Chemistry Pack 6

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Chapter 6 of AS level Chemistry
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Chemistry Pack 6
1 Shapes of molecules
1.1 The electron pair repulsion theory predicts the shape and the polyatomic ions.
1.1.1 The e- pairs ,both bonded and lone, surrounding the central atom give the shape.
1.1.1.1 Different numbers of electron pairs result in different shapes
1.1.1.1.1 The lone pair affects the shape but isn't shown as a bond or anything in the structure
1.1.2 The e- pairs repel each other so arrange themselves around the central atom so they are as far apart of possible.
1.1.3 This arrangement minimises repulsion and the bonded atoms are held in a definate shape.
1.2 The shapes are shown in a 2D way but are 3D. Different signs can be used to show their 3D shape.
1.2.1
1.3 -3 bonded pairs - 120 - Trigonal planar e.g. BF3 -5 bonded pairs - 120,90 - triangular pyramidal e.g. PCl5 -6 bonded pairs - 90 - octahedral e.g. SF6
1.4 When giving a shape of an ion, you draw the dot and cross so the charge affects the first element in the compound but the rest is the same.
1.4.1 You tend to show the gained electron as a star and you also tend to put the charge on the diagram e.g. NO3 2-
2 Bond Angles
2.1 It is the angles between bonds.
2.2 It is affected by the number of bonds and lone pairs of the central atom.
2.2.1 e.g. there are 4 bonded pairs around the carbon atom which repel eachother.
2.2.1.1 4 bonded pairs give a bond angle of 109.5 and a tetrahedral shape.
2.3 The shape of the solid wedge comes out infront of the plane in the 3D structure. it is shaped like an isosceles triangle on its side and filled in.
2.4 The dotted wedge is when the bond goes behind the plane and is also shaped like a sideways triangle but is made out of lines.
2.5 A staright line just represents the bond.
2.6 Each lone pair repels the bonded pairs closer together meaning the bond angle decreases by 2.5 for each lone pair
2.6.1 e.g. NH3 has 3 bonded pairs and 1 lone pair. the bond angle is 107.
2.6.2 The lone pair affects the shape but isn't shown as a bond or anything in the structure.
2.6.3 H20 forms a non-linear molecule with 2 pairs of lone pairs. This reduced the bond angle from 109.5 to 104. The lone pairs are what make it non-linear.
2.6.4 Lone pairs repel more strongly than bonded pairs as they are closer to the central atom and occupy more space.
2.6.4.1 Bonded pair to bonded pair<bonded pair to lone pair<lone pair to lone pair.
2.6.4.2 Increasing repulsion
2.7 3 bonded pairs without lone pairs bond angle= 120
2.8 If e.g. 2 double bonds are around the central atom, they are called bonded regions rather than onded pairs.
2.8.1 The 2 bonded regions repel eachother as far as possible giving a bond angle of 180 and a linear structure. e.g. co2.
3 How to Predict the shape and Bond angle: 1. Draw a dot and cross diagram
3.1 2. Count up the number of regions of electro denisty around the central atom -a bonded pair -a lone pair -a double bond -a triple bond. These all count as regions.
3.1.1 3.Decide on the bond angle: -6 regions is 90 -5 regions is 120,90 -4 regions is 109.5 -3 regions is 120 -2 regions is 180 each will decrease by 2.5 when a lone pair is present.
3.1.1.1 4. Decide on the shape (Remember the lone pairs aren't seen/shown:
3.1.1.1.1 2 regions = linear
3.1.1.1.2 3 regions: 3 bonded pairs = trigonal planar 2 bonded pairs + a lone pairs = non-linear
3.1.1.1.3 4 regions: 4 bonded pairs = tetrahedral 3 bonded pairs + a lone pair = triangular pyramidal 2 bonded pairs + 2 lone pairs = non-linear.
3.1.1.1.4 5 regions = trigonal bipyramidal
3.1.1.1.5 6 regions = octaedral
4 Simple Molecular Forces
4.1 When solid, simple molecules form a simple molecular lattice - a regular structure where the molecules are held close together by weak intermolecular forces. The covalent bonds are very strong.
4.1.1 When heat energy is given to a molecule, the intermolecular forces break down NOT THE COVALENT BONDS.
4.1.1.1 Weak intermolecular forces require less energy to break them.
4.2 Simple molecules: They all have strong covalent bonding between atoms which aren't broke n by melting or boiling. They also have weak IMFs whether they are London forces or permanent dipole-dipoles. As they are weak, melting and boiling points are low.
4.2.1 Solubility: Non-polar molecules are soluble in non- polar solvents e.g.F2,CL2.IMFs between the molecules in the simple molecular lattice break and new IMFs form between the molecules and the solvent molecules. Non-polar molecules are insoluble in polar solvents (water) as IMFs between the H20 molecules are too strong to be broken by non-polar molecules.
4.2.1.1 H20 is a polar solvent so many ionic compound can dissolve in it e.g. NaCl.
4.2.1.1.1 The H20 molecules are attracted to the Na+ and Cl- ions. The ionic lattice breaks down as it dissolves. When the ions are in solutions, the ions are surrounded by water molecules. The ions then become hydrated ions.
4.3 Polar Molecules
4.3.1 They are soluble in polar solvents as the polar solute molecules and the polar solvent molecules attract each other.
4.3.1.1 A permanent dipole-dipole interaction is formed between the Cl and H doe to polar charges.
4.4 They don't conduct electricity as there are no mobile charged particles(no ions or delocalised electrons) so molecules don't conduct electricity.
