Transcript for:
Overview of Atoms, Molecules, and Ions

Hello everybody, my name is Iman. Welcome back to my YouTube channel. Today we're going to start chapter 2 for General Chemistry 1. This chapter is titled Atoms, Molecules, and Ions.

Now in this chapter we're going to cover the following objectives. First we're going to start with a discussion on fundamental chemical laws. Here we're going to define the law of conservation of mass, the law of definite proportions, and the law of multiple proportions. Next, we're going to learn about Dalton's atomic theory.

So in 1808, Dalton published a new system of chemical philosophy in which he presented his theory of atoms that we're going to go over. Then our third objective is a continuation of some of the earlier ideas of the basic structure of an atom. We're going to compare that to our current understanding of an atom before we move into objective four. which will further build on the modern view of atomic structure. Here we're going to discuss what is an atom made of, and how do the atoms of the various elements differ.

In addition, we're going to cover a lot of important terminology here in Objective 4. We're going to talk about atomic number, mass number, atomic weight, atomic mass unit, moles, and molar mass. Then, the fifth objective is on molecules and ions, and here we're also going to learn the difference between covalent and ionic bonding. Then we're going to move to the sixth objective which is an introduction on the periodic table.

We will also cover periodic table trends. This is going to require us to understand atomic radius, effective charge, ionization energy, electronegativity, as well as electron affinity. And then last but not least, we get to move on to our last and final objective.

We're going to learn how to name simple. compounds. Now with that, let's go ahead and get started.

But before we even dive into objective one, I really want to set the stage for this chapter. I really like to do this because having a little bit of a background story really helps to hopefully pique your interest into the topics we're going to cover and why we want to cover them. Chemistry has been important since ancient times.

The Greeks were the first to try to explain why chemical changes occur. And by about 400 BC, they had proposed that all matter was composed of four fundamental substances fire, earth, water, and air. Only the avatar mastered all four elements, but when the world needed him most, he vanished and so did this idea. But the Greeks, they had considered a lot of questions that really lays the basis for some of the for chemistry.

They asked whether matter is continuous. and whether it was infinitely divisible into smaller pieces or not. They had no experiments that supported their claims, but they had these questions, they had these thoughts and ideas.

And actually the next 2000 years of chemical history were dominated by a pseudoscience called alchemy. Some alchemists were fakes who were just really obsessed with the idea of just turning cheap metal into gold, but Many alchemists were indeed serious scientists, and this period did see some important advances. The alchemists discovered several elements, and they learned to prepare the mineral acids, etc. Now, the first quote-unquote chemist to perform truly quantitative experiments was Robert Boyle, who measured the relationship between the pressure and volume of air.

And we're going to talk about... Robert Boyle a lot more when we get to our chapter on gases. Now, all his ideas weren't correct, but he was an excellent scientist, and he did lay some groundwork that was extremely important. And after him, a lot of interest in the 17th and 18th century actually centered around the phenomena of combustion.

The phenomena of combustion really evoked intense interest. in the 17th and 18th century. And by the late 18th century, combustion had been studied extensively, and the gases like carbon dioxide, nitrogen, hydrogen, and oxygen had been discovered. And it was Antoine Lavoisier, forgive my pronunciation, I'm not French and I suck at pronouncing French names, but he was a French chemist who finally explained the true nature of combustion.

And his precise measurements in chemical reactions demonstrated a really important concept that's going to allow us to get started here with our first objective on fundamental chemical laws. And that was the law of conservation of mass. He showed that combustion and respiration involved oxygen, a gas he named. Conservation of mass is really important to understand, and it essentially states that in a closed system, the total mass of the reactants, which are the substances that start a chemical reaction, equals the total mass of the products. Products are the substances that are formed by the reaction.

And so in simple terms, mass is neither created nor destroyed. In chemical reactions, it's simply transformed. We can use this example right here to kind of motivate this idea, right? What we see here is methane, CH4. It reacts with oxygen gas, O2, to form carbon dioxide, CO2, and water, H2O.

All right, this is going to be used as an example to illustrate the law of conservation of mass. So we have one molecule of methane, CH4. That's going to be 16 grams.

This is just adding up the individual weights of all the atoms that make up CH4 here. Carbon 12, hydrogen 1, and there's four of them, so total of 16 grams. Then we have two O2 molecules, all right? O2, that's two oxygens.

Each oxygen is 16. For a total of 32, that's for one O2 molecule. And if you have two, that's 64 grams. So for the reactants, we have a total of... 80 grams. The total mass before the reaction is 80 grams.

That's 16 grams of methane plus 64 grams of oxygen. After the reaction is complete, these reactants are going to form carbon dioxide and two water molecules. Now, if you add up the weights of those, you're going to get 44 grams for 1 CO2 and 36 grams for 2 H2O molecules, and lo and behold, 44 plus 36 equals 80 grams.

The mass of these products combined will also be 80 grams, demonstrating the law of conservation of mass. Now, this law holds for all chemical reactions, showing that mass is conserved throughout the process. All right, no mass lost or gained in the reaction. Now, building on... Antoine's principles, Joseph Proust, another French chemist, established the law of definite proportions.

How? Let's talk about that next. He established the law of definite proportion by demonstrating that compounds always contain the same proportion of elements by mass.

His findings actually influenced John Dalton, who proposed that atoms As indivisible particles make up elements, this atomic theory explained why compounds always have consistent relative masses of their constituent elements. But Dalton discovered another principle that convinced him even more of the existence of atoms, and that was the law of multiple proportions. The law of multiple proportions. This said, when two elements form a series of compounds, the ratios of the masses of the second element that combine with one gram of the first element can always be reduced to small whole numbers.

So let's kind of demonstrate this with a problem. Here we have data that was collected for several compounds of nitrogen and oxygen. So you have compound A, B, and C, and here you have the mass of nitrogen that combines with one gram of oxygen. So in compound A, you have 1.75 grams of nitrogen that combines with one gram of oxygen.

In compound B, you have 0.875 grams of nitrogen that combines with one gram of oxygen. And C, you have 0.4375 grams of nitrogen that combines with one gram of oxygen. Now to demonstrate the law of multiple proportions, we can start comparing these compounds. We can compare compound A to B, for example. The mass of nitrogen combining with one gram of oxygen in compound A is 1.750 grams, and that for B is 0.8750 grams.

All right, this is exactly two to one. All right, this gives us a two over one relationship. And so the mass of nitrogen combining with one gram of oxygen in compound A is exactly double that in compound. B. Now we could do the same.

