Periodic Trends (3)
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Q: 1 2 3
electronic configuration of the atom of an element "E"is:
1s2 2s2 2p6 3s2 3p6 3d10 4s2 4p3
which Periodic Table Group does "E" belong?
Write the formula of a chloride of "E".
Write the formula of an oxoacid of "E".
For the elements of Periodic
Group 15, write the name and symbol of:
element which has the most acidic oxide
the element which has the highest density
the element which has the lowest ionization enthalpy
For the elements of Periodic
Group 16 (excluding element 84),
which element has the lowest ionization enthalpy?
compare the boiling points of their hydrides and account for the
unexpectedly high boiling point of the hydride of oxygen, H2O.
For the elements of Periodic
Group 17 (the first four elements only), write the formula of each
of the following -
element of lowest boiling point
the hydride of lowest boiling point
the least acidic hydride
element of lowest electron affinity
For the set of elements
with atomic numbers 11 to 18 inclusive, write the formula or the
symbol of each of the following -
element of lowest boiling point
the most acidic hydride
the element of highest first ionisation enthalpy
the element able to exhibit the highest positive oxidation number
in its compounds
fluoride with the greatest partial ionic character
element which is the strongest reductant
element which is the strongest oxidant
An element "Q"
has the electronic configuration
1s2 2s2 2p6 3s2 3p6
3d10 4s2 4p6 4d7 5s2
From this information
make reasonable predictions regarding the following -
the physical state of the element at room temperature
the appearance of a sample of the element
the physical state of QCl3 at room temperature
the tendency of Q to have multiple oxidation states
the tendency of Q to form coloured compounds
the tendency of Q to form complex ions.
Trends (3) (Answers)
This is shown by the outer shell arrangement of 4s2 (filled),
4p3 giving a total of 5 electrons in the 4s and 4p orbitals.
ECl3 or ECl5
In ECl3, the central E atom would have one lone pair of electrons
and three bonding pairs in the valence level. In ECl5, there would
be five bonding pairs and no lone pairs in the valence level of
the E atom.
Metallic character increases down a Group so the least metallic is
at the top. Non-metals have acidic oxides, so the most non-metallic
would have the most acidic oxide, as measured by the strength of the
acid derived from it.
Density increases down any Group due to the increasing numbers of
protons and neutrons in the nucleus of the constituent atoms.
Ionization energy decreases down any Group due to the weaker attraction
between outer electrons and the nucleus resulting from the increasing
distance between them, and the increased screening of the outer
electrons by the inner electrons.
Ionization energy decreases down any Group of the Table due to the
weaker attraction between outer electrons and the nucleus resulting
from the increasing distance between them, and the increased screening
of the outer electrons by the inner electrons. Thus the lowest element
in the Group, tellurium, has the smallest first ionization enthalpy.
The boiling points of the hydrides, H2X, where X is O, S, Se and
Te would be expected to increase down the Group. This is due to
the increasing strength of van der Waals forces and in particular
the dispersion forces between the molecules arising from the increasing
numbers of electrons on the X atom. However, this trend is broken
in the case of H2O which in addition has the much stronger hydrogen
bonding forces operating between molecules and consequently considerably
more energy is needed to separate the molecules from the liquid
phase to the gas phase.
Group 17 contains the elements fluorine, chlorine, bromine and iodine.
These are molecular covalent species which means that they exist as
discrete molecules, in this case as diatomic molecules of formulas
F2, Cl2, Br2 and I2 respectively. In the solid and liquid states,
only dispersion or London forces are operating. The strength of dispersion
forces is dependent on the number of electrons in the atoms and molecules,
so the larger the atom or molecule, the stronger are the attractions
between them and thus the higher the MPt and BPt. As F2 has the least
electrons in its molecules, it will experience the weakest intermolecular
forces and have the lowest boiling point.
Hydrogen chloride, HCl.
The hydrides are HF, HCl, HBr and HI. Based on just the van der
Waals intermolecular forces which consist of predominantly dispersion
forces plus the weaker dipole/dipole and dipole/induced dipole forces,
one would expect HF to have the lowest boiling point as its dispersion
force contribution to the van der Waals forces between the molecules
is the least. However, apart from van der Waals forces, HF molecules
also experience a much stronger intermolecular force, hydrogen bonding,
between HF molecules and consequently HF has the highest BPt. This
leaves HCl which does not experience hydrogen bonding as the hydride
with the weakest intermolecular forces operating and therefore the
Hydrogen fluoride, HF.
