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Let's start with physical properties of alcohols. And so we're going to compare, in this case, alcohols to alkanes. In this alkane on the left here, two carbons, so this is, of course, ethane. On the right, if we take off on those hydrogens and replace it with an OH, we of course have ethanol right here. So let's start with the boiling point. So the boiling point of ethane is approximately negative 89 degrees Celsius. And since room temperature's somewhere around 20 to 25 degrees Celsius, we're-- At room temperature, we are much higher than the boiling point of ethane which means it's already boiled. It's already turned into a gas. So at room temperature and room pressure, ethane is a gas. Ethanol however, has a much higher boiling point somewhere around 78 degrees Celsius. And once again, since room temperature is somewhere around 20 to 25, the boiling point of ethanol is much higher than room temperature. So at room temperature and pressure, ethanol is a liquid. It hasn't boiled, yet. And these a large differences and boiling points between these two molecules can be attributed to the intermolecular forces that are present. So if two molecules of ethane are interacting, really the only intermolecular forces holding those molecules together would be London dispersion forces which are the weakest of the intermolecular forces. So it's relatively easy to pull part two ethane molecules. And that accounts for the very low boiling point, doesn't take a lot of energy to pull them apart. So it's easy for it to turn into a gas. Ethanol, however, has a much higher boiling point which means it's much harder to pull those molecules apart. It takes more energy. So let's look at why ethanol such a higher boiling point. So if I show two ethanol molecules interacting-- So here's one ethanol molecule. And let's go ahead and draw another ethanol molecule right here. And if I think about the oxygen-hydrogen bond, I know that's a polarized covalent bond. I know that there's a large difference in electronegativity between the oxygen and the hydrogen. Oxygen's much more electronegative which means the electrons in the bond between oxygen and hydrogen are going to be much closer to the oxygen atom, giving the oxygen atom a partial negative charge. So these electrons in this bond between oxygen and hydrogen are going to move away from the hydrogens. So hydrogen is going to lose a little bit of electron density, leaving it relatively positive. So we give it partial positive charge. It's the same thing for the other ethanol molecule, partially negative oxygen, partially positive hydrogen. And we know that opposite charges attract. So the partially positive hydrogen is attracted to the partially negative oxygen. And so there's a strong intermolecular forces that holds those two molecules together and then, of course, is hydrogen bonding. So there's some hydrogen bonding. So there's hydrogen bonding between alcohol molecules. So hydrogen bonding. And since hydrogen bonding is the strongest intermolecular force, it's relatively difficult to pull those molecules apart. It takes a lot of energy, takes a lot of heat. And that's why the boiling point of ethanol is so much higher than the boiling point of ethane and also accounts for the state of matter. What about solubility? So is ethanol soluble in water? And of course, it is. And the reason why is hydrogen bonding, once again. So if we draw a water molecule in here, I know that the water molecule is polarized in the same way that the alcohol molecule is. So the hydrogen is partially positive, and the oxygen, right over here, is partially negative. And so once again, opposite charges attract. The hydrogen is attracted to this oxygen here. And so because of hydrogen bonding, there's interaction between the water molecule in between the alcohol molecule. So the water molecule is polar. So if you want to think about it in terms of polarity, because of the difference in electronegativity, water is a polar molecule. Ethanol is a polar molecule. And like dissolves like, so these two molecules will be soluble in each other. So if I look at the structure of ethanol, the reason why it is soluble in water is because of this portion of the molecule, this hydroxyl group, this OH. It's the differences in electronegativity that allow the hydrogen bonding. So this portion of the molecule is the polar portion of the molecule. And this portion of molecules the part that loves water which is why it is soluble. So if it loves water, we say it's hydrophilic, hydro meaning water, phil meaning love, so hydrophilic. Whereas this portion over here on the left, this is more of an alkane-type environment, a nonpolar type of environment. So this part of the molecule scared of water. So it's hydrophobic. So we have the hydrophobic portion of our alcohol molecule. And we have the hydrophilic portion of the alcohol molecules. Now, we know that like dissolves like, so nonpolar will not dissolve in polar. But as long as we have, in a relatively small number of carbon atoms in our alcule group, the OH group is enough-- is polar enough for the alcohol to be soluble in water. Now if you have a large number of carbon atoms, your molecule's more nonpolar than polar. And so alcohols will stop being soluble in water if they have a lot of carbon atoms on them. So let's look at, now, the preparation of alkoxides. So let's look at an alcohol. So here we have our alcohol. And if we react our alcohol with a strong base-- So we'll give it a lone pair of electrons, a negative 1 formal charge. So we have a strong base here. And our strong base is going to take this proton and leave these electrons behind on this oxygen. So now we have an oxygen that used to have two lone pairs of electrons and now has three lone pairs. That gives it a negative 1 formal charge. And the base is going to form a bond with that proton like that. So this is an acid base reaction. So if we react an alcohol with a strong base-- So this is a strong base here-- we're going to form the conjugate base to an alcohol which is called an alkoxides. So this is an alkoxides ion right here. So it's a chemical property of alcohols. They are acidic if you use a strong enough base. And the conjugate base to an alcohol is called alkoxides. Let's look at an example. So let's take ethanol. So here I have my ethanol molecule. And we'll react that with a strong base something like sodium hydride, so NaH. So Na+ and H with 2. Hydrogen with two electrons around it which makes it a negatively charged ion, so that's called the hydride anion. So we have the basic portion, the negatively charged hydrogen. It's going to function as a base. It's going to take these two electrons. It's going to take that proton right there. So the acidic proton on alcohols is the one on the oxygen. And the electrons in here are going to kick off onto our oxygen like that. So we're going to get for our product an alkoxide with three lone pairs of electrons around it, giving it a negative 1 formal charge. The sodium is floating around positively charged. So it's going to electrostatically, ionically, interact with our alkoxide anion. And the hydride anion picked up a proton. So those two hydrogens combine to form hydrogen gas which will, of course, bubble out of your solution. So the formation of hydrogen gas will be observed this reaction. And this is how you form an alkoxide. This molecule is called sodium ethoxide. And so we have sodium ethoxide over here on the right which is a relatively strong base that is used a lot of organic chemistry reactions. And let's see we used a strong base to form it. We used sodium hydride over here to form that molecule from ethanol. So there's another way to form alkoxides. So let's take a look at a generic way to form alkoxides from Group One, Alkali Metals. So here we have our alcohol like that. And if we react our alcohol with a Group One metal-- So an alkali metal, those all have one valence electron being in Group One on the Periodic Table, so something like lithium or sodium or potassium. We are going to form an alkoxide. So we're going to form-- Let's see, three lone pairs of electrons, a negative 1 formal charge. And the mechanism, the metal is going to donate its one valence electron, leaving it with a plus-1 charge. So it's going to interact with your alkoxide like that. And this also releases hydrogen gas like that. So that's your general reaction. Let's look at an example where we'll react-- Let's use cyclohexanol. So we're going to react cyclohexanol with sodium. So let's actually-- Let's go ahead and redraw that cyclohexanol molecule here because I want to show a little bit of the mechanisms. So let's go ahead and draw it like that. And put our lone pairs of electrons on the oxygen. So sodium metal has one valence electron like that. So if we think about what happens, sodium is-- Sodium will donate it's one valence electron very easily because it will then have the stable electron configuration of a noble gas. So the first step of the mechanism is donation of this one valence electron. So we're going to show the movement of one electron. So we're going to use a half-headed arrow. And then the two electrons in the bond between oxygen and hydrogen, we're going to use a two-headed arrow to show the movement of those two electrons off onto that oxygen. So let's go ahead and draw what we have now. We have our cyclohexane ring. And we now have three lone pairs of electrons around my oxygen which makes it negatively charged. And the sodium donated its one valence electron. So now it has a plus-1 charge. So it's going to interact with that negatively charged oxygen. And what happened to the hydrogen? That hydrogen there is going to pick up one electron. So now we have hydrogen with one electron around it. That is extremely reactive. Hydrogen prefers to have two electrons around it. So when two of those hydrogen atoms get close to each other, they're going to, of course, react and share their electrons to form hydrogen gas. So I could draw it like that. So that's where those two electrons are. The two electrons are here. So each-- One from one hydrogen and one from the other hydrogen. So hydrogen gas is going to be produced in this reaction as well. So that's an overview of physical properties and also the preparation of alkoxide anions.