Practice: Which of the following compound(s) cannot be classified as an acid?
E. All are acids
|Ch.1 - Intro to General Chemistry||4hrs & 4mins||0% complete|
|Ch.2 - Atoms & Elements||4hrs & 21mins||0% complete|
|Ch.3 - Chemical Reactions||4hrs & 18mins||0% complete|
|BONUS: Lab Techniques and Procedures||1hr & 38mins||0% complete|
|BONUS: Mathematical Operations and Functions||47mins||0% complete|
|Ch.4 - Chemical Quantities & Aqueous Reactions||3hrs & 54mins||0% complete|
|Ch.5 - Gases||3hrs & 22mins||0% complete|
|Ch.6 - Thermochemistry||2hrs & 26mins||0% complete|
|Ch.7 - Quantum Mechanics||2hrs & 17mins||0% complete|
|Ch.8 - Periodic Properties of the Elements||3hrs||0% complete|
|Ch.9 - Bonding & Molecular Structure||3hrs & 20mins||0% complete|
|Ch.10 - Molecular Shapes & Valence Bond Theory||1hr & 53mins||0% complete|
|Ch.11 - Liquids, Solids & Intermolecular Forces||2hrs & 21mins||0% complete|
|Ch.12 - Solutions||2hrs & 49mins||0% complete|
|Ch.13 - Chemical Kinetics||2hrs & 46mins||0% complete|
|Ch.14 - Chemical Equilibrium||2hrs & 26mins||0% complete|
|Ch.15 - Acid and Base Equilibrium||4hrs & 42mins||0% complete|
|Ch.16 - Aqueous Equilibrium||3hrs & 48mins||0% complete|
|Ch. 17 - Chemical Thermodynamics||1hr & 44mins||0% complete|
|Ch.18 - Electrochemistry||2hrs & 21mins||0% complete|
|Ch.19 - Nuclear Chemistry||1hr & 31mins||0% complete|
|Ch.20 - Organic Chemistry||3hrs||0% complete|
|Ch.22 - Chemistry of the Nonmetals||2hrs & 1min||0% complete|
|Ch.23 - Transition Metals and Coordination Compounds||1hr & 54mins||0% complete|
|Identifying Acids and Bases||52 mins||0 completed|
|Arrhenius Acid and Base||7 mins||0 completed|
|Bronsted Lowry Acid and Base||18 mins||0 completed|
|Amphoteric||6 mins||0 completed|
|Lewis Acid and Base||14 mins||0 completed|
|pH and pOH||64 mins||0 completed|
|Ka and Kb||21 mins||0 completed|
|Ionic Salts||46 mins||0 completed|
|Diprotic Acid||31 mins||0 completed|
|Polyprotic Acid||24 mins||0 completed|
|Strong Acids and Strong Bases (IGNORE)|
|Conjugate Acids and Bases|
Most acids have one common feature and that is the presence of the hydronium ion, which is represented by H+ or H3O+.
Concept #1: Identifying Binary Acids.
Hey guys, in this new video, we’re going to take a look at acids. We’re going to see what are the key features that make acids acids, and more importantly, how do we distinguish one type of acid from another.
So if we take a look at the beginning it says the most common feature of an acid is that it possesses an H+ ion. Now here’s a thing, not all acids have H+ but a majority of them do. Or we’re going to say the name of this H+ ion is hydronium ion. Later on we’re going to learn that certain types of acids don’t possess H+ at all but they’re still acids. They’re just a special type of acid which we call louise acids. When we get to that section we’ll cover that in detail, but for right now just remember most acids have H+. This H+ ion is called hydronium ion or hydrogen ion. And we’re also going to say here is H+ is the same thing as H3O+. You’re going to see that your professor uses them interchangeably. That’s because those same exact thing, they’re just a different rate way of writing the same exact ion. So just remember H+ is the same thing as H3O+.
Now we’re going to say when it comes to acids there are two major forms. Now the simpler of the two are called binary acids. Now binary acid possesses H+ and that H+ is connected to some electronegative element. Now we’re gonna say the electronegative element would be Nitrogen, Sulphur, Phosphorus, or some type of Halogen, Fluorine, Chlorine, Bromine iodide.
Now the key things that binary acids don’t have is, were going to say no binary acid has oxygen. If you have oxygen on an acid that will be different type of acid. So all binary acids none of them oxygen present. Now the second thing is, they usually possess no metals. Usually if we have a metal within a compound, it’s ionic. And usually it will be a base, if a metal is involved, it’s usually a base. Of course there are exceptions to this and we’ll talk about that later on.
