all right engineer so what we're going to do in this video is we're going to talk about acids and bases so we're going to go through first off before we get into any calculations because that's what we're going to spend a lot of time on kind of the key to chemistry at getting better at it is doing tons and tons of practice problems and so that's our goal what we're going to do here at ninine nered science is help you as much as we can with a lot of practice problems before we dive in let's go ahead and get some just Baseline uh introduction introductory material out of the way so first off let's go ahead and Define what an acid and what a base is all right so we're going to do that then we're going to talk about all the strong acids and strong base we're going to talk about conjugate acid conjugate base pH scale and we're going to talk about this KW K KB pkp KB all right so let's go ahead and dive in first with definition of an acid and a base in order for us to do this though we have to go through three different types of definitions all right so the first one we're going to talk about is Bron Lowry okay so Bron stead Lowry so Bron said Lowry acid and when we're talking about a bronze dead Lowry acid he said that when you take an acid and you put it into a solution an acid gives up protons so how does he Define an acid he defines an acid as a proton donator okay or we can even draw this like this he loves to give give up H+ ions because we're going to see this a lot okay whenever we're doing a lot of these practice problems so again what is a bron stad Lowry acid by definition a proton donator well if an acid is a proton donator someone has to be able to gain that or accept that proton who accepts the proton the Bron stead Lowry base so the Bron stead Lowry base is going to be the guy who is accepting the proton so he is a proton acceptor all right so that's pretty easy right so it's not too bad there so for example if I were to ask you if I were to draw right here let's say I did this h a uh plus water and it yields two products let's say one of them is a negative and the other one is h3o+ if we were to go based on this definition of Bron Lowry what happened to this ha it turned into a negatives so what did it do it gave up a proton so he must be by definition the Bron said Lowry acid and then what happened to this H2O he turned into h3o plus hydronium so he must have accepted that proton and if he accepted the proton that the ha gave he must be the base now while we're here let's just go ahead and get these terms out of the way since we're already here conjugate acid conjugate base let's go ahead and get those terms out of the way because we're already here so an a when he gives up the proton so he donated that proton to the H2O he turns into a different term now that we can call so he was an acid but now he's a base by definition if he's a base usually bases have negative charges right so they're really good at having negative charges that's kind of the way that we can see bases is they're negatively charged uh molecules or atoms so we're going to say that this is a base but it's another type of Base it's the conjugate base so this acid it gets converted into a negative which we're going to call the conjugate base and again why is he the conjugate base because the acid gives up a proton and turns into a conjugate base well if he turns into the conjugate base then this must be the acid why is he the acid because this water when he accepts that proton because he's the base he turns into an acid so he is the conjugate acid so if we were to combine the pairs this acid and this conjugate base this is the acid conjugate base pair right there and this right here is the base conjugate acid pair here okay so not too bad there right so that's that's just giving you an understanding of bronze stead Lowry what an acid and what a base is and then what is an acid what is its conjugate base what is a base and what is it conjugate acid all right let's go into the next next individual who came up with this another theory this guy is Lewis so Lewis came up with a theory of acids and bases but he looked at it in a different way bonstead looked at protons Lewis looked at electrons so Lewis said if an acid is donating a proton it has to be accepting electrons that's his definition he says by definition a Lewis acid is a molecule who does what it accepts electrons so he is a electron acceptor okay so while this guy is's donating protons he's also accepting electrons Okay cool so then a base what was a base doing a base was accepting a proton so he must be donating electrons so he is a electron donator all right not too bad here then right so this is if if we were to go based off this we could apply the same concept here with the Lewis acid and Lewis base there is two exceptions we're not going to talk about them so oranus he was the other guy that came up with this Theory here right and he said that an acid is a little bit more specific than just a proton donator so arenus said an acid is actually going to be a guy who gives up this guy has to give up a proton so he is a proton donator okay so he does give up these actual protons but more specifically he gives up these protons in an aquous solution so we'll say in water in an aquous solution so he even kind of expanded more on that he said an uranous acid is a proton donator in an aquous solution okay and the definition of a uranous base right so an acid is the one