2 There are three types of strong bonds: Ionic Covalent Metallic Some substances contain both covalent and ionic bonding or an intermediate.
4 4.1 Ionic bonding Ionic bonding is an electrostatic attraction between oppositely charged ions.
5 One ore more electrons are transferred from the outer shell of one atom to the outer shell of another atom. The charge of an ion depends on the number of electrons the atom needed to loose or gain to achieve a full outer shell. 2 Na(s) + F2 (g) 2 NaF (s)
6 The electrons are shown in pairs, because each pair of electrons occupies an orbital. The successive energy levels in the atoms and ions are shown getting closer together. The radius of a sodium atom is approximately twice that of a chlorine atom. The radius of a sodium ion is approximately half that of a sodium atom. The radius of a chlorine ion is approximately twice that of a chlorine atom.
7 Cations If an atom loses e-, it becomes a positively charged ion. Group 1: Group 2: Group 3 Transition metals can form more than one ion, for example Cu+ and Cu2+, Fe2+ and Fe3+
8 Anions If an atom gains one or more e-, it becomes a negatively charged ion. Group 15: Group 16: Group 17:
9 Polyatomic ions
10 Ionic compounds Between metals (electropositive elements) and nonmetals (elements with high electronegativity). The difference in electronegativity values needs to be greater than about 1.8.
11 Formulas of ionic compounds The overall charge of the compound must be zero. Ex. CaF2
12 Lattice When an ionic compound is formed, the ions are packed in an organized crystalline structure, a lattice. The sum of all the electrostatic attractions between the oppositely charged ions is called the lattice energy.
13 The lattice energy has a high value and this energy is released when the ionic compound is formed. e.g. the formation of NaCl from Na(s) and Cl2(g) is an exothermic reaction. The value of lattice energy depends on: The charge of the ions The size of the ions The higher the value of lattice energy, the more stable is the ionic compound.
14 Physical properties Melting: The crystal structure is broken down, but there are still some attractive forces between the particles. Boiling: The attractive forces between the particles are completely broken. The stronger the bonds, the higher the boiling point.
15 Properties of ionic compounds High melting and boiling points because of strong attractive forces between the ions in the lattice (mp of Na 801º C) Conducts electricity when molten or dissolved in water. When a salt dissolves, new bonds are formed between the water molecules and the ions. This process is called hydration and the ions are said to be hydrated.
16 4.2 Covalent bonding Covalent bonding is the electrostatic attraction between a pair of electrons and positively charged nuclei.
17 Multiple covalent bonds Single bond: One shared electron pair with one electron from each atom. Double bond: Two shared electron pairs with two electrons from each atom. Triple bond: Three shared electron pairs with three electrons from each atom.
18 The more pairs of electrons there are in a covalent bond: - the shorter the bond length - the stronger the bond
19 Polarity of molecules
20 Molecules with polar bonds can be non-polar if they are symmetrical, that is if the central atom is symmetrically surrounded by identical atoms. In carbon dioxide the dipoles are exactly opposite in direction and cancel each other. O=C=O
21 Non-polar molecules In a chlorine molecule, the difference in electronegativities of the atoms is 0. This means that the electronpair in the covalent bond is on average shared EQUALLY between the 2 chlorine atoms. The bond is called a non-polar bond, thus making the molecule a non-polar molecule.
22 Polar molecules In hydrochloric acid, the difference in electronegativities is 1.0. The more electronegative chlorine atom draws the bonding pair of electrons towards itself and becomes negatively charged. The hydrogen atom then becomes positively charged. The bond is polar and the molecule has a dipole moment.
24 4.3 Covalent structures Lewis symbols show the number of valence electrons of an element represented ass either dots or crosses.
