Periodic Trends 2

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Atomic Properties .

Electronic Structures of Atoms and Ions

Shortcut to Questions

Q: 1 2 3 4 5 6 7 8 9 10

1

(a) Does metallic character increase or decrease across a row in the periodic table from left to right?

(b) Does metallic character increase or decrease down a group in the periodic table.

(c) Therefore the best metals tend to be found in what region of the Periodic Table?

(d) The most non-metallic elements tend to be found in what region of the Table?

2

Give three characteristic properties of strongly metallic elements.

3

Give three characteristic properties of strongly non-metallic elements.

4

List some typical characteristics of basic oxides.

5

List some typical characteristics of acidic oxides.

6

Give an example of an amphoteric oxide and give two typical reactions.

7

Define the term "first ionisation enthalpy".

8

Define the term "electron affinity".

9

  What are the general trends of the above within the periodic table?

10

  What is meant by the term "electronegativity"?

Periodic Trends 2 (Answers)  

1

The metallic character of an element refers to the extent to which that element displays the characteristics of a metal. Chemical characteristics of metals include the following:
* form cations in ionic compounds with non-metals
* have ionic halides
* have ionic hydrides containing the H- ion
* have basic oxides

The underlying reason why metals have these properties is the relative ease with which electrons can be removed from metal atoms, allowing them to form cations. This is expressed quantitatively via the ionization energy of any given element, defined as the energy required to remove an electron from a gas phase atom of that element. Contrary to the statement that is sometimes encountered, atoms are not "keen to lose electrons" - just the opposite as the ionization energy is always energy that must be supplied in order to remove electrons from the attraction of the positively charged nucleus. However, for metals this energy requirement is much less than for non-metals in the same Period of the Table. For example, the first ionization energy of the metal sodium is 502 kJ/mol while that of the non-metal chlorine from the same Period is 1257 kJ/mol. The energy needed to remove a second electron from any atom is higher than that for the first as there is a greater attraction between the remaining electrons and the excess +1 charge the nucleus bears after the loss of the first electron. Even more energy is required to remove a third electron and so on.

Metallic character can be related to Periodic Table position in the same way as the first ionization energies of the elements. Thus the following generalizations are valid:
Across a Period: Metallic character decreases from left to right across each Period of the Table as the outer electrons experience increasing effective nuclear charge and thus require more energy for their removal.


Down a Group. Metallic character increases down each Group of the Table as the outer electrons are further from the nucleus and require less energy for their removal.

(a) metallic character decreases from left to right in each Period.

(b) Metallic character increases down a Group.

(c) The most metallic elements are in the bottom left hand corner of the Table.

(d) The most non-metallic elements are in the top right hand corner of the Table.

2 Strongly metallic elements would have:
*low ionization energy for the valence electrons
*basic oxides containing the O2- ion.
* ionic halides
* strongly basic hydrides because they contain the H- ion

3

Strongly non-metallic elements would have:
*very high ionization energies, even for the first ionization energy.
* acidic oxides
* covalent halides
* covalent hydrides which may be acidic, neutral or weakly basic but not strongly basic as they lack the H- ion.

4

Basic oxides are those that react with H+ ions to form a salt and water. They can do this because basic oxides contain the O2- ion which reacts with the H+. Such compounds containing the O2- anion must also contain a cation. It is a defining property of metals that they form cations in compounds because their outer electrons are relatively easy to remove in reactions. Thus basic oxides are usually the oxides of metals.

e.g. sodium is a Group 1 element, all of which are metals. Hence Na2O is a basic oxide and it reacts with H+ as follows:

Na2O + 2H+ ®2Na+ + H2O

The salt would be sodium chloride if the acid used were hydrochloric acid.

