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§3 Periodicity & Periodic Table
During the 19th century the chase was on to discover new elements. Ever since the pre-Greek era humans knew of and used gold, silver, copper, iron and lead. Subsequently other elements, metals in the main, were also discovered, eg tin, mercury and the non-metal sulphur. It wasn't until more modern chemical techniques became available that several gaseous non-metals were identified. The most important of these was oxygen which is the key to understanding the combustion process. The discovery of oxygen is an intriguing story in itself. Here is an extract from http://www.science.ca/askascientist/viewquestion.php?qID=127 ...
"The discovery of oxygen happened in the 1770s, not the 1850s. It is normally credited to Joseph Priestley (1733–1804), an English clergyman and self-taught chemist. In those days chemistry was more like alchemy, or as it was called apothecary. Nobody really knew what anything was made of. They thought everything was made of air, earth, water, and fire. There was no such thing as the periodic table of elements. The concept of "elements" was still evolving. Same with the idea of a gas. In 1774, during an experiment, Priestley noticed that mercuric oxide (they used to call it "mercury calx") when heated, yielded a "special air" that made a candle burn much faster, and that it would support respiration or life. He put a mouse in a vessel with his "special air" and it lived four times longer than with regular air, and was revived afterwards. Priestley published these experiments and observations in 1774.
A couple years before, in 1772 Karl W. Scheele, a Swedish apothecary discovered oxygen in a similar way, but never published his findings until 1777.
It was shortly after Priestley's discovery that he went on a trip to France and met Antoine Lavoisier, a French lawyer who was conducting scientific experiments trying to figure out what air was made of. Priestley told Lavoisier of his experiments, and this was the key Lavoisier needed to make sense of his own findings that ordinary air was made up of two major components in a ratio of about 3 to 1. Lavoiser named Priestley's special air “oxygen,” which comes from the Greek word roots that mean “acid-former.”
In summary, oxygen was probably first discovered by Scheele, but first reported by Priestley. It was not named until a bit later by the Frenchman Lavoisier. Who should get the credit?"
During the 19th century there was a real spurt in the discovery of elements. Sweden, and even the Stockholm region, played a major role in this story. At the mine area near to Ytterby, a village in Sweden on the island of Resarö, close to Vaxholm (east of Stockholm), a deposit of many unusual minerals, containing rare earth (RE) and other elements was discovered. Here four of the rare earth elements (yttrium, ytterbium, terbium and erbium), were extracted from the minerals gadolinite, (Y,RE)2FeBeSi2O10 and yttrotantalite, YTaO4. The Ytterby Mine was a source of feldspar and quartz for the porcelain trade in Great Britain and Poland. The mine
opened in the 1750s and was closed in 1933.
But as the list of elements and their properties grew certain similarities became evident. Around the middle of the century Dmitri Mendeleev dreamed of a table that could reflect these regularities.
Here, to the right, is a representation of Mendeleev's table ...
What is most remarkable is how dissimilar Mendeleev's table is to the modern periodic table.
In fact Mendeleev, in his original table, organized the information about the elements, among others atomic mass, in a different way.
What was brilliant about his breakthrough was that he could see gaps in the patterns (periodicity) which he proposed should be filled by, as yet, undiscovered elements with predictable properties.
Lo and behold, those elements were ultimately discovered with properties very close to those predicted by Mendeleev - the Scientific Method in practice.
3.1 The Periodic Table
Assessment Statement Obj Teaching notes 3.1.1 Describe the arrangement of
elements in the periodic table in order of increasing atomic number.
2 Names and symbols of the elements are given in the Chemistry data booklet. The history of the periodic table will not be assessed.
3.1.2 Distinguish between the terms group and period.
2 The numbering system for groups in the periodic table is shown in the Chemistry data booklet. Students should also be aware of the position of the transition elements in the periodic table.
3.1.3 Apply the relationship between the electron arrangement of elements and their position in the periodic table up to Z = 20
2
3.1.4 Apply the relationship between the number of electrons in the highest occupied energy level for an element and its position in the periodic table.