Now we’re going to say here, the most common types of these particular acids are called your haloacid. Halo means that we have a halogen, and of course acid is H+. So remember the halogens are in group 7A, and if you’re in a group 7A what’s your charge? Hopefully you guys remember the charge distribution in the periodic table. So remember if you’re in a group 7A your charge is -1, so we have F-, CL-, BR-, I-. So when we have positive H+ connecting with one of these negative ions, our haloacids would be HF, HCL, HBR, HI. So these are your four haloacids, hydrogen connected to a halogen.
Now we could talk about other types of common binary acids based on groups. So we just looked at group 7A. So that would be HF, HCL, HBR, HI. But remember we also got group 6A involved. Remember when you’re group 6A, your charges are -2. So we’re going to skip Oxygen in group 6A because again, binary acids don’t possess any oxygen. So if we go down groups 6A and you’re looking at the periodic table, the next would be S2-, Se2-, Te2-. And again this guys are combining with H+ to give us a binary acid.
So let’s just remember what we learn about ionic compounds. How do we make them into compounds? Remember the charges, the number from the charges crisscross to give us our compound. So we’d have H2S, H2Se, H2Te, so these will be your binary acids that are common to group 6A.
And then we could say that we have certain ones that are a little bit different. So these are the odd ones, that their technically still binary acids, they’re just a little bit different. So here we can have HN3 and here we can have HCN, so those would be the odd ones out. They’re a little bit different from group 7A and 6A but they’re still binary acids because in here we have H+ connected to a negative ion (N3-), and here we have H+ connected to a polyatomic ion CN-, cyanide ion.
Now we’ll learn how to name these later on in coming videos, but for right now, this are our typical common types of binary acids from group 6A, group 7A and also the odd ones. Now that we’ve seen these first type of acids, it’s time to talk about the second type.
Binary acids can be identified by the fact that they all possess an H+ ion attached to an electronegative element.
Oxyacids can be identified by the fact that they all possess an H+ ion connected to a nonmetal and oxygen, hence the prefix “oxy”.
Concept #2: Identifying Oxyacids.
Now were going to say the second type, they’re called oxy acids. And from the name oxy you can tell that it has one particular element in there, Oxygen. So we’re going to say oxyacids are acids that contains H+ ion of course, they contain Oxygen, and they contain some type of none metal. So the difference between binary and oxy is that, oxy has Oxygen involved binary does not.
Something else they both share in common is that, they usually don’t possess any metals. Again we’re going to see, usually when we have a metal involved that compound will be a base. Now we’re going to say that oxy acids are formed when we hydrate none metal oxides. So with this example here, our none metal oxide would be Carbon dioxide here, and here would Sulphur trioxide. So none metal oxide just means you have a none metal connected to Oxygen. Now all what’s going to happen is the water is going attach itself to this whole compound. And what we get are two types of oxyacids, and they’re oxyacids because both have H, both have Oxygen, remember that H is really H+, they both have Oxygen and they both have some none metal involved. And that’s what makes them both oxyacids. So here we would have Carbonic acid, and here we have Sulphuric acid. And again we’ll learn how to name these particular types of acids later on. But for right now just realize, how do we create oxyacids? We just add H+ to some type of none metal oxide, CO2, SO2, SO3, those types of compounds. PO3, those kind of compounds you add water to, to make your new types of oxyacids.
Now that you guys have seen the two types of acids, binary and oxy, I want you guys to try to answer the next following question. Here I say, “Which the following compound or compounds cannot be classified as an acid?” I gave you just brief description on both, but try to use the best judgement to determine which ones could be, which ones could not be. Then once you’re done with that, click on the explanation button, you’ll see a video of me explaining how to approach this problem, guys. Goodluck!
They are created by the hydration of nonmetal oxides.
Conversely, metal oxides create bases when hydrated by water.
Practice: Which of the following compound(s) cannot be classified as an acid?
