orbital (i.e. penetrates the shield set up by ls e-‘s). Hence,
the Li 2s electron would feel a nuclear charge of somewhat > +1.
 For p-orbitals, the probability of finding an electron in a p-orbitals near the nucleus is small.
Hence, electrons in a 2p orbital, for example, are well shielded by the ls2 (and even the 2s2)
electrons.
 Since the electron energy is directly proportional to the electron nuclear attraction attractive
forces, if an electron is more shielded, it is attracted less by the nucleus, and its energy is
higher!
Remember
1. Electrons in a given shell are shielded by electrons in an inner shell but not by an outer shell!
2. Inner filled shells shield electrons more effectively then electrons in the same subshell shield
one another!
3
e.g.
for the reaction He (g)  He+ (g), we need to supply 2373kJ of energy. However, for the
reaction He+ (g)  He2+ (g), we need to supply need 5248kJ of energy. Why?
 Due to electron-electron repulsions in ground state, the first electron is slightly shielded.
Hence, the electron nuclear attractive forces are somewhat lessened for the first electron.
When the first electron is removed, the 2nd electron then feels the full force of the 2+ nuclear
charge.
Trends in Chemical Properties
Atomic Radius
 Atomic size is one of the factors that helps determine atomic reactivity.
 Note: quantum mechanics predicts that the greatest volume of the atoms is occupied the
electron cloud. The cloud (2) gradually thins out as the radial distance, r, approaches infinity
 we can’t measure atomic size, but we can estimate some effective size by seeing how close
together atoms are in molecules (i.e., in bonding situations).
Example
Cl-Cl distance between the Cl nuclei in Cl2 = 0.198 nm
 atomic radius of Cl atom = 0.198/2 = 0.099 nm.
Similarly  the C-Cl bond distance = 0.190 nm
 C atomic radius = 0.190- 0.099 = 0.091nm.
(Note l nm = 1000pm).
 How does the atomic radius (calculated in this fashion) vary in the periodic table?
1. Across a row  decrease from left to right. Note that as we are adding an electron to the
same energy level (same value of n), we are adding a proton to the nucleus. This results in
an increase in the effective nuclear charge (Zeff) as we proceed from left to right in the
periodic table a nuclear, since the screening constant () remains essentially constant as we
proceed across a row. This results in an increase in the electron-nuclear attractive forces with
the overall effect being the electron cloud is drawn closer to the nucleus.
2. Down a group  the atomic size increases from top to bottom. Note we are going to a
larger shell (an energy level with a higher value of n).  the size has to increase (i.e.
remember the slide of the ls2s3s shells)
Ionic Radius
Ionic Radius  radius of a cation or an anion.
 Cl (g) +e-  Cl- (g)  size increases  we are adding an electron to the electron shell,
whereas the charge of the nucleus is staying the same. Therefore, the electron-nuclear
attractive forces are diminished and the size of the anion is greater than the size of the
neutral atom.
4
 Na (g) Cl- (g) +e-  size decreases  we are removing an electron from valance shell.,
and again the nuclear charge remains constant. Hence, the magnitude of the electronnuclear attractive forces is increased and the electron cloud is drawn in closer to the
nucleus, resulting in the size of the cation being smaller than the size of the neutral atom.
A comparison of Atomic and Ionic Radii
Atomic Species
Na
Cl
K
Br
Atomic Radius / nm
0.186
0.099
0.227
0.114
Ionic Species
Na+
ClK+
Br -
Ionic Radius / nm
0.095
0.181
0.133
0.195
What about the trends in the ionic radii?
1. Across a row  Example: Let’s examine the following isoelectronic species
F-, Na+, Mg2+, Al3+
They all have 10 electrons, but each species has a different atomic number (i.e., they all
contain different number’s of protons. Since in Al3+, the number of protons is greatest, the
magnitude of the electron-nuclear attractive forces is largest, the ionic radius would be the
smallest in this series. For F-, we have 9 protons and 10 electrons. Therefore, the magnitude
of the electron-nuclear attractive forces is the least in this series, and hence, the radius is the
largest! In general, for an isoelectronic series, the ionic radius decreases in the direction of
increasing effective nuclear charge.
Al3+, Mg2+, Na+, FZeff
Ionic radius
2. Down a group  as is the case with the atomic radius, the ionic size increases from top to
bottom. As we descend the group, the electrons have been added to shells with a higher
value of n.
Ionization Energy
 Another factor that determines the physical and chemical properties of elements is how
strongly the valance electrons are held in the space surrounding the nucleus. We need some
way of determining this.