2
What should be evident is the unwieldy layout of Mendeleev's table. The modern form of the Periodic Table takes into account the electron structure of atoms, especially the s, p, d, f subshells.
You can view several forms of the Periodic Table ... see the PT menu bar on the left of the screen.
What you should know is the first 20 elements by name, symbol and number ... H, He followed by "LiBeBCNOFNe" then "NaMgAlSiPSClAr" followed by K, Ca.
The rough breakdown of the elements into metals, semi-metals or metalloids and non-metals.
Remember ...
the rows are called PERIODS - they indicate the number of full and partly filled shells
the columns are called GROUPS - they indicate how many electrons are in the outer shell
You should also be able identify certain elements as being in special groups known as alkali metals (Group 1 metals), alkaline earth metals (Group 2 metals), (transition metals in the d-block, see below), halogens (Group 7) and noble gases (Group 8).
Although not indicated in the course plan you should also know the s-block, p-block, d-block and f-block elements.
Chemists now identify ...1 x s-orbital containing max 2 electrons (Pauli exclusion principle),3 x p-orbitals each containing max 2 electrons, giving max 6 electron "states" then 5 x d-orbitals with max 2 electrons giving a combined total of 10 electron "states" and finally 7 x f-orbitals giving up to 14 possible electron "states".
Note: I use the word "state" entirely liberally
3.2 Physical Properties
Assessment Statement Obj Teaching notes 3.2.1 Define the terms first ionization energy
and electronegativity. 1
3.2.2 Describe and explain the trends in atomic radii, ionic radii, first ionization energies, electronegativities and melting points for the alkali metals ( Li --> Cs ) and the halogens ( F -->I ).
3 Data for all these properties is listed in the Chemistry data booklet. Explanations for the first four trends should be given in terms of the balance between the attraction of the nucleus for the electrons and the repulsion between electrons. Explanations based on effective nuclear charge are not required.
3.2.3 Describe and explain the trends in atomic radii, ionic radii, first ionization energies and electronegativities for elements across period 3.
3 Aim 7: Databases and simulations can be used here.
3.2.4 Compare the relative electronegativity values of two or more elements based on their positions in the periodic table.
3
We will mostly be concerned about the physical properties of atoms of the elements i) atomic radii ii) ionic radii iii) electronegativity and iv) 1st ionization energy ... 3.2.1 Some definitions
1. Effective nuclear charge:"Unshielded protons" mentioned below refers to the fact that full shells of electrons seem to blank out the attractive force of an equal number of protons.
The unshielded protons whose force attracts the valence electrons are equal in number to the number of outer shell or valence electrons.
2. Ionization energy:
"Ionization energy" - is the energy necessary to remove an electron from the neutral atom. The "first ionization energy" electron and so on.
It is a minimum for the alkali metals and generally increases across periods being a maximum for the noble gases with full outer shells. This is explained by the greater
attraction the positive nucleus has for the outermost electron the nearer it is to the nucleus.
Two closer at this graph shows more detail which illuminates further the electron shell (energy level) model of atoms ...
3. Electronegativity:
"Electronegativity" - is the measure of the tendency of an atom to attract (a bonding pair of) electrons. The Paulingmost electronegative and is assigned a value of 4.0. Cesium and Francium are the least electronegative with a value of 0.7.A clear trend of higher electronegativity the further right and up an element is found is evident. Of course this corresponds to the greater force of attraction for any nearby electron to form a bonding pair the greater the number of unshielded protons there are in the nucleus.
Another graph showing electronegativity trends ...
Notice the extreme values for fluorine and oxygen. They are the most electronegative elements of them all and, given fluorine's relative rarity, then oxygen is by far the element that has most attraction for electrons. We will see that this has consequences when considering oxidation numbers.
4. Atomic & Ionic radii:
The trends in atomic and ionic radii follow the same reasoning.
The greater the attracting force of unshielded electrons the less the atomic radius is. The more shells of electrons the greater the radius is.