We can compare B and C. For example, this is going to be 0.8750 grams over 0.4375 grams. And this also gives us a two over one relationship. So It is also double here.

The mass of nitrogen combining with one gram of oxygen in compound B is exactly double that in compound C. We can do another comparison. We can do A and C. This is going to be 1.750 grams over 0.4375 grams.

This is going to give us a 4 over 1 relationship. So the mass of nitrogen combining with one gram of oxygen in compound A is exactly four times that in compound C. And what you notice is these ratios between the different compounds, two to one, two to one, four to one, are all small whole numbers, which perfectly illustrates the law of multiple proportions.

And just to kind of summarize that, I want to restate the definition once more. The law of multiple proportions is a fundamental principle in chemistry. Stating that when two elements combine to form more than one compound, the masses of one element that combine with a fixed mass of the other are in a ratio of small whole numbers.

Now with that discussion, we can go ahead and move into objective two, which is on Dalton's atomic theory. In 1808, Dalton published a system of chemical philosophy in which he presented his theory. theory of atoms and this is what he wrote.

First, he said each element is made up of tiny particles called atoms. Second, the atoms of a given element are identical and the atoms of different elements are different in some fundamental way or ways. Three, chemical compounds are formed when atoms of different elements combine with each other. A given compound always has the same relative number of atoms.

and types of atoms. And last but not least, chemical reactions involve reorganization of the atoms, so changes in the way they are bound together. The atoms themselves are not changed in a chemical reaction. John Dalton really viewed atoms as the indivisible building blocks of matter.

He viewed them kind of like billiard balls, atoms of different elements, simply differed from each other by mass and nothing else. they could not be subdivided any further. It clearly did not account for the subatomic particles that make up atoms that was discovered later. And this is the perfect segue into objective three, where we want to talk about all the early ideas and experiments that aim to understand and characterize atoms. The concept of atoms is really interesting and inevitably scientists began to wonder.

about the nature of the atom. What is an atom made of? How do the atoms of the various elements differ? So let's start that conversation.

We now know that an atom is the smallest unit of ordinary matter that forms a chemical element. It's composed of a nucleus that's made out of protons, positively charged, and neutrons, which have no charge. Electrons which are negatively charged, orbit the nucleus. The protons and neutrons define the mass of the atom, while the arrangement and movement of electrons determines the atom's chemical properties and reactivity. How did we reach this conclusion, right?

It was the cumulative work of many scientists that has allowed us to understand the basic structure of an atom. We just need mention. of John Dalton who proposed his atomic theory and how he viewed atoms as sort of like billiard balls.

Atoms of different elements simply differed from each other by mass and nothing else and they could not be subdivided any further. So this is kind of our starting point here. It wasn't until J.J. Thomson's work with cathode ray tubes in 1898 that negatively charged particles called electrons were found to be present in atoms. And so Thompson's model is usually called the plum pudding model.

You have these scattered negative charges around in an atom. Then in 1911, Ernest Rutherford further refined our view of the atom by showing that although atoms consist of mostly empty space, at the center of every atom is an extremely small and dense nucleus which is positively charged. He was able to make these conclusions based off of his gold foil experiment. So then we had Rutherford model. Center, nucleus, positively charged, and then the negative charges, the electrons, kind of surrounded this extremely small and dense nucleus.

Then Bohr came along and he developed his model, which suggests that negative charge orbits the nucleus of an atom, which is made out of protons and neutrons. There is more to that, truly, as we can't really... pinpoint where electrons are.

We can only really talk about electron density probability, where the electron's most likely to be found, but we're going to dive into that later. So essentially, we know that atoms have a nucleus made out of protons and neutrons, and surrounding that are electrons. In our quantum mechanical model, our latest model, the electrons kind of exist in a cloud of probability.

around the nucleus. We don't really want to think about electrons as quote-unquote orbiting the nucleus, but more so as existing as a cloud of probability, electron density probability around the nucleus. We're not going to get too much into that until maybe later, because it is a complicated topic.

But now that we understand that, I do want to focus on the subatomic particles in more detail. We've said that an atom the smallest identifiable unit of an element. So in our fourth objective, we're going to focus on that modern view of atomic structure.

We have an atom. We talked about how it contains a nucleus with protons. These are subatomic particles with a positive charge.

Neutrons, subatomic particles with no electric charge, they make up the nucleus. And then we have electrons, subatomic particles with negative charge. Now, in addition to that, we should know some quantifiable information about protons, neutrons, and electrons. Here we have a table with some of that important information. Proton, the mass in kilograms, all right, 1.672 times 10 to the minus 27 kilograms, or the mass in AMU, atomic mass units, we'll talk about that in a second, is about 1.00727.

The charge for proton is 1.602 times 10 to the minus 19. And the relative charge is plus 1. So this is the charge in coulombs. Neutrons, here is the mass in kilograms. The mass in AMU, it has zero charge.

And then we have the electron. It's the smallest of the three, as you can obviously tell by the mass. 9.109 times 10 to the minus 31 kilograms. In AMUs, it's 0.000548.

The charge is the same as the proton, but with a negative sign. So protons had positive charge, electrons have a negative charge. Now, something to know about electrons is that they move around the nucleus at varying distances away from the nucleus. We're going to get into a lot of the details about this a little later on. But some few points that I want you to keep in mind, let them marinate in your subconscious until we get to it later on.

But the electrons that are closer to the nucleus are at lower energy levels, and those that are further out in higher shells are going to have higher energy. And while we're going to revisit this in more detail, we should know that the electrons furthest away from the nucleus are going to have the strongest interactions with the surrounding environment. You can think about the electrons further away from the nucleus as being kids that are not as close to their parents, and therefore they have a stronger interaction with their surroundings, their friends, their environment, their school, etc. Just as an analogy, if you will.

They're going to have the strongest interactions with the surrounding environment and the weakest interactions with the nucleus since it's further away. These electrons are called valence electrons. They are the electrons that are going to participate in chemical reactions, and they are more likely to participate in bonds. And this is a central idea that you're going to carry even into organic chemistry. Now that we understand the atom, let's talk about a couple of important terms related to the atom.

So atomic structure is defined by three critical numerical descriptors. We have atomic number, mass number, and atomic weight. So typically if you're looking at an element on the periodic table for example you'll see something like this. You have the elemental symbol for oxygen for example, the name of the element at the top here is your atomic number, at the bottom here is your mass number.

