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INTERMOLECULAR FORCES. Introduction: The physical properties of melting point, boiling point, vapor pressure, evaporation, viscosity, surface tension, and solubility are related to the strength of attractive forces between molecules.
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INTERMOLECULAR FORCES Introduction: The physical properties of melting point, boiling point, vapor pressure, evaporation, viscosity, surface tension, and solubility are related to the strength of attractive forces between molecules. These attractive forces are called Intermolecular Forces. The amount of "stick togetherness" is important in the interpretation of the various properties listed above. There are four types of intermolecular forces. Most of the intermolecular forces are identical to bonding between atoms in a single molecule. Intermolecular forces just extend the thinking to forces between molecules and follows the patterns already set by the bonding within molecules.
1. IONIC FORCES • The forces holding ions together in ionic solids are electrostatic forces. Opposite charges attract each other. • These are the strongest intermolecular forces. Ionic forces hold many ions in a crystal lattice structure.
Ion - Ion Interactions Two oppositely-charged particles flying about in a vacuum will be attracted toward each other, and the force becomes stronger and stronger as they approach until eventually they will stick together and a considerable amount of energy will be required to separate them. They form an ion-pair, a new particle which has a positively-charged area and a negatively-charged area. There are fairly strong interactions between these ion pairs and free ions, so that these the clusters tend to grow, and they will eventually fall out of the gas phase as a liquid or solid (depending on the temperature).
Ionic Bonding • Ionic bonding is best treated using a simple electrostatic model . The electrostatic model is simply an application of the charge principles that opposite charges attract and similar charges repel. • An ionic compound results from the interaction of a positive and negative ion, such as sodium and chloride in common salt.The IONIC BOND results as a balance between the force of attraction between opposite plus and minus charges of the ions and the force of repulsion between similar negative charges in the electron clouds. • In crystalline compounds this net balance of forces is called the LATTICE ENERGY. Lattice energy is the energy released in the formation of an ionic compound.
Ionic Bonding • DEFINITION: The formation of an IONIC BOND is the result of the transfer of one or more electrons from a metal onto a non-metal.Metals, with only a few electrons in the outer energy level, tend to lose electrons most readily. The energy required to remove an electron from a neutral atom is called the IONIZATION POTENTIAL. • Energy + Metal Atom ---> Metal (+) ion + e- • Non-metals, which lack only one or two electrons in the outer energy level have little tendency to lose electrons - the ionization potential would be very high. • Instead non-metals have a tendency to gain electrons. The ELECTRON AFFINITY is the energy given off by an atom when it gains electrons. • Non-metal Atom + e- --- Non-metal (-) ion + energy • The energy required to produce positive ions (ionization potential) is roughly balanced by the energy given off to produce negative ions (electron affinity). The energy released by the net force of attraction by the ions provides the overall stabilizing energy of the compound.
METALLIC STRUCTURES • What is a metallic bond? • Metals tend to have high melting points and boiling points suggesting strong bonds between the atoms. • Even a metal like sodium (melting point 97.8°C) melts at a considerably higher temperature than the element (neon) which precedes it in the Periodic Table.
Metallic bonding in sodium • Sodium has the electronic structure 1s22s22p63s1. When sodium atoms come together, the electron in the 3s atomic orbital of one sodium atom shares space with the corresponding electron on a neighbouring atom to form a molecular orbital - in much the same sort of way that a covalent bond is formed. • The difference, however, is that each sodium atom is being touched by eight other sodium atoms - and the sharing occurs between the central atom and the 3s orbitals on all of the eight other atoms. And each of these eight is in turn being touched by eight sodium atoms, which in turn are touched by eight atoms - and so on and so on, until you have taken in all the atoms in that lump of sodium.
Metallic bonding in sodium • All of the 3s orbitals on all of the atoms overlap to give a vast number of molecular orbitals which extend over the whole piece of metal. There have to be huge numbers of molecular orbitals, of course, because any orbital can only hold two electrons. • The electrons can move freely within these molecular orbitals, and so each electron becomes detached from its parent atom. The electrons are said to be delocalised. The metal is held together by the strong forces of attraction between the positive nuclei and the delocalised electrons.
