Covalent Bonds

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Covalent Bonding

'A covalent bond is the attraction of a shared pair of electrons for the nuclei of two atoms'

  • When the nuclei of two atoms are simultaneously in close attraction to the same pair of electrons, the two atoms will be locked together
  • A covalent bond is often shown by a line between two symbols
  • Some atoms can also share two pairs of electrons and this is shown by two lines
  • Hydrogen only has one electron so can only form one covalent bond
  • For other non-metallic elements, the atoms often form bonds so the total number of electrons in the outer shell, including shared and unshared electrons adds up to 8

Lone Pairs

  • Outer shell electrons not involved in bonding also form pairs. These are called lone pairs of electrons
  • They are very important in determining the chemical reactivity and the shapes of molecules
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Covalent Bond

Exceptions to the '8 electrons' guide

There are two important exceptions;

  • Group 3 elements have only 3 outer shell electrons so can only form 3 normal covalent bonds. They have to therefore settle for less than 8 electrons in the outer shell
  • There are many compounds where dot and cross diagrams show more than 8 electrons in the outer shell of an atom but only if the element is in period 3 or below

Bonding in compound ions

  • Hydroxide ions are compound ions; the oxygen and hydrogen are covalently bonded together and an electron is gained from elsewhere resulting in a charged ion

Dative Covalent Bond

'A dative covalent bond is one in which both of the electrons in the covalent bond are provided by the same atom'

  • In displayed formula a dative covalent bond is represented by an arrow
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Polar and Non-Polar Covalent Bonds

  • In a diatomic molecule such as hydrogen, the pair of electrons making up the covalent bond is equally shared between the two identical atoms
  • However when the bonded atoms are different, the pair of electrons are not equally shared
  • This is due to differing electronegativities of the atoms in the bond. We use symbols to indicate the consequent charges on the atoms showing that the charge is less than a full +1 or -1 charge. We call these partial charges
  • When this is the case, the covalent bond is said to be polar and the entire molecule can be described as polar, or as having a permanent dipole
  • In order to predict the polarity of covalent bonds and hence of the molecules that contain them, we need to know about the abilities of different elements to attract electrons when they are shared in a covalent bond. We measure this as electronegativity

'Electronegativity is the ability of an atom to attract the bonding electrons in a covalent bond'

  • Electronegativity increases from left to right across the periodic table
  • Electronegativity decreases on going down the group in the periodic table
  • The three most electronegative elements are fluorine, oxygen and nitrogen
  • The other halogens also have high electronegativity
  • Hydrogen has about the same electronegativity as carbon
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Shapes of Molecules

'Electron pair repulsion theory states that the electron pairs in the outer shell of the central atom in a molecule repel each other and therefore arrange each other to be as far away from each other as possible'

The theory includes;

  • The shape of a molecule or ion depends on the number of electron pairs (bonding pairs and lone pairs) in the outer shell of an atom
  • The pairs of electrons repel each other and are arranged as far away from each other as possible
  • The presence of one or more lone pairs in the outer shell modifies the shape slightly because a lone pair repels slightly more than a bonding pair
  • Two pairs of electrons in a double bond exerts about the same repulsion as one pair of electrons

To explain or predict a shape you need to first determine the number of regions of electron density in the outer shell of the middle atom

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Shapes of Molecules

Linear

  • Two regions of electron density
  • Two bonded pairs and no lone pairs
  • Molecule is linear
  • Bond angle is 180 degrees

Trigonal Planar

  • Three regions of electron density
  • Three bonded pairs and no lone pairs
  • Flat shape 
  • Bond angles of 120 degrees

Tetrahedral

  • Four regions of electron density
  • Four bonded pairs and no lone pairs
  • Three dimensional molecule
  • Bond angle of 109.5 degrees
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Shapes of Molecules

Octahedral

  • Six regions of electron density
  • Six bonded pairs and no lone pairs
  • Three dimensional molecule
  • Bond angles are 90 degrees

Pyramidal

  • Four regions of electron density
  • Three bonded pairs and one lone pair
  • Three dimensional molecule
  • Bond angles are 107 degrees

Non-Linear

  • Four regions of electron density
  • Two bonded pairs and two lone pairs
  • Three dimensional molecule
  • Bond angles are 104.5 degrees
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