E. All are acids
Concept #3: Acids and Electrolytes
In this video we're going to take a look at the strength of binary acids. So, remember we talked about there being two classes of acids, there are binary acids and then there are oxyacids, for binary acids we'll go over the different roles that are associated with determining the strengths of different binary acids, binary acids have their own set of rules, oxyacids, which are a different type of acid, will have their own separate list of rules that we follow for them and we're trying to compare the strengths of different oxyacids. Now, regardless if you're looking at a oxyacid or binary acids, we're going to say strong acids in general are strong electrolytes. So, they ionize completely in water, here we have HCl, which is a strong binary acid, we have a single arrow going forward to show that it completely ionized it, it goes to completion, so that means we're going to make 100% of both of these ions, weak acids are weak electrolytes and again whether it be weak oxyacids or weak binary acids, weak acids in general are weak electrolytes. So, they don't completely ionize in water, here we don't have a solid arrow going forward, we have double arrows to say that instead of going to completion our reaction reaches an equilibrium. So, we're going to say an equilibrium is established, so that means we're going to make less than 100% of these ions, and in fact, we tend to make a very small amount of these particular ions. So, when it comes to weak acids we're going to have to basically set up what are called ice charts later on in order to determine the pH of these different types of weak acids. Now, what we're going to talk about binary acids, keep in mind again there are two classes of acids, these are just the rules for binary acids, when we talk about oxyacids they have their own separate list of rules.
Strong Acids are strong electrolytes that completely dissociate into ions when dissolved in water.
Weak acids are weak electrolytes that don’t completely ionize, but instead reach a state of equilibrium.
Concept #4: Strength of Binary Acids.
Now, the strength of a binary acid is based on two major factors. First, it’s based on the electronegativity or the size of the non-metal that’s connected to my hydrogen. In this case, this non-metal in HCl that we’re talking about is Cl. And in HF, the non-metal that we’re talking about is fluorine. So the strength of a binary acid is based on those two factors, electronegativity as well as the size of the non-metal. Now, we’re going to say for elements in the same period and remember when I say period—we have a periodic table here. This is period one, two, three, four, five, six, and seven. Here we’d have boron, carbon, nitrogen, oxygen, fluorine and neon. They’re all on period two with each other. We’re going to say, elements in the same period, we’re going to look at their electronegativities in order to determine acidity. And basically, what you’re going to see here, the more electronegative the nonmetal, the more acidic.
If we’re taking a look at certain examples, let’s say we’re looking at CH4, HF and BH3. So, we’re going to say that boron, carbon and fluorine are all in the same period. In terms of electronegativity, remember, it increases going from left to right. And because the increase is going from left to right, we’re going to say fluorine is the most electronegative out of these three non-metals that I’ve underlined. And because fluorine is more electronegative than boron and carbon, HF will be a stronger acid than CH4 or BH3. That’s the way we apply this. If they’re on the same period, we look at electronegativity.
However, for elements in the same group, remember groups go vertical. This is: 1A, 2A, 3A, 4A, 5A, 6A, 7A. We’re going to say here—in this case, we’re going to say we look at their size when they’re in the same group. We’re going to say, the bigger, the more acidic. Remember, when we’re talking about size, we’re really talking about atomic radius. And we’re going to say, your atomic radius increases when you go down a group. So we’re going to say, let’s say we’re looking at different types of acids. Here we have: Cl, Br, and I. Let’s say, you wanted to compare HF, HCl, HBr, HI. We’re going to say that, all four of these non-metals are all in the same group. They’re all in group 7A. We’re no longer going to look at electronegativity instead we’re going to look at size, atomic radius. Going down the group, it’s going to increase. We’re going to say here, “I” would be the largest. And because “I” is the largest, HI would be strongest acid out of this group. So that’s how we look at electronegativity and how we look at atomic radius or size to determine which binary acid is the strongest.
Now, one last thing before I’ll let you guys attempt to do the two practice questions at the bottom. The most common types of binary acids--we’ve already seen them. When H+ is connected to HF, we have HCl, HBr, HI. Group 6A will also be another common type of binary acids: H2S, H2Se, H2Te. For all purposes, we’re going to say that, these are the most common types of binary acids but only three of them are really strong. The only three strong binary acids you’re going to know this semester are these three: HCl, HBr, HI. Those are your three strong binary acids. For those three, we never have to do an ICE chart in order to find pH. For all the other binary acids: HF, HCN, H2S, H2Se, H2Te. All those others are weak binary acids. So we would need to do an ICE chart in order to figure out their pH. Just remember that these are the things that we use electronegativity and size to break ties between different binary acids but there are only three strong binary acids you need to know, HCl, HBr, HI.
When looking at a binary acid we look at both electronegativity and size to determine their strength.
Practice: Which is the weakest acid from the following?
Practice: Which of the following acids would be classified as the strongest?
