Slide 1:BONDING S. Francis Part 2 Metallic Bonding & Physical properties


Slide 2:Metallic Bonds Metallic Bonds occur between metal atoms. The valence electrons are delocalised (not attached to a particular atom) – known as a “sea of electrons”. They are free to move throughout the metal. The rest of the atoms (actually positive ions) are packed together as tightly as possible – called a lattice because it extends in all directions.


Slide 3:Metallic Bonds A metallic bond = the attraction between the positive ions and the mobile valence electrons.


Slide 4:Strength of Metallic Bonds Metallic bonds increase in strength with an increase in the charge on the positive ion, and with an increase in the number of valence electrons in the electron sea. It also depends on ionic radius. Sodium (large ionic radius & 1 valence electron) has weak bonding. Sodium’s melting point & boiling point is low & sodium is soft. Most metals have strong metallic bonding, so mp and bp are high, and they are hard.


Slide 5:Properties of Metals Metals are malleable (can be bent & reshaped) because the layers of positive ions can slide past each other without breaking any bonds.


Slide 6:Properties of Metals Metals are ductile (can be pulled into a wire) because the layers of positive ions can slide past each other without breaking any bonds. Metals also conduct heat & electricity well. Electrons are free to move from one side of the lattice to the other, so they can carry an electric current. The mobile electrons also transfer heat easily as well. A = Random movement of electrons in delocalised sea B = Movement of electrons in an electric field (conductivity) C = Conduction of heat energy by electrons D = Malleability


Slide 7:Physical Properties related to Bonding Melting & Boiling Points Ionic Bonding In an ionic compound, all the attractive forces between the ions are strong & act in all directions, so a lot of energy is needed to break the regular structure of a solid to turn it into a liquid. Melting Points are high. The melting point of Sodium chloride is 801°C. All ionic compounds are solids at room temperature.


Slide 8:Melting & Boiling Points Covalent Bonding – Giant Macromolecules Covalent Macromolecules (like diamond or silicon dioxide) have extremely high melting & boiling points. The melting point of Diamond is over 4000°C. Diamond structure (all carbon atoms in a tetrahedral lattice)


Slide 9:Melting & Boiling Points Covalent Bonding – Simple Molecules Although covalent compounds are held together by strong bonds, the forces BETWEEN the molecules are involved in keeping the molecules in an ordered array as a solid, or close together as a liquid. Melting and boiling breaks these intermolecular forces, not the covalent bonds inside the molecule. Iodine ( I 2 ) as a solid in a regular array – held together by weak intermolecular forces.


Slide 10:Covalent Bonding – Simple Molecules Liquid nitrogen (N2) – molecules held close together by intermolecular forces. Nitrogen (N2) gas -intermolecular forces broken, no attraction between molecules.


Slide 11:Melting & Boiling Points Covalent Bonding – Simple Molecules Strength of Intermolecular Forces: Hydrogen bond (strongest) = about 1/10th as strong as covalent bonds. Van der Waals’ forces (weakest) = less than 1/100th as strong as covalent bonds. Dipole – dipole forces are weaker than hydrogen bonds but much stronger than van der Waals’ forces.


Slide 12:Melting & Boiling Points Covalent Bonding – Simple Molecules Strength of Intermolecular Forces: C H H H C O H Propane Mass = 44 Mp = -42.2°C Ethanal Mass = 44 Mp = 20.8°C Ethanol Mass = 46 Mp = 78.5°C Non-polar Polar Polar Van der Waals’ Dipole - dipole Hydrogen bond d+ d- d- d+


Slide 13:Volatility Volatility = will evaporate easily The weaker the intermolecular forces, the more volatile the liquid will be. Small molecules & those with weak van der Waals’ forces will tend to be volatile. Ionic compounds and giant covalent compounds are not volatile.


