Periodic Trends (3)

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Q: 1 2 3 4 5 6


The electronic configuration of the atom of an element "E"is:
1s2 2s2 2p6 3s2 3p6 3d10 4s2 4p3

(a) To which Periodic Table Group does "E" belong?

(b) Write the formula of a chloride of "E".

(c) Write the formula of an oxoacid of "E".


For the elements of Periodic Group 15, write the name and symbol of:

(a) the element which has the most acidic oxide

(b) the element which has the highest density

(c) the element which has the lowest ionization enthalpy


For the elements of Periodic Group 16 (excluding element 84),

(a) which element has the lowest ionization enthalpy?

(b) 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 -

(a) the element of lowest boiling point

(b) the hydride of lowest boiling point

(c) the least acidic hydride

(d) the 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 -

(a) the element of lowest boiling point

(b) the most acidic hydride

(c) the element of highest first ionisation enthalpy

(d) the element able to exhibit the highest positive oxidation number in its compounds

(e) the fluoride with the greatest partial ionic character

(f) the element which is the strongest reductant

(g) the 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 -

(a) the physical state of the element at room temperature

(b) the appearance of a sample of the element

(c) the physical state of QCl3 at room temperature

(d) the tendency of Q to have multiple oxidation states

(e) the tendency of Q to form coloured compounds

(f) the tendency of Q to form complex ions.

Periodic Trends (3) (Answers)  


(a) Group 15.
This is shown by the outer shell arrangement of 4s2 (filled), 4p3 giving a total of 5 electrons in the 4s and 4p orbitals.

(b) 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.

(c) H3EO3, H3EO4

2 (a) Nitrogen, N.
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.

(b) Bismuth, Bi.
Density increases down any Group due to the increasing numbers of protons and neutrons in the nucleus of the constituent atoms.

(c) Bismuth, Bi.
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.

3 (a) Tellurium, Te.
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.

(b) 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.

4 (a) Fluorine, F2.
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.

(b) 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 lowest BPt.

(c) 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 density.]

(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.

5 The elements are those in the Period sodium to argon.
(a) Ar.
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.

(b) HCl.
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 H- ion.

(c) Ar.
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.

(d) Cl.
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, perchloric acid.

(e) NaF.
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.

(f) Na.
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 reductant.

(g) Cl.
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.

6 From the 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:

(a) Solid at room temperature (but note, mercury is a liquid d-block element).

(b) Metallic appearance - shiny when freshly cut.

(c) Solid chloride QCl3 at room temperature due it being an ionic compound.

(d) 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+).

(e) 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.

(f) complex formation is normal for d-block elements.