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=Task 1: Trends in the Periodic Table- Electronegativity = Electronegativity is a characteristic of atoms (or groups of atoms), and it describes their tendency to attract electrons. This is in contrast with electropositivity, the tendency to donate electrons. In order to represent this tendency, atoms (and groups of atoms) are assigned an electronegativity number, with the electronegativity increasing with the associated number. Unlike many other chemical characteristics, electronegativity cannot be directly measured, and therefore must be calculated from other known properties. This has given rise to multiple methods of calculation, between which the calculated value varies. Despite this variation, all methods show that certain trends exist in the periodic table.

With a few exceptions, electronegativity increases up groups and across periods. This makes fluorine the most electronegative element, while francium or caesium is the least. This varies between methods of calculation. The diagram to the left shows the increasing electronegativity of elements going diagonally up the periodic table. The noble gases (including helium) are shown with no electronegativity, which is due to their complete (8) valence electron shell. Since electronegativity is the measure of an atom's tendency to attract electrons to itself, and noble gases do not attract electrons to themselves (due to their complete valence shell), they have no electronegativity. Data is not available for the electronegativity of some elements, and as such, they are not shown on the diagram.

There are many methods of calculating electronegativity, of which the Pauling method is the most prominent. Linus Pauling first proposed the concept of electronegativity in 1932, and his method provides results on a scale from about 0.7 to 3.98. Other methods commonly provide their answers on a scale with a similar range, although not always. This is known as electronegativity in Pauling units.

Atoms are capable of attracting negatively charged electrons because of their positively charged nucleus. This means that, moving across a period, the increasingly positive nucleus causes electronegativity to increase. Moving down groups however, electronegativity decreases. This is because of the increased number of electron shells. The valence shell (where electrons will move to) is more shielded from the effects of the positive nucleus by the electron shells between them. The combination of these two properties is what results in the diagonally increasing electronegativity of elements.

It is this property that causes ions to form. For instance, a sodium atom will lose its single valence electron to the more electronegative chlorine atom, forming the ionic bond between a sodium ion, and a chloride ion. This is because chlorine's increased electronegativity 'pulls' the electron away from the sodium atom. Non-polar covalent bonds form when two atoms of similar (or equal) electronegativity bond. Neither atom attracts the electrons to itself, so they are shared between both atoms. Polar covalent bonds result when two atoms of different electronegativity bond. For example, when hydrogen and oxygen bond, the more electronegative oxygen cannot completely remove hydrogen's single valence electron. The resulting shared electron is attracted more to the oxygen atom, creating a polar bond- a bond more negatively charged at the oxygen end than the hydrogen end.

For a full list of the electronegativity values of elements, including the values calculated using different methods, click here. A useful YouTube video on the effects of electronegativity in bonding can be found here.

References: http://www.docstoc.com/docs/47375143/The-Periodic-Table-and-Electronegativity http://en.wikipedia.org/wiki/Electronegativity http://en.wikipedia.org/wiki/Electronegativities_of_the_elements_(data_page) http://www.youtube.com/watch?v=Kj3o0XvhVqQ http://www.chemhume.co.uk/ASCHEM/Unit%201/Ch3IMF/Chemical%20Struct.htm

=Task 3: Trends in the Periodic Table- Test =

1.
 * **Element** ||  **Proton Number (Z)**  ||  **First Ionisation Energy (kJmol-1)**  ||
 * **Li** ||  3  ||  520  ||
 * **Na** ||  11  ||  496  ||
 * **K** ||  19  ||  419  ||
 * **Rb** ||  37  ||  403  ||
 * **Cs** ||  55  ||  376  ||

2.

3. The graph shows that the first ionisation energy decreases as you move down a group. For example, Lithium in period 2 has a much higher ionisation energy (520 kJmol-1) than Caesium in period 6 (376 kJmol-1). This is because of the increased number of electron shells present in the atom. These electrons more effectively shield the valence electrons from the positive attractive forces of the nucleus. They also serve to position the valence shell further from the nucleus, combining to give the effect of decreased ionisation energy. This is because ionisation energy is the energy necessary to remove an electron, which decreases with the lower attractive forces.

