Structure, Bonding and Main Group Chemistry

Structure, Bonding and Main Group Chemistry

Atomic Structure:

Definitions

Atomic Number is the number of protons in the nucleus of the atom.

Mass number is the number of protons and neutrons in the nucleus of the atom.

Isotopes are atoms of the same element with the same number of protons and different numbers

of neutrons.

Relative Atomic Mass is the average mass of the atom divided by 1/12th of the mass of an atom

of carbon-12.

Relative Isotopic Mass is the mass of an isotope divided by 1/12th the mass of an atom of carbon-12.

Relative Molecular Mass is the sum of the relative atomic masses of the atoms in it.

Molar Mass is the mass of one mole of a substance. It is the same as the relative molecular mass.1st Ionisation Energy is the energy required to turn one mole of gaseous atoms of an element into one

mole of gaseous 1+ ions by removing a mole of electrons.

1st Electron Affinity is the energy required to turn one mole of gaseous atoms of an element into one

mole of gaseous 1- ions by adding one mole of electrons.

Mass Spectroscopy:

A source is vapourised and then ionised by firing electrons at it which remove electrons from the source atom's shells forming positive ions. These are accelerated through an electric field and deflected according to their masses using a magnetic field on a curved path, and detected. They are then plotted on a graph and the data is used to calculate the source's average relative atomic mass. All species detected are positive because of ionisation, so that they can be detected and accelerated.Electron Structure:

Two electrons can occupy one orbital The first shell has one orbital (1s)

The second shell has four orbitals (2s, three 2p)

The third shell has nine orbitals (3s, three 3p, five 3d)

The fourth shell has sixteen orbitals (4s, three 4p)

Due to stability, the 4s orbital is occupied before the 3d orbitals.

Example of electron structure notation:

Phosphorus: 1s22s22p63s23p3, or [Ne]3s23p3Calcium: 1s22s22p63s23p64s2, or [Ar]4s2Trends of sizes of atoms and ions:

Atoms are smaller moving right to left because the nuclear charge increases, pulling each atom towards it more. The radius of a chlorine atom is smaller than that of a magnesium atom for this reason.

Atoms are larger moving down because there are more shells of electrons. A negative ion is larger than its neutral atom because the extra electron repels the others, so the shell gets slightly wider.

Ionisation Energies:

In general, the ionisation energies going from left to right in a period increases because of the increasing nuclear charge without an increasing amount of electron shielding.

Ionisation energies going down the table decrease because the electrons are further from the nucleus, so the attraction is weaker and less energy is required to break it.

The second ionisation energy is always larger than the first because the second means removing an electron from an already positive ion.

Large jumps in successive ionisation energies indicate that electron is coming from a lower shell. This helps understand what group the element is in; if the third ionisation energy is much higher than the first two, then the third removed electron is from a lower shell and so the element must have had 2 electrons in the outer shell - is was therefore in group 2.

Electron Affinity:

The first electron affinity energy is always negative because it is an electron being attracted to a positive nucleus.

The seconds and subsequent affinities are positive because it is an electron being forced onto an already negative ion.

Structure and Bonding:Intramolecular Forces (chemical bonds inside molecules):

Ionic - electrostatic attraction between ions of opposite charge after an electron transfer

Polar covelant - covelant bonds where the shared electron pair is closer to one atom than the other.

Pure covelant - covelant bonds where the shared electron pair is in the centre of the covelant bond. Metallic - delocalised electrons moving around the solid due to overlapping electron shellsIntermolecular Forces (bonds between molecules):

Hydrogen bonds - bonds between a + H from one molecule and a - N, O or F from another.

Dipole-dipole bonds - bonds between any + atoms and any other - atoms from different molecules. Hydrogen bonds are specific dipole-dipole bonds because of their strength. Van der Waals bonds - bonds existing due to dispersion forces (electron positions in atoms)Definitions:

Ionic Bonds are electrostatic attractions between cations and anions that results from an electron transfer of one or more electrons.

Covelant Bonds are atoms with a shared bonding electron pair. It results from the atoms' orbitals overlapping, either head on, causing bonds, or side on, causing bonds. A double covelant bond is one bond and one bond.

Dative Covelant Bonds are covelant bonds where both electrons in a pair come from the same atom.

A Metallic Bond is the force of attraction between the sea of delocalised electrons and the positive ions which are arranged in a regular lattice.

Electronegativity is the tendency of an atom in a covelant bond to attract the bonding electron pair.Van der Waals bonds are bonds caused by the repulsions between electrons of neighbouring atoms or ions. When an atom or ion's electrons are unevenly dispersed in their shells, they can repel atoms to make more atoms unevenly dispersed, causing an induced dipole.

Ionic Bonding:

In general, there must be a large enough gap in atom electronegativities to cause an ionic bond. (around 1.5) This gives the atom with the high electronegativity enough relative electronegativity to completely remove the other atom's electron.