Another property of basic oxides is that if they are soluble in water, they form a basic solution. Most oxides are not soluble, but Group 1 oxides dissolve.

e.g. Na2O + H2O ® 2Na+ + 2OH-

5

Acidic oxides are those that react with OH- ions to form a salt and water. They can do this because the central atom is a non-metal and it can form covalent bonds to the O atom of the OH- ion. It is a defining property of non-metals that they can form covalent compounds with other non-metals. Thus acidic oxides are generally the oxides of non-metals.

e.g.. CO2 + 2OH- ® CO32- + H2O

The salt would be sodium carbonate if the hydroxide used were sodium hydroxide.

Another property of acidic oxides is that if they are soluble in water, they form an acidic solution.

e.g. CO2 + H2O ® H2CO3 H+ + HCO3-

6

Amphoteric oxides are able to react with both H+ and OH- and such compounds are usually found as oxides of elements located near the middle of the Periodic Table. These elements typically exhibit some properties of metals as well as some of non-metals. Thus they have sufficient ionic character for their oxides to react with H+ but also the ability to form covalent bonds to the O atoms of OH- ions.

e.g. PbO + 2H+ ® Pb2+ + H2O

The salt would be lead(II) chloride if the acid used were hydrochloric acid.

PbO + H2O + 2OH- ® [Pb(OH)4]2-

The salt formed would be sodium tetrahydroxoplumbate(II) if excess sodium hydroxide were used.

Other commonly encountered amphoteric oxides are Al2O3 and ZnO.

7

The first ionization enthalpy (energy) of an element is defined as the energy required to remove the first electron from a gas phase atom of an element, and can be represented by the equation:

X(g) ® X+(g) + e-

Ionization enthalpy is listed as kJ/mole of electrons removed.

Note that even for the most metallic elements, the sign of this enthalpy is +ve, indicating that energy must be supplied and that the often seen statement that "metals are keen to lose electrons" is incorrect.

8

The electron affinity of an atom is the energy change for the reaction

X(g) + e- ® X-(g) in which one mole of electrons is added to a mole of isolated, gas phase atoms.

Usually this reaction is exothermic, so the sign of DH would be negative. Consequently compilations of electron affinity are usually listed as -DH.

9 (i) First Ionisation Enthalpy
The following generalizations apply to trends in atom's first ionization energy with Periodic Table position.
Down a Group: As the outer electrons are progressively occupying orbits further from the nucleus and screened by increasing numbers of inner electrons, the attraction they experience to the nucleus decreases and therefore the energy needed to remove an electron (i.e. to ionize the atom) also decreases.
Across a Period: Recall that the reason that atomic radius decreases from left to right across any Period is the increasing effective nuclear charge experienced by the outer electrons as more protons are added to the nucleus while the accompanying additional electrons are placed in the same outer orbitals. The increased attraction to the nucleus that results not only decreases the atomic radius, but also leads to an increase in the amount of energy needed to remove an electron - i.e. to ionize the atom.

(ii) Electron Affinity
The following generalizations apply to trends in atom's electron affinity with Periodic Table position.
Down a Group: Electron affinity usually decreases down a Group due to the increasing distance between the nucleus and the added electron in the atom's outer shell, an effect enhanced through screening by the inner electrons. This reduced attraction between the added electron and the nucleus results in less energy being released when the electron is captured by the atom.
Across a Period: Electron affinity generally increases from left to right across any given Period of the Table. This arises for the same reasons as atomic radius decreases and first ionization energy increases - namely the increasing effective nuclear charge experienced by the outer electrons as progressively more protons are added to the nucleus.
However, note that these trends are not as consistent as other trends in properties with Periodic Table position examined. For example, the electron affinity of F = 339 kJ/mol while the value for Cl = 355 kJ/mol is larger.

10

The electronegativity of an atom is the power of that atom to attract electrons to itself. Electronegativity is not a directly measurable quantity like ionization energy or electron affinity but instead is calculated from measured bond energies. A number calculated from the bond energy is assigned as the electronegativity of the element. The larger the number, the greater is the electronegativity. The most electronegative atom is F with an electronegativity value of 3.99 while the least electronegative is Fr with a value of 0.7. It is useful to know the order of electronegativities for the three most electronegative elements is F > O > Cl.