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Chemistry: Unit F321: Atoms, Bonds & Groups Covalent Bonding o Covalent bonding occurs between non metal atoms. They overlap their outer shells to give each a noble gas structure, with a stable octet of 8 outer electrons. The resulting species is a molecule o A covalent bond is a shared pair of electrons. There is a force of attraction between the bonding pair of electrons and the nuclei of the atoms at each side. o Hydrogen atoms need only 2 electrons in their outer shell to reach the noble gas structure of Helium. The covalent bond holds the two atoms together because the pair of shared electrons is attracted to both nuclei. The electrons will be held in the exact centre of the bond because each atom is the same A single covalent – or sigma (σ) – bond is and therefore they both have the same formed when one pair of electrons is electronegativity. shared. o Some atoms make double covalent bonds: oxygen and carbon share 2 pairs of electrons in this double covalent bond in carbon dioxide. o Some atoms make triple covalent bonds: nitrogen combines with another nitrogen atom, sharing three pairs of electrons to make a triple covalent bond. o There are some exceptions to the octet rule: Boron for example has a depleted octet, having only 6 electrons in its outer shell. o Sulfur can have an expanded octet: in the compound sulphur hexafluoride, the sulphur atom has 12 electrons in its outer shell o A dative covalent bond (or coordinate bond) is a shared pair of electrons which has been provided by one of the bonding atoms only. A dative covalent bond can be formed between an ammonia molecule and an H+ ion (which is electron deficient); both of the bonding electrons are donated by the nitrogen atom. o Boron triflouride-ammonia is a second example of a dative covalent bond being formed; both of the bonding electrons are donated by the nitrogen atom. o The shape of a molecule is determined by the repulsion of electron pairs around a central atom. The electron pair repulsion theory states that; the four pairs of electrons around a central atom repel each other the pairs of electrons arrange themselves so that they are as far apart as possible the repulsion between lone pairs is greater than the repulsion between bonding pairs o There are 6 shapes for molecules with up to 6 electron pairs surrounding a central atom: Shape: Tetrahedral Bond Angle: 109.5 Fischer projection: Shape: Trigonal Planar Bond Angle: 120 Fischer projection: Example: Methane, CH4 – 4 bonding pairs of electrons Example: BCl3 – 3 bonding pairs, no lone pairs Shape: Non linear Bond Angle: 104.5 Fischer projection: Shape: Pyramidal Bond Angle: 107 Fischer projection: Example: Water, H2O – 2 bonding pairs, two lone pairs Example: Ammonia, NH3 – 3 bonding pairs, 1 lone pair Shape: Linear Bond Angle: 180 Fischer projection: Shape: Octahedral Bond Angle: 90 Fischer projection: Example: CO2 – 2 double bonding pairs, no lone pairs Example: SF6 – 6 bonding pairs o A spectrum of bonding: As the difference in electronegativity increases the bonds become more polar, then become ionic. o Electronegativity is the ability of an atom to attract the bonding electrons in a covalent bond. o Pauling’s Scale of Electronegativity gives each element a value for electronegativity. Li 1.0 Be 1.5 B 2.0 C 2.5 N 3.0 O 3.5 F 4.0 Na 0.9 Mg 1.2 Al 1.5 Si 1.8 P 2.1 S 2.5 Cl 3.0 As we move from left to right across a period, there is an increased nuclear charge – more protons in the nucleus. Electrons are added to the same shell, so there is no increase in distance between the bonding electrons and the nucleus, and no additional electron shielding. The electronegativity therefore increases. As we move down a group, there are a greater number of electron shells. This results in increased electron shielding, so there are more layers between the bonding electrons and the nucleus. There is a greater distance between the bonding electrons and the nucleus. Although there are an increased number of protons in the nucleus of the atoms as we go down the group, this is outweighed by the increased distance and electron shielding. o In a non polar covalent bond, the bonding electrons are shared equally between atoms which have the same electronegativity values. o In a polar covalent bond, a difference in electronegativity results in unequal sharing of bonding electrons and a permanent dipole. A permanent dipole is therefore a small charge difference across a bond that results from the difference in the electronegativities of the bonded atoms. The greater the difference in electronegativity, the greater the permanent dipole. o Overall Polarity – not all molecules containing polar bonds are polar overall. Sometimes, the dipoles can cancel out, resulting in a net no overall dipole, making the molecule as a whole non polar. Net dipole – polar No Net dipole – non polar Net dipole – polar o Intermolecular Forces – Acting between molecules 1. Van der Waals’ forces – non polar molecules have small, temporary, fluctuating dipoles due to random movement of electrons. This instantaneous dipole induces a dipole on another neighbouring molecule. The larger the molecule, the greater the number of electrons, the larger the temporary dipoles and so the greater the attractive force. 2. Permanent Dipole-Dipole Interaction – polar molecules have permanent dipoles. The permanent dipole on one molecule attracts the permanent dipole on another molecule to form the permanent dipoledipole interaction. Molecules with permanent dipoles have much stronger intermolecular forces than molecules with induced dipoles. 3. Hydrogen Bonding – a hydrogen bond is a strong dipole-dipole interaction between an electron deficient hydrogen atom attached to an electronegative atom and the lone pair of electrons on a highly electronegative atom (O, N or F) on a different molecule. The atoms involved in hydrogen bonds are small; they therefore have a high charge density, creating greater intermolecular forces, resulting in higher boiling points. The peculiar nature of water… 1. Ice is less dense than water Each water molecule is bonded to 4 others in an open lattice structure. This means that the volume of the ice is larger than the volume of the liquid, making it less dense. When ice melts, the hydrogen bonds break and the structure breaks down, so water molecules can then move closer together, filling all of the gaps. 2. Water has a higher boiling and freezing point than expected Water has a much higher boiling point than H2S, H2Se and H2Te. This is because of the hydrogen bonding between water molecules; the other compounds have only weak Van der Waals forces, whereas water, with hydrogen bonds, must be heated to a higher temperature before the hydrogen bonds can be overcome.