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Types of bonding. There are three types of bond that can occur between atoms:. an ionic bond occurs between a metal and non-metal atom (e.g. NaCl). a covalent bond occurs between two non-metal atoms (e.g. I 2 , CH 4 ). a metallic bond occurs between atoms in a metal (e.g. Cu).
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Types of bonding There are three types of bond that can occur between atoms: • an ionic bond occursbetween a metal and non-metal atom (e.g. NaCl) • a covalent bond occursbetween two non-metal atoms (e.g. I2, CH4) • a metallic bond occursbetween atoms in a metal (e.g. Cu)
Charge on the ions Group Charge Example Metals lose electrons to form positive ions while non-metals gain electrons to form negative ions. The number of electrons gained or lost by an atom is related to the group in which the element is found. 1 2 3 4 5 6 7 8/0 1+ 2+ 3+ N/A 3- 2- 1- N/A Na+ Mg2+ Al3+ N/A N3- O2- F- N/A The elements in groups 4 and 8 (also called group 0) do not gain or lose electrons to form ionic compounds.
Strength of metallic bonding: ion charge Element Charge on ion Melting point (K) The strength of metallic bonding depends on two factors: 1. the charge on the metal ions 2. the size of the metal ions. 1. The charge on the metal ions The greater the charge on the metal ions, the greater the attraction between the ions and the delocalized electrons, and the stronger the metallic bonds. A higher melting point is evidence of stronger bonds in the substance. Al Na Mg 3+ 1+ 2+ 933 371 923
Element Ionic radius (nm) Melting point (K) Strength of metallic bonding: ion size 2. The size of the metal ions The smaller the metal ion, the closer the positive nucleus is to the delocalized electrons. This means there is a greater attraction between the two, which creates a stronger metallic bond. Li Na K Rb Cs 0.076 0.102 0.138 0.152 0.167 454 371 337 312 302
What is electronegativity? In a covalent bond between two different elements, the electron density is not shared equally. This is because different elements have differing abilities to attract the bonding electron pair. This ability is called an element’s electronegativity. Electronegativity values for some common elements. Values given here are measured on the Pauling scale.
Electronegativity and atomic radius The electronegativity of an element depends on a combination of two factors: 1. Atomic radius As radius of an atom increases, the bonding pair of electrons become furtherfrom the nucleus. They are therefore less attracted to the positive charge of the nucleus, resulting in a lower electronegativity. higher electronegativity lower electronegativity
Electronegativity, protons and shielding 2. The number of unshielded protons The greater the number of protons in a nucleus, the greaterthe attraction to the electrons in the covalent bond, resulting in higher electronegativity. However,full energy levels of electrons shield the electrons in the bond from the increased attraction of the greater nuclear charge, thus reducing electronegativity. greater nuclear charge increases electronegativity… …but extra shell of electrons increases shielding.
Electronegativity trends: across a period Electronegativity increases across a period because: 1. The atomic radius decreases. 2. The charge on the nucleus increases without significant extra shielding. New electrons do not contribute much to shielding because they are added to the same principal energy level across the period.
Electronegativity trends: down a group Electronegativity decreases down a group because: 1. The atomic radius increases. 2. Although the charge on the nucleus increases, shielding also increases significantly. This is because electrons added down the group fill new principal energy levels.
Non-polar bonds If the electronegativity of both atoms in a covalent bond is identical, the electrons in the bond will be equally attracted to both of them. cloud of electron density This results in a symmetrical distribution of electron density around the two atoms. Bonding in elements (for example O2 or Cl2) is always non-polar because the electronegativity of the atoms in each molecule is the same. both atoms are equally good at attracting the electron density
Effect of electronegativity on polarization Element Pauling elecronegativities Molecule Electronegativity difference between atoms The greater the electronegativity difference between the two atoms in a bond the greater the polarization of the bond. This can be illustrated by looking at the hydrogen halides: H F Cl Br I 2.2 4.0 3.2 3.0 2.7 H–F H–Cl H–Br H–I 1.8 1.0 0.8 0.5 decreasing polarization
Ionic or covalent? Rather than saying that ionic and covalent are two distinct types of bonding, it is more accurate to say that they are at the two extremes of a scale. Less polar bonds have more covalent character. More polar bonds have more ionic character. The more electronegative atom attracts the electrons in the bond enough to ionize the other atom. increasing polarization
Polar molecules Molecules containing polar bonds are not always polar. Non-polar molecules Polar molecules If the polar bonds are arranged symmetrically, the partial charges cancel out and the molecule is non-polar. If the polar bonds are arranged asymmetrically, the partial charges do not cancel out and the molecule is polar.
Types of intermolecular force The molecules in simple covalent substances are not entirely isolated from one another. There are forces of attraction between them. These are called intermolecular forces. There are three main types of intermolecular force: • van der Waals forces – for example, found between I2 molecules in iodine crystals. • permanent dipole–dipole forces – for example, found between HCl molecules in hydrogen chloride. • hydrogen bonds – for example, found between H2O molecules in water.
Strength of van der Waals forces 200 The strength of van der Waals forces increases as molecular size increases. 150 100 50 0 boiling point (°C) -50 This is illustrated by the boiling points of group 7 elements. -100 -150 -200 F2 Cl2 Br2 I2 element Atomic radius increases down the group, so the outer electrons become further from the nucleus. They are attracted less strongly by the nucleus and so temporary dipoles are easier to induce.
Strength of van der Waals forces The points of contact between molecules also affects the strength of van der Waals forces. butane (C4H10) boiling point = 272K 2-methylpropane (C4H10) boiling point = 261K Straight chain alkanes can pack closer together than branched alkanes, creating more points of contact between molecules. This results in stronger van der Waals forces.
Permanent dipole–dipole forces If molecules contain bonds with a permanent dipole, the molecules may align so there is electrostatic attraction between the opposite charges on neighbouring molecules. Permanent dipole–dipole forces (dotted lines) occur in hydrogen chloride (HCl) gas. The permanent dipole–dipole forces are approximately one hundredth the strength of a covalent bond.
What is hydrogen bonding? When hydrogen bonds to nitrogen, oxygen or fluorine, a larger dipole occurs than in other polar bonds. This is because these atoms are highly electronegative due to their high nuclear charge and small size. When these atoms bond to hydrogen, electrons are withdrawn from the H atom, making it slightly positive. The H atom is very small so the positive charge is more concentrated, making it easier to link with other molecules. Hydrogen bonds are therefore particularly strong examples of permanent dipole–dipole forces.
Hydrogen bonding In molecules with OH or NH groups, a lone pair of electrons on nitrogen or oxygen is attracted to the slight positive charge on the hydrogen on a neighbouring molecule. hydrogen bond lone pair Hydrogen bonding makes the melting and boiling points of water higher than might be expected. It also means that alcohols have much higher boiling points than alkanes of a similar size.
Boiling points of the hydrogen halides 40 The boiling point of hydrogen fluoride is much higher than that of other hydrogen halides, due to fluorine’s high electronegativity. 20 0 -20 boiling point (°C) -40 -60 -80 -100 HF HCl HBr HI The means that hydrogen bonding between molecules of hydrogen fluoride is much stronger than the permanent dipole–dipole forces between molecules of other hydrogen halides. More energy is therefore required to separate the molecules of hydrogen fluoride.