But in addition, you might see this information communicated in other ways. For example, in this format where you have your elemental symbol, and then at the top left corner, you have your mass number and your bottom left corner, your atomic number. So oxygen could look like this on the periodic table, but sometimes that same information can be communicated like this in writing. Regardless of the format, I want to define some of these terms. So atomic number.

denoted by the symbol Z is the count of protons in the nucleus of an atom. This number is fundamental because it determines the chemical identity of the atom. Each element on the periodic table is characterized by a unique atomic number.

Then you have the mass number. So let's do this in a different color. The mass number denoted by A.

So you could see that here highlighted in red. This is the total number of protons and neutrons in an atom's nucleus. Neutrons and protons are collectively known as nucleons, and their sum gives us the mass number. Now, the difference between the mass number and the atomic number is the number of neutrons in the nucleus.

It's important to note that the mass number is not the same as the atomic weight. Atomic weight a more nuanced concept. It is the weighted average of the masses of all naturally occurring isotopes of an element and it's measured in atomic mass units. Since isotopes of an element have same number of protons but varying numbers of neutrons, they also have different mass numbers but of course they'll have the same atomic number.

So Atomic weight takes into account the relative abundance of each isotope in nature in the calculation. And that kind of makes it important for us to really define what an isotope is. So let's quickly define that and then let's go back to this topic of atomic weight to make sure that we are all on the same page. Isotopes are variants of elements that have the same number of protons.

but different number of neutrons and that leads to different mass numbers. They're significant in both natural processes and technological applications. So for instance, stable isotopes are used in medical diagnostics, whereas radioactive isotopes can be applied in treatment or to help date archaeological finds or as tracers in biochemical research. Understanding atomic structure and isotopes is really important for grasping the bigger picture of chemical behavior. So with that being said, let's go ahead and go back to atomic weight now that we have defined isotopes.

All right, we defined atomic weight and we have it right here for us for our convenience as the weighted average of the masses of all. naturally occurring isotopes of an element measured in atomic mass units. So let's really break this down because this calculation is not super straightforward.

It takes into account, all right, it has to take into account the fact that most elements exist naturally as a mixture of isotopes, each with its own mass number and natural abundance. And so to calculate atomic weight, Scientists consider both the mass and the natural abundance of each isotope. And here's how that works.

You have to consider one. All right, we're going to make a little list over here. One, you have to consider the isotope masses. All right, first, the exact mass of each isotope of the element is determined. These masses are close to whole numbers, but are slightly different.

All right, then you have to consider two, natural abundance, natural abundance. You have to know each isotope's natural abundance. This is the percentage of each isotope that occurs naturally on Earth.

And then you can go ahead and calculate that weighted. average. The atomic weight is then calculated by multiplying the mass of each isotope by its natural abundance and then adding these values together.

The result is a weighted average that represents the average mass of all isotopes that occur naturally. Now the best way to understand this is by doing a quick example. So what I'm going to do is I'm going to scroll down here really quickly and I have a problem.

that relates to what we just discussed about atomic weight that we're going to do together. So in this problem, we say we have an element. This element has two isotopes. Isotope A has a mass of 10 amu, and we're going to cover what amu is in a second, but let's just say that this is a measure of mass here.

All right, isotope A with a mass of 10 amu and a natural abundance of 90%, and then we have isotope B with a mass of 11 amu. and a natural abundance of 10%. What is the atomic weight? So based off of our verbal description of atomic weight, we said that it would be calculated like so. We have our first isotope, isotope A, okay?

It's 10 amu in mass, and it has a natural abundance of 90%, and we want to express 90% as a decimal, so that's 0.90, okay? That's for isotope A, and we're going to have to add this to isotope b all right isotope b 11 amu and it has a natural abundance of 10 so 0.10 all right the atomic weight can then be calculated as follows we just follow through with this mathematical expression right here plug this into a calculator follow pemdas rules okay and what we're going to get is that this is going to equal to 9 plus let me use my pen 9 you plus 1.1, and that gives us 10.1 amu. And so the atomic weight of this element, all right, that has two isotopes is 10.1 amu, all right? So again, the key to understanding atomic weight is to recognize that it's an average that reflects isotopic composition and abundance.

It's not the mass of a single atom, but an average mass of all atoms. of that element as they are found in nature. Wonderful. Now let's scroll back up because we skipped a few things to get to that problem that we still want to go ahead and cover.

So we've talked about atomic number, mass number, and atomic weight. There are a couple more concepts that are going to be super crucial for us to understand before we move on. And those are the concepts of mass, atomic mass unit, and moles. So let's go ahead and cover those now.

Let's start with atomic mass unit. The atomic mass unit is a standard unit of mass that quantifies the mass of atoms or molecules, and it's defined as one twelfth the mass of a carbon-12 atom. So let me repeat that. It's defined as one twelfth the mass of a carbon-12 atom. This is an isotope of carbon with six protons and six neutrons.

It's stable. and its abundance makes it a standard for measuring atomic masses. One atomic mass unit is approximately 1.66 times 10 to the minus 27. Minus 27, I want to make that clear, kilograms. Now the atomic mass of an element that's usually found on the periodic table is expressed in atomic mass unit, and it represents the average mass of all the isotopes of that element, taking into account their relative abundance. The next thing I want to talk about is moles.

A mole is a unit that measures the amount of substance. So like one dozen signals that you have 12 things, right? If you say you have one dozen eggs, you have 12 eggs.

You have one dozen books, you have 12 books. One dozen magazines, you have 12 magazines. Same thing for a mole.

It is a unit that measures the amount of a substance. One mole of any substance. contains exactly 6.022 times 10 to the 23 entities.

This could be atoms, molecules, ions, etc. This number is known as Avogadro's number. The mole allows chemists to count particles by weighing them and the mass of one mole of a substance is equal to its molecular or atomic mass in grams.

And last but not least, I want to talk about molar mass. This is The mass of one mole of a substance, it's expressed in units of grams per mole. Now that we've defined these three terms, Atomic mass unit, moles, molar mass. How do we connect them all together?

And by that, I mean, how do we convert from one form into another? This is going to be really important in general chemistry. So let's talk about it.

Let's talk about how we can connect these concepts. Let's say to start off, you know the number of particles. You know the number of particles, and you want to convert that to moles.

How do you go? from number of particles to moles. You do this by dividing by Avogadro's number.

All right, you do this by dividing by Avogadro's number. So let me demonstrate this a little bit better. All right, you divide, divide by Avogadro's number.

Okay, that's going from number of particles to moles and that's kind of written up here at the top as well. Okay, what about if you know the number of moles but now you want number of particles, right? So you're going in the opposite direction. Now you know moles, you want number of particles.