The acidity of the hydrides can be measured in terms of their Ka
values. However, the hydrides HCl, HBr and HI are all completely
ionized in water, so all three are strong acids in water. The hydride
HF is a weak acid in water because of the presence of hydrogen bonding
between HF molecules. The energy required to break these intermolecular
hydrogen bonds when added to the energy needed to ionize the separated
HF molecules to H+ and F- ions is the reason that HF is a weak acid
while the other hydrides, having no hydrogen bonds between their
molecules, are all strong acids.
[Using other solvents, the Ka values for HCl, HBr and HI can be
measured and the order is HCl < HBr < HI. Usually hydrides
of the same general formula increase in acidity down a Group of
the Table due to the increasing ionic radius of their conjugate
base which makes the base more stable through a lower ionic charge
(d) Iodine, I.
Electron affinity generally decreases down a Group, although there
are many departures from this generalization, e.g. in this Group,
Cl actually has a higher EA value than F. However, the lowest EA
value is as expected, for I.
are those in the Period sodium to argon.
Argon is a noble gas and exists as single atoms. The other elements
in the Period exist as molecules containing at least two atoms or
as aggregates containing many atoms. Consequently argon has the least
intermolecular forces operating between its constituent particles
and therefore requires the least energy to undergo the transition
from liquid to gas phase.
The hydrides of non-metals are covalently bonded and range from
being neutral (e.g. H2O) through weakly basic (e.g. NH3), weakly
acidic (e.g. H2S) to strongly acidic (e.g. HCl). However, they are
never strongly basic like the hydrides of metals which are ionic
and contain the H- ion which reacts with water to form OH- ions.
The hydrides which are the most acidic in any Period are those of
the halogens, Group 17, in this case HCl. This occurs because the
conjugate base of the hydride is most stable when the ion is large
and has a small charge, giving it the least ionic charge density.
Having the most stable conjugate base means that HCl (a strong acid)
tends to ionize more than H2S (a weak acid) while the hydrides of
Group 15 are basic, their central atom being able to use a lone
pair of electrons to accept an H+. The metals to the left of the
Period have strongly basic hydrides e.g. NaH which contains the
Ionization enthalpy increases from left to right across any Period
due to the increasing effective nuclear charge that the outer electrons
experience as the number of nuclear protons increases. Effective
nuclear charge attains a maximum at the extreme right of the Period
with a noble gas. The next element after a noble gas has its outer
electron in an orbit that is further out and this last electron
is screened by inner electrons so it experiences a weaker attraction
to the nucleus.
The oxidation number that an element can exhibit increases regularly
across any Period. This was one of the factors that lead Mendeleef
to propose his version of the Periodic Table. Thus in this Period,
as noble gases do not form compounds, chlorine has the highest positive
oxidation number of +VII shown for example in the compound HClO4,
Partial ionic character is greatest for compounds between the most
metallic element (Group 1) and the most non-metallic (Group 17).
For this period, the most metallic element is sodium, so NaF would
have the largest partial ionic character.
Reducing power of an element is related to the ease with which electrons
can be removed from that element. This can be quantified by the
first ionization enthalpy which is least for metals and largest
for non-metals. In this Period, sodium from Group 1 would require
the least energy for removal of an electron so it is the strongest
Oxidizing power is determined by the ability of an atom to attract
electrons to itself and can be gauged by examining the electron
affinity values for the elements. Non-metals have the largest electron
affinity values and in this Period, chlorine has the greatest.
structure given, Q has a completed 5s orbital and an incomplete 4d
orbital, so it is a d-block element. Based on the general properties
of d-block elements, the following would be predicted:
at room temperature (but note, mercury is a liquid d-block element).
Metallic appearance - shiny when freshly cut.
chloride QCl3 at room temperature due it being an ionic compound.
Multiple oxidation states are common among the d-block elements
due to the availability of d orbitals to participate in forming
covalent bonds as well as the ability of some elements to form cations
of different charges (e.g. Fe2+, Fe3+).
Coloured compounds are common among d-block elements. The d-orbitals
in transition element compounds cease to be degenerate and electron
transitions between the two energy levels within the d orbitals
absorb energy in the visible region.
formation is normal for d-block elements.