Bonding summary

Bonding and structure. Here's a summary of the main types of bonding that's used in matter and the consequent effects it has on the macroscopic structures that we see. So we have two sides of things. We have the microscopic structure and the macroscopic properties. There are four fundamental types of structure. Simple covalent, giant covalent, giant ionic and giant metallic. So you see, there's only one of these four structures is actually simple, and that's the simple covalent structure. The simple covalent structure, we just have covalent bonding that's shared pairs of electrons between nuclear centers. For example, hydrogen joins together to make a hydrogen molecule. This structure, the hydrogen molecule, is relatively stable, far more stable than the individual atoms. Simple covalent properties. But because the particles are not held together very tightly, I.e. the molecules are not held together tightly, they have fairly common properties between them. Covalent substances have low melting and boiling points. They often the smaller molecules will be gases at room temperature, certainly possibly liquids at room temperature. They are low melting and low boiling point. We call that volatile. They don't tend to conduct electricity because there's no way for electrons or ions. There are no ions there and there's no way for electrons to move from one molecule to the next. The molecules are separate. They're usually soluble in non polar solvents. That's because the bonding type in a non polar solvent is usually London dispersion forces. And London dispersion forces can also interact with small molecules or larger molecules and carry them off into solution. When we do get solutions of simple covalent substances, they don't conduct electricity because once again, there are no charged particles in the solution to do so. If the covalent substance does form a solid, such as iodine, there's a picture of iodine there on the screen. If they do form a solid, well, the solids are usually very weak in structure. The crystals are weak or soft crystals. They're easily shattered. So a simple covalent structure consists of molecules held together by very weak forces. These weak forces are London dispersion forces. The London dispersion force can become relatively strong if the molecule is large, such as iodine. Iodine has got a relative mass, the ideal molecule of 254 relative mass units. That makes it relatively large. And so the London dispersion forces acting between the molecules, the intermolecular forces, are strong enough to make it into a solid at room temperature. Iodine dissolves easily in non polar solvents such as hexane, but it's only slightly soluble in water. And in reality its solubility in water is actually a chemical reaction. It's not really soluble, it's a non conductor, and its solution does not conduct electricity. Now, occasionally we have covalent bonding, which is continuous. It goes from atom to atom to atom to atom and it never seems to stop. In fact, it only stops when it runs out of atoms. And the two most common examples of this are diamond and graphite. I'm sorry, the most common examples, diamond, graphite and silicon dioxide. Silicon dioxide is a compound, obviously, and diamond and graphite are both elements. But you can see in the diamond structure, each carbon atom is tetrahedrally bonded to four other carbon atoms. In the silicon dioxide structure, it's like the diamond structure, only with an oxygen atom between the two silicon atoms. So if we count up the actual numbers of particles in the structure of silicon dioxide, we see that there are two oxygen atoms for every silicon atom. And that gives the formula SiO2. The silicon atoms are tetrahedral and the oxygen atoms are based on an electronic tetrahedron. But of course, the bond angle is angular or bent about 104 degrees. Graphite is a bit special in that the atoms are not three. Sorry, the structure is not three dimensional. It's actually a two dimensional layer of graphites, each on top of another two dimensional layer of graphite. In fact, the single layer is called graphene. You may have heard of graphene. It's a relatively new substance. It's a, it's a mono layer of graphite. Each individual layer of graphite is only held by London dispersion forces. And so the layers can slide over one another. And this gives graphite its characteristic slippery feel. So a giant covalent properties. Giant covalent properties are they're usually hard, brittle crystalline solids, except for graphite, which is soft because of the slipperiness. The layers able to slide over one another, they're very high melting point and boiling point insoluble in water. And they will only conduct if there are valence electrons not used for bonding. Now, graphite is a conductor, but it only conducts in two dimensions. It conducts along each layer. In fact, graphene is a conductor along its layer. But each individual graphite layer doesn't pass electrons one to the other. The structure of silicon dioxide shown here, another structure. All