who's going to be able to give up or produce an H+ you know he's producing a proton in an aquous solution a base is just producing hydroxides so what is it doing it's producing or it's we could say even donating hydroxides so we could say hydroxide donator if you'd want to say that or it's producing hydroxides in an aquous solution okay so that's the definition of an uranous acid and base an acid is going to be one who gives up protons in an aquous solution whereas a base is going to be producing hydroxides in an aquous solution or donating those hydroxides in an aquous solution okay so now we have an understanding of what Bron lowy acid leis acid and base uranous acid and base and we even know what conjugate acid and conjugate base pairs are so let's go ahead and check those off so what do we do we did the definition of an acid and a base and we did what conjugate acids and conjugate bases are all right let's talk about what strong acids and strong bases are all right well first off how how would you define a strong acid a strong acid is defined as something that can deprotonate or disassociate completely so for example I'm going to write down all of our strong acids here so let's go ahead and put them right here strong acids these guys are going to be hydrochloric acid Hydro iodic acid Hydro bromic acid hno3 which is nitric acid hcl4 which is perchloric acid h2so4 which is sulfuric acid so how many do we have 1 2 3 four five six these are our six strong acids and again what does it mean that they're a strong acid that means if you guys if you put this into an aquous solution they're going to disassociate so it's a one Arrow movement if you notice I did not have an equilibrium Arrow where did I have an equilibrium Arrow here here I will not because all of the HCL will deprotonate into the form of H+ and cl negative and again a strong acid is defined as a proton donator so it's going to give up that proton and it completely disassociates in an aquous solution so with this one look like hi would give way to H+ and I negative h H BR would give way to H+ and BR negative H3 would give way to H+ and nitrate hcl4 would give way to H+ and cl4 negative h2so4 is different so if you look here h2so4 it has an H2 this is defined as a d pric acid meaning it can give up two protons so if we were to run this reaction he would deprotonate and we turn into h s so4 negative and then he would react even further and get converted into S so4 2 negative right so what we give up we should actually show the H+ in there so again it would react here and give you H+ and this guy then this would react even further and produce sulfate and H+ okay so these are six strong acids and they the reason why they're strong acids is because they disassociate completely all right let's do strong bases so strong bases same thing the reason they're called strong bases is because they disassociate completely when put in an aquous solution so what are this what are these strong bases these ones are easy cuz if you take all the group one and group two metals there is a couple exceptions but if you take almost all of those so for example we take sodium hydroxide we take lithium hydroxide we take potassium hydroxide we take rubidium hydroxide we could even keep going we could even do cesium okay cesium hydroxide we can't do francium though all right that's a big mofo and not a good one to do right all right now what's another one we could do group two metals so what are the group two metals the group two metals we could do uh magnesium isn't red necessarily soluble um calcium is too but it's kind of slightly soluble but we'll put that one in there CO2 um we could also do strontium hydroxide and we could even do barium hydroxide but not radium okay so in general these are our strong bases magnesium isn't really considered to be a strong base because it's not really soluble okay so if we were to kind of look at all of these guys here this would pretty much give us all of our strong bases and what are they strong why are they strong bases because they can give up or donate all of these hydroxides all of them will disassociate into solution okay so if you were to put all of these guys into a solution I'm only going to do one here if I were to take sodium hydroxide here put him into an aquous solution what would happen it would give up or completely disassociate into hydroxide ions okay so that's the definition of a strong base and that's the definition of a strong acid they completely disassociate it's not a reversible reaction One More relationship I want to develop here because this is really important when we're going through a lot of these calculations there is a a a relationship here and it says the stronger the acid whenever the reaction occurs the weaker it's conjugate base the weaker it's conjugate base and this would be the vice versa if you start off with a strong base the weaker it's conjugate acid it's either way let me give you an example here if I take something like HCL I react this in water what would I get I would get H+ which is a strong acid and I would get chloride well this by definition is a strong acid this right here is a very very weak conjugate base that's going to be important whenever we talk about the acidity of of uh and basicity of salts okay all right so again remember that relationship stronger the acid the weaker it's conjugate base and like I said you can do the same thing with the other one you could say stronger