4.5 Intermolecular Forces (IMFs)
4.5.1 Covalent and ionic bonds are very strong but when heated, it is the intermoleculr forces that break.
4.5.2 Intermolecular forces are weak attractive forces vetween dipoles of different molecules.
4.5.2.1 They effect the physical properties. Covalent bonds determine the chemical reactions.
4.5.3 There are 3 types of intermolecular forces
4.5.3.1 Induced dipole-dipole interaction: These exist between all molecules and in inert gas atoms whether they are polar or non-polar. They can also be called london forces or dispersion forces. They occur due to an instantaneous dipole which induces neighbouring dipoles causing them to attract eachother. They are only temporary and disappear or reappear in a different direction.
4.5.3.1.1 The strength of these forces depend on the circulating electrons.
4.5.3.1.1.1 The more electrons in a molecule: The larger the instantaneous and induced dipoles, the greater the induced dipole-dipole interactions, the stronger the attractive forces between the molecules, the more energy needed to break the IMFs, the higher the MP and BP.
4.5.3.2 Permanent dipole-dipole interactions: Molecules with polar bonds that have permanent dipoles which give rise rise to permanent dipole-dipole interactions between the molecules.
4.5.3.2.1 These are the bonds that break and require the extra energy to break.
4.5.3.3 Hydrogen Bonding
4.5.3.3.1 This is a special type of permanent dipole-dipole interactions between molecules containing: a very electronegative atom with a lone pair of electrons, either O,N or F. Bonded to a H atom so either a O-H, N-H, F-H bond. The bond forms between a lone pair of eectrons on one molecule and the H atom on a different one.
4.5.3.3.1.1 It is therefore the strongest type of intermolecular attraction.
4.5.3.3.2 It is shown with a dotted line:
4.5.3.3.2.1 The bonding gives water anomalous properties e.g. ice is less dense than liquid water as the H bonds hold the H20 molecules apart in an open lattice and in ice there are 4 H bonds in a molecule.
4.5.3.3.2.1.1 This means ice floats on water. On melting some H bonds break and the strcuture collapses.
4.5.3.3.2.1.2 Water has quite high MP+BP. If the only IMf is london forces, the BP would be -75 without the H bonds meaning there would be no life on earth. The H bonds water does have are however strong and require lots of energy to break them so the actual BP is much higher at 100.
4.5.3.3.2.1.3 It also has a high viscosity (thick) which is the ability of molevules to move past each other. It also has a high surface tension.
4.5.3.4 Simple Covalent Bonding
4.5.3.4.1 They are made up of simple molecules. They are particles with a definate number of atoms with a definate molecular fomrmula e.g. H2
4.5.4 Strongest to weakest: Induced dipole-dipole interaction, permanent dipole-dipole interaction, hydrogen bonding, single covalent bonfing.
5 Electronegativity
5.1 The bonded pair is evenly shared when the shared pair electrons are attracted to the nuclei of the same elemnt.
5.1.1 The bonded pair will be shared unevenly when: - The nuclear charges are different -The atoms are different sizes - The shared pair are closer to one nucleus that another.
5.1.1.1 The result will be a polar covalent bond.
5.1.1.2 The attraction of a bonded atom for the pair of electrons in a covalent bond is called the electronegativity.
5.1.1.2.1 The higher the attraction the higher the value of electronegativity (0-4), there are no units.
5.1.1.2.1.1 In the periodic table, (across to the right form the left and upwards) the nuclear charge increases (no. of protons). The atomic radius decreases (as electrons are pulled closer in by increases in the number of protons.)
5.1.1.2.1.2 Non- metals are most electronegative e.g. Cl
5.1.1.2.1.2.1 Group 1 metals are least electronegative e.g. K
5.2 The difference in electronegativity of a compound shows if its ionic, covalent or polar.
5.2.1 The greater the electronegativity difference between the bonded atoms, the more ionic the bond will be.
5.2.1.1 Covalent = 0 ,Polar covalent = 0<x>1.8 Ionic = >1.8
5.2.1.1.1 Covalent - The electrons are equally shared.
5.2.1.1.2 Polar covalent - The electrons are unequally shared.
5.2.1.1.3 Ionic - The electrons are transferred.
5.2.1.1.4 Polar bonds bonded electron pair are unequally shared as the atoms have different electronegativities so the bond is POLARISED.
5.2.1.1.4.1 e.g. HCl - H has a small amount of positive charge and Cl has a small bit of negative which is shown with a delta positive/negative.
5.2.1.1.4.1.1 The charge separation is called a dipole. This dipole doesn't change so is called a permanent dipole. This is now a polar molecule.
5.2.1.1.4.1.1.1 The direction of the dipole is positive to negative.
5.2.1.1.4.1.1.1.1 e.g. H20. It has 2 permanent dipoles and as H20 is a symmetrical, non-linear molecule, the overall one is in the middle. this makes water a very polar molecule.
5.2.1.1.4.1.1.2 Polar molecules with one permanent dipole. For molecules with more than 2 atoms, there may be polar bonds. The shape fo the molecule determines if the dipoles can add together to produce a larger dipole over the whole molecule o0r if they cancel eachtoher outif they are in opposite directions.
5.2.1.1.4.1.1.3 CO2 has 2 permanent dipoles but the shape is linear so they cancel each other out sot theres no overall dipole meaning the molecule is non-polar.
5.2.1.2 Non- polar bonds are where shared electrons are equally shared between bonded atoms (covalent) so have little or no electronegativity.
5.2.1.2.1 Non- polar bonds produce non- polar solvents.
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