Here now you're going to have to multiply by Avogadro's number. Fantastic. So we got that conversion, number of particles to moles or moles to number of particles.

We know that the secret lies in... Avogadro's number. Phenomenal. What if you know moles, all right, and you want to convert to mass?

You want to go from moles to mass. How do you do that? In order to do that, you're going to multiply by molar mass. So you're going to multiply by grams per mole. And if you want to go the other way, if you know the mass and you want the moles, you're going to divide by molar mass.

So you're going to think mole over grams. Good. So far, so good.

Here's something else. What if you know the mass and you want to convert, all right, you want to go from mass to number of particles. What do you do now?

All right, this is going to involve two steps, and it's the two steps that we've talked about going in the same direction. First, you're going to divide by molar mass. All right, so step one is divide by molar mass.

And then the second step is to multiply by Avogadro's number. Multiply, I can't spell, by Avogadro's number. So if you're going from mass to number of particles, two steps, you divide by molar mass, then you multiply by Avogadro's number.

Okay, last thing that we want to cover now is, you guessed it, what if we want to go from number of particles to mass? Again, two steps. First step, you... divide by Avogadro's number, and then the second step you multiply by molar mass.

So this is kind of like a flow chart, a roadmap of how to convert between number of particles, moles, and mass. Let's do a quick example to make sure that we understand this concept fully. So this is the problem right here. Suppose you have 3.011 times 10 to the 23. particles of argon, what is the mass of the argon sample?

So here's our goal. We know the number of particles. We want to go to mass. So this is the path we want to follow. This one right here, this is going to require us to divide by Avogadro's number and then multiply by molar mass.

Now the molar mass of argon is going to be 39.95. grams per mole. And we're just going to go ahead and kind of set up this dimensional analysis. We're starting off. What do we know?

We know that we have, let's pick a color. Let's do green. 3.011 times 10 to the 23 particles of argon. All right.

We want to convert this to mass. Ultimately, first step is divide by Avogadro's number. So Avogadro's number is 6.022 times 10 to the 23 particles, all right, is equal to one mole.

Phenomenal. So this is our first conversion, okay? Notice how mole by Avogadro's number, that's exactly what we have here. Mole by Avogadro's number.

We have particles here because that's what we're dealing with, particles of argon. And the units here cancel out. Good.

Now, what's the next step? The next step is to multiply by molar mass. So what's the molar mass of argon? It's just 39.95 grams per mole.

And notice again, mole cancels out. And that means our final answer is going to have units of... grams.

And that's exactly what we want because we're trying to figure out what the mass of this argon sample is. And then you just go ahead and calculate this. You should get about 19.98 grams of argon. And there we have it.

We did it. Now, in summary, we've talked about a lot here in objective four. I want to scroll back up and just show you how much we've talked about.

We talked about protons, neutrons, and electrons, the subatomic particles that make up an atom. We talked about atomic number and mass number and how to identify it in different forms, in different writing styles, if you will. And then we had a little bit of a tangent on how mass number is not equal to atomic weight. And so we were motivated to talk about atomic weight, which also required us to talk about isotopes. And we learned how to calculate atomic weight and that atomic weight is defined as the weighted average.

of the isotope masses of the elements that are naturally occurring isotopes. And then we took that and we continued to talk about a couple of other important terminology, atomic mass unit, moles, and molar mass. And we learned how to convert between different units, going from number of particles to moles to mass and back and forth and in between.

And in the process, we did two example problems as well. Now, We can confidently move into our fifth objective. Our fifth objective is about molecules and ions. Something important to remember is that in ordinary chemical reactions, the nucleus doesn't change, but the number of electrons that an atom possesses can change.

The only time that the number of protons or neutrons possessed by an atom change is during nuclear reactions. We're not talking about that here. Now it's safe to say that the movement of electrons constitute reactions, and this will be the focus of most discussions with reactions in general and organic chemistry.

Now it's important to pay close attention to whether a particle is charged or not, if a particle is a charged ion or a neutral atom, because an ion possesses properties that are quite different from those of a neutral molecule. An ion is a charged particle resulting from the fact that the number of protons is not equal to the number of electrons. So when the number of protons equals the number of electrons, we're dealing with a neutral molecule. But when the number of particles does not equal the number of electrons, then we are dealing with ions. There are two types of ions.

We have cations and anions. A positively charged atom is called a cation and a negatively charged atom is called an anion. So for cations, they're positively charged because the number of protons is greater than the number of electrons.

So you've lost an electron. So hence cations are positively charged. Anions on the other hand, you have a greater number of electrons than protons. You've gained an electron. So they are negatively charged.

Again, it's important to pay attention to whether a particle or molecule is charged or not because an ion possesses properties that are quite different from those of a neutral atom. Understanding ions allows us to discuss molecules and ionic compounds. Molecules are distinct. electrically neutral groups of atoms that are held together by chemical bonds, specifically covalent bonds where atoms share electrons.

We'll get into the details of that here just in the next page. The structure and properties of a molecule, they're determined by the specific atoms it contains and how those atoms are bonded. A molecule's formula indicates the types and numbers of atoms present using element symbols. and subscripts to denote their ratios.

On the other hand, in contrast, ionic compounds, they consist mainly of oppositely charged ions that are held together by ionic bonds. These are electrostatic forces between the ions. Again, we're also going to get into that here in the next page.

These compounds typically form highly ordered three-dimensional crystal lattice structures. The formula of an ionic compound represents the smallest neutral unit of a compound, indicating the ratio of cations to anions. So again, the ratio of cations, positively charged ions, to anions, negatively charged ions.

So for example, a good example is table salt. Table salt is sodium chloride, NaCl. This is an ionic compound comprising of a one-to-one ratio of sodium cations to chloride anions. And they don't exist like this, like a typical traditional covalent bond.

This is an ionic compound. And so it exists in a three-dimensional crystal lattice where you have these alternating bonds between sodium cations and chloride anions. to form a specific crystal lattice shape, but that is an inorganic topic that we're not going to get into here.

Now, molecules form when atoms are held together by covalent bonds. This involves the sharing of electron pairs between atoms. On the other hand, ionic compounds form ionic bonds, which occur when electrons are transferred from one atom to another, resulting in positively charged ions, cations, and negatively charged ions, anions, that attract each other due to their opposite charges. How can we distinguish between the two?

