General Chemistry Chemical Reactions II

Hi, welcome back to Educator.com.0000

Today we are going to be talking about additional chemical reactions; especially those that happen in aqueous solution.0003

Let's go ahead and go over the lesson overview; we are going to cover a few more areas.0013

We are going to build upon what we discussed last time on chemical reactions.0019

The very first type of reactions that we are going talk about are going to be involving what we call acids and bases.0024

That is going to be our first really general category.0035

We are going to be then jumping into another category of acids and bases.0040

These are going to be nontraditional ones; something that we couldn't look at just by formula; I will build on that.0046

After that, we are going to then go on to what we call redox reactions.0058

We are going to learn how to identify when redox reactions occur and exactly what is occurring during the reaction.0064

Finally we go on to our summary and sample problems.0071

We are first going to tackle acids.0078

The first interpretation of what an acid and a base is is called the Arrhenius definition.0081

The Arrhenius definition tells us the following.0090

That acids are going to dissolve in water and dissociate into a proton and anion.0093

For example, we can take hydrochloric acid; this is going to dissociate in water.0100

This is going to form H⁺ which we call proton and its anion, Cl^1-.0109

This is very familiar to what we talked about in the previous session on the first part of chemical reactions.0119

If you recall, we talked about what was labeled an electrolyte, strong and weak.0127

This is pretty much an electrolyte; the Arrhenius acids essentially behave as the electrolytes.0135

We can go ahead and look at another example.0144

HNO₃ aqueous can also break up into H⁺ aqueous and NO₃^1- aqueous.0147

Just want to do a reminder for strong electrolytes.0162

Remember we are going to use a single arrow pointing in the forward direction0165

to show we get full dissociation--that we completely break up into ions.0168

However there are also weak electrolytes.0175

There is going to be Arrhenius acids that are going to be weak electrolytes.0179

For example, hydrofluoric acid is the typical example.0183

If everybody tried to recall, remember this was called an equilibrium arrow.0187

This is going to form H⁺ aqueous and F^- aqueous.0193

Now that we have covered the Arrhenius definition of an acid,0201

let's go ahead and cover the Arrhenius definition of a base.0204

According to Arrhenius, bases dissolve in water and dissociate into a cation and hydroxide.0208

In other words, Arrhenius bases are also electrolytes.0217

The typical example will be sodium hydroxide.0221

Sodium hydroxide is going to be a strong electrolyte; we use a single arrow.0228

That is going to dissociate and break up into its cation, Na¹⁺, and hydroxide, OH^1-.0233

There is also going to be other types of Arrhenius bases that don't completely dissociate.0243

Again we are going to use a double arrow for them.0249

We can imagine maybe the following; magnesium hydroxide aqueous.0253

This one again is not going to completely dissociate; we use an equilibrium arrow.0259

That is going to give us Mg²⁺ aqueous and two hydroxides aqueous.0264

The Arrhenius interpretation of an acid and a base is relatively straightforward.0273

Once again they simply are ionic compounds that break up into their respective cations and anions when in aqueous solution.0278

The Arrhenius definition is fine for most purposes.0289

However what we are going to cover now is called the Bronsted-Lowry definition.0292

The Bronsted-Lowry definition, the main difference is the following.0297

It always takes to in count the rule of water.0301

Let's first start off with the definition of acids.0306

Acids dissolve in water; what they do is the following.0309

They are going to donate a proton to water.0313

The first example that we talked about was hydrochloric acid.0318

Let's go ahead and talk about that once again.0321

This time interpreting it from the Bronsted-Lowry perspective; HCl aqueous.0324

This time we are going to follow the definition.0330

Instead of just dissociating into cation and anion, the acid is going to react with water.0333

It is going to donate a proton to water.0339

So HCl aqueous plus H₂O liquid goes to the following.0342

If HCl loses its hydrogen, all that is left is going to be Cl^1- aqueous.0347

If H₂O is accepting the hydrogen from HCl, H₂O becomes the following; H₃O¹⁺ aqueous.0356

There is a couple of things I want to point out here.0368

Number one, we give H₃O¹⁺ a unique name.0370

This is what we call the hydronium cation.0376

The hydronium cation, anytime you see it, it is always indicative that an acid is present in water; acid in water.0382