the base the weaker it's conjugate acid okay so now we have that relationship taken care of so what have we done so far we've talked about definitions of acids and bases with Bron said Lewis and oranus we talked about all the strong acids and strong bases and we tied in a relationship with conjugate acid conjugate bases and the relationship of strength all right let's get over here to the pH scale all right so let's go ahead and talk about pH scale so pH scale is a is basically a logarithmic scale here and we're going to talk about that in a little bit but just know that it's a scale that goes from zero all the way to 14 and what it does it helps to measure from zero all the way to seven measures how acidic the solution is and then going from Seven all the way up to 14 measures how basic a solution is now again what do I mean by acidic and what do I mean by basic so when we talk about acidic we're talking about specifically we're talking about how many protons we have or how many hydronium ions we have so again we talk about acidic environments we're talking about the concentration of H+ and we put these brackets in there to signify concentration so the more you go closer to zero the more acidic Your solution is so if you have a solution of that's uh in between one and zero that's that's pretty freaking acidic all right so when we talk about a acidity we're talking about H+ concentration when we're talking about basicity we're talking about hydroxide ion concentration so again if I were to say okay I got a solution that can have H+ is and I got solution that can have hydroxides how do I know specifically that it's acidic well that means that the H+ concentration in the acidic environment is greater than the concentration of the Hydrox oxide ions all right specifically in the acidic environment because the more H+ is the more acidic in the basic solution the basic solution we want it to be measured by hydroxide ion concentration so the greater the hydroxide ion concentration is than the h plus concentration that means that we're in a basic solution what if it's neutral though then these two are equal so then the H+ concentration is equal to the hydroxide ion concentration and this is an important relationship when we talk about KW here in a second okay before we do that I want to talk about another thing so I said before that this is a logarithmic scale right so in other words if I wanted to go from zero all the way to five most people would say okay how many uh what was the pH change well they'd be like oh it was five that's wrong and the reason why is each number you go up from 0 to one 1 to two uh 3 to four and then four to five is a 10 times change it's a tenfold change so for example if I'm going up here I'm actually going up what 10 * 10 time 10 * 10 time 10 okay so what is that 10 * 10 * 10 * 10 * 10 okay well that's going to be what that's going to be a 100,000 fold change that is a 100,000 th000 fold change so how much of a change was it it was a 100,000 times change so remember that whenever we're saying okay what's the pH change going from zero to five well each change is a tenfold increase so if I go 10 and 10 and 10 and 10 and 10 that is a fivefold increase but specifically by 10 * 10 * 10 * 10 right that's just 100,000 times change there same thing if I was going from uh 14 to 9 if I'm going all the way from 14 to 9 what am I doing it's changing by a 10-fold scale so 10 * 10 time 10 * 10 and then over here to this guy times 10 so same situation how many was that we went from 14 all the way to nine most people be like oh that's a five times difference no it's a 100,000 times difference because it's a 10 times each one okay now another thing when because it's this this uh base 10 if you will change they came up with a way to be able to calculate a specific number which we call this this pH right and so what they said is okay well we can technically figure out what the pH of a solution is since we have this scale now and what they said is we could actually say that pH is equal to we put a negative in there for uh specifically mathematical sake and we said it's a log base 10 relationship it's a logarithmic scale so pH is equal to the negative log of the H+ concentration okay and that's how we derive this formula to solve for pH we said it's based on this logarithmic scale which is a base 10 right 10 full change same thing that's pH so that would give you the H+ concentration P would give you the hydroxide ion concentration so it's the negative log of the hydroxide on concentration okay well that's cool because now I have this ability to calculate P I have the ability to calculate pH what if I add both of these two together so then I take and I add P plus pH that gives me a specific formula pH plus P wait isn't that the whole scale it is so it equals 14 and that's how we got a lot of these formulas right so what do we have so far we have equal negative log of the hydroxide ion phal negative log of the hydronium ion and then pH plus P = 14 in the same way if you know your algebra if I wanted to solve for hydrogen ion concentration and they gave me the pH well knowing my logarithmic relationship it's a base 10 so it's going to be 10 to the ne pH so if I wanted to solve for hydrogen ion concentration that formula would be H+ concentration is equal to 10 to the ne pH if I wanted to solve for hydroxide on concentration all I would do is I'd take 10 the Nega P so I can move all the way around this