So now, this is the perfect segue into talking about the different kinds of bonding, covalent and ionic bonding. In order to do this, we're going to have to define a really important... important topic called electronegativity. Electronegativity is a chemical property that describes the ability of an atom to attract electrons toward itself when it is part of a compound.

It's a dimensionless quantity, often represented on the periodic table with values that typically range anywhere between 0.7 to 4.0, which is a scale of lower to higher electronegative elements respectively. So a lower number, less electronegative, higher number, more electronegative. Now understanding electronegativity is vital because it helps us predict the type and the strength of bonds that are going to form between atoms.

Now depending on your course here, you might be required to memorize some values for electronegativities. But in addition here in the next objective... We're going to talk about periodic table and periodic table trends, and we're going to learn the trend for electronegativity. So sometimes you can typically follow the trends to be able to follow the flowchart that we're going to discuss next to determine what kind of bonding is occurring between two atoms. Now, the concept of bonding is really central to chemistry.

It explains how atoms bind together to form molecules. Bonds can be... broadly classified as ionic or as covalent.

And this depends on the electronegativity difference between the bonding atoms. So let's go ahead and get started first with this branch, the covalent bonding branch. Covalent bonding occurs when two atoms share a pair of electrons and that shared electrons allows each atom to attain the electron configuration of a noble gas leading to a more stable molecule.

Covalent bonds, they can be polar or nonpolar. So let's talk about each of these. Nonpolar covalent bonds, these bonds form between atoms with similar electronegativities. Usually the electronegativity difference is less than 0.5. Since the electron pull is almost equal, the electrons here in nonpolar covalent bonds are shared fairly equally.

So here's a good example. If you have a carbon bonded to a carbon, if you're trying to decide what kind of bond that is, you would take the electronegativity value for carbon and the electronegativity value for carbon, and you would subtract the two from each other. Now, the electronegativity value for carbon is 2.6.

essentially what you're doing here is 2.6 minus 2.6. That equals zero. That is less than 0.5.

So that is one way you can conclude that the bond between a carbon and a carbon atom is a non-polar covalent bond. Another one is whenever you have bonds between the same atoms, it makes sense that their electronegativity difference is going to be zero. And if it's zero, then what you have is a non-polar covalent bond. Let's do another example.

What about carbon and hydrogen? Carbon is 2.6, and hydrogen, let's go back to our table, hydrogen should have a, it's not on here, but I'll tell you, hydrogen should have an electronegativity value of about 2. 2.1. All right, so about 2.1. So what we have is 2.6 minus 2.1. That's about 0.5.

Traditionally, something on the border like this typically falls under nonpolar covalent bonding here. Anyways, carbon-hydrogen bonds, those are nonpolar covalent bonds. Then we have polar covalent bonds. Okay, these bonds occur when there's a moderate difference in electronegativities. greater than 0.5, but less than 1.7.

All right, if the electronegativity differences falls in this range, then what you have is a polar covalent bond. The more electronegative atom attracts the shared electrons more strongly. What this leads to is a partial negative charge on that atom and a partial positive charge on the other.

One good example is Carbon and oxygen. Carbon has an electronegativity value of 2.6. Oxygen, let's scroll back up here, 3.5.

So 3.5 minus 2.6, this is equal to 0.9. This falls in this range between 0.5 and 1.7. So the bond between carbon and oxygen is a polar covalent bond.

Now, which one's more electronegative? Oxygen is because... its electronegativity value is 3.5.

What does that mean? That means oxygen pulls those electrons, those shared electrons, a little more towards it because it's more electronegative. It likes those electrons just a little more than carbon. What that results in is a partial negative charge on that oxygen since it's pulling those electrons towards it and a partial positive charge on the carbon. So that is polar covalent bonds.

Okay. So that is our covalent branch. What about ionic bonding? Ionic bonds, when the difference in electronegativity is significant, so think electronegativity difference greater than or equal to 1.7, ionic bonding is likely. In an ionic bond, one atom, usually a metal, loses one or more electrons, it becomes a positively charged ion, while the other atom, usually a nonmetal, gains those electrons and becomes a negatively charged ion.

And this electrostatic attraction between the oppositely charged ions forms a strong ionic bond. And sodium chloride is a classic example of an ionic compound. So if we look up here, sodium has electronegativity value of 1 and chloride 3.2.

So 3.2 minus 1, that's a difference of... 2.2. That is greater than 1.7. This is an ionic bond.

Again, sodium chloride, great example, sodium, positive cation, chloride, and ion. So that is bonding for us. We talked about electronegativity, measure of the ability of an atom to attract electrons.

And we use the difference in electronegativity values between two atoms to determine what kind of bonding is. occurring. All right, look at us learning so much chemistry. Let's move on to the sixth objective.

This objective is an introduction to the periodic table. Here we see a chemist's favorite tool, the periodic table of elements. We're going to continuously refer to the periodic table in this course, and we're going to build our knowledge. But to start, let's talk a little bit of history.

In 1869, the Russian chemist Dmitry Mendeleev published the first version of the periodic table of elements and he showed that ordering the known elements according to atomic weight revealed a pattern of periodically recurring physical and chemical properties. Since then the periodic table has been revised using the work of physicist Henry Mosley to organize the elements based on increasing atomic number, which is the number of protons in an element rather than atomic weight. And using this revised table Many properties of elements that had not yet been discovered could be predicted. The periodic table created a visual representation of the periodic law, which states the chemical and physical properties of the elements are dependent in a periodic way upon their atomic numbers.

Now, the modern periodic table arranges the elements in two periods or rows and in two columns. or groups. Now, each period, all right, talking about periods, each period.

is filled sequentially and each element in a given period has one more proton and one more electron than the element to its left. So this atom right here has this element right here has one more proton and one more electron than the element before it and so on and so forth. Now as for groups, groups contain elements that have the same electron configuration in their valence shell.

and they share similar chemical properties. We're going to talk about electron configuration in a lot more detail in a different chapter. But with this preliminary information, we want to talk a little bit about the different element types that we're going to see and discuss. And there are three element types we're concerned about here, metals, nonmetals, and metalloids.

Metals are found on the left side and in the middle of the periodic. table. They include the active metals, the transition metals, and also don't forget these two lonely rows at the bottom, the lanthanides and the actinides. Metals are shiny solids, except for mercury, by the way, which is a liquid under standard conditions.

They generally have high melting points and densities. There are exceptions like lithium, which has the density about half that of water. Metals have the ability to be deformed without breaking.