However, the electronegativity gap might be small enough for there to exist polarisation in the ionic bond. Polarising cations are relatively electronegative (have a high polarising power) due to their high nuclear charges and small radius'. Cations with high polarising powers are Be2+ and Al3+. Ionic bonding leads to lattices because each ion can be attracted to as many other (oppositely charged) ions as the repulsion between them will allow.

Ionic substances will be soluable in water if the energy required to split the ionic lattices up is less than the hydration energy. This is why some strong ionic compounds with high lattice energies do not dissolve in water.

Covelant Bonding:

There are four types of covelant bonding:

Giant covelant networks, such as diamond and graphite, where each atom can covelantly bond enough times to make large 3d structures.

Simple molecular, such as I2 and CH4, where there are a specific number of bonds.

Hydrogen bonded molecular, such as ice and ethanol, where simple molecular covelantly bonded molecules have hydrogen bonds altering their properties.

Non-crystalline, such as polymers.

Some covelant bonds are polar, such as water, because of their structures, although linear, planar, tetrahedral or octahedral covelant molecules are not because of symmetry - the charges cancel out.Intermolecular Forces:

The strongest intermolecular force is the hydrogen bond.

The next strongest is usually van der Waals bonds. Generally, the more electrons in the atoms, the stronger the van der Waals forces, which explains why the pure covelant simple molecule I2 is a solid and Cl2 is a gas, both at room temperature and pressure, as Iodine has more electrons than Chlorine. The same can be said about the noble gasses - the boiling temperature increases down the group because there are stronger vdW forces between atoms.

In general, going down the table will increase the van der Waals forces.

If the number of electrons of atoms in comparison is roughly equal, and there are no hydrogen bonds in either, then the differences in biolong temperatures can be put down to dipole-dipole attractions.

Melting:

Melting occurs when the kinetic energies of the vibrating solid particles are large enough to be transferred into potential energies of separate molecules, and the lattice is broken up.

When ionic lattices such as NaCl melt there is a lot of kinetic energy needed by the ions to overcome the bonds and split up, as the ionic bonds are very strong. Therefore NaCl has a high melting point.

In covelantly bonded simple molecular structures such as ice, not nearly as much kinetic energy is needed to overcome the bond strength and be transferred to potential energy because the bonds between water molecules (hydrogen bonds) are weaker than ionic bonds.

In giant covelantly bonded atomic structures, such as diamond, as with ionic lattices, lots of kinetic energy is needed to break the covelant bonds, so the structure has a high melting point.Shapes of molecules:

Electron Pair Repulsion Theory states that (i) the electron pairs arrange themselves as far apart from each other as possible in order to minimise repulsion, and that (ii) the repulsion between lone pairs is greater than that between a lone pair and a bond pair, which is greater than that between two bond pairs.The shapes of molecules is defined by how the bond electron pairs and lone electron pairs repel.

bonds should not be counted when working out molecular shapes.The Periodic Table:Valence Electrons are outer shell electrons.

Electron shells are filled in the following way: 1s,2s,2p,3s,3p,4s,3d,4p,5s,4d,5p,6s

Trends:

Melting temperatures of the period 3 elements:

The melting temperatures depend on the bonds in the solids. Goind across a period means going between types of bonds - metallic, giant covelant structures, and simply bonded covelant molecules with intermolecular bonds holding them together. In metallic bonds, the higher charge density of the ions means the stronger the metallic bond.

Na, Mg and Al are the three period 3 atoms that bond metallically. Because of the last point, you would expect their melting temperatures to increase, staring high as would be expected from metallic bonding. Silicon bonds in a giant atomic lattice and so has a very high melting temperature as you need to break the covelant bonds themselves, and the others - P, S, Cl and Ar - all bond in covelant simple molecules (P4, S8, Cl2, and Ar) the only forces of attraction keeping these as solids are dispersion forces (vdW), and are weak in comparison to the first 4.

Electrical conductivity trends:

Conductivity is the ability to allow a flow of delocalised electrons through a substance. Metals are very good because of their already delocalised electrons, and some other molecules such as graphite can due to its bonds leaving electrons free to move above and below the hexagonal planes.Groups I and II:

Going down the table, elements become more metallic, meaning:

Their oxides become more basic

They form covelant bonds less readily

They form positive ions more readily.

Group I are all metals and their melting temperatures decrease down the group, because the metallic bonds are further apart due to more electron shells. Group II are all metals as well, and their melting temperatures decrease down the group as well, in general. Their melting temperatures are higher than group I elements though, as the nuclei in the metallic bonds have a stronger charge and so form stronger metallic bonds.

Reactions with Oxygen:

All group I metals react with oxygen (with increasing vigour down the group) to form oxides, Na forms a peroxide and K and below form superoxides.

All group II metals react with oxygen to form oxides, and Ba forms a peroxide.

The reactions with oxygen become increasingly vigorous down the groups because oxygen is an oxidising agent and it can gain electrons easier further down the group when the distance from the outer electron is greater and therefore the pull of the nucleus is weaker.