Metallic bonding in sodium • This is sometimes described as "an array of positive ions in a sea of electrons". • If you are going to use this view, beware! Is a metal made up of atoms or ions? It is made of atoms. • Each positive centre in the diagram represents all the rest of the atom apart from the outer electron, but that electron hasn't been lost - it may no longer have an attachment to a particular atom, but it's still there in the structure. Sodium metal is therefore written as Na - not Na+ or potassium metal is written as K but not K+.
Metallic bonding in magnesium • If you work through the same argument with magnesium, you end up with stronger bonds and so a higher melting point. • Magnesium has the outer electronic structure 3s2. Both of these electrons become delocalised, so the "sea" has twice the electron density as it does in sodium. The remaining "ions" also have twice the charge (if you are going to use this particular view of the metal bond) and so there will be more attraction between "ions" and "sea". • More realistically, each magnesium atom has one more proton in the nucleus than a sodium atom has, and so not only will there be a greater number of delocalised electrons, but there will also be a greater attraction for them. • Magnesium atoms have a slightly smaller radius than sodium atoms, and so the delocalised electrons are closer to the nuclei. Each magnesium atom also has twelve near neighbours rather than sodium's eight. Both of these factors increase the strength of the bond still further.
Metallic bonding in transition elements • Transition metals tend to have particularly high melting points and boiling points. • The reason is that they can involve the 3d electrons in the delocalisation as well as the 4s. • The more electrons you can involve, the stronger the attractions tend to be.
d-block elements • Remember that the 4s orbital has a lower energy than the 3d orbitals and so fills first. Once the 3d orbitals have filled up, the next electrons go into the 4p orbitals as you would expect. • d-block elements are elements in which the last electron to be added to the atom is in a d orbital. The first series of these contains the elements from scandium to zinc, which you probably called transition elements or transition metals. • The terms "transition element" and "d-block element" don't quite have the same meaning, but it doesn't matter in the present context.
d-block elements • d electrons are almost always described as, for example, d5 or d8 - and not written as separate orbitals. • Remember that there are five d orbitals, and that the electrons will inhabit them singly as far as possible. Up to 5 electrons will occupy orbitals on their own. After that they will have to pair up. d5 means: d8 means :
d-block elements • Notice in what follows that all the 3-level orbitals are written together, even though the 3d electrons are added to the atom after the 4s. • Sc: 1s22s22p63s23p63d14s2 • Ti:1s22s22p63s23p63d24s2 • V :1s22s22p63s23p63d34s2 • Cr:1s22s22p63s23p63d54s1 Whoops! Chromium breaks the sequence. In chromium, the electrons in the 3d and 4s orbitals rearrange so that there is one electron in each orbital. It would be convenient if the sequence was tidy - but it's not!
d-block elements • Mn:1s22s22p63s23p63d54s2 (back to being tidy again) • Fe : 1s22s22p63s23p63d64s2 • Co: 1s22s22p63s23p63d74s2 • Ni : 1s22s22p63s23p63d84s2 • Cu: 1s22s22p63s23p63d104s1 (another awkward one!) • Zn: 1s22s22p63s23p63d104s2 (And at zinc the process of filling the d orbitals is complete.) A transition element is defined as one which has partially filled d orbitals either in the element or any of its compounds. Zinc (at the right-hand end of the d-block) always has a completely full 3d level (3d10) and so doesn't count as a transition element.
The metallic bond in molten metals • In a molten metal, the metallic bond is still present, although the ordered structure has been broken down. • The metallic bond isn't fully broken until the metal boils. That means that boiling point is actually a better guide to the strength of the metallic bond than melting point is. • On melting, the bond is loosened, not broken.
The physical properties of metals • Melting points and boiling points • Metals tend to have high melting and boiling points because of the strength of the metallic bond. The strength of the bond varies from metal to metal and depends on the number of electrons which each atom delocalises into the sea of electrons, and on the packing. • Group 1 metals like sodium and potassium have relatively low melting and boiling points mainly because each atom only has one electron to contribute to the bond. • They have relatively large atoms (meaning that the nuclei are some distance from the delocalised electrons) which also weakens the bond.
The physical properties of metals • Electrical conductivity • Metals conduct electricity. The delocalised electrons are free to move throughout the structure in 3-dimensions. They can cross grain boundaries. Even though the pattern may be disrupted at the boundary, as long as atoms are touching each other, the metallic bond is still present. • Liquid metals also conduct electricity, showing that although the metal atoms may be free to move, the delocalisation remains in force until the metal boils.