The ability of metal to be hammered into shapes is called malleability and the ability to pull metals or draw them into wires is called ductility. Many of the transition metals, which are these metals right here by the way, are going to have two or more oxidation states, so charges when forming bonds with other atoms. And because the valence electrons of all metals are are only loosely held to their atoms, they're free to move, which makes metals really good conductors of heat and electricity. So that's metals for us. What about non-metals?

Non-metals are found predominantly on the upper right side of the periodic table. So over here, they have, they're generally brittle. in the solid state, they show little or no metallic luster. They are poor conductors of heat and electricity.

And all of these characteristics are manifestations of the inability of non-metals to easily give up electrons. Non-metals are also less unified in their chemical and physical properties than metals are. Separating the metals and non-metals in this stair-step group right here that I've attempted to draw in black. So the elements that surround this stair step right here, these are called metalloids. The metalloids are also called semi-metals because they share characteristics with both metals and non-metals.

The electronegativities and ionization energies of the metalloids, they're going to lie between metals and non-metals and their physical properties like density, melting point, boiling point, They vary widely and can be a combination of metallic and non-metallic characteristics. Now what we're kind of equipped to do is begin our discussion on periodic trends. But to have a productive conversation on periodic trends, we're going to have to define a couple of really important words. We're going to have to define atomic radius, effective charge.

ionization energy, electronegativity, and electron affinity. First, let's talk about atomic radius. The atomic radius is a term that's used to describe the size of an atom. It's defined as the average distance from the center of the nucleus to the outermost boundary of the electron cloud.

This average distance, it's not a fixed physical boundary, because remember,... electrons exist in a region of probability rather than in a specific orbital path. However, the definition I want you to keep in mind for now is that the atomic radius is the distance from the center of the nucleus to the outermost boundary of the electron cloud.

Second, we want to talk about effective nuclear charge. Effective nuclear charge is the actual amount of positive charge experienced by an electron in a multi-electron atom. It is the charge that's felt by an electron when both the actual nuclear charge and the repulsive effects of other electrons are considered.

This is a really important concept. It's quantified. as the nuclear charge minus the shielding effect due to inner shell electrons.

Now this can be estimated using Slater's rules. We might talk about this a little more down the line, but in short, the effective nuclear charge is the net positive charge that valence electrons experience from the nucleus. It is less than the actual number of protons in the nucleus because those inner electrons partially shield.

the outer electrons from the full positive charge. So in simple terms, I suppose you can think of it as number of protons minus the core electrons that are getting in the way. So that gives you kind of that actual amount of positive charge that's experienced by this electron in particular.

All the number of protons in here and all the electrons in the way. Then next up, we have ionization energy. This is a critical atomic property that quantifies the energy required to remove an electron from a gaseous atom or ion. This is an endothermic process, meaning that energy is absorbed to overcome the electrostatic attraction between that negatively charged electron and the positively charged nucleus. Ionization energy is a key indicator of an element's reactivity.

and it is influenced by atomic radius as well as effective nuclear charge. Again, it is the energy required to remove an electron. Next up is electronegativity.

We've briefly defined this earlier. It is a measure of an atom's ability to attract and retain shared electrons in a chemical bond. Now, electronegativity, chemical property.

that describes the tendency of an atom to attract electrons towards itself. It's influenced by elements atomic number, as well as the distance that its valence electrons reside from the charged nucleus. So again, measure of the atom's ability to attract and bind with electrons in a chemical bond.

And then the last thing that we want to talk about is electron affinity. This is the measure of the energy change when an electron is added to a neutral atom, forming a negative ion. So it's kind of the opposite of ionization energy.

This is the energy required to remove an electron. Electron affinity is the amount of energy released when an electron is added to a neutral atom or molecule to form a negative ion. Now these... Concepts are extremely important in the study of chemical periodicity.

They each describe different but interrelated aspects of atomic behavior. And now what we want to do is take these ideas and dive into periodic trends. What is the trend for each of those properties going from left to right?

What about going from bottom to top? Let's start with atomic radius first. What is the trend for atomic radius? Going from left to right on the periodic table. The atomic radius is going to decrease.

This is due to the increase in the number of protons, which enhances the effective nuclear charge, thereby drawing those valence electrons closer to the nucleus. What about atomic radius going from bottom to top on the periodic table? Here it's going to decrease. decrease as you ascend a group. Fewer electron shells are present, which reduces the shielding effect and it allows the nucleus to exert a stronger pull on the remaining electron shells.

Next up, effective nuclear charge. What's the trend for effective nuclear charge going from left to right? It is going to increase. As you move across a period, The number of protons in the nucleus increases. That should increase the nuclear charge.

However, electrons are also being added to the same energy level, so their ability to shield each other from the nucleus, that doesn't increase proportionally. Therefore, the actual increase in positive charge felt by the electrons is going to be greater, thereby enhancing the effective nuclear charge. So it increases going from left to right. What about from bottom?

to top. The effective nuclear charge is also going to increase. going from bottom to top, so moving up a group.

Electrons in lower shells are going to shield the outer electrons less effectively as you move up, because the inner electron shells are going to be closer to the nucleus and thus exert a stronger pull on the outer electrons. And so this results in an increase in the effective nuclear charge for the valence electrons since they experience a more significant attraction to the nucleus. Next up, ionization energy.

What's the trend for ionization energy moving from left to right? It is going to increase. The electrons are held more tightly by the stronger effective nuclear charge, and so it requires more energy for electrons to be removed.

And what about going from bottom to top for ionization energy? It is going to increase moving upward. Fewer electron shells are present. The electrons are going to be held closer to the nucleus, and so increasing the energy is needed to ionize them. Then we have electronegativity.

What is the trend for electronegativity going from left to right? It is going to increase due to the greater effective nuclear charge, which allows atoms to attract bonding electrons more strongly. And then for electronegativity going from bottom to top, it is also going to increase because atoms higher up in the group have fewer electron shells, resulting in a stronger attraction between nucleus and the valence electrons.

And then last but certainly not least, we have electron affinity. What is the trend going from left to right for electron affinity? It is going to increase.

This means the atoms are more likely to accept additional electrons, and this trend is because the atomic nucleus becomes more positively charged due to the addition of protons, which increases the effective nuclear charge without a proportional increase in shielding effects, since those additional electrons are being added to the same energy level, and so as a result, the added electrons feel a stronger attraction to the nucleus. What about going from bottom to top? Electron affinity increases going from bottom to top. So as the valence shell gets closer to the nucleus. So with that, we've covered our periodic table trends.

And now we've actually made it to our last and final objective for this chapter, naming. In this section, we're going to specify the most important. rules for naming compounds other than organic compounds.

