In this module we will be looking at various types of acids including binary an oxyacidss. Our objective is to be able to identify different types of acids. As well as to look at periodic trends that determine the relative strength have those acids. The first thing we want to distinguish between are binary verses oxyacids. Binary acids are those that contain 2 elements. Not necessarily to atoms. We can have acids that have to different elements but have three or more atoms. Oxyacids are those that contain an oxygen atom and those that form for many of the polyatomic ions that we learned earlier in our chemistry studies. HClO_4 as an example. There are many different examples are these oxyacids including many different options with chlorine and oxygen. We can also have HClO_3 HClO_2 and HClO. When we look at our binary acids we see a periodic trend that makes it easy to rank these acids in order of increasing strength or increasing acidity. As I go from the top to the bottom I see an increase in acidity due to a decrease in the bond a strength. Because that bond gets weaker it is easier for those hydrogens to come off and therefore we will get more dissociation. So HF is a weak acid, HCl, HBr ,and HI are all considered strong acids but have those three HI is considered to be the strongest. When I look at the trend from left to right what I see is that as I go from water to HF from H_2S to Cl and so on, I see it increase in the electronegativity my anion. Because it that increase in electronegativity I also see an increase in the acidity. Because if these two periodic trends I can generally ranks a set of acids in order increasing or decreasing acid strength. If we look at the bond energies for a few of these binary acids we see that as we go from HF to HCl tp HBr the bond energy decreases. The bond is getting weaker and the acid is becoming stronger. Because it easier for the H+ two separate from the anion. When we look at oxyacids we have two issues to worry about. We have to look at the identity of the element in the act oxyacid other than the hydrogen and oxygen. Many times this is actually a halogen so we're going to focus on those here for example. The other issue we have to deal with for oxyacids is the number of oxygen. What happens when we look at something like HClO_4 HCl_3 and so on and so we have to separate these two issues to determine the relative acid strength. Here I'm looking at HOI, HOBr, and HOCl. What I notice is that as I go from iodine bromine to chlorine my electronegativity is increasing and when my electronegativity increases my acid strength also increases. We can tell that by looking at our K_a values, 10 ^ -11 10 ^ -9, and 10 ^ -8. The greater the K_a value the stronger the acid the more \ionization that we will see from our molecule. If I look at a generic molecule HOX, the increased electronegativity of X will pull electron density towards the X and therefore shift the electron density of the molecule such that a weakens this bond here and therefore will be easier for that hydrogen to be removed. Now we can look at what happens when I increase the number of Oxygen. Oxyacids increase strength with the number of oxygens. Perchloric or HClO_4 is one of our six strong acid that we have memorized. We know that that is a strong acids so its K_a value will be much much greater than 1. When I look at HClO_3 we have a K_a value of approximately 1 for HClO_2, 1.1 x 10 ^-2 and HClO only 2.9 times 10 ^-8, much weaker acid. When we compare these molecules for example HOCl versus HClO_2 where we have two oxygens. We are looking at the strength of the bond between the hydrogen and the oxygen. As we go from HClO to HClO_2 the strength this bond is weakened. It is easier to remove the H+ and therefore we have a stronger acid. Every time we add an additional oxygen into our structure we see the same trend continuing. So let's summarize what we know about acid strength. For binary acids those with just two elements regardless of how many atoms there are,
we see the same trend continuing.
So let's summarize what we know about acid strength. For binary acids those with just two elements regardless of how many atoms there are, we see the increases as we go to the right and down the column. For our oxy-anions we have to trends to look at. We have HOX where we have the same number of oxygen in each but the identity of X changed and that increases with the increasing electronegativity of X. Typically we're looking at this with the halogens other we can do it with other elements as well. Then we can look at what happens as we increase the number of oxygen atoms. So we have HAO where Y is increasing. And as Y gets bigger our acid strength also increases. So let's look at an example which is the following is the strongest acid? Remember our periodic trend is that for oxyacids increasing the number of oxygens will increase the acid strength. The hydrogen oxygen bond gets weaker in therefore it is easier to remove that hydrogen. Now let's look at another example. Which of the following is the strongest acid. Now we're looking at our trend for binary acids. Remember when we go down the problem we see increasing acidic strength. Now remember that we look at three separate trends. It very difficult without having quantitative data such as K_a values to compare say HF and HClO versus some other acid. So when we look at these periodic trends within the acid we want to look at a analogis group where we have the same general structure we are just changing one specific property. Either the identity of the one of the elements or the number of oxygen atoms. Going back to what we look at earlier in this unit about Bronsted-Lowry acids and bases. We talked about transferring protons. Where we have proton donors and proton acceptors behaving as either an acid or a base. Now what we wanna look at as a slightly different definition of acids and bases that applies to different substances where proton donor acceptor model doesn't really do a good job explaining it. And that is the Lewis acid-base theory. So when we look at electron pair transfer instead of proton transfer that we did in the Bronsted-Lowry acids and bases we have to define what we mean by an acid or base. Here for an acid we mean an electron pair acceptor. Compared to our Bronsted-Lowry definition
Here for an acid we mean an electron
pair acceptor. Compared to our Bronsted-Lowry definition where we had a proton donor and for our Lewis base we have an electron pair donor where in Bronsted-Lowry we had a proton acceptor. So what we wanna look at for a electron pair acceptor if some place for those electrons to go. In a base for an electron pair donor we need to have some lone pair some unbonded pair of electrons they can actually donate those electrons. Once this transfers occurs we have a coordinate covalent bond. Lets look at Lewis acids and bases. Here we have a reaction representing that Lewis acid-base reaction. Note, that we have a Lewis base that is going to be acting as an electron pair donor and a Lewis acid which is going to be acting as an electron pair acceptor. One thing that we look for in a Lewis acid is to look for a hole, or some place for those electrons to go. Here we see that aluminum doesn't have any electrons in this fourth spot and as a result it can accept electrons there which is what makes it a Lewis acid. If we look at the chlorine atom here . We know that we have three lone pairs of electrons around that chlorine that can serve as electron a pair donors. Once they react we form a complex that shows a coordinate covalent bond between the chlorine and the aluminum. This reaction could continue until we get to a Lewis acid Where we have a positively charged species that now can behave as an electron pair acceptor and a negative species that they can now serve as an electron pair donor. So much in the same way that we talked about conjugate acids and bases with Bronsted-Lowry acids and bases, we see a similar approach here looking at Lewis acids and bases. Let's look at an example. Which is the following could act as a Louis base. Remember that Lewis base are electron pair donors. For this question we need to look at the Lewis structures are electron pair donors. For this question we need to look at the Lewis structures which is the same Lewis, but we look at the Lewis structures and what we see is that PCl_3 has this non-binding pair of electrons. If I look at aluminum note that I have no non-binding electrons. It has nothing to contribute. So it only can act as a Lewis acid as an electron pair acceptor.
note that I have no non-binding
electrons. It has nothing to contribute. So it only can act as a Lewis acid as an electron pair acceptor. When I look at AlCl_3 I also see this kind of unoccupied space where this can act as an electron pair acceptor. Making it a Louis acid. where this can act as an electron pair acceptor. Making it a Louis acid. Now we completed our unit on acid-base equilibria. There are other issues surrounding this which we further explored in the unit on
Now we completed our unit on acid-base
equilibria. There are other issues surrounding this which we further explored in the unit on aqueous equilibria which will include titrations, buffers and solubility equilibria.
aqueous equilibria which will include titrations, buffers and solubility equilibria.