Basically a Bronsted-Lowry acid will always form hydronium in water; Bronsted-Lowry acids form H₃O¹⁺.0395

This is the Bronsted-Lowry interpretation of what an acid is.0408

Let's go ahead and look at another example.0414

Let's go ahead and look at HClO₄; HClO₄ aqueous.0421

Once again this is going to go ahead and react with water liquid.0428

That goes on to form ClO₄^1- aqueous plus H₃O¹⁺ aqueous.0433

Always remember everyone that you always want to pay attention to balancing; don't forget.0444

You always make sure your equation is balanced; not only the coefficients but also the charges.0449

My overall charge on the reactant side is 0 and 0 here.0454

Overall here it is -1 and +1 on the right side giving me a net sum of 0 too.0461

By breaking up into hydronium and the anion, we are going to get the balanced chemical equation also.0471

There is a couple of different types of acids that we want to talk about right now.0480

Right now when you just see one hydrogen in the formula of the acid, that is what we call monoprotic acids.0485

But of course there are going to be many other types of acids that contain more than one hydrogen.0497

These are what we call polyprotic acids.0504

The typical example of a polyprotic acid would be sulfuric acid, H₂SO₄ aqueous.0510

H₂SO₄ specifically, this is called a diprotic acid; the prefix di meaning there is two of them.0520

Let's go ahead and react this with water; here is the rule of thumb.0527

For polyprotic acids, anytime you write them reacting with water, you never take off both of the hydrogens at once.0535

You always do one deprotonation at a time; for polyprotic acids, again one deprotonation at a time.0544

That means the following; I am not going to get just sulfate.0563

I am going to get just one proton is going to be taken off.0568

That is going to be HSO₄^1- aqueous; again we always form H₃O¹⁺ aqueous.0571

The rule of thumb again, don't forget, it is going to be one deprotonation at a time.0586

Let's go ahead and take HSO₄^- and deprotonate it one more time.0597

That is going to be HSO₄^1- aqueous plus H₂O liquid.0603

We have another rule of thumb that we also have to go over.0610

For polyprotic acids, only the first step is going to be significantly strong.0614

Everything else is going to be considered a weak step.0621

In other words, HSO₄^1- will be functioning as a weak electrolyte.0624

We are going to use therefore an equilibrium arrow.0635

We are going to get now sulfate plus H₃O¹⁺ aqueous.0639

This is the basic definition of Bronsted-Lowry acids.0648

Once again a couple rules to remember.0654

That Bronsted-Lowry acids are going to react with water and are going to donate a proton to water.0657

If we are dealing with polyprotic acids, you only write their deprotonations one step at a time.0662

Finally for the polyprotic acids, only the first step is considered a strong electrolyte step.0671

Everything else after that is going to be considered weak.0679

You may be wondering: how do I know what are the strong electrolytes for acids?0683

How do I know which ones are the weak electrolytes for acids?0688

We are going to go over that right now.0691

Basically the strong electrolytes are what we call strong acids.0693

Traditionally there are seven strong acids that are normally taught to a freshman level chemistry course.0700

Your instructor may give you more or less; so definitely pay attention to him or her.0706

But the basic seven that are always strong are the following.0712

Hydrochloric acid, HBr, HI, HNO₃, H₂SO₄, HClO₄, and HClO₃.0716

Pretty much it is pretty safe to assume that any acid that is not on this list you can assume to be weak.0734

Again these are the seven strong acids that I would recommend definitely committing to memory.0742

You will always use a single arrow for them.0747

But remember, make a note, for sulfuric acid which is what we just looked at.0750

That is the polyprotic acid; only the first step is strong.0756

Now that we have talked about the Bronsted-Lowry definition of an acid,0762

let's go ahead and talk about how Bronsted-Lowry defines a base.0766

Bases dissolve in water and what they do, instead of donating a proton, they accept a proton from water.0773

The typical example of a Bronsted-Lowry base is going to be ammonia.0779

Ammonia... this is definitely something you should just have it memorized... is NH₃.0785

Ammonia is going to react with water, H₂O liquid.0790

It turns out that ammonia is going to be what we call a weak base.0794

Weak bases are going to be weak electrolytes; we are going to use an equilibrium arrow for them.0800

