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Hybridisation and angle suggested by LINEAR arrangement?
- sp hybridisation: s + p -> 2 sp
- 180 degrees
Hybridisation and angle suggested by TRIGONAL PLANAR arrangement?
- sp2 hybridisation: s + 2 p -> 3 sp2
- 120 degrees
Hybridisation and angle suggested by TETRAHEDRAL arrangement?
- sp3 hybridisation: s + 3 p -> 4 sp3
- 109.5 degrees
Hybridisation and angle suggested by TRIGONAL BIPYRAMIDAL arrangement?
- sp3d hybridisation: s + 3 p + d -> 5 sp3d
- 120 degrees equatorial, 90 degrees axial
Hybridisation and angle suggested by OCTAHEDRAL arrangement?
- sp3d2 hybridisation: s + 3 p + 2 d -> 6 sp3d2
- 90 degrees
The relationship between quantum number n and the size of the orbital?
As n increases, the size of the orbital increases
The relationship between the quantum number l and the electron distribution of an orbital?
- l increases, the electron distribution of the orbital becomes more elaborate.
- nodal planes or radial nodes
In every orbital how many nodes, nodal planes and radial nodes are there?
- These electron distributions depend on the quantum numbers n and l
- For every orbital, there are n - 1 nodes, of which l are nodal planesthe remaining are radial (spherical) nodes
What is lattice energy?
Lattice energy is the amount of energy requires to break 1 mole of a compound into its constituent gaseous ions.
What does PAN stand for?
- peroxyacetyle nitrate
- produced when NO2 reacts with air contaminated with unburned hydrocarbons
What is bond energy? What are the three trends in bond energies?
- the amount of energy required to break a particular chemical bond
- 1. Bond energies increase as more leectrons are shared between the atoms. Shared electrons are the 'glue' of chemica lbonding, so sharing more electrons strengthens the bond.
- 2. Bond energies increase as the electronegativity difference between bonded atoms increases.
- 3. Bond energies decrease as bonds become longer. As atoms become larger, the electron density of a bond is spread over a wider region. This decreases the net attraction between the electrons and the nuclei. (N.B. electronegativity difference also changes, supporting the trend)
There are also numerous exceptions these three trends
What is the polarisability of a molecule referring to? What is its connection with intermolecular forces?
- The strength of dispersion forces depends upon the polarisability of a molecule - the ease with which the electron cloud of a molecule can be distorted.
- Large atoms with many electrons are easily polarised and produce higher dispersion forces. Hence dispersion forces increase with increased molar mass of a molecule.
What are the relative magnitude of forces (intramolecular & intermolecular)
- Intramolecular - force that holds the atoms or ions together INSIDE the molecule
- Intermolecular - force that acts BETWEEN molecules
- Intramolecular forces (ionic, covalent and metallic forces) are of similar strength - generally much stronger than van der Waals forces
- Of the intermolecular forces (dispersion or 'London forces', dipole-dipole and hydrogen bonds), hydrogen bonding is strongest (especially for small molecules - dominant effect in determining molecular properties such as m.p. and b.p.)
What is a dimer?
- The name given when two identical or very similar monomers join together held by bonds that can be strong or weak, covalent or intermolecular.
- E.g. Two units of acaetic acid forms a very stable dimer due to the presence of two hydrogen bonds between the pair of CH3COOH molecules.
What is the molar volume at STP?
- Standard Temperature and Pressure are 0 degrees Celcius and 1 atm.
- One mole of ANY GAS occupies 22.4 L at STP
- can use V = nRT/P to calculate the molar volume where n = 1 mol, R = 0.0821 L atm K-1 mol-1, T = 273 K, and P = 1 atm)
Under what conditions does the Ideal Gas Equation apply? What is the limitation for the Ideal Gas Equation?
- Under normal temepratures and pressures most real gases behave close to idea and obey the Ideal Gas Equation.
- At LOW TEMPERATURES (when intermolecular forces are important and become more stable) or HIGH PRESSURES (molecules are forced close together and attract one-another - real gas molecules stick together momentarily upon collision - molecular volume becomes a significant fraction of container volume), real gases deviate from the Ideal Gas Equation. Hence modifications are made to it to cover these situations.
TWO properties of an ideal gas
An ideal gas is an imaginary substance whose molecules have ZERO MOLECULAR VOLUME and NO ATTRACTIONS BETWEEN MOLECULES
What is the van der Waals Equation?
- A modification of the Ideal Gas Equation to include compensations for molecular attractions (low temp) and finite molecular volumes (high pressures) of real gases - fits the behaviour of real molecules better.
- where a and b are t he van der Waals constants - specific values for each gas.
- a constant is related to the no. of electrons in the gas molecule - influences magnitude of the intermolecular attractions
- b constant is related to the intrinsic volume of the gas molecule
What is the compression factor and its relation with ideal and real gases?