As a metal loses electrons to form an ion the radius shrinks
As a non-metal gains electrons to form an ion the radius grows (sometimes seemingly out of proportion)
So in summary
3.3 Chemical Properties
Assessment Statement Obj Teaching notes 3.3.1 Discuss the similarities and
differences in the chemical properties of elements in the same group.
3 The following reactions should be covered. • Alkali metals (Li, Na and K) with water • Alkali metals (Li, Na and K) with halogens (Cl2, Br2 and I2) • Halogens (Cl2, Br2 and I2) with halide ions (Cl–, Br– and I–)
3.3.2 Discuss the changes in nature from ionic to covalent and from basic to acidic, of the oxides across period 3. Equations are required for the reactions of Na2O, MgO, P4O10 and SO3 with water.
3
It should not be surprising that the chemical behavior of elements in the same group is similar. The main difference being in "strength" of the reaction.
3.2.1 Reactions with water:
In Group 1, the alkali metals, the reaction get stronger as you move down the group from lithium to cesium. A good example is their reactions with water. Viewing the Brainiac videos in the "modern Periodic Table" menu shows some graphical evidence of this.Here are the reaction equations for all alkali metals A = Li, Na, K, Cs, Rb, Fr
2A(s) + H2O(l) ----> 2AOH(aq) + H2(g)
The heat given out in the reaction is enough to explosively set alight the hydrogen gas generated for K to Cs. The AOH is, of course, the hydroxide, commonly called an alkali, of the metal hence the name for this group.
Similarly the Group 2 alkaline earth metals react with water to form an alkaline solution. E = Be, Mg, Ca, Sr, Ba, Ra
E(s) + 2H2O(l) ----> E(OH)2(aq) + H2(g)
These reactions are much "gentler" than for the alkali metals. In the case of magnesium small bubbles of hydrogen slowly form on the surface of the metal. The reaction get "stronger as we move down the group.
3.2.1 Reactions with halides:
Similar reaction equations are possible for the reactions of group 1 and 2 metals with the halogens, X = F, Cl, Br, I, At
2A(s) + X2(g,l or s) ----> 2AX(s)
E(s) + X2(g,l or s) ----> EX2(s)
The products here are all ionic salts, e.g. sodium chloride, barium iodide etc.
3.2.1 Reactions among the halogens:
The difference in reactivity between the halogens, here in reverse to the groups 1 and 2 (i.e. more reactive the higher in the group one goes) can be neatly illustrated by reacting them with halide solutions.
F2(g) + 2X-(aq) ----> 2F-
(aq) + X2(g, l or s) for all other halides
Cl2(g) + 2X-(aq) ----> 2Cl-
(aq) + X2(l or s) for the other halides except fluorine
Br2(l) + 2X-(aq) ----> 2Br-
(aq) + X2(s) for iodine and astatine only
Considering trends across the periodic table we focus on period 3
3.2.2 Reactions with oxygen, nature of the oxide:
4Na(s) + O2(g) ----> 2Na2O(s) and then Na2O(s) + H2O ----> 2NaOH(aq) ; sodium hydroxide is a strong alkali
2Mg(s) + O2(g) ----> 2MgO(s) and then MgO(s) + H2O ----> 2Mg(OH)2(aq) ; magnesium hydroxide is a
4Al(s) + 3O2(g) ----> 2Al2O3(s) aluminum oxide does not react with water, rather it is amphoteric - acting as an acid with an alkali and as a base with an acid
Si(s) + O2(g) ----> SiO2(s) silicon dioxide is commonly known as quartz or sand and does not react with water but is
4P(s) + 5O2(g) ----> P4O10(g) then P4O10(g) + H2O ----> H3PO4(aq) , phosphoric acid - a strong acid
S(s) + O2(g) ----> SO2(g) , SO2(g) + O2(g) ----> SO3(g) then SO3(g)+ H2O ----> H2SO4(aq), sulfuric acid -a
2Cl2(g) + 7O2(g) ----> 2Cl2O7(g) which reacts with H2O ----> 2HClO4(aq) - called perchloric acid - also a
A nice link explaining can be found at periodic trends.