NH₃ aqueous plus H₂O liquid, it is going to accept the proton from water.0806

NH₃ is going to become... remember your polyatomic ions... NH₄¹⁺ aqueous.0812

H₂O is going to lose a hydrogen this time forming hydroxide.0818

We see it pretty clear cut; Bronsted-Lowry acids always generate hydronium.0826

And Bronsted-Lowry bases are always going to generate hydroxide when reacting with water.0831

What are some other typical bases, Bronsted-Lowry bases, that you want to be on the lookout for?0840

Pretty much anything that is related to ammonia is going to be a weak base.0847

Specifically the molecules that I am talking about that are related to ammonia, these are going to be called amines.0861

Amines contain nitrogen and usually with carbon and hydrogen.0871

NH₂R would be an example where R is a molecule that contains hydrogen and carbon; what we call a hydrocarbon.0890

Not only NH₂R but NHR₂; not only NHR₂ but NR₃.0904

These are the again generic formulas for amines; also Bronsted-Lowry bases.0913

Just by looking at the name, you can usually tell if it is going to be a base or not.0922

A lot of the organic bases that are also biomolecules, they all end in the same suffix.0926

What you want to look for is a ?ine suffix.0935

You actually know a lot already of examples that are weak Bronsted-Lowry bases that are also bimolecules.0942

For example, here is one of my favorite.0951

Of course it is going to be caffeine; you see this caffeine ends in ?ine.0956

Adrenaline is another one that ends in ?ine.0961

A lot of the drugs also end in ?ine; they are also bases and amines.0968

For example, morphine, cocaine, and heroine; the list goes on and on.0975

I am sure you will encounter these in some of your biology classes.0983

But just by looking at the name again, the ?ine suffix usually indicates if something is an amine.0986

What are the strong bases and what are the weak bases?0998

Pretty much the strong bases are going to be group 1 metal hydroxides and some group 2.1002

Sodium hydroxide, potassium hydroxide, lithium hydroxide, etc; those are the group 1 metal hydroxides.1023

Some of the group 2 metal hydroxides that are also considered strong are going to be calcium hydroxide and magnesium hydroxide.1032

Pretty much it is safe to assume... again you should check with your instructor about this.1045

It is safe to assume that anything that is not on this list of strong bases can be considered to be weak.1050

Especially ammonia and its derivatives which we call amines.1057

Now that we have talked about the Bronsted-Lowry interpretation of acids and bases...1065

We saw that they have one thing in common; that they all react with water.1071

But in one situation, that is with acids, water is accepting the hydrogen.1076

In the second situation, that is with bases, water is actually donating the hydrogen.1083

Let's take a look more closely at that.1089

Water has the unique ability of functioning as both a Bronsted-Lowry acid and also as a Bronsted-Lowry base.1093

This is what we call the amphiprotic nature of water.1099

Because of this, let's go ahead and further scrutinize water reacting with itself.1105

H₂O liquid plus H₂O liquid; this is going to be a weak process.1111

If we focus just on maybe the left water molecule... let's call that the acid.1124

Let's call the other water molecule the base.1135

If the acid is going to function as an acid, it is going to donate its hydrogen.1143

But if it donates a hydrogen, it becomes hydroxide.1149

The second water molecule which we are labeling as the base is going to accept the hydrogen.1157

All of a sudden it becomes H₃O¹⁺.1163

This process is known as the autoionization of water.1168

It is physically representative of what happens in normal water.1172

The water that you get out of your tap, it is not just H₂O.1178

There is a little bit of hydroxide and there is a little bit of hydronium.1184

But the majority is clearly going to be water.1187

Again this is what we call the autoionization of water.1192

When we go ahead and look at quantitative acid-base chemistry,1200

down the line, this is going to come into play a lot more.1204

Now that we have covered the traditional acids and bases, Arrhenius and Bronsted-Lowry,1210

we are now going to look at some more acids and bases1218

that are going break the trend from Bronsted-Lowry and Arrhenius.1224

So far we learned that acids can be identified as basically having one or more hydrogens.1229

Just by looking at the formula; the hydrogens are usually at the beginning.1235

Bases were pretty easy to spot; they either contained hydroxide or they were amines related to ammonia.1240