The term PV/nRT is called the compression factor
- For an ideal gas, PV = nRT and for one mole of gas, the compression factor (PV/nRT) is equal to 1.
- For real gases, pV does NOT EQUAL nRT, and so the value will be something different to 1. The compression factor (pV/nRT) varies as you change temperature and pressure of a real gas.
What causes non-ideal behaviour of gases?
With the compression factor, pV/nRT, the variables in the denominator can be determined quite accurately but it is not so true for pressure and volume.
With assumptions we make about ideal gases (zero molecular volume & no attractions between molecules) this is not true for real gases and these have an effect on both pressure and volume.
- THE VOLUME PROBLEM
- The kinetic molecular theory for ideal gases assumes the particles have negligent intrinsic volume. The distance between particles is so large compared with the particle size that particles are treated as point objects.
- For a real gas, this assumption is not true. The molecules themselves do take up a proportion of space in the container. The space in the container available for gas particles to move around in, is less than the measured volume of the container.
- The problem gets worse as more gas is compressed into the container - as the gas gets more compressed, the proportion of the total volume that the molecules themselves take up increases. So if the pressure is low, the volume taken up by the actual molecules IS insignificant compared with the total volume of the container.
- THE PRESSURE PROBLEMAnother assumption of the Kinetic Theory for ideal gases is there are no intermolecular forces between molecules. This is wrong for every real gas.
- If there weren't any intermolecular forces, it would be impossible to condense the gas into a liquid. Even helium, with the weakest of all intermolecular forces, can be turned into a liquid if the temperature is low enough.The effect of intermolecular forces For a gas molecule located in the middle of the gas mixture, there is no net effect. It will be attracted to some extent to all other molecules around it, but, on avg, those attractions will cancel each other out. Attractions from behind a molecular, tending to slow it down, will be cancelled by attractions from in front of it, tending to speed it up.
- The gas molecule in the centre here (marked in green) is experiencing intermolecular forces but it will just continue to move in the same direction as the same speed.
- However if the molecule is about to hit the wall of the container, there aren't as many gas molecules in front of it, and the net attraction is backwards into the body of the gas. The molecule will slow down just before it collides into the wall. If it slows down, it will not the wall as hard, and so exerts less pressure.
- Overall effect - measured pressure is less than it would if the gas was ideal. Hence if you put the measured pressure into the expression PV/nRT, the compression factor you calculate will be less than it would be if the gas was ideal. This is why, under some conditions, graphs of compression factors drop below the ideal value of 1. This effect is noticeable at lower temperatures.
- At lower temperatures, the molecules are moving more slowly on average. Any pull they experience back into the gas will have relatively more effect on a slow moving particle than a faster one. i.e. Intermolecular attractions are more dominant than the molceular volume effect. PV/RT ratio < 1.
- At higher temperatures, where the molecules are moving a lot faster due to great kinetic energy, any small pull back into the body of the gas doesn't do much and is hardly going to be noticeable on the molecule. At high temperatures, the effect of intermolecular forces is negligible and hence high temperatures gives a higher pressure for real gases than an ideal gas would - due to molecular volume contribution to pressure outweighing intermolecular forces. PV/RT ratio > 1.
Another effect concerning intermolecular forces - as pressure increases
, the molecules are forced more closely together - intermolecular forces become more important - as pressure increases, at first the value of the compression factor falls but it soon rises up again. Why? Because at this point, the effect of the size of the molecules starts to become more important - and as pressure increases even more, this effect becomes dominant.
For a real gas: effect of moderate temperature on pressure?
- Considerable intermolecular bonding during collisions of gas molecules
- Thus at any instant, the number of moles of particles if effectively less
- If V & T are kept constant, then P is proportional to n, so fewer 'n' means pressure of real gas tends to be lower than ideal.
- Hence intermolecular bonding factor prevails at low temp, so Preal < Pideal
For a real gas: effect of high temperature on pressure?
- gas molecules have too much kinetic energy to form effective intermolecular bonds during collisions of gas molecules.
- The pressure is quite high at higher temperatures, and intrinsic molecular volume is a significant fraction of the total volume of the container.
- Thus, the actual volume of free space avilable for molecules to move around in is somewhat less than the container volume
- n moles in a smaller V gives high P
- Hence intrinsic molecular volume contribution to pressure outweighs the weak intermolecular forces, so Preal > Pideal
What is surface tension?
The resistance of a liquid to increase its surface area
Energy process when a solute dissolves in a solvent?
- 1. solute-solute bonds need energy to break so solute particles can disperse and dissolve
- 2. solvent-solvent bonds need to be broken when a solute particle inserts between them
- 3. solute-solvent bond formation releases energy
If the total energy input of 1 & 2 > release of energy from 3
, then the solute is not very soluble