of the atoms on the. Sorry, all of the electrons in the outer shell of the silicon are used in bonding. And so there's no way for electrons to pass around the structure. It's an insulator, silicon dioxide. To continue, it's the most common giant molecular compound. Once again, Although it looks complicated, you can think about it as silicon atoms arranged tetrahedrally, with oxygen atoms acting as bridges between the silicon atoms. And the Average formula is SiO2. Ionic bonding is when electrons are transferred from a metal to a nonmetal. And this creates ions which are then held together by strong non directional electrostatic forces. And that means that the positive ions can gather negative ions around them and the negative ions can gather positive ions around them. So once again, it's a giant structure and there are many strong electrostatic forces that need to be broken if we want to disrupt the structure. So this gives ionic substances their unique properties. They're hard, brittle crystals. They've got a high melting point and boiling point. They're non conductors in the solid state, but conductors when molten or in solution, these are called electrolytes. And they're often soluble in water, but insoluble in non polar solvents. So our ionic lattice, then you've got negative ions surrounded by positive ions. Positive ions surrounded by negative ions. The strength of a lattice is determined by the magnitude of the charge and the and the actual size or the atomic or the ionic radius of the ions. There is a third factor which is the packing of the ions. But this isn't covered by the IB syllabus, so I'm not going to go into it. So our ionic structure then is this giant structure of particles. Giant structure of ions. Any kind of stress brings oppositely charged ions into close proximity, causing the crystal structure to fracture rather than bending. So ionic compounds are brittle. It's the opposite of malleable. Polar water molecules can surround the ions and carry them off into solution. The ions in the solid are not free to move other than vibrate. So in the solid state, it's an insulator. But once molten or dissolved, it becomes an electrolyte. Solubility of ionic compounds. Ionic compounds dissolve in water because the water molecules can surround the ions and carry them off into the solution. This is our brittle property. A bit of stress and similar charges come close together. And so the compound, the ionic crystal, breaks along a cleavage plane. Metal structures once again, giant structures. Many, many, many metal ions, all surrounded by a sea of electrons. And it's the attraction between the negative sea of electron electrons and the positive ions that keep the whole structure together. The direction of the bonding is non directional, so metals can be bent and shaped and they still don't lose their strength. The actual strength of the metallic bonding will depend upon the number of available valence electrons and the magnitude of the charge on the metal. Ions. These numbers 1 and 2, of course, are effectively different, different ways of saying the same thing. The ionic range, the metal ions, as they get bigger, the metallic bonding gets weaker. And the arrangement of the ions in the lattice, again, this is something we're not going to go into. This is something that's not covered by the ib. You get periodic trends. Metallic bonding strength increases from Group 1 to Group 13. And that's because of the increasing number of electrons in the, in the valence shell and the increasing charge on the ion that we do get some occasionally anomalous melting points. Alright. And this is always caused by the arrangement of ions in the lattice. So metallic properties then are fairly well known. We know that metals conduct electricity and heat. Metals are usually hard and strong. They usually have a high melting point. They are malleable. That means they can be knocked into shape, bent into shape, and they're quite often shiny and sonorous. They ring like a bell because the vibrations are easily passed through the whole structure by the electrons holding on to the ions. You can see there the periodicity of the of metallic character. We've got increasing strength of lattice going up and increasing strength of lattice going to the right. So our summary. Elements and compounds may be metals or non metals. The bonding may be metallic then, or covalent. Either simple covalent in the case of groups 17 and 16, or giant covalent in the case of groups 13 and 14. Metallic structures, giant lasses. It's a giant structure. Metal ions surrounded by delocalized valence electrons. Covalent structures. The smallest particles are molecules. They're usually soft, volatile materials. But we may get giant covalent structures. And giant covalent structures have got very different properties to simple covalent structures. Ionic structures, of course, only occur in compounds. We have to have positive ions and negative ions and they're held together in a giant structure of many, many, many electrostatic attractions. It.