There are molecules that fit neither profile.1249

Specifically we are going to now look at oxides.1254

Basically oxides are anything that contain oxygen.1257

Certain oxides of metals and nonmetals, they contain neither hydrogen or hydroxide.1260

However they can still react with water to produce hydronium or hydroxide.1266

What we are going to do now, we are going to spend some time on going over1275

what are called the acidic oxides and what are called the basic oxides.1279

Certain oxides of nonmetals... remember certain oxides of nonmetals are actually acidic.1286

That is they will dissolve in water to form hydronium.1294

Here I present to you three examples.1301

They are going to be the oxides of carbon, of sulfur, and of nitrogen.1305

Let's first look at the oxide of carbon, carbon dioxide.1315

Carbon dioxide can go ahead and react with water to form H₂CO₃; this is called carbonic acid.1318

Carbonic acid remember is a compound we learned before; it is a relatively unstable product.1329

It can go on to react with water to form bicarbonate and hydronium.1337

Essentially carbonic acid is going to function as a Bronsted-Lowry acid.1343

This process is very interesting; this is actually what is happening in your body.1351

If you take in too much carbon dioxide, the carbon dioxide is going to react with the water in our bodies.1356

It is going to form H₂CO₃; H₂CO₃ in turn is going to form hydroxide.1363

If we get too much carbon dioxide in our body, that is actually make our blood and etc more acidic.1367

That is going to be very very dangerous.1378

The next oxide is that of sulfur; that is sulfur dioxide.1381

Sulfur dioxide is a slightly different reaction.1387

Sulfur dioxide is going to react with molecular oxygen, O₂, to form sulfur trioxide.1390

Sulfur trioxide is going to go ahead and react with water to form sulfuric acid.1397

You know sulfuric acid is one of the seven strong acids that we covered.1402

That is definitely going to go on to form hydronium.1408

Finally the last oxide is that of nitrogen; that is going to be nitrogen dioxide.1415

Nitrogen dioxide is going to go ahead and react with water.1421

We are going to go ahead and form a strong acid, nitric acid, and a weak acid, nitrous acid.1424

Again both of them are going to go on and form hydronium.1431

Again it is the oxides of carbon, sulfur, and nitrogen that are going to react with water somehow to give us hydronium.1436

That is what we call acidic oxides.1447

Now that we have covered acidic oxides, let's go on to basic oxides.1451

Basic oxides are going to be oxides of certain metals1455

because they are going to go on and react with water and to form hydroxide.1459

The typical examples are going to be group 1 and group 2 oxides; group 1 and group 2 oxides.1465

Here I want to present to you calcium oxide.1477

Calcium oxide is going to react with water to go ahead and form calcium hydroxide.1480

Remember calcium hydroxide was one of those relatively strong bases that we talked about.1486

That is going to go ahead and form Ca²⁺ and two hydroxides.1494

The other example is the group 1 oxides.1499

Sodium oxide is going to go ahead and react with water to form sodium hydroxide.1502

We saw that sodium hydroxide was one of those strong bases also.1506

It is going to go ahead and go on to form Na¹⁺ and hydroxide.1511

We have learned the Arrhenius interpretation of an acid and base.1517

We have learned the Bronsted-Lowry interpretation.1520

Now we have learned that acids and bases will also include oxides of metals and nonmetals.1523

We now move on to the second half of this presentation on chemical reactions.1535

This type of chemical reaction is what we call oxidation-reduction reactions; which are also known as redox reactions.1542

In an oxidation-reduction reaction, basically one reactant is going to get oxidized and a second reactant is going to get reduced.1551

The most basic, the simplest interpretation of an oxidation-reduction reaction is one that involves the loss or gain of oxygen atoms.1564

That is pretty easy to remember; when you hear the word oxidation, it pretty much sounds like oxygen.1576

Let's go ahead and look at the following example.1583

We are going to take iron(III) oxide; we are going to react it with carbon monoxide.1585

We are going to form iron solid and carbon dioxide gas.1592

You see that one compound has gained oxygens and one compound has completely lost them.1597

Fe₂O₃ has lost oxygens; CO has gained oxygens.1605

Now to build upon our terminology, if a reactant gains oxygens, we say that it has become oxidized.1619

