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Covalent Bonds. Electronegativity Polar/non-polar covalent compounds Lewis dot structures Resonance. Covalent Bonding – What?.
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Covalent Bonds Electronegativity Polar/non-polar covalent compounds Lewis dot structures Resonance
Covalent Bonding – What? • A covalent bond is a form of chemical bonding that is characterized by the sharing of pairs of electrons between atoms, or between atoms and other covalent bonds. • Different from Ionic bonding because the atom does notgive up electrons, orgain them.
What does covalent mean? • The term covalent bond dates from 1939.The prefix co- means jointly, associated in action, etc.; thus a “co-valent bond”, essentially, means that the atoms share “valence.” In the molecule CH4, the hydrogen atoms share their electrons with those of carbon via covalent bonding.
Electronegativity • Covalency is greatest between atoms of similar electronegativities. • Covalent bonding does not necessarily require the two atoms be of the same elements, only that they be of comparable electronegativity.
Electronegativity cont. • Electronegativity, symbol χ, is a chemical property that describes the ability of an atom to attract electrons towards itself. • Electronegativity cannot be directly measured and must be calculated from other atomic or molecular properties.
Pauling Method • There are several ways to calculate electronegativity but the most common is to use the Pauling Method named after the man who first proposed electronegativity – Linus Pauling – in 1932
The Pauling Scale • Using the Pauling Method gives dimensionless quantities placed on the “Pauling scale”, which runs from 0.7 to 4.0 (hydrogen = 2.2).
Pauling’s method • Electronegativity is calculated by looking at the difference between two atoms joined in a covalent bond. (XA – XB) • Dissociation energies, Ed are expressed in Electron Volts (eV)
Example. • The difference in Pauling electronegativity between hydrogen and bromine is 0.73 (Ed: H–Br, 3.79 eV; H–H, 4.52 eV; Br–Br 2.00 eV)
The scale needs a base! • As only differences in electronegativity are defined, it is necessary to choose an arbitrary reference point in order to construct a scale. Hydrogen was chosen as the reference, as it forms covalent bonds with a large variety of elements. • Electronegativity of hydrogen is 2.20 • Every other electronegativity is based of hydrogen's. (0.7 – 4.0)
More or less than Hydrogen? • It is also necessary to decide which of the two elements is the more electronegative (equivalent to choosing one of the two possible signs for the square root). • This is done by "chemical intuition": • Example: Hydrogen bromide dissolves in water to form H+ and Br− ions, so it may be assumed that bromine is more electronegative than hydrogen.
How it affects a bond. • Same electronegativity. If the atoms are equally electronegative, both have the same tendency to attract the bonding pair of electrons, and so it will be found on average half way between the two atoms. To get a bond like this, A and B would usually have to be the same atom. You will find this sort of bond in, for example, H2 or Cl2 molecules. Note: It's important to realise that this is an average picture. The electrons are actually in a molecular orbital, and are moving around all the time within that orbital.
What if one is more e-negative. • B is slightly more electronegative. That means that the B end of the bond while sharing with A, has slightly more time with the electrons, so it becomes slightly negative. At the same time, the A end becomes slightly positive. In the diagram, " " (read as "delta") means "slightly“ - so + means "slightly positive". This is also called a Polar Bond
And if one is a lot more? • B is a lot more electronegative than A In this case, the electron pair is dragged over to B's end of the bond. To all intents and purposes, A has lost control of its electron, and B has complete control over both electrons. This is now no longer covalent but an ionic bond.
So when do the bonds change? • The implication of all this is that there is no clear-cut division between covalent and ionic bonds. In a pure covalent bond, the electrons are held on average exactly half way between the atoms. In a polar bond, the electrons have been dragged slightly towards one end. • How far does this dragging have to go before the bond counts as ionic? There is no real answer to that. You normally think of sodium chloride as being a typically ionic solid, but even here the sodium hasn't completely lost control of its electron. Because of the properties of sodium chloride, however, we tend to count it as if it were purely ionic.
Electronegativity difference predictions • Non-polar covalent bond: < 0.5 • Slightly polar bond: 0.4-0.9 • Moderately polar bond: 1-1.3 • Highly polar bond: 1.4-1.7 • Slightly ionic bond: 1.8-2.2 • Ionic Bond: 2.3+
Summary • No electronegativity difference between two atoms leads to a pure non-polar covalent bond. • A small electronegativity difference leads to a polar covalent bond. • A large electronegativity difference leads to an ionic bond.
Polar and Non-polar • Polar bonds occur when one atom attracts electrons more strongly in a covalent bond. • This causes atoms to have a slightly negative and slightly positive “end”
Polar substances are soluble in water • An example of a polar compound is also water - the electrons of water's hydrogen atoms are strongly attracted to the oxygen atom, and are actually closer to oxygen's nucleus than to the hydrogen nuclei; thus, water has a relatively strong negative charge in the middle (red shade), and a positive charge at the ends (blue shade).
Non-Polar • Non-Polar compounds are formed when the electrons are shared equally between two atoms.
Non-Polar • Compounds that are non-polar are insoluble in water. • Due to this, there are no positive or negative “ends” to the compound. • Examples of Non-polar compounds include: fats, oil and petrol.
Lewis Structures G.N. Lewis
Lewis Dot Structures • Lewis Structures were originally created by Gilbert Newton Lewis, who introduced them in his 1916 article The Atom and the Molecule. • Also called Lewis dot diagrams, and are similar to electron dot diagrams. • He developed them to assist him in teaching chemical bonding, by putting dots in the place of valence electrons. Because of this, Lewis structures only deal with the valence electrons.
Cont. • Atoms always want to have 8 electrons in the outer shell, or in the case of hydrogen and helium, 2. • The goal in all bonding is to gain a full valence shell, or to obtain a noble structure. This is called the Octet rule.
Covalent Bonds – Lewis Diagrams Here we are only dealing with covalent bonding, all of the Lewis structures we deal with will be sharing electrons. This means that the valence electrons will be shared between both participating atoms. This is represented in Lewis structures by either a line, or a circle encompassing both of the shared electrons.
Steps to creating a Lewis Structure • Draw the atoms, and then draw the appropriate number of valence electrons around the atoms. • Because we are only dealing with covalent bonding, automatically assume that it is covalent and pair up the electrons. • Apply the Octet rule. • Determine if it is polar or non polar, and state which atom is slightly negative, and which atom is slightly positive. This step is not necessary for Lewis Structures, but you should be able to do it.
Resonance • Many bonding situations can be described with more than one valid Lewis Dot Structure • (for example, ozone, O3). In an LDS diagram of O3, the center atom will have a single bond with one atom and a double bond with the other. The LDS diagram cannot tell us which atom has the double bond • These two possible structures are called resonance structures. In reality, the structure of ozone is a resonance hybrid between its two possible resonance structures. Instead of having one double bond and one single bond, there are actually two 1.5 bonds with approximately three electrons in each at all times.
Resonance • Lewis Dot Structures for molecules with resonance are shown by creating the dot structure for every possible form, placing brackets around each structure, and connecting the boxes with double-headed arrows.