Ionization energy  the energy required to remove an electron completely from an atom (or
ion) in the gaseous state.
e.g.,
Na(g)  Na+(g) +e - 1st ionization energy (I1) = energy = 495.9kJ/mol
Na+(g)  Na2+ (g) + eI2 = 4560 kJ/mol
5
Why the big difference? Na+ (g) has a noble gas electron configuration (i.e., a closed valance
shell), which is very energetically stable!
Note: this reaction is endothermic  energy is needed
Ca (g)  Ca+ (g) + eI1 = 589.5 kJ/ mol.
Ca+ (g) Ca2+(g) + eI2 = 1145 kJ/mol
What about the trends in the ionization energy?
1. Across a Row  I (ionization energy) increases as we proceed from left to right in the
periodic table. This is due to the increase in Zeff as we move across a row. Hence, the
electron-nuclear attractive forces increase meaning that the electron is ‘more tightly held’ by
the nucleus, making it harder to remove the electron from the valence shell.
2. Within a group  The ionization energy decreases from top to bottom. This is due to the
fact the increased Zeff is largely cancelled out by the shielding of inner shell electrons.
Hence, the electron-nuclear attractive forces decrease and the amount of energy needed to
completely remove the electron from the valence shell decreases.
What about the transition metals?
A slight increase in the first ionization energy is observed as we work across the transition
metals. The electrons are being added to inner shells (e.g., the 3d and the 4d level). Since we
most remove the ns electron,ans the elctrons are being added to the (n-1)d level, this slight
increase in the shielding almost cancels out the increased nuclear charge!
Electron Affinity

Electron Affinity - the energy charge when an electron is captured by a mutal atom in the gas
phase
e.g.,
Cl (g) + e-  Cl- (g) 1st electron - affinity (EA1) = -349 kJ/mol
The large - value indicates a greater tendency for the Cl atom to accept an electron. Note that
we are using the thermodynamic sign convention for the electron affinity!
6
How does the electron affinity vary in the periodic table?
1. Across a row  The electron affinity generally decreases from left to right. The reason for
this is that the increasing effective nuclear charge leads strong electron-nuclear attractions.
This indicates that the ability of the gas-phase atom to capture an electron increases!
2. Down a group  The electron affinity increases as we descend the group. Since the atomic
radius increases as we go down a group, the electron-nuclear attractive forces become
weaker. Hence, the ability of the gas phase atom to capture the electron to decreases.
Variation in Chemical Properties

General Trends in Chemical properties
We can use our knowledge of the trends in the electron affinities and ionization potentials to
understand the variation in chemical properties. The first thing that we should be aware of is
that the first member of a group is usually a bit different then the other members of the
group.
e.g. C is different than Si, Ge etc.
 diagonal relationships  the similarities that exist in the properties of an element and another
element located diagonally below it.
e.g. Li
B
Mg
Si
Trends in the Chemical Properties of Elements
 Recall that groups have been arranged by chemical properties. The first (top) members are
not very representative of the rest of the groups due to the rather small size of these species.
Therefore, we will look at the second element in the group to obtain representative properties
Hydrogen ls1 
can form H+ like group 1 A
can form H- like group V A
 Distinct from other members of the series in that it forms H2 (g), a diatomic gas like halogens
(group 7A). Its most important (and best-known) compound is water, H20.
Group 1A