Therefore carbon monoxide became carbon dioxide; it was oxidized.1632

If a reactant is going to lose oxygen, we say that it has become reduced.1642

In this example, it is the iron(III) oxide that has become reduced.1648

Two more vocabulary terms; the oxidized reactant is what we call the reducing agent.1656

The reduced reactant is what we call the oxidizing agent.1663

Let's go ahead and move on then.1672

Again the simplest interpretation of a redox reaction is involving a loss or gain of oxygens.1677

Another interpretation involves a transfer of electrons.1683

The reducing agent is going to be the reactant that loses electrons while the oxidizing agent gains electrons.1691

This is very important that electron transfer is always from reduced agent to oxidizing agent in a redox reaction.1699

Let's go ahead and take a look at the following example.1714

Zinc solid plus Pb²⁺ is going to form is going to form Zn²⁺ and lead solid.1717

Look what happened here; zinc solid became Zn²⁺; Pb²⁺ became lead solid.1726

In other words, zinc gave up electrons and became a cation; Pb²⁺ gained electrons to get reduced.1737

In other words, look at the charge.1752

Zinc started off with a charge of 0 and it became 2+.1753

Lead 2+ started out here on the reactant side and it became 0.1760

We can easily tie this in to the activity series that we talked about in chemical reactions from the previous section.1769

Basically good reducing agents include those at the top of the activity series that are above hydrogen.1779

Basically the higher on the list, the better its reducing ability.1788

Now that we saw that we can interpret a redox reaction as a transfer of electrons, let's go ahead and talk about another one.1796

This is what we call a change in oxidation number.1806

This is going to be the less obvious one.1811

It is easy to see the change in number of oxygen atoms.1815

It is easy to see a change in ionic charge.1819

So this one we are going to spend a little more time on, takes a little getting used to.1824

Let's go ahead and identify and define what we mean by an oxidation number first.1829

Basically an oxidation number is going to be a pseudo charge.1833

It has no physical meaning and don't think of it as an ionic charge.1839

But it is a convenient way for us to keep track of electrons.1842

What you want to remember is the following.1849

During the redox reaction, an oxidation number is going to change for the compound that is involved.1851

If the oxidation number has increased, we say that the compound the element was a part of was oxidized.1859

If the oxidation number has gone down, then reduction has occurred.1868

There is a couple rules that we have to know in order to do this.1879

This is in order of priority; please remember that.1885

Rule number one: for all homonuclear species including monatomic, diatomic, and polyatomic, the oxidation number is 0 for each element.1891

For example, if you just see sodium solid, that is going to be a neutral atom sodium; its oxidation number is 0.1906

Let's look at some diatomic homonuclear species.1919

That would be Cl₂ gas, Br₂ liquid, I₂ solid, N₂ gas, H₂ gas, and F₂ gas.1923

These are what we call homonuclear diatomic compounds.1941

Each element in the compound has an oxidation number of 0.1946

There are some homonuclear polyatomic species; for example, sulfur.1953

The natural form of sulfur is actually S₈ solid.1960

In S₈ solid, all eight sulfurs, each have an oxidation number of 0.1966

That is the first rule; not too bad to remember.1975

The second rule is actually very straightforward too.1979

It tells us that for monatomic ions, the oxidation number is identical to the charge.1982

When we look at Na¹⁺ aqueous, N^3- aqueous, and O^2- aqueous, these are what we mean by monatomic ions.1989

The ionic charge is equal to the oxidation state.2004

Next rule, this deals with fluorine.2010

Fluorine is the first element that gets very high priority.2013

Fluorine is going to be having an oxidation number of -1 when part of a heteronuclear compound.2017

For example, if you look at the molecule HF, this is a heteronuclear compound.2025

The fluorine here is going to have an oxidation number of -1; pretty straightforward.2034

The next element which gets a pretty big priority is going to be oxygen.2040

Oxygen is going to have an oxidation number of -2 which is very very common.2045

For example, in water, oxygen has an oxidation number of -2.2051

However there is going to be one main exception for oxygen; this is in what we call peroxides.2063

Remember peroxides we had you memorize for polyatomic ions; an example would be Na₂O₂.2069

In this case, for sodium peroxide, each oxygen is -1 oxidation number.2090

Once again peroxides is a big exception for oxygen.2103

Going on, after fluorine and oxygen comes the remainder of the halogens.2110

The halogens, chlorine, bromine, and iodine, are also pretty high on the list of rules here.2117

