Week 1 Lecture Notes Review: Jan. 9 – Jan. 13
Key Terms
Intramolecular Forces – the forces within single molecules that are large responsible for chemical properties
Intermolecular Forces – the forces between 2 or more molecules that are responsible for keeping the molecules together (largely responsible for physical properties)
Ideal gas – a gas in which all collisions between molecules are elastic and all atoms take up no volume
Intermolecular Forces:
Ion-dipole – the strongest IMF. Defined as the interaction of an ion (either positive or negative) with the oppositely charged area of the dipole of a polar molecule.
Hydrogen Bonding – the interaction of a hydrogen which is itself attached to a F, O or N with another F, O or N
Dipole-dipole (or Dipolar) – the interaction of the positive pole of one polar molecule with the negative pole of another.
Dispersion (or London) – the weakest IMF. Defined as the negative end of a distorted electron cloud interacting with the positive end of another distorted electron cloud.
Boiling point – the point at which a substance changes phases from a liquid to a gas (or vapor). Occurs when vapor pressure equals external pressure.
Normal boiling point – the boiling point at 1 atm.
Melting point – the point at which a substance changes phases from a solid to a liquid.
Freezing point – the point at which a substance changes phases from a liquid to a solid.
Cohesive forces – forces of attraction between molecules in the same phase
Adhesive forces – the attractive force between a liquid and the sides of it’s container.
Meniscus – the curvy surface on top of a liquid in a cylinder
Surface Tension – the resistance of a liquid to an increase in surface area
Capillary Action – the upward movement of a liquid in a narrow tube against the force of gravity
Viscosity – the resistance a liquid has to flow
Vapor Pressure – the partial pressure of a particular vapor in equilibrium with the condensed phase (or the amount of gas sitting on top of a sample of liquid)
Dynamic equilibrium – the point when the rate of a forward reaction exactly equals the rate of a back reaction.
Bonding in solids:
Ionic – metal and nonmetal (electron transfer)
Covalent – nonmetal and nonmetal (electron sharing)
Metallic – metal with metal or IMFs (electron pooling)
Types of Solids:
Molecular Solids – discrete molecules in an ordered array (no actual chemical bonds present)
Network Solids – atoms covalently bonded in an infinite array
Ionic Solids – alternating positive and negative charges set up in strict arrays
Metallic Solids – bonding between metals that share an electron sea
Malleability – the ability to be hit by a hammer and not shatter.
Notes
-The IMFs in a gas are very weak, which is why the molecules are far apart and can move so fast. They are stronger in a liquid, which causes the molecules to be close together, but still able to move. In a solid, they are strong enough to both pull the molecules as close together as possible and restrict movement.
-Very few gases behave ideally. This is because the collisions between the molecules do mean something and are affected by IMFs and that the volume of the molecules take up some amount of space. If the IMFs of a sample are playing a big role if the pressure of the sample is less than the pressure of the ideal gas. If the pressure is greater, then the volume of the molecules is playing a role (see graph below). To determine if a gas is behaving ideally, you have to do work with both the Ideal gas equation and the van der Waals equation (see below). (Know how to calculate this and how whatever answer you may get deals with the graph below).
-Gases exhibit the least ideal behavior at high pressure and low temperature.
-The van der Waals equation is a version of the Ideal Gas equation that corrects for the molecular attractions (n2a/V2) and molecular volumes (nb) (see below).
-IMFs are not actually bonds. They are long-range electrostatic attractions. This is important to keep in mind because they are broken much more easily than a bond would be.
-Ion-diople forces are responsible for the dissolving of salt in water. The Na+ is attracted to the negative O on the water and the Cl- is attracted to the positive H on the water. Any ion will be attracted to a dipole in a similar manner. This interaction occurs until the solid is used up.
-Please remember that not everything that has a F, O or N can hydrogen bond with itself. One of those three has to have an H on it (HF vs. CH3F, for instance). However, as long as there is a molecule present that has an H on an F, O or N, then any available F,O or N should be able to H bond to it (HF can H-bond to water, but so can CH3F or CH3OCH3). Just remember that you have two rules (1. H on F, O or N and 2. H must bond to another F, O or N), and as long as both of them are satisfied, life is good.
-If something can hydrogen bond, don’t write both dipole-dipole and H-bonding down on the exam. H-bonding is a very specific type of dipole interaction, so you would actually just be redundant.
-Absolutely every sample of everything has dispersion forces. If any other force is present, though, it will win because these are super weak. They are the only force that is present in any non-polar or mono-atomic sample. These occur and disappear super fast (which is part of the reason they are so weak).
-IMFs can affect physical properties. Take boiling for instance. You have to add some amount of energy to get most samples to boil. The amount of energy necessary is dependant upon the strength of the forces holding the sample together. So, boiling point and melting point would both increase with increasing IMFs. Freezing point would decrease with increasing IMFs since it is the opposite phase change to melting.
-If you’re asked to rank a bunch of substances in terms of boiling point or some other property, the first thing you should do is determine the intermolecular forces. Next, you should pick depending on either size or shape.
Ex. A – if you have to pick between H2O and H2S, then you would pick H2O because H-bonding is a stronger force than dipole-dipole.
Ex. B – if you have C2H6 or C5H12, they both only have dispersion forces, but C5H12 is both longer so it has more surface area, and weighs more, so it has more electrons.
Ex. C – if you have n-pentane or 2-methyl butane, both only have dispersion forces and both have the formula C5H12. However, one is long and one is branched (draw the structures out if you need to), so in this case only surface area would determine the winner.
-Water forms droplets because of the cohesive forces within water. They are trying to get the molecules as close together as possible to reduce the surface area of the molecules. The surface tension also has to do with the formation of droplets.
-If a liquid is in a cylinder that has forces similar to the forces present in it (say H2O in SiO2 – both have a dipole), then the liquid will try to interact with the cylinder as much as possible and will climb up the sides, making the meniscus concave. If the forces aren’t similar (H2O and Pt – H2O has a dipole and Pt has a perfectly distributed electron sea) – then the liquid will be repelled by the sides of the cylinder and you’ll end up with a convex meniscus.
-the stronger the forces, the stronger the surface tension, capillary action and viscosity will be and the weaker the vapor pressure will be.
-Vapor Pressure is inversely proportional to boiling point. If something boils at a low temperature, that means that we won’t have to add as much energy in order to get enough of the sample to be a vapor that the vapor pressure equals the external pressure and the sample boils (yes, that’s a run-on sentence). So, high boiling samples will have a low vapor pressure (or a low number of moles already a gas) and low boiling samples will have a high vapor pressure (or a high number of moles already a gas).
-When an equilibrium is reached, then there is no net change in either reactant or product. This doesn’t mean that there is no change in the system, it just means that the change is one direction is equal to the change in the opposite direction (or if you make something on one side, something on the other side is being made at the same time.)
-In a molecular solid, the molecules are put in their place by IMFs. Generally there is one best way for atoms or molecules to line up to form a solid, and the IMFs will maximize this attraction. Network solids are similar, but tend to have higher melting points that molecular solids.
-Ionic solids are held together by electrostatic forces. This is why the (+) and (-) charges alternate. If you hit an ionic solid with a hammer, it will shatter, because you’re moving the charges around and you’re getting the (+) one to close to another (+) one and they end up repelling and shattering you crystal. Metallic solids won’t break if you hit them with a hammer because there is an evenly shared electron sea, and there isn’t much of a way to get like charges too close together. Both of these have high melting points, but ionic solids are much more brittle.
Equations
Ideal Gas Equation:
PV = nRT
van der Waals Equation:
R = 0.08206 L*atm/mol*K
T = Temp in K
n = moles
P = Pressure in atm
V = Volume in L
a = a constant
b = a constant