The more electronegative an element the greater its ability to attract electrons and so the stronger an oxidising agent it will be.
ii) Displacement reaction:
In simple terms a displacement reaction is what it states it is. An element- (strictly speaking an ion), within a compound is literally displaced from that compound within a reaction, and is replaced (displaced) by another (more reactive) element, (this normally occurs when the displaced element of the compound is in an aqueous state). Essentially more reactive metals can compete with others , and them from solution.
Displacement reactions are characterised by one element displacing another out of its compound and physically replacing it, one substances pushes out another from a compound and takes it place (in that compound). This happens when an element is more reactive than the element it is replacing in a compound and relates to the position of elements in the ‘reactivity series’.
Lister and Renshaw, define a displacement reaction as, ‘a chemical reaction in which one atom or group of atoms replaces another in a compound.’ (2008, p238).
The typical reaction demonstrative of a displacement reaction is the example of (metal) zinc reacting with (blue) copper sulphate, this reaction sees the zinc displacing the copper ions from the aqueous compound of copper sulphate, resulting in the products of: zinc sulphate and elemental copper, see equation below.
Zn(s) + CuSO4(aq) = ZnSO4 (aq) + Cu(s)
0 +2 -2 +2 -2 0
Zinc atoms undergo oxidation: forming cations
Zn (s) = Zn2+ (aq) + 2e-
Copper atoms undergo reduction: returning to an elemental state.
Cu2+ (aq) + 2e- = Cu (s)
The ionic equation for this reaction=
Zn(s) + Cu2+ (aq) = Zn2+ (aq) + Cu (aq)
Sulphate (SO4) as the spectator ion neither undergoes oxidation nor reduction.
This typical ‘redox’ reaction sees the displacement of one element by another due to the superior reactivity of zinc over copper -in this case, as can be seen when referring to the metal reactivity series; which indeed sees zinc sitting higher up the series than copper.
If an attempt was made to reverse the reaction, where elemental copper was added to zinc sulphate, no reaction would occur. The copper wouldn’t replace the zinc in the compound, this is because as stated, zinc is higher up in the reactivity series compared to copper, meaning the copper cannot displace it from the compound.
If a reaction was attempted where a less reactive element is placed with a compound which contained an element which was more reactive than that element then no displacement-(replacement) would occur, no reaction would take place. The reason for this would be that the element in the compound would be higher up in the reactivity series than the other element attempting to react with the compound, so will keep its position, bound in the compound. In the case of the non-metal halogens, reactivity would be higher ascending further up the group. In general, a halogen will always displace the ion of a halogen below it in the group. This will be further considered in due course.
In other words, generally a more reactive element will displace a less reactive element from it compound, and when this happens a displacement reaction would have occurred.
iii) Electronegativity
Electronegativity is the ability of an atom to attract bonding electrons in a covalent bond; it is a measure of the tendency of an atom to attract a bonding pair of electrons. Electronegativity increases towards the top right of the periodic table, with the halogen of group 7, fluorine, exhibiting the most electronegative atoms, nearing 4.0 on the Pauling scale, (Noble gases due to their inert, unreactive properties generally are not assigned a value as they generally never form covalent bonds, with the exceptions of Krypton and Xenon). The lowest electronegative value are assigned to the large metals seen at the bottom left of the periodic table, these elements tend to have large atomic radii, with francium exhibiting a value of 0.7 on the Pauling scale.
The forces that hold atoms together are all to do with the attraction of positive charges to negative charges. The complete transfer of electrons from one atom to another occurs in ionic bonding, which results from a vast difference in electronegativity between elements which sees the greater electronegative element effectively capturing both the bonding electrons in the reaction, due to this the result is an ionic bond being formed, but in covalent bonding, electrons are shared by atoms, though in these shared pairs if one atom is better disposed at attracting electrons than the other atom, the sharing becomes uneven, and this atom which possesses a stronger attractive force on the bonding pair is termed to be more electronegative than the other atom in the bond. Its power to attract the bonding pair is greater due to its stronger electronegative properties. Essentially the shared electrons are drawn closer to the more electronegative atoms nucleus.
Atoms which are covalently bonded which show a discrepancy, which show a charge difference in the covalent bond are said to form polar bonds, where one side of the bond is assigned a partially positive figure and the other side a partially negative figure. It is the side of the bond which is assigned the partially negative figure which is the more electronegative of the two bonding electrons.
We can use the example of hydrogen bonding with fluorine, from consultation with data books and from using our knowledge and from expected trends in the periodic table we know fluorine to be the most electronegative atom, and when bonded with hydrogen, which has a much lower electronegative figure, it will be seen that the electron in the covalent bond, (hydrogen only having 1 electron), will be more attracted towards the fluorine than the hydrogen, this in turn then distorts the electron cloud towards the fluorine. Due to the fluorine’s greater-stronger attractive forces on the bonding pair of electrons, it is the more electronegative of the two atoms in the covalent bond and results in a clear polar bond being apparent. Polarity is all about the unequal sharing of the electrons between the covalently bonded atoms in a bond, and the greater the difference in electronegativity between atoms in a bond, the greater the polarity is.
However, when the electronegativity of covalently bonded atoms is equal, as is seen for diatomic elements such as the group 7 elements, the halogens (e.g. F2, Cl2), a different form of bond is created. When bonds form between atoms which possess an equal level of electronegativity there will be no discrepancy in the bond to one side over the other, each will exhibit the same forces of attraction on the bonding pair of electrons, which means that the bonding electrons are shared equally between atoms, which results in the bond exhibiting non-polar properties. The bonding electrons will reside ‘in between’ the atoms with a uniformly formed electron cloud being apparent as both atoms of the bond have the same tendency to attract the bonding pair of electrons. In a chlorine molecule, which possesses two atoms covalently bonded, a pure covalent bond is created, which is generally formed when two of the same atoms combine covalently. Two chlorine atoms each with a value of 3.0 electronegativity on the Pauling scale, on average, have an equal ability to attract the bonding pair of electrons.
Essentially the take home message is that: no electronegativity difference between two atoms leads to a pure non-polar covalent bond. A small electronegativity difference leads to a polar covalent bond and a large electronegativity difference between two atoms leads to an ionic bond.
Electronegativity is determined by the same general factors which determine ionisation energy. It increases as the number of protons in the atoms increases, it decreases when the number of shielding electron increases, and it decreases when the atomic radius, (the distance of outer electrons to the nucleus) increases.
Bibliography:
Angelosanto. A et al (2008) AS-Level Chemistry The Revision Guide
Coordination Group Publishing, Newcastle UK
Gent. D, & Ritchie. R (200 ) OCR Chemistry AS
Heinermann, UK
Hill. G, & Holman. J (2000) Chemistry in Context
Nelson Publishing, Surrey UK
Lister. T, & Renshaw. J (2008) AQA Chemistry AS level
Nelson Thornes Ltd, Cheltenham UK
Zumdahl. S, & Decoste. D (2010) Introduction to Chemistry a Foundation 7th International Edition
Charles Hartford Publishing, USA
Class notes 2010/2011.
Websites:
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