They are going to get oxidation numbers of -1.2127

Coming up is hydrogen; hydrogen is almost always a +1 oxidation number when part of a heteronuclear compound.2133

For example, if you look at HF which we just dealt with.2143

Heteronuclear compound so hydrogen is going to get a +1 oxidation number.2147

There is an exception however; this is when hydrogen is going to be in a compound that contains a metal.2151

When it is directly bonded to a metal; for example, NaH.2159

In this case, hydrogen is going to get an oxidation number of -1.2165

These types of compounds are what we call metal hydrides.2171

When hydrogen is part of a compound with a metal, what we call metal hydrides.2176

The final rule is going to be something very familiar to us.2183

Remember when we learned how to name ionic compounds.2189

We said that the overall charge, the net charge, must be equal to the overall charge of the compound itself.2193

Along a parallel track, the net sum of all oxidation numbers in the compound must add up to the overall charge.2200

This is very familiar; let's look at those examples we just talked about; HF.2212

We said that hydrogen is going to get +1 and fluorine is going to get -1.2220

HF is neutral overall so we get +1 and -1 adding up to 0.2227

If you look at NaH, Na is going to get a +1 charge.2232

Remember this is an ionic compound so we go by group 1; we just go by the charge.2241

Hydrogen we said is going to get -1; once again +1 and -1 giving us 0.2246

Let's do one final look; that is going to be H₂O.2253

What get priority?--is it hydrogen or oxygen?2256

Going back to your rules, oxygen is going to get the priority.2260

That is going to have a -2 oxidation state.2263

Hydrogen here we said is going to get a +1 oxidation state when part of a heteronuclear compound.2266

However how many hydrogens are there?--there is two of them.2274

When we do the arithmetic, we are going to get +2 plus -2 adding up to 0 again.2278

This is again very familiar territory where the net sum of the oxidation numbers must equal to the overall charge of the compound.2290

Now that we have talked about the rules...2302

Unfortunately, these are rules you are just going to have to commit time to memorizing.2306

Let's go ahead and apply them to do some problems.2310

We are going to now assign oxidation numbers to all elements in each of these compounds.2315

Ammonia, nitrogen and three hydrogens; if you look at your list, hydrogen is going to get the priority.2324

Each hydrogen is +1; there is three of them giving me +3 overall.2334

Ammonia is neutral; that means nitrogen has to be high or low enough of an oxidation number to balance the +3 from hydrogen.2341

Therefore nitrogen is going to be -3 oxidation number.2352

Next one is a barium chloride.2356

We recognize this as a metal and a nonmetal; this is an ionic compound.2359

Barium is column 2; that is going to be +2.2364

Each chlorine, group 7, is -1; there is two of them.2368

2 plus -2 is going to give us 0; so we are good to go.2373

C₆H₁₂O₆, this is a molecular compound, nothing but nonmetals.2379

This is actually the formula for glucose; let's take it one by one.2384

Oxygen gets priority; each oxygen is -2; there is six of them giving me -12 overall.2391

After oxygen, hydrogen is going to get the next priority.2401

There is +1 oxidation number times 12 hydrogens giving me +12 overall.2406

What does that mean for carbon then?2418

If this whole molecule is neutral, that means carbon has an oxidation number of 0 in this compound.2423

When I add all the numbers together, 0 times 6 added to +12 added to -12, I add up to 0.2431

Finally, oxygen difluoride, this is a molecular compound; oxygen difluoride.2440

Which one is going to get priority?--fluorine is going to get priority.2449

Each fluorine is -1; there is two of them giving me -2 overall.2452

That means the oxygen has to be +2 to balance the -2 oxidation state.2459

Now that we know how to assign oxidation numbers to elements within a compound, let's apply this now to the following chemical reaction.2469

We have CH₄ plus 2O₂ goes on to form carbon dioxide and water.2482

Let's go ahead and assign oxidation numbers to each element.2489

And see if there is any changes from the reactant side to the product side.2492

Carbon is in CH₄; hydrogen is going to get priority.2498

The hydrogen is going to be +1; there is four of them which means carbon has to be -4.2504