Intermolecular Forces PDF

Structure of Water and Other Simple Compounds: Electron Dot structures do not tell us anything about the shape of molecules, they only tell us how many covalent bonds there are. H H O H H N H...

Structure of Water and Other Simple Compounds: Electron Dot structures do not tell us anything about the shape of molecules, they only tell us how many covalent bonds there are. H H O H H N H H C H H H Water (H2O) Ammonia (NH3) Methane (CH4) The VSEPR (Valence Shell Electron Repulsion) theory tells us that because electron pairs repel, molecules adjust their shapes so that the valence electron pairs are as far apart as possible. This gives a tetrahedral type of bonding, where the angle between electron pairs is called the tetrahedral angle (approx. 109.5) Methane (CH4) Polar Covalent Bonds: When pairs of electrons are shared between the same type of atom as in H 2 , O2 or Cl2 they are shared evenly. However in many heteroatomic molecules such as HCl , H2O and NH3 the electrons are not shared evenly. δ+ H Cl δ- The δ (delta) symbol means a small amount charge is present at that end of the molecule. Covalent bonds in which the electrons are unequally shared are called polar covalent bonds. A pair of equal and opposite charges separated in space as in the H—Cl molecule is called a dipole. Polar molecules are molecules that have a net dipole.  net zero dipole net zero dipole In a tetrahedrally shaped molecule the slightly polar bonds combine to cancel each other out, this means that the molecule is non-polar Electronegativity: The electron attracting power of an atom is called its Electronegativity, this increases from left to right across the periodic table and decreases from top to bottom. In a polar molecule the atom with the highest Electronegativity will have the slightly negative charge. Electronegativity Values for Atoms of Selected Elements H 2.1 Li Be B C N O F 1.0 1.5 2.0 2.5 3.0 3.5 4.0 Na Mg Al Si P S Cl 0.9 1.2 1.5 1.8 2.1 2.5 3.0 K Ca Ga Ge As Se Br 0.8 1.0 1.6 1.8 2.0 2.4 2.8 Electronegativity Differences and Bond Types Electronegativity Difference (approx) Type of Bond Example 0.0 – 0.4 Covalent (nonpolar) H – H (0.0) 0.4 – 1.0 Covalent (slightly polar) H – Cl (0.9) 1.0 – 2.0 Covalent (very polar) H – F (1.9) ≥ 2.0 Ionic NaCl (2.1) Dipole-Dipole Forces: Polar molecules are able to line as shown below, this has an effect on some of the properties of these substances. Dispersion Forces: These are weak intermolecular forces between pairs of molecules, they arise from the attraction of the nucleus of one molecule to the electron cloud of a neighbouring molecule. This force of attraction is opposed by the repulsion between the two electron clouds, so the resulting force is very small. The strength of dispersion forces increases as the molecular mass increases. Hydrogen Bonding: A hydrogen bond is a type of intermolecular attraction which occurs when a slightly positively charged hydrogen atom (in a molecule) is attracted to an unshared electron pair in a nearby molecule. δ+ δ- δ+ δ- H Cl H Cl Hydrogen bond The strength of a hydrogen bond is about one-tenth of a normal covalent bond. Predicting Shapes of Molecules: Compounds with Double or Triple Bonds: When we are referring to double and triple bonds to determine the shape of a molecule we treat it as a single bond. Eg Carbon dioxide CO2 O C O (linear) Hydrogen Cyanide HCN H C N (linear) Methanal CH2O H C O (Trigonal Planar) H