Electrical conductivity of metals Electric current is the flow of electrons in a wire. In metals, the outer electrons of the atoms belong to a ‘cloud’ of delocalised electrons. They are no longer firmly held by a specific atom, but instead they can move freely through the lattice of positive metal ions. Normally they move randomly. However, when the wire is connected to a cell, they are pushed away from the negative terminal and drawn to the positive one. The cloud of electrons drifts through the wire. The drift velocity of the cloud is about 3 mm s-1. The electrons within the cloud are still moving randomly (at much higher speeds) - rather like a swarm of bees leaving a hive. Animation showing electrons moving randomly and then the movement of electrons through a wire
Thermal conductivity • Metals are good conductors of heat. • Heat energy is picked up by the electrons as additional kinetic energy (it makes them move faster). • The energy is transferred throughout the rest of the metal by the moving electrons.
Ionic vibrations The positive metal ions in a metal structure are packed closely together in a symmetrical geometric arrangement. They don’t move from their position in the lattice but they are constantly vibrating. If a metal is heated, the positive metal ions vibrate more vigorously. These ions collide with neighbouring ions and make them vibrate more vigorously too. In this way, the energy is passed, or conducted, through the metal. A cool lattice. If we heat the left hand end, then the energy will be carried along by conduction.
Free electrons • However, metals are particularly good conductors of heat. In general, they are better than ionic compounds which also have strong bonds. So we need another mechanism to explain their especially good conductivity. It is their free electrons. • Free electronsThe ions in the lattice are vibrating . The ions at the hot end of a piece of metal vibrate more. Let's look at just a few electrons. • The electrons at the hot end will speed up – they gain kinetic energy from the vigorously vibrating ions. • Some of them will move down to the cooler end and collide with ions that are vibrating less vigorously than those at the hot end. • In these collisions, the electrons will lose kinetic energy and make the ions vibrate more vigorously. • In effect, the electrons have carried the vibrational energy from the hot end to the cold end. And, because they are free to move through the lattice, they are able to do this more quickly than the bonds between the ions in the lattice
Thermal conductivity of metals How a metal conducts by the movement of free electrons. Metals are good conductors of heat. There are two reasons for this: • the close packing of the metal ions in the lattice • the delocalised electrons can carry kinetic energy through the lattice
Strength and workability • Malleability and ductility • Metals are described as malleable (can be beaten into sheets) and ductile (can be pulled out into wires). • This is because of the ability of the atoms to roll over each other into new positions without breaking the metallic bond. • If a small stress is put onto the metal, the layers of atoms will start to roll over each other. • If the stress is released again, they will fall back to their original positions. Under these circumstances, the metal is said to be elastic.
Strength and workability • If a larger stress is put on, the atoms roll over each other into a new position, and the metal is permanently changed.
The hardness of metals • This rolling of layers of atoms over each other is hindered by grain boundaries because the rows of atoms don't line up properly. It follows that the more grain boundaries there are (the smaller the individual crystal grains), the harder the metal becomes. • Offsetting this, because the grain boundaries are areas where the atoms aren't in such good contact with each other, metals tend to fracture at grain boundaries. Increasing the number of grain boundaries not only makes the metal harder, but also makes it more brittle.
Controlling the size of the crystal grains • If you have a pure piece of metal, you can control the size of the grains by heat treatment or by working the metal. • Heating a metal tends to shake the atoms into a more regular arrangement - decreasing the number of grain boundaries, and so making the metal softer. Banging the metal around when it is cold tends to produce lots of small grains. Cold working therefore makes a metal harder. To restore its workability, you would need to reheat it. • You can also break up the regular arrangement of the atoms by inserting atoms of a slightly different size into the structure. Alloys such as brass (a mixture of copper and zinc) are harder than the original metals because the irregularity in the structure helps to stop rows of atoms from slipping over each other.
Metallic Bonding A. Outermost electrons wander freely through metal. Metal consists of cations held together by negatively-charged electron "glue.“ B. Free electrons can move rapidly in response to electric fields, hence metals are a good conductor of electricity. C. Free electrons can transmit kinetic energy rapidly, hence metals are good conductors of heat. D. The layers of atoms in metal are hard to pull apart because of the electrons holding them together, hence metals are tough. But individual atoms are not held to any other specific atoms, hence atoms slip easily past one another. Thus metals are ductile. Metallic Bonding is the basis of our industrial civilization.