The strength of an oxyacid depends on the number of oxygen atoms and the electronegativity of the nonmetal.
Concept #5: The Strength of Oxyacids.
In this video we're going to take a look at oxyacid strength. So, let's say here, that the strength of oxyacids is based on the number of oxygens or the electronegativity of the nonmetal, we're going to go over when do we look at either one. So, here we're going to say the rule is, the number one rule for oxyacids is, if my oxyacid has two or more oxygens and then hydrogens, then my oxyacid is a strong acid. So, if we take a look at the first one, we have here, HNO3, we're going to face an oxyacid because it contains hydrogen, a nonmetal and an oxygen, here's the number of oxygens is 3 the number of hydrogen's is just 1, when we subtract them we have 2 oxygens remaining, you must have a minimum of 2 oxygens remaining in order to be considered a strong oxyacid, here it meets that requirement. So, it's strong, next one, we have phenyl. So, C6H5OH, here it has one oxygen and then it has five plus one more that six hydrogens. So, in this one we don't have any oxygens remaining, remember, we need a minimum of two oxygens remaining to be considered a strong oxyacid. So, this one doesn't meet the requirements. So, it's weak. Then here we have perbromic acid. So, here we have four oxygens, we have one hydrogen. So, we have three options remaining. So, we've met the minimum requirement of having at least two oxygens left. So, this one would be strong. So, when it comes to oxyacid, this is the first tool that you need to remember, you need to have a minimum of at least two more oxygens than hydrogen's to be considered a strong oxyacid, we're going to move on to the next portion when we talk about this topic of oxyacids but keep in mind this first important.
If your oxyacid has 2 or more oxygens than hydrogens then your oxyacid is a strong acid.
Concept #6: Comparing the Strength of Oxyacids.
So, when comparing the strengths of oxyacids, we need to remember a few things. Alright, so when comparing the strengths of different oxyacids, remember, if they have different number of oxygens then the more oxygens you have the more acidic. So, for example, if we're comparing HClO versus HClO4, if we did the math of oxygens minus hydrogen's, this one has nothing left, this one here would have three oxygens left. So, this one would be stronger, if they have the same number of oxygens though then the more electronegative the nonmetal the more acidic, for example, if I was comparing HBRO3 versus HClO3, here in our periodic table we have fluorine, chlorine, bromine and iodine here and we're going to say here, when it comes to electronegativity. Remember, electronegativity increases, when we go from left to right and it decreases when we go down a group. So, when we do the number of oxygens minus the number of hydrogen's here, we're going to have two remaining for both, two remaining oxygens for this one, two remaining oxygens for this one, but chlorine is not as far down as bromine, since bromine is further down into less electronegative, and remember we just said, the more electronegative the nonmetal the more acidic it would be. So, when it comes to oxyacid strength, first look at the number of oxygens remaining, the one with more oxygens remaining will be stronger but if the number of oxygens remaining are the same then we look at the nonmetal, the more electronegative one will be the stronger oxyacid. Now, it's all rules of course there are exceptions that exist and here we have these three exceptions. So, this first compound here this is oxalic acid, this one here is ionic acid and this one here represents amphoteric species, which we'll talk about.
Alright, so if we take a look at the first two, we have four oxygens in this one and two hydrogens, three oxygens and one hydrogen, both of them have the minimum requirement of oxygens left of two, but they're both weak, the reason they're weak is because the carbon in them and the iodine in them are not very electronegative. So, this weakens the acid overall. So, although they've met the minimum requirement of having at least two more oxygens and hydrogen's they are still weak because their nonmetals are not very electronegative. Now, amphoteric species are species that connect acids or bases, here this connect as an acid because it possesses an H but also can be a base because it's negatively charged, other ones that could fit this criteria H2PO4-, HCO3-, HPO4 2-, all of them have hydrogen so they can act as acid and they have a negative charge, so it doesn't matter if you do the math and they have a minimum of two oxygens remaining because they're amphoteric they're going to be considered weak as well. So, I just remember these are the three major types of exceptions when it comes to determining the strength of oxyacids and now that we've gone over that, let's see if you guys can answer the practice question left at the bottom of the page, attempt it on your own but if you get stuck don't worry just come back and take a look at how I explained the best approach of this question.
When comparing the strengths of different oxyacids remember:
If they have different number of oxygens then the more oxygens the more acidic
If they have the same number of oxygens then the more electronegative the nonmetal the more acidic.