pH scale right so again how many formulas have we've been able to uh come up with here we came up with H+ equals 10 Negative PH pH plus P equal 14 pH equal negative log of the hydronium ion P equal negative log of the hydroxide ion and Hy oxide is equal to the 10 NE P one more thing I want to talk about here okay water has the ability to act as uh what's called amphoteric meaning Can it can act like an acid and a base and there's other molecules that can do that so when water actually disassociates right so let's say here's water water disassociates let's say it reacts with two two waters react and when it or we can say it auto ionizes when it Auto ionizes let's actually just do that say it's Auto ionizing right when it auto ionizes it produces an H+ and a hydroxide ion okay that's cool right because I I just kind of like I know what H+ is I know what hydroxide is and I know what water is we actually were able to calculate a rate at which this actually occurs with the rate at which water Auto ionizes and we said it's defined as KW and KW is equal to 1 * 10 to the4 so that's that rate that's theate and that takes a long time so that's a very you know it's a very very small small number right so very very very small number so with this what's happening is the auto ionization of water you have a specific rate at which this reaction occurs to produce H+ and hydroxide ions so we can actually say one more thing okay well I know that KW because I'm producing these guys is equal to the H+ concentration times the hydroxide ion concentration so at any point in time here I could rearrange this formula and can get three other formulas from this I could have this formula and I could have H+ equal KW over hydroxide or I could have hydroxide equal KW over H+ so there's a lot of different formulas I could drive from that but this gives me the relationship of KW all right now let's talk about one other thing here um and that's going to be talking about weak acids and weak bases very briefly and then we're going to go into a lot of calculations all right so we finished up talking about the pH scale we talked about P pH H+ hydroxide KW so what have we finished so far we've talked about the pH scale okay again we talked about the logarithmic scale H+ hydroxide we talked about uh acidity basicity we talked about um KW right so we talked about KW now we got to talk about Ka and KB PK and pkb so right now we've only been talking about strong acids and strong bases so we have to get just a baseline definition of what's what's up with these weak acids and weak bases so I'm going to give you an example first off let's just do a weak acid so an example of a weak acid could be I could pick acetic acid so I'm going to pick acetic acid I'm going to say hc2 h32 acetic acid is a weak acid why is it a weak acid what was a strong acid that means that the strong acids are defined as completely disassociating giving all of their protons up so if I had put in a 0.1 molar solution of hydrochloric acid all of that 0.1 molar will go over to forming H+ and cl if I put a 0.1 M solution of acetic acid into a solution not all of it not all that 0.1 molar will get converted over into to H+ and acetate why because not all of it dis Associates so now what we say is let's say that we take this reaction it's a double arrowed reaction so in other words this can occur it can actually be an equilibrium so whenever he gives up his H+ he'll give up his H+ and he gives up and he forms these two structures and again what would this be weak acid this would be its conjugate base so that's the weak acid conjugate base pair and if I did this in water water would be the base and H+ would be the conjugate acid all right so now because he didn't disassociate completely we have to have some type of number or rate if you will to determine how much of that H+ disassociates so what do we use for that that's where Ka comes into play so Ka is defined as the acid ionization constant in other words it gives us how much of this acetic acid dis Associates into a solution okay how much of it goes into H+ and now here's what we need to know out of this the greater the KA all right so the more that Ka the larger it is the larger the KA the stronger the acid okay in other words the bigger this is the more H+ are going to disassociate okay so if I were to do something like this a low Ka a low Ka implies a a weaker acid okay cool while I'm here let's just talk about PKA so PKA you can actually calculate it because PKA is the negative log of the KA okay so now if I said that a Ka if it was really really big and I wanted to try to be able to figure out PKA PKA is the actual opposite so now we say the lower the pka the stronger the acid the higher the KA the lower the pka the stronger the acid because it's going to give you this log amount and again we can actually solve for KA how 10 to the negative PKA so Ka is equal to 10 to the pka so you could solve for that relationship so again higher the KA stronger the acid lower the pka stronger the acid lower the KA the weaker the acid the higher the pka the weaker the acid okay now and again weaker acid would imply pH so you could say this would be a low PH this would be a higher pH now let's go ahead and talk about weak bases so if I talk about weak bases here let's say I use ammonia so I use ammonia so NH3 and I take NH3 and I react him with water what's the definition of a base a base accepts proton so in this case he has to accept