2. DIPOLE FORCES • Polar covalent molecules are sometimes described as "dipoles", meaning that the molecule has two "poles". One end (pole) of the molecule has a partial positive charge while the other end has a partial negative charge. The molecules will orientate themselves so that the opposite charges attract principle operates effectively.
FORCES BETWEEN MOLECULES • There are in fact three basic types of interaction between molecules which are, in order of increasing strength: • van der Waals interactions or dispersion forces, • dipole-dipole interactions, • hydrogen bonds, • and these secondary bonding interactions, like the primary bonding interactions, involve the electrons (atomic glue!).
DIPOLE FORCES • In the example on the right,hydrochloric acid is a polar molecule with the partial positive charge on the hydrogen and the partial negative charge on the chlorine. • A network of partial + and - charges attract molecules to each other.
Polar Covalent Compounds • Introduction to Covalent Bonding: • Bonding between non-metals consists of two electrons shared between two atoms. In covalent bonding, the two electrons shared by the atoms are attracted to the nucleus of both atoms. Neither atom completely loses or gains electrons as in ionic bonding. • There are two types of covalent bonding: • 1. Non-polar bonding with an equal sharing of electrons. • 2. Polar bonding with an unequal sharing of electrons. The number of shared electrons depends on the number of electrons needed to complete the octet.
Polar Covalent Compounds • POLAR BONDING results when two different non-metals unequally share electrons between them. One well known exception to the identical atom rule is the combination of carbon and hydrogen in all organic compounds. • The non-metal closer to fluorine in the Periodic Table has a greater tendency to keep its own electron and also draw away the other atom's electron. It is NOT completely successful. • As a result only partial charges are established. One atom becomes partially positive since it has lost control of its electron some of the time. The other atom becomes partially negative since it gains electron some of the time.
WATER • Water, the most universal compound on all of the earth, has the property of being a polar molecule. • As a result of this property, the physical and chemical properties of the compound are fairly unique.Hydrogen Oxide or water forms a polar covalent molecule. • The graphic on the right shows that oxygen has 6 electrons in the outer shell. Hydrogen has one electron in its outer energy shell. • Since 8 electrons are needed for an octet, they share the electrons.However, oxygen gets an unequal share of the two electrons from both hydrogen atoms. • Again, the electrons are still shared (not transferred as in ionic bonding), the sharing is unequal. • The electrons spends more of the time closer to oxygen. As a result, the oxygen acquires a "partial" negative charge. • At the same time, since hydrogen loses the electron most - but not all of the time, it acquires a "partial" charge. The partial charge is denoted with a small Greek symbol for delta.
If the difference in electronegativity is not so great, however, there will be some degree of sharing of the electrons between the two atoms. The result is the same whether two ions come together or two atoms come together: Polar Molecule
Polar molecules can interact with ions: • Ion - Dipole Interactions
or with other polar molecules: Dipole - Dipole Interactions
HYDROGEN BONDING • To recognize the possibility of hydrogen bonding, examine the Lewis structure of the molecule. • The electronegative atom must have one or more unshared electron pairs as in the case of oxygen and nitrogen, and has a negative partial charge. • The hydrogen, which has a partial positive charge tries to find another atom of oxygen or nitrogen with excess electrons to share and is attracted to the partial negative charge. This forms the basis for the hydrogen bond.
3. HYDROGEN BONDING • The hydrogen bond is really a special case of dipole forces. A hydrogen bond is the attractive force between the hydrogen attached to an electronegative atom of one molecule and an electronegative atom of a different molecule. Usually the electronegative atom is oxygen, nitrogen, or fluorine. • In other words - The hydrogen on one molecule attached to O or N that is attracted to an O or N of a different molecule.
HYDROGEN BONDING • In the graphhic on the right, the hydrogen is partially positive and attracted to the partially negative charge on the oxygen or nitrogen. Because oxygen has two lone pairs, two different hydrogen bonds can be made to each oxygen. • This is a very specific bond as indicated. Some combinations which are not hydrogen bonds include: hydrogen to another hydrogen or hydrogen to a carbon.
HYDROGEN BONDING • Hydrogen bonding is usually stronger than normal dipole forces between molecules. Of course hydrogen bonding is not nearly as strong as normal covalent bonds within a molecule - it is only about 1/10 as strong. • This is still strong enough to have many important ramifications on the properties of water.