With some rules there are exceptions. Oxalic acid and Iodic acid have two more oxygens than hydrogens but are weak acids because carbon are iodine have low electronegativity.
Amphoteric species (compounds that can act as both an acid or a base) are also an exception to the rule for oxyacids.
Practice: Rank the following oxyacids in terms of increasing acidity.
Concept #7: Bases and Electrolytes.
Hey guys! In this new video, we’re finally going to take a look at the strength of bases. So up to this point we’ve talked about acids. We’ve talked about binary acids and oxyacids. And we should know their strengths at this point as well. How do we determine when they’re week and when they’re strong? Again, this is important because weak acids, we have to use an ice chart in order to find Ph, strong acids we don’t have to do that.
So what we’re going to say now let’s talk about bases. We’re going to say that strong bases are considered to be strong electrolytes. And what does a strong electrolyte mean? It means that it completely breaks up or dissociates when dissolved in water. Here you’re going to notice that you have a solid arrow going forward. Again this means, reaction goes to completion. So all of the reactant will break down basically to 100% so we get those two ions, we get NH+, we got OH-.
Now weak acids are considered to be weak electrolytes. They don’t completely break down, they don’t completely ionize. Here again we’ll notice that there’s a double arrow, meaning that equilibrium is established. That means our reaction goes into forward direction as well as the reverse direction. Now weak acids, weak bases, favor the reactants. All that means is that our reactants really don’t break down that much. So a majority of our compounds will be in the form of a reactant, very little of it will be made into product. In the opposite way, strong acids and strong bases, we’re going to say they favor the products, which means nearly all of our reactants break down to give us products. So again remember, weak acids, weak bases, we need to use an ice chart to find Ph or POH. We’ll talk about those later down the road. It’s imperative that you guys remember how to identify things.
Strong Bases are strong electrolytes that completely dissociate into ions when dissolved in water.
Weak bases are weak electrolytes that don’t completely ionize, but instead reach a state of equilibrium.
Bases are characterized by THREE major features: they may possess metals, they may have a negative charge or they may be an amine.
Concept #8: Bases and Group 1A.
Now if we take a look here, we’re going to say when it comes to bases, there are three features that we’re supposed to look out for. Generally, bases posses metals. If a metal is involved, then it’s most likely going to be a base. We’re going to see a couple of exceptions of that when we get to amphoteric species. But for right now, let’s not worry about that.
Next, we’re going to say bases tend to have a negative charge. So bases will have metals or a negative charge, or they’re going to get classified as another compound which we call amines. Now amines are an organic type of compound. But we get to see them for the first time in chemistry. So before you guys move to organic, you got a small taste of what amines are, and how exactly do they relate to bases. So if we take a look at the rules, we’re going to say any group 1A metal, when they combine with any four of these ions (OH-, H-, O2- or NH2-) make a strong base. Now these four ions, we call them hydroxide ion. Hydrogen can be positive or negative. When it’s positive, it’s called hydronium ion. But when it’s negative, it’s called hydride ion. O2- is oxide and NH2- is amide. So notice how all of them end with ide because they have a negative charge present. So if one of these negative ion group with any group 1A metal, it’s going to create a strong base. Now recall, what are the charges of group 1A metals? Group 1A metals are +1. So examples, we could have NA+, Li+, Cs+ and K+. Any of these ions could react with any one of those four ions to give me a strong base.
So here, Na+ OH- combined together to give me NAOH. This (Li+ H-) would give me LiH. The 2 from here (O2-) would come here (Cs+), the 1 from here (Cs+1) would come here (O2-) Cs2O. Then here (K+ NH2-) we’d have KNH2. Remember you could mix and match any group 1A ion with any one of those four. He we get Sodium hydroxide, Lithium hydride, Cesium oxide, Potassium amide. So again, any group 1A metals with any of those four ions will create a strong base. And strong bases don’t require an ice chart in order to find Ph or POH.
Any Group 1A metal when combined with OH –, H –, O2– or NH2 – makes a STRONG BASE.
Concept #9: Bases and Group 2A.