26 Drawing Lewis structures of molecules
27 Draw the Lewis structures for: a) O2 b) N2 c) CO2 d) HCN
28 Shapes of molecules and ions The shape of a molecule or ion can be predicted by the valence shell electron pair repulsion theory (VSEPR). The theory states that electron pairs (= electron domains) repel each other, and are therefore located as far away from each other as possible. The order of repulsion strength is: lone pair-lone-pair > lone pair-bond pair > bond pair-bond pair
29 If one or more of the negative charge centres is a nonbonding pair, this will influence the final shape of the molecule. e.g NH3 and H2O
31 Resonance structures For some molecules it is possible to write more than one correct Lewis structure. These structures are called resonance structures and true structure is an intermediate form known as a resonance hybrid.
32 Ex. All of the C-C bonds in benzene have the same bond length:
33 Coordinate covalent bonds In coordinate covalent bonds (dative covalent bonds) the shared pair of electrons comes from the same atom.
34 Covalent network solids Pure carbon has several different structural forms: These forms have different physical properties and they are called allotropes. Allotropes are crystalline forms of the same element, in which the atoms are bonded differently.
36 Silicon Tetrahedral arrangement
37 Silicon dioxide, SiO2 (quartz) Strong Insoluble in water High melting point Non-conductor of electricity
38 A common impure form of silicon dioxide is sand, which is colored yellow by the presence of iron (III) oxide.
39 Metallic bonding Delocalized valence electrons move freely through the metal. The attraction between these electrons and the cations holds the piece of metal intact.
40 Electrical conductivity The delocalization electrons enables free movement in response to electric fields.
41 Thermal conductivity Tight packing of cations and delocalized electrons transmit kinetic energy rapidly.
42 Malleability Individual atoms are not held to any other specific atoms, hence atoms slip easily past one another.
44 Intramolecular forces: - holds the atoms together within a molecule - affects molecular geometry and reactivity Intermolecular forces: - between the molecules within a compound - affects melting and boiling points
46 London forces (dispersion forces) Attractive forces that exist between ALL atoms and molecules. These forces are only temporary and very weak. Compounds that only have London forces have very low boiling points (they are gaseous at room temperature)
47 Factors that affect the magnitude of the London forces 1. Number of electrons in an atom The more electrons, the stronger the London forces. The more electrons, the further they are from the nucleus = less attraction the electron cloud is more easily polarized
48 2. Size of the electron cloud - The longer the carbon chain, the larger the electron cloud the stronger the London forces and the higher the boiling point
49 3. Shapes of molecules - The more contact area for the molecules, the stronger the forces.
50 Van der Waal s forces are due to the motions of electrons, which causes temporary dipoles. These forces generally increase in strength as the number of electrons in a molecule increases or if the surface area between the molecules increases. These forces are so weak that non-polar molecules have low boiling-points (many of them are gases at room temperature).
51 Dipole- dipole bonding Between permanent dipoles The negative pole of one polar molecule is attracted to the positive pole of another polar molecule.
52 Hydrogen bonding In molecules where hydrogen is directly bonded to a small highly electronegative element such as oxygen, nitrogen or fluorine.
53 Small molecules can have surprisingly high boiling points due to hydrogen bonds.
54 The lattice structure of ice
55 14.1 Further aspects of covalent bonding and structure (HL) The octet is the most common electron arrangement because of its stability. Exceptions: a) Fewer electrons (incomplete octet) if the central atom is a small atoms, e.g. Be and B b) More than eight electrons (expanded octet) if the central atom is a 3rd row element or below, e.g. P and S
56 Species with five negative charge centres If a molecule has five charge centres and they all are bonding electrons, the shape is triangular bipyramidal.
57 If one or more of these five negative charge centres is a non-bonding pair, this will influence the final shape of the molecule. One: Tetrahedron Two: T-shaped Three: Linear ClF3 I3-
58 Species with six negative charge centres Molecules with six charged centres that are all bonding have an octahedral shape, e.g. SF6. One non-bonding pair: square pyramidal BrF 5 Two non-bonding pairs: square planar XeF4
59 Formal Charge Formal charges are assigned to atoms that have an abnormal number of bonds. Formal charge Ex. For the nitrogen in ammonium: formal charge = 5-8/2 0 = +1
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