Atomic structure and bonding

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Electronic Structure

Electron shells are made up of sub-shells and orbitals, and the electrons have fixed energies. They orbit in energy levels. Each shell has a quantum number and the further away from the nucleus it is the higher its energy and the larger the quantum number. Not all electrons in a shell have the same energy. The shells are split into sub-shells with different energies. The sub-shells have orbitals containing up to two electrons. These are called the s, p, d and f orbitals. The s orbitals are spherical in size, whilst p orbitals are dumb-bell shaped, orientated in three different directions at right angles to one another. The orbital is where the electron orbits the nucleus in. Electron configurations show the amount of electrons in an atom and what energy levels they are in. Electrons fill out the lowest energy level first, starting at 1s with two electrons, 2s with two electrons, 2p with 6, 3s with two, 3p with 6, 4s with two, 4d with 6, and the list continues. Electrons fill orbitals singly before sharing. Inert gas symbols may be used in electron configurations, where the inert gas symbol is shown in brackets and the energy levels not in the inert gas are shown after it. Chromium and copper donate an electron from the 4s shell to the 3d shell because they are more stable with a full/half-full 3d sub-shell. The electronic structure decides the chemical properties of an element. The s-block elements (Groups 1 and 2) have 1 or 2 outer shell electrons. These can easily form positive ions with a 1+ or 2+ charge to get noble gas configurations. The elements in group 5, 6 and 7 can gain 1, 2 or 3 electrons and get a noble gas electronic configuration. In group 0 the atoms are inert. The d block elements tend to lose s and d electrons to form positive ions.

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Ionisation energies

Ionisation is the removal of one or more electrons. The energy needed to remove the first electron from its atom is the first ionisation energy, for the second it is the second ionisation energy and so on. The first ionisation energy is the energy needed to remove 1 electron from each atom in 1 mole of gaseous atoms to form 1 mole of gaseous 1+ ions. To ionise an atom you put energy in, so it is an endothermic process. The equation for the first ionisation energy would be 1 mole of a gaseous atom makes 1 mole of 1+ ions and an electron. For second it is the 2+ ion add one electron, and so on. The factors affecting ionisation energy are: Nuclear charge - The more protons in the nucleus, the stronger the positive charge and so the electrons are attracted more. Distance from nucleus - The further the electron is from the nucleus the less attracted it is to the nucleus. Shielding - The electrons in the inner shells shield the effect of the positive charge pulling the electrons in. A high ionisation energy shows that there is a strong attraction between the nucleus and the electron. The ionisation energy decreases down a group but generally increases across a period due to stronger nuclear charge. Sucessive ionisation energies involve removing more electrons. The more you remove the harder it is to remove any more, due to the fact that it is being pulled by an increasingly positive ion. The big jumps in ionisation energies show shell structure. In sodium for example the first big jump is after one electron is removed - so it is in group 1. You can predict the electronic structure of an element by counting all the points before each jump to work out the amount of electrons in each shell.

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Periodic properties

All elements within a period have the same number of electron shells and all elements in a group have the same number of electrons in the outer shell - so all have similat properties. You can use the periodic table to work out electron configuration by labelling all the shells and sub-shells, then reading across and down until you get to your element. Atomic radius decreases across a period because the positive charge in the nucleus increases and so pulls electrons in more. Look at 2. Ionisation energies for a refresher of what affects ionisation energies. The general trend is ionisation energies increase across a period because the number of protons is increasing giving a stronger nuclear attraction. All the electrons are in the same energy level so there is little shielding to affect the ionisation energies. However, there are small drops inbetween Groups 2 and 3 and 5 and 6. For Beryllium, its outer electron is in its 2s sub-shell, but boron's outer electron is in the 2p sub-shell. This sub-shell has a higher energy level so it is a little bit further from the nucleus, and it has the 2s shells giving some shielding. The combination of these two factors overrides the effect of the increased nuclear charge. For nitrogen and oxygen, the key factor here is that nitrogen has one electron in each orbital, but oxygen has two electrons sharing one orbital. These repel each other and make it easier to remove one of them. Melting points increase across a period for metals because the metal-metal bonds get stronger due to more delocalised electrons and a decreasing radius. This makes the ions attract more strongly. The elements with macromolecular structures have strong covalent bonds linking their atoms, and a lot of energy is needed to break these. Simple molecular structures have only London forces between them and these are easily broken. More atoms means stronger London forces so sulphur has a slightly higher melting/boiling point than phosphorous or chlorine. The noble gases have lowest melting/boiling points as they exist as individual atoms.