Reactions with Chlorine: All group I metals react vigorously with chlorine to produce metal salts in the form MCl, and all these dissolve in water. All group II metals react vigorously with chlorine to produce ionic metal salts except Be which is covelant when anhydrous.

Reaction with Water:

Group I metals react with cold water to form an alkaline hydroxide solution and hydrogen. Group II metals from Ca down react with cold water to give hydroxides and hydrogen, Mg reacts with steam to give a metal oxide and hydrogen.

The rates of reactions with water increase down the group because the reaction with water involves giving up outer shell electrons. Further down the group, the elements give up their outer shell electrons with increasing ease.Soluabilities of Group II Sulphates and Hydroxides: Group II sulphates decrease in soluability down the group. The first 2 are soluable, the third is slightly soluable and further down no sulphates are soluable.

Group II hydroxides increase in soluability down the group. The first 2 are insoluable, the next 2 are slightly sluable and the rest are soluable.

Thermal stabilities of Group I and II nitrates and carbonates:

All increase in thermal stability, because as you go down the groups, the electrons have less attraction to the nucleus, and so make more stable ionic bonds to the anions, and so more heat is needed to break the bonds.Oxidation and Reduction (REDOX):Redox Reactions are reactions which involve a transfer of electrons.Defenitions:

Oxidation occurs when a substance loses a number of electrons, and its oxidation number increases.

Reduction occurs when a substance gains a number of electrons and its oxidation number decreases.

An Oxidising Agent is an agent which causes an oxisation of the substance it is reacting with, itself being reduced. It accepts electrons.

A Reducing Agent is an agent which causes reduction of the substance it reacts with, and it realidy donates electrons.

Oxidation numbers:

An oxidation number, simply, is the charge on each atom in a compound. For example, in:

H2O, the most electronegative, the oxygen, is given a 2- because it forms 2 bonds. The H must have (and usually has) +1. 2(+1) + (-2) = 0, which is the overall charge on the atom. If the water were to be electrolysed: 2H2O ---> 2H2 + O2, the oxygen's oxidation number would rise and the hydrogen's oxidation number would fall - both going to 0, as they are both separate molecules with no net charge. In this reaction, the oxygen is bein oxidised, and the hydrogen is being reduced.

Half Equations:These are written as either oxidations or reductions:

Oxidation: 2O2- + 2e- ---> O2Reduction:2H2+ ---> 2e- + 2H2Which can be added together once the 'coefficients' of electrons one side of one half equation and the opposite side of the other are equal. They can then be cancelled out, giving the original equation.

Oxidising agents with oxygen in sometimes need H+ ions on the left hand side to 'mop up' spare oxygens, giving H2O on the right hand side. The reduction of permanganate is a good example: MnO4 + 8H+ + 5e- ---> 4H2O + Mn2+ Oxidising agents in alkaline solutions sometimes need OH- on one side and H2O on the other, such as the reduction of chromate: Cr3+ + 8OH- ---> Cr2O42- + 4H2O + 3e-Group 7:

As you go down the group, the outer shell electrons are less attracted to the nucleus so it is easier for them to be lost. Therefore, the halogens become better reducing agents down the group (reducing agents are oxidised and oxidation is the loss of electrons). Chlorine holds onto its electrons the strongest and is therefore the strongest oxidising agent (best for gaining electrons). This means that the reaction Cl2 + 2HI ---> 2HCl + I2 can happen, as chlorine has a stronger attraction to the hydrogen's electron, but the reaction, for example, Br2 + 2MgCl ---> 2MgBr + Cl2, cannot happen as MgCl is more stable than MgBr, again because the chlorine has a stronger attraction to the magnesium's lost electron because that electron is in a shell closer to the nucleus.

Defenitions:

Disproportination is where one element is simultaneously oxidised and reduced. Cl2 can be oxidised and reduced in an alkali solution to oxygen chloride, and a chlorine ion and water.

Chlorine is a yellowy green gas at room temperature

Bromine is a brown gas at room temperature

Iodine is a dark gray lustrous solid with a violet vapour

Addition of concentrated sulphuric acid to a solid halide:

The reaction of sodium chloride to concentrated supluric acid:2NaCl + H2SO4 ---> 2HCl + Na2SO4The reaction of sodium bromide to concentrated sulphuric acid:

2NaBr + H2SO4 ---> 2HBr + Na2SO42HBr + H2SO4 ---> Br2 + 2H2O + SO2

The Hydrogen bromide is a powerful enough reducing agent to reduce the concentrated sulphuric acid, forming bromine gas, water, and sulphur dioxide.

Redox with Chlorine:

Chlorine can come as chlorates - mixtures of chlorine and oxygen. Because the chlorine is less electronegative than the oxygen it is assigned a positive oxidation number. Chlorates are very strong oxidising agents, because they will very readily accept electrons. Once reduced to Cl2 the chlorine isn't as strong an oxidising agent but still will readily gain electrons better than Br2 or I2.

The oxidation states of chlorine are:

OCl3- - oxygen chlorate (V)

OCl- - oxygen chlorate (III)

Cl2 - chlorine

Cl- - chloride