a proton from who water so water gives up a proton and puts it onto ammonia and converts him into ammonium which is a weak acid so there's our base conjugate acid pair and then H2O gives up a hydrogen so it turns into hydroxide so there is your acid conjugate base base pair now same thing ammonia is not a strong base so it doesn't accept protons as rily and it doesn't produce hydroxides as readily as something like sodium hydroxide so we have to get another type of constant if you will to determine how much of this hydroxide is being produced or how much hydrogens is he accepting so we give that a k b okay and KB is the base ionization constant okay and it's the same thing we did up here with the KA so the higher the KB but remember we got to be careful now the higher the KB the stronger the base okay stronger the base the lower the KB the weaker the base okay and the same thing is applied with the PK pkbs and um so pkb here if I were to say for this one if it has a high KB it has to have a low pkb and if it has a low KB it has to have a high pkb all righty so again these are your inverse relationships here so the higher the KB again Lower the pkb so that's going to be a stronger base in this condition and if there's a low KB it's it's a weaker base and if it's a weaker base it's going to have a high pkb all right so we finished off talking about Ka and KB now right so let's go ahead and say one other thing here before we go into PKA and pkbs so I talked about Ka and I talked about KB right and again the higher the KA the stronger the acid the higher the KB the stronger the base higher the KA the lower the PK the higher the KB the lower the pkb right and again that's just a measure of acidity or basicity now what we can do is we can say KW this is actually working perfectly KW is equal to k a times KB so now why is that because if you remember from before what did we say KW was equal to if you remember from before we said KW which was 1.0 * 10us 14 is equal to the H+ concentration times the hydroxide ion concentration right well we talked only about that for strong acids and strong bases the formula still applies here but now we just have to take into consideration the amount of the acid that ionizes and then KB the amount that the base ionizes so then for this case we can actually say H+ we could apply here the KA so KW equals Ka a and then hydroxide we can say how much of the hydroxide disassociates times KB and so we could use this formula once again to be able to determine some of these calculations that we're going to do okay so again K wal K * KB is an important one now let's finish up with PK and pkb we already know the relationship right that PKA is equal to the negative log of the KA and pkb is equal to the negative log of the KB well we can actually say PKA plus pkb equal 14 doesn't that sound familiar like pH plus P equal 14 same concept here because here here's one thing I want to finish up with PKA and pkbs we're going to talk about buffers in our final videos and a buffer is basically a mixture so a buffer is designed to resist any abrupt changes in the pH so a buffer is made up of an acid a weaker acid so a weak acid and it's Associated conjugate base that's what a buffer is so when we're picking a buffer for maybe some type of uh reaction in a chemistry lab or within the actual human body we want to make sure that the buffer that we're applying is pretty darn equal to the pH that we desire so for example my blood is naturally going to keep a pH of about 7.35 to 7.45 so let's say that's the pH let's just say 7.4 to keep it right in the middle right keep it a a nice number here 7.4 if I had to pick between a whole bunch of different buffer systems that I'd want in my bloodstream to keep it at that part part right there I'd want the pka of a buffer to be as close to 7.4 as possible but we even give it a nice little cushion there so we say the best buffer is plus or minus one of the pka of the pH so anything that is plus or minus one of the pH is a good buffer so for example if I had two buff offers here one of them had a PKA that is equal to 9.3 the other one had a PKA that's equal to let's say 6.5 and I had another PKA that's equal to uh 3.2 which one of these am I going to pick I'm going to pick the one that's closest to the pH because that's going to give me the best buffer which one's plus or minus one oh this guy right there he's our best choice so he's going to be the best buffer in this case Okay so again what do we talked about we talked a lot a lot of stuff in this video guys we talked about the definition of what an acid and a base is according to three different theories we talked about all of our strong acids and strong bases and what that means to be a strong acid and strong base we discussed conjugate acid base pairs we talked about the pH scale and its relationships and formulas we talked about KW Auto ionization of water and Ka and KB and their relationship to the pH scale and then we talked about PKA and pkb with respect to buffers and also with respect to Ka so I I hope you guys learned a lot out of this video it's going to help you guys when we start doing some practice problems all right guys in the next video we're going to put what we did into action so the next video we're going to talk about specifically being able to calculate the pH of specifically strong acids and strong bases so if you guys want to click up in the right corner there and I'll see you guys there