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Ionic Bonding

Compounds are atoms of different elements bonded together. They contain either ionic or covalent bonds. Ionic bonding is when ions are stuck together by electroscatic attraction. Ions are formed when electrons are transferred from one atom to another. The simplest ions are those that have lost or gained 1, 2 or 3 electrons. Positive ions are called cations and negative ones are called anions. Sodium chloride and Magnesium oxide are ionic compounds. Dot and cross diagrams are used to show how ionic bonding works. Dots represent the electrons of one element and crosses are the electrons of the other one. The total overall charge is zero. Ionic crystals are giant lattices of ions, made up of the same unit repeated over and over again. Sodium chloride, for example is a cubic structure. All ionic compounds have differently shaped structures.  Elements in the same group form ions with the same charge. The theory of ionic bonding fits the evidence from physical properties such as they have high melting points, which tells you the atoms are held together by a strong attraction. They are soluble in water but not in non-polar solvents - this tells you the particles are charged. The ions are pulled apart by polar molecules but not by non-polar molecules. Ionic compounds only conduct electricity when molten or dissolved - this supports the idea that there are ions but they are free to move. These all support the ionic model. Not all ions are made from single atoms. Lots of ions are made up of a group of atoms with a total overall charge, such as nitrate, with a -1 charge, carbonate and sulfate with a -2 charge, and ammonium with a +1 charge.

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More on Ions

Ions are smaller than atoms for metals but larger for non-metals. The ionic radius of a metallic element is smaller than the atomic radius because removing an electron means the positive charge is higher than the negative charge, and are pulled closer to the nucleus. There is also less electron shielding. For non-metals, they gain electrons, so there is a bigger negative charge in the electron cloud and therefore there is greater repulsion between the electrons and the cloud expands a bit. The size of an ion depends on its atomic number and charge. The ionic radius decreases as you go down a group, because extra electron shells are added. Isoelectronic ions are ions of different atoms with the same number of electrons, like flourine with a -1 charge and sodium with a +1 charge. But sodium has more protons, so the electrons are attracted to the nucleus more strongly. There are different kinds of evidence for the existence of ions, like the migration of ions on wet filter paper. When you electrolyse a green solution of copper(II) chromate(VI) the filter paper turns blue at the cathode and yellow at the anode. Copper(II) ions are blue in solution and chromate(VI) ions are yellow. Together they make a green colour. When you pass a current through the solution, the positive ions (copper) move to the cathode while the negative ions (chromate) move to the anode. Electron density maps made using X-ray crystallography show that there are spaces between the ions where the density of electrons is zero. This shows that the atoms in an ionic crystal have no shared electrons - the bonding electrons have moved from one atom to the other.

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Ions and Born-Haber Cycles

Born-Haber cycles can be used to calculate lattice energies. Born-Haber cycles show enthalpy changes when a solid ionic compound is formed from its elements in their standard stats. They show two routes, much like Hess Law diagrams - one direct and one indirect. Both routes have the same enthalpy change. Born-Haber cycles are used to calculate lattice energies as they can't be found directly from experiments. To make sodium chloride from sodium and chloride, you need to atomise the sodium, then ionise it by removing one electron, then you atomise the chlorine, then you need the electron affinity of chlorine, then you need to add a certain lattice energy. Or you can form it directly with a standard enthalpy change equal to that of the standard enthalpy change of formation. To work out the lattice energy you take the sum of the enthalpy of atomisation of sodium, the 1st ionisation energy of sodium, then the enthalpy of atomisation of chlorine, then the electron affinity of chlorine, then the lattice energy and put it all equal to the standard enthalpy of formation. Then you rearrange it, being careful about the signs. Born-Haber cycles can show why some compounds don't exist. If a lot of energy is released during the formation of a compound the compound is stable. But some compounds either don't form or break up into more stable compounds. For example, sodium chloride forms because a negative enthalpy of formation. But for sodium(II) chloride, the enthalpy change is now positive because you need the second ionisation energy for sodium. It is energetically unfavourable. Similarly, magnesium chloride(III) doesn't form because the standard enthalpy change of formation is positive. Both magnesium chloride and magnesium chloride(I) both have negative enthaply changes of formation, but magnesium chloride releases more energy. So any magnesium chloride(I) forming disproportionates to magnesium chloride and magnesium.