Next we’re going to say group 2A ions, and here there’s a caveat. It’s not all group 2A. It’s certain metals in group 2A. So we’re going to say, any group 2A metal from Calcium to Barium, when combined with OH-, H-, O2- or NH2- will create a strong base. So remember, if you look at your periodic table, group 2A, we’re going to see Beryllium, Magnesium, Calcium, Strontium, Barium. And we have Arum at the bottom as well and usually we don’t see bases with that so we’re not going to worry too much with that. So what we’re seeing from Calcium and lower, so the cut off is Calcium. These two (Be, Mg) will get left out. So here we’re going to say, group 2A is +2 in charge. So we could have these (Ca2+, Ba2+, Sr2+, Ca2+) reacting with any one of these ions (OH-, H-, O2-, NH2-). So the 2 from here (Ca2+) will come over here (OH-), the 1 from here (Ca2+) will move over here(OH-) and cause nothing, so this would be Ca(OH)2. Here we’re going to get the 2 from here (Ba2+) coming here (H-), BaH2. This (Sr2+) is +2, this (O2-) is -2 so when they combine they cancel out (SrO). And then the 2 from here (Ca2-) can come down here (NH2-), Ca(NH2)2. So those are all examples of strong bases.
Now we’re going to say here, that Magnesium is above Calcium. And we’re going to say Magnesium hydroxide is a base. Now technically it’s a weak base, it’s not as strong as the others. The others break up 100%. But here’s a thing about Magnesium hydroxide. We’re going to say that a small portion of it will dissolve, let’s say about 10% of it dissolves. That small portion of it that does dissolve, it dissolves all the way. And when it comes to weak acids and bases, sometimes a piece of it will dissolve, but not completely. For Magnesium hydroxide, a small portion of it dissolves, and it dissolves completely. Because of that weirdness, we’re going to say Mg(OH)2 is not a strong base, but it’s strong enough that we don’t need to do an ice chart. So we’re going to say that Magnesium hydroxide is a weird special case. Technically it’s weak, but it’s strong enough that it may not require an ICE chart. Just make a little note of that.
Group 2A metals, from Ca2+ to Ba2+, when combined with OH –, H –, O2– or NH2 – makes a STRONG BASE.
Amines are compounds with only nitrogen or hydrogen (i,e. NH3) or with carbon, nitrogen and hydrogen (i,e. CH3NH2).
Concept #10: Amines.
Then finally the last group we talked about, we’re going to say amines. And since we’re going to need some room guys to do this, I’m going to remove myself from the image so we have more room to work with.
So amines, what the heck is an amine? Well we’re going to say an amine is a compound that has N and H, Nitrogen and Hydrogen, such as example, NH3. Or it is a compound that has Carbon, Hydrogen and Nitrogen. Example of that, we could have CH3NH2. So just remember that. Amines are compounds with Nitrogen and Hydrogen or Carbon, Hydrogen and Nitrogen. We’re going to say here that neutral amines, ones without charges, are considered to be weak bases. Examples, the two we just wrote, NH3 or CH3NH2. Obviously there are more than just these, but these are two good examples. We have Ammonia and then Methylamine. The name even tells you that it’s an amine.
And we’re going to say next that positively charged amines are considered to be weak acids. So here we could have NH4+ which is Ammonium. This (NH3) is Ammonia. So Ammonium is positive so it is a weak acid. Or CH3NH3+, this would be Methyl ammonium.
Now, we saw compound earlier when we talked about binary acids. That compound was HCN. Now HCN has the characteristic of an amine right? I said it was Carbon, Hydrogen and Nitrogen. But remember, you can’t be a binary acid and amine at the same time. So HCN, even though it has those three elements is not an amine. It’s not an amine, it is a binary acid, so make note of that. This is Hydrocyanic acid, not an amine. So just make note of that, what an amine is. It usually has just N and H, or it could have H, C and N. HCN though, is not an amine, it’s a binary acid. So look out for those characteristics. And if they give you the name, usually you’ll see an amine in the name in some way. That would be a clear giveaway that you’re dealing with an amine.
Now hopefully you guys have learned how to identify acids and bases now, because again, it’s imperative that you know how to identify them as either weak or strong because if you don’t know that, you won’t know how to calculate PH or anything else dealing with acid base chemistry.
Neutral Amines are considered weak bases.
Positive Amines are considered weak acids.
Example #1: Classify each of the following as a strong acid, weak acid, strong base or weak base:
a) HCHO2 c) H2NNH2
b) (CH3CH2)3NH+ d) HBrO3
Practice: Classify each of the following as a strong acid, weak acid, strong base or weak base.
Practice: Classify each of the following as a strong acid, weak acid, Strong base, or weak base.
Practice: Classify each of the following as a strong acid, weak acid, strong base or weak base.
Join thousands of students and gain free access to 46 hours of Chemistry videos that follow the topics your textbook covers.
Enter your friends' email addresses to invite them: