- Describe Liquids at a molecular level
- Describe the effects of energy application to Liquids
- Explain the effect of temperature on Vapor Pressure
- Contrast the stuctures of Metals and Salts(ionic network lattices)
- Describe the relationship between structure and properties in Metals and Salts
The Nature of Liquids
Liquid particles slide past one another and flow like gases. However, the difference between a liquid and a gas is that the particles of a liquid are more strongly attracted to each other than the particles in a gas. This attractive force between water molecules is called an intermolecular force. Because of intermolecular forces, liquid particles vibrate and spin in fixed positions while moving freely, but they do not have enough kinetic energy to overcome the intermolecular forces and become a gas. Kinetic energy is the energy an object has because of its motions. To overcome the intermolecular forces and become a gas, commonly called evaporation or vaporization, a liquid particle must have enough kinetic energy.
Liquids also have something called vapor pressure. Vapor pressure is basically the pressure exerted on the surface to propel itself up. Think about it, if you push off something really really hard, you go up much faster, than say, if you just give a feeble kick. This means that the higher the vapor pressure the liquid has, the easier it evaporates, and this may sometimes lead to sublimination. Sublimination is basically when there is virtually no liquid stage from the solid to gas phase. Iodine and "dry ice" are all examples of molecules with a high vapor pressure. To have a high vapor pressure, the intermolecular bonds have to be extremely weak too. This is because then it'll take less energy(heat) to break them and it'll evaporate faster. So basically, in summary, high vapor pressure=low intermolecular forces=low boiling point.
METALS VS SALT aka Metallic Compounds VS Ionic Compounds
So, lets start talking about metals first, shall we? Basically, metals are metallic compounds, which means that electrons are free-flowing throughout the entire structure. The reason? As you (should) know, metals are on the left side of the periodic table. This means that it has a low electronegativity, so it has a hard time keeping its electrons. So, since its a bunch of the same cations together, the bonds are extremely weak and are easily broken, so the electrons are always shared around, so that why they are free-flowing. There are some charactersitics special to metals because of these metallic bonds:
1. Dense-its denser than ionic compounds because the electrons are free-flowing, so when you "squish" it, the cations don't repel because the negative electrons can flow through little spaces in between the cations, forming a sort of cushion.
2. Malleable-Once again, this has to do with the free-flowing electrons. If you add pressure to the metal, instead of shattering, the electrons cushion the cations, so they just basically float away and mush their way away from the point. So thats why metals bend, while ionic compounds don't, which I'll explain later.
3. Conductive-So, guess why metals are conductive. If you don't get it in one guess....nothing'll happen...but once AGAIN, its because of the free-flowing electrons. If you send an electric shock through the metal, technically, the electricity is also a bunch of electrons. So, once they come in one way, the same amount of electrons goes out the other. So thats how they conduct electricity.