electron configuration - ns1 (n> 2)
 “Alkali-metals”
 They all tend to lose one electron to get a noble gas configuration (most stable). They easily
7
form the M+ cations (low ionization energy).
 extremely reactive (good reducing agents, i.e., easily oxidized)
 never exist in nature in pure form
2 M + 2 H2O (l)  2 MOH (g) + H2 (g) + lots of heat
Group IIA

electron configuration - ns2 (n > 2)
 “alkaline-earth metals”
 Less reactive than the alkali metals.
 Their metallic character increases from Be to Ra.
 In general
M (alkaline earth) + 2 H2O (l)  M (OH)2 + H2 (g)
Be  does not react with water.
Mg  reacts with c steam (has to be “pushed”)
Sr, Ba  react with cold water
Group IIIA 
configuration - ns2 np1 (n > 2)

May lose 3e- to get the noble gas electronic configuration.

The M+ cation becomes more preferred as we descend the group since the I1 values for the
M+  M2+ are too high for e.g. Tl and In (inert-pair effect).

Metallic elements in group 3 tend to form molecular compounds.
Group IVA
electron configuration - ns2 np2 (n >2)

C is a nonmetal; Si and Ge are metalloids.


The metallic nature of these elements increases as we descend the group.
The upper members of this series do not form ionic compounds (dominated by covalent
compounds).
Sn, Pb lend to form ionic 2+ compounds
The stability of the divalent compounds increases down the group (i.e., compounds of Sn2+
and Pb2+ are common).



All the members of this series can form compounds in the +2 and +4 oxidation states, e.g.,
CO, CO2
Group IVA electron configuration - ns2 np3 (n > 2)
 N and P are nonmetals, whereas As, Sb are metalloids. Bi is metallic in nature.
 The metallic nature of these elements increases as we proceed down the group.
 All these elements tend to form compounds via covalent binding!
8
 N exists in nature as N2, P as P4, As, Sb  exists in 3-D networks.
Group VIA (ns2 np4, n > 2)
 O, S, and Se are nonmetals, whereas Te and Po are metalloids.
 The chemistry of these elements tends to be denominated by the formation of a large number
of molecular compounds with non-metals (esp. O). They can also accept two electrons to
form the A2- species.
 O in nature exists as O2. S and Se exist as covalent molecules (e.g., S8).
Group VIIA
-
electron configuration - ns2np5 (n > 2).
 The Halogen family. All the halogens are nonmetals, and they exist as the discrete species X2
in nature (e.g., F2, Cl2, Br2). Astatine is radioactive.
 F2 is very reactive; reacts with water to generate O2.
2 F2 (g) + 2 H2O (l)  4 HF (g) + O2 (g)
 The halogens tend to form anions of the type X - from X +e-; these anions are isoelectronic
with the noble gases.
 Interhalogen compounds exist, e.g. BrF3, ClF3.
 React with H2
H2 (g) + X2 (g)  HX (g) (X = F, Br, Cl, I).
This reaction is explosive with F2!
 HF is a weak acid, whereas HCl, HBr, and HI are strong acids in water.
Group VIIIA  noble gases (ns2np6, n > 2)
 All noble gases are monatomic species
 The noble gases are very unreactive, due to their closed valance shells.
 Compounds of Xe and Kr have been synthesized.
Comparison of Some Properties of the third - period elements
Na2O

Cl2O7
Ionic in nature
molecular compounds (discrete molecules)
Acidic vs. basic oxides
Na2O(s) + H2O (l)  2NaOH (aq) basic oxide
Cl2O7 (s) + H2O (l)  2 HClO4 (aq) acidic oxide
 Al2O3  amphoteric oxide (can act as either an acid or a base).
Al2O3(s) + 6 HCl (aq)  2 AlCl3 (aq) + 3 H2O (l) (basic)
Al2O3 (s) + 2 NaOH (aq) + 3 H2O (l)  2 NaAl(OH)4 (acidic oxide)