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Lattice energies and Polarisation of ions

Theoretical lattice energies are based on the ionic model. You can work out lattice energies with a Born-Haber cycle or you can do some calulations based on the purely ionic model of a lattice. To work out a theoretical lattice energy you assume that all the ions are spherical and have their charge evenly distributed around them - a purely ionic lattice. Then you work out how strongly the ions are attracted to one another based on their charges, the distance between them and so on. Comparing lattice energies can tell you 'how ionic' an ionic lattice is. For any compound the experimental and theoretical lattice energies are usually different. If they are very similar with less than 1% difference between them the compound fits the purely ionic model very well. However the bigger the difference the more covalent character it has - i.e. not purely ionic, and are more polarised. Polarisation of ionic bonds leads to covalent character in ionic lattices. The reason has to do with the charge on the positive ion. With a +1 charge it doesn't pull electrons that much, but with a +2 or higher charge the positive charge pulls electrons towards it, polarising the bond. Small cations are very polarising. What normally happens in ionic compounds is that the positive charge on the cation attracts electrons towards it from the anion - this is polarisation. Small cations with a large charge are very polarising because they have a high charge density - it can pull electrons towards it. Large anions are polarised more easily because their electrons are further from the nucleus. So the electrons on large anions can be pulled away more easily towards cations. If a compound contains a cation with a high polarising ability and an anion that is easily polarised, some of the anions's electron charge cloud is dragged towards the positive cation. If the compound is polarised enough a partially covalent bond is formed. 

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Covalent bonding

Molecules are groups of atoms bonded together. Molecules are held together by strong covalent bonds, where two atoms share electrons. Electron density maps give evidence for covalent bonding. There is an area of high electron density between the two atoms - this shows that they are sharing electrons. Covalent bonds can be sigma bonds, where two s orbitals overlap in a straight line to form an area of high electron density. This is a single covalent bond. Or they can be pi bonds, when two electrons in p orbitals overlap. It has two parts to it - one above and one below the molecular axis, because the pi orbitals which overlap are dumb-bell shaped. Atoms can form double and triple covalent bonds. Double bonds are made up of a sigma bond plus a pi bond. The pi bond makes ethene more reactive than ethane, which only has sigma bonds. You need to be able to show electrons in the outer shells bonded to other atoms, using dots and crosses to show electrons. The electrons in the outer shells is the same as the group number the atom is in. Dative covalent bonding is where both electrons come from one atom, such as the ammonium ion. It's formed by dative covalent bonding. It forms when the nitrogen atom in an ammonia molecule donates a pair of electrons to a proton (hydrogen with a +1 charge.)

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Giant Covalent Structures and Metallic Bonding

Covalent bonds can form giant molecular structures, which have a huge network of covalently bonded atoms, sometimes called macromolecular structures. Not all form this type of giant structure, but carbon and silicon do because they can form four strong covalent bonds. Diamond is one example of a giant molecular structure, made up of carbon. Silicon dioxide has a tetrahedral arrangement. Silicon atoms form a giant covalent structure when they bond with oxygen to form silicon dioxide (quartz). Each silicon bonds with four oxygen atoms to form a big crystal lattice, but the structure isn't the same as diamond as oxygen can only bond with two silicon atoms. The properties of giant structures provide evidence for covalent bonding, such as they are insoluble in polar solvents, which shows they don't contain ions. They form hard crystals with very high melting points, due to the network of very strong covalent bonds. They don't conduct electricity (except for graphite), because all the bonding electrons are used to form covalent bonds. Graphite has delocalised electrons within its sheets of atoms. Metals have giant structures too. The outermost shells of electrons are delocalised - the electrons are free to move. The positive metal ions are attracted to these electrons, forming a lattice of closely packed positive ions in a sea of delocalised electrons. Metallic bonding explains the properties of metals. The number of delocalised electrons per atom affects the melting point. The more there are the stronger the bonding. The size of the metal ion affects charge density. A high charge density means a high melting point. They are malleable and ductile as the metal ions can easily slide over each other. The delocalised electrons can pass kinetic energy to each other, making them good thermal conductors. The delocalised electrons allow a current to pass through so they are good electrical conductors. They are insoluble except in liqiud metals due to the strength of the metallic bonds. Ionic compounds have high melting/boiling points, are solids at RTP (room temperature and pressure) do not conduct electricity as solids but will as liquids, and dissolve in water. Simple molecular compounds have low melting/boiling points, are usually liquid or gas at RTP, do not conduct electricity, and their solubility in water depends on how polarised the molecule is. Giant molecular and metallic compounds both have high melting/boiling points, are solids at RTP, and do not dissolve in water. But whilst giant molecular structures do not conduct electricity, metallic compounds do conduct electricity.

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