Organic compounds is a worry for organic chemistry. So we're not going to concern ourselves with that here. Instead, we're going to begin with the systems for naming inorganic binary compounds. These are compounds that are composed of two elements, and we're going to classify this into various types for easier recognition.

Now, the first thing to kind of set the stage is... Let's talk about how do we determine if a substance is ionic or molecular because that's going to be really important for naming chemical compounds and knowing which rules to follow. So if you're looking at a binary compound and both elements are non-metals, then the compound is molecular. Water is a great example of this.

But if one element is a metal and the other is a non-metal, then the compound is ionic. And here sodium chloride is a really good example. Now in addition to that, I also want you to make yourself accommodated.

Please know these common monatomic cations and anions and their names. So here we have a list of cations and their names, H+, hydrogen, Li+, that's lithium, Na+, that's sodium, so on and so forth, Ag+, that's silver. And then on the other side we have anions, H-, hydride, Cl-chloride, so on and so forth. Now, something else that's important to keep in mind that'll make remembering some of these names easier is making flashcards and or recognizing that, at least for the charges that are associated with each of these cations, you can kind of extract that information from the periodic table.

So what I'm going to do is I'm going to scroll back up here where we had an image of the periodic table. And what you noticed in the common names for monatomic cations specifically is any of the elements that you find here in group one are going to have plus one charges. The metals that you're going to find here in group two are going to have plus two charges and usually some of these elements in group three are going to have plus three charges like aluminum and gallium. So keeping that in mind, that at least makes it easier to associate the charge with the ion. But of course, also, please know the names and the same goes for anions.

There's one more I want to add here just to keep in mind as we go through the lecture and we'll go over it again. OH minus, this is called hydroxide. So also keep that in mind through this lecture because we're going to make mention of this anion.

Now, let's go ahead and get started with all that preliminary fluff we can get to. talking first about binary compounds. Now, type one, binary ionic compounds, they contain a positive ion, a cation, that's always written first in the formula, and then it's followed by a negative ion or an anion.

In naming these compounds, we're going to follow a couple of rules. The first rule is that the cation is always named first, and then the anion. The anion is second.

Now, how do we name those different parts? How do we name the cation? How do we name the anion? That's what the next rules are about. The second rule says a monatomic cation.

It takes its name from the name of the element. So for example, Na plus, that's just called sodium, and it's called sodium in the names of compounds that contain this ion. The third rule, now it addresses monatomic anions.

They're named by taking the root of the element name and then adding ide so the cl minus ion is called chloride right so cl is chlorine and then this anion is now we take that root name and we add ide so chloride and we can see a couple of these examples right here here's a compound sodium chloride we have the cation sodium so we keep it just like that as per rule two and then here we have chlorine it's the anion we take the root of the element name, and then we add I-D-E, and that's why the name is sodium chloride. Now, I want to interject right here and have a little bit of a tangent on something important that will reappear in just a few moments. But here, if we're looking at sodium chloride, what makes up sodium chloride? We have our sodium cation and our chlorine or chloride anion.

How do we know that by taking... sodium, and chloride that we can get NaCl and in this kind of ratio, one sodium to one chloride. Well, one way to do this is by crisscrossing the charges to give you the ratio of how many cations and how many anions you need to give the chemical compound that you see here. All right. And the purpose of this is because...

For ionic compounds, the overall formula of an ionic compound is charge neutral. So you want to balance the charges and you can do that by crisscrossing the value that's associated in the charge with the other ion that makes up the compound. So here we take this one, we assign it to the chloride ion, and we take this one from the chloride ion, assign it to sodium, and so now we know we have one sodium and one chloride. Let's do another example of compounds and naming them and then also working through this kind of ratio and how the overall formula of an ionic compound is charge neutral.

How we assure that that's true when we're working through this in either direction, going from the anions to the compound, going from the compound to the name, etc. So let's do this one. This one's interesting.

Li3N. Okay, let's break this up. What are the ions that make up this compound?

Well, we obviously have lithium. Lithium has a plus one charge, and nitrogen, and this nitrogen anion has a three minus charge. Okay, now we can go ahead and name this, right? Off the bat, we do the first part, which is the cation.

We name the cation first, and we just use the name of the element, so lithium, and then we can Name the anion. We have nitrogen, but remember, we take the root of the element name, we add ide, so it's nitride. And so the name of this molecule is lithium nitride.

Now, if you were only given these ions, how do you know what you would get as the chemical formula? Again, you would crisscross these charges to tell you the ratio of the anions that you need. So this one charge on lithium tells us that we need one nitrogen, and the nitrogen 3-tells us we need... 3 lithium, so it's lithium 3N. And the reason for that is if we have 3 lithiums with a plus 1 charge, we have a plus 3 total charge from the lithium 3, and the total charge for the nitrogen is minus 3, and so notice how we get...

these charges to cancel out. And remember, the overall formula of the ionic compound is charge neutral. Now, let's go ahead and do some example problems.

So here, we're given the binary compound, and we want to name each binary compound. So let's look at A. Let's break this into the respective cation and anion. We have CS.

All right, CS, this is cesium. All right, it has a plus one charge. And then we have fluorine.

We have it as an anion, F minus one. All right, and so it's gonna be called fluoride, right? So we broke it into the anions it makes and we named it following the rules two and three for how to name cations and anions respectively.

And we can just put this together. The name of this molecule is cesium fluoride. Fluoride, missing a U.

All right, let's do another one. All right, here we have AlCl3. Okay, so let's write aluminum.

All right, Al, that's aluminum. It has a three plus charge. And remember, cations are just named after the element. So this is aluminum. And then we have Cl, Cl minus one here.

This is chlorine, but for naming anions, you take the root name, you add Ide. So this is chloride. And the name of this molecule is aluminum chloride. Beautiful.

Let's do the last one. We have lithium. Lithium has a plus one charge.

We write lithium because that is the cation. It keeps its name. And then H minus, this is a hydrogen anion.

How do we name this? We take the root element name and then we add IDE. So this is hydride.

The name of this binary compound is lithium hydride. Wonderful. So that is for naming type one binary compounds.

Very simple rules. You name the cation first, then the anion. Monatomic cation, it takes the name from the name of the element.

And for the monatomic anion, you take the root of the element name and add I-D-E. So that's it for binary compounds type one. Now, something else that we need to have a conversation about is, well, we've started with the chemical formula of a compound. We learned how to decide its systematic name, but the reverse process is equally important. So for example, if we were given the name calcium hydroxide, can we write the chemical formula?