Slide 14:Conductivity Conductivity = can allow electrons or heat to move through the substance. To conduct, the substance must have electrons or ions that are free to move. Metals have delocalised electrons, which are free to move. Metals are excellent conductors of heat & electricity. If one end of the metal is made positive, and the other negative, the valence electrons in the metal will move towards the positive end (& are replaced by more electrons from the electricity supply).


Slide 15:Conductivity Ionic compounds have no free ions when solid. Solid salt does not conduct electricity. When ionic compounds are molten (liquid), the individual ions are free to move around. They will conduct electricity, but will also be chemically broken down into their original atoms (because the positive ions will move to the negative end, and the negative ions will move to the positive end).


Slide 16:Conductivity Ionic compounds don’t conduct when solid, but do when molten. Molten Sodium chloride conducts electricity


Slide 17:Graphite (another form of carbon) is an exception – it’s structure has covalent layers separated by delocalised electrons that are free to move. Graphite can conduct heat & electricity. Conductivity Covalent compounds (giant & simple) do not have free ions or electrons. They do not conduct electricity as either solids or liquids. graphite diamond Graphite structure delocalised electrons


Slide 18:Solubility “Like tends to dissolve like”. Metals do not dissolve in other substances, but they can dissolve in other metals to form alloys (metal mixtures). Giant covalent molecules do not dissolve in any solvents. (Because the forces between each particle is so strong). Solubility involves the complete mixing together of the particles in the two substances. For dissolving to occur, the forces between the two types of particle in the mixture must be as strong, or stronger, than the forces between each type of particle in the two pure substances.


Slide 19:Solubility “Like tends to dissolve like”. Ionic compounds have strong forces between the particles, so are insoluble in most solvents. Water is a very polar molecule, so it can bind to the ions in the compound, allowing it to broken up and dissolved. Note – if the forces between the ions are very strong, it won’t even dissolve in water.


Slide 20:Solid sodium chloride is added to water Sodium chloride starts to dissolve – from where it is next to the water The NaCl is fully dissolved – each ion is surrounded by many water molecules


Slide 21:Solubility “Like tends to dissolve like”. Simple covalent molecules dissolve in similar solvents. Non-polar molecules tend to dissolve in non-polar solvents, (both have weak van der Waals’ forces between particles) but are insoluble in polar solvents, (they would need to break the stronger bonds holding the solvent together to dissolve). Eg iodine (I2) dissolves completely in hexane, but does not dissolve well in water.


Slide 22:Solubility “Like tends to dissolve like”. Simple covalent molecules dissolve in similar solvents. Polar molecules tend to dissolve in polar solvents, but are insoluble in non-polar solvents, (the weak bonds holding the solvent together can’t overcome the stronger bonds of the solute). Eg water can’t dissolve in oil. methanol (CH3OH) can dissolve in water, but hexane (C6H14) won’t.


Slide 23:Solubility “Like tends to dissolve like”. Ethanol (C2H5OH) is a good solvent, because it has both a polar end and a non-polar end. Organic molecules such as the alkanols (alcohols) can form hydrogen bonds (& dissolve in water) but as the chain gets larger, it interferes with the hydrogen bonding, and they become increasingly insoluble in water. methanol CH3OH dissolves completely in water ethanol C2H5OH dissolves completely in water propanol C3H7OH dissolves completely in water butanol C4H9OH barely dissolves in water pentanol C5H11OH doesn’t dissolve in water


Slide 24:Conductivity of Solutions When ionic compounds dissolve in water, the ions are free to move, so they will be able to conduct electricity. Solutions of covalent molecules do not conduct electricity.


Slide 25:Hardness Metals are hard because metallic bonds are strong. The stronger the bonds between particles in a substance, the harder it will be. Ionic compounds are hard, but if one layer moves a little (because it is hit) ions of the same charge will be next to each other & repulsion will occur. So ionic solids are also brittle. Giant covalent compounds are very hard. Molecular covalent compounds form soft solids because the intermolecular forces are weak. (eg wax, fat). Molecules with hydrogen bonds will be harder but also brittle (eg ice, sugar).


Slide 27:The end