4.
 * **Element** ||  **Proton Number (Z)**  ||  **First Ionisation Energy (kJmol-1)**  ||  **Atomic Radius (pm)**  ||
 * **Li** ||  3  ||  520  ||  145  ||
 * **Na** ||  11  ||  496  ||  180  ||
 * **K** ||  19  ||  419  ||  220  ||
 * **Rb** ||  37  ||  403  ||  235  ||
 * **Cs** ||  55  ||  376  ||  260  ||

5. Atomic radius increases down a group. This is because the presence of an extra electron shell enlarges the size of the atom.

6. The decreasing ionisation energy of elements down the group means that they will have increased reactivity. This is because they are more easily able to ‘lose’ electrons. For example, when forming ions, Lithium requires 520 kJmol-1, whereas Caesium requires only 376 kJmol-1.

Li + 520kJ à Li + + e -

Cs + 376kJ à Cs + + e -

This is obvious when these elements react with water. While Lithium will burn in water, producing a bright light, Caesium will produce a much more violent reaction.

7. Yes, moving down group 2, the same characteristics can be observed (increasing number of electron shells). This means that the same trends will be apparent in atomic radius and first ionisation energy.


 * **Element** ||  **Proton Number (Z)**  ||  **First Ionisation Energy (kJmol-1)**  ||  **Atomic Radius (pm)**  ||
 * **Be** ||  4  ||  900  ||  125  ||
 * **Mg** ||  12  ||  738  ||  160  ||
 * **Ca** ||  20  ||  590  ||  174  ||
 * **Sr** ||  38  ||  550  ||  191  ||
 * **Ba** ||  56  ||  503  ||  198  ||

8.
 * **Element** ||  **Proton Number (Z)**  ||  **First Ionisation Energy (kJmol-1)**  ||
 * **Li** ||  3  ||  520  ||
 * **Be** ||  4  ||  900  ||
 * **B** ||  5  ||  801  ||
 * **C** ||  6  ||  1806  ||
 * **N** ||  7  ||  1402  ||
 * **O** ||  8  ||  1314  ||
 * **F** ||  9  ||  1681  ||
 * **Ne** ||  10  ||  2081  ||

9.

10. As a general trend, ionisation energy increases, but this is not always the case. The graph shows that Boron (5) has a lower than expected ionisation energy, and Carbon (6) has a much higher than expected ionisation energy. The general trend (increasing) is due to the increasingly positive nucleus. This causes the valence electrons to be more strongly attracted to the nucleus, requiring more energy to be removed.

11.
 * = **Element**  ||=  **Proton Number (Z)**  ||=  **Electronegativity (Pauling)**  ||
 * = **Li**  ||=  3  ||=  0.98  ||
 * = **Be**  ||=  4  ||=  1.57  ||
 * = **B**  ||=  5  ||=  2.04  ||
 * = **C**  ||=  6  ||=  2.55  ||
 * = **N**  ||=  7  ||=  3.04  ||
 * = **O**  ||=  8  ||=  3.44  ||
 * = **F**  ||=  9  ||=  3.98  ||
 * = **Ne**  ||=  10  ||=  0  ||

12.

13. Moving left to right across the periodic table, electronegativity increases. The only exception to this trend is Neon (10). This is because electronegativity describes the tendency of an atom to attract electrons. Having a complete valence shell, Neon does not attract electrons, giving it no electronegativity. The reason that a trend of increasing electronegativity is visible, is to do with the increasingly positive nucleus. This causes the atom to attract electrons more strongly. Additionally, the decreasing atomic radius (due to increasingly positive nucleus) causes the valence shell of electrons to be closer to the attractive forces of the nucleus, increasing electronegativity.

14. Reactivity initially decreases across a period, but eventually begins to increase again (with the exception of the noble gases, which are non reactive). This is because of the increasing electronegativity (more reactive) and ionisation energy (less reactive). The combination of these two characteristics causes this trend. For example, Sodium is reactive, more so than Oxygen, but so is Fluorine.

Na > O < F

15. Neon is generally non-reactive. This is because of its complete valence electron shell. This means that it does not interact with other elements easily, as it does not attract or donate electrons. (it has high ionisation energy and no electronegativity)