Well, calcium hydroxide gives us a lot of information. Calcium is Ca2+, and hydroxide is OH-1. And we know that the overall formula of an ionic compound is charge neutral.

And to figure out how much of each ion we need, we crisscross the numbers associated with charges. And we figure out that the chemical formula is CaOH-1. 2. All right. Calcium forms only Ca2 plus ions and hydroxide is OH minus. And we need two of these ions.

They're required to give us a neutral compound. So let's work through this even more by doing a couple more examples because very important to be able to work in either direction for these kinds of problems. A says potassium iodide.

All right. Let's look at the cation. Cation is named first.

So potassium is our cation. That's K plus. And then we have iodide.

That's our anion. Okay. And we know iodide. That is I minus.

All right. We crisscross the charges and what we get is KI. So that is the formula for potassium iodide.

Let's do this next one. Gallium bromide. Gallium is our cation. It is GA3 plus.

And then bromide is our anion. It's Br minus one. All right.

And so again, we crisscross the charges because the overall formula of an ionic compound is charge neutral. And since gallium has a three plus charge, we're going to need three bromide anions to balance the charge. And so our formula, our chemical formula for gallium bromide is G-A-B-R-3, G-A-B-R-3.

Wonderful. Now that we've covered that, we can move into talking about binary ionic compounds type 2. So in binary ionic compounds type 1, the metal that's present forms only a single type of cation. So we know sodium only forms Na+, calcium only forms Ca2+, and so on. However, what we're going to see now is that there are many metals, specifically the transition metals, that form more than one type of positive ion and thus form more than one type of ionic compound with a given anion.

So for example, we can consider FeCl2. and FeCl3. FeCl2 is formed with the Fe2 plus cation, and FeCl3 is formed with the Fe3 plus cation. In a case like this, the charge on the metal ion has to be specified, and so that needs to be considered in the systematic name.

FeCl2 is actually going to be named iron 2 chloride, and FeCl3 is going to be named FeCl3. three would have the systematic name iron, three chloride. And so you notice that the Roman numerals here are used to indicate the charge of the cation.

And here you see a list of examples of certain metals that could have different charges that can form more than one type of positive ion. So we see iron here, copper, cobalt, tin, lead, and so on, mercury. You'll see this a lot, again, with the transition metals.

They're typically able to form more than one type of positive ion. So we need to get used to being able to look at a chemical formula of a binary ionic compound type 2 and be able to give the systematic name and also be able to work in the other direction, being given a systematic name for a binary ionic compound type 2 and be able to write the chemical formula. So that's exactly what we're going to work through now. So here in this first example, we want to be able to give the systematic name for each of the following compounds here.

So what we're going to do is look at A first, okay? And let's figure out what we have for the cation and for the anion. So for the cation, we have copper. This is one of those metals that can form multiple different positive ions. Okay, so we need to keep that in mind while we work through this problem.

So copper is our cation. Chloride is our anion. It has a minus one charge always, okay?

Now, we can use this information to figure out what the charge of copper has to be here to give us CuCl. It's gonna have to be plus one so that the charges balance and we would get back exactly what we are looking at, CuCl. And that helps us because now we can name it.

The name of this molecule is we name. The cation first, copper. Here, it's a metal that can form different positive ions. So we need to write the Roman numeral to indicate the charge that's associated with the copper in this formula right here. And then we can name the anion.

It's just going to be chloride. Wonderful. Let's do B now. Okay, again, same logic.

What do we have here? We have mercury. All right, mercury is one of those elements that can have different positive, it could form different positive ions, okay?

And what is our anion? We have oxide, all right? And the anion oxide has a charge of 2 minus. This helps us determine what the charge of mercury is because the charges need to balance to give us HGO.

And so the charge of... Mercury here is going to be 2 plus and this is good. We figured that out. So now we can write the systematic theme and This is going to give us if we cross charges HG 202 and that's okay because the reduced form of this is HG Oh, and now we can name it. It's going to be mercury for the cation We have to use Roman numerals to indicate the charge that's going to be 2 and then we can now write the anion Which is oxide wonderful Now, I want you to attempt C and let me know in the comments what you got for the systematic name of Fe2O3.

This should be a 3, my bad. So let me fix that for you. It should be Fe2O3.

So let me know what the systematic name of this is going to be. But we're going to work through this next example now together, which says write the formula, the chemical formula, for each compound. Okay, A tells us we have manganese.

4 oxide. This is nice though the cation is listed first. Manganese is our cation and we're told 4. So that's the charge that's associated with manganese.

So it's 4 plus and then we know that our anion is oxide. Oxide has a charge of 2 minus. Now we have to be careful here when we write the chemical formula. Remember the overall formula of an ionic compound is is charge neutral. So what we can do here is try to figure out what does that look like?

What we can do is we can crisscross these charges. All right. You can crisscross these charges. It gets us MN204, but we want to reduce this down, right?

Because these are factors of each other. And so it's MN204. MnO2.

And that makes sense. If I had two oxides, I would have two minus and two of those. So that's going to give me a total of four minus, four minus and four plus.

Those balance each other out for us to get a ionic compound that is charge neutral. And so the formula for manganese four oxide is MnO2. Let's also do this next one, lead two chloride.

Okay. So again, lead is Pb and it's, we're given the charge two plus, and then we have chloride, which is a minus one, and then we can crisscross these charges. And what we're going to get is that it's PbCl2, and that is the chemical formula for lead to chloride.

Now we can move into the next objective in nomenclature here. But before moving on, I want to make note that this table has a lot of common polyatomic ions and their names. This is another table that I highly recommend getting familiar with because it is super important and you're going to see these kinds of ions reappear throughout general chemistry.

And it will make general chemistry a lot easier if we're able to understand when we're given a chemical formula, what is the systematic name for it? Or if we're given a systematic name in a problem that we know what the chemical formula for that is, because as we... dive deeper into general chemistry, those things are going to have a huge effect on so many concepts down the line. We won't be able to do any of the other concepts if we don't have nomenclature down. So all that to say that we should familiarize ourselves with these ions.

So some common ones that are definitely important, ammonia, NH4 plus. We have nitrite, that's NO2 minus, and nitrate, which is NO3 minus. We have sulfite. SO3 2-and sulfate SO4 2-. This is hydrogen sulfate, HSO4-.

Again, our hydroxide reappears, OH-. We also have cyanide, CN-. We have PO4 3-, that's phosphate.

And when there's one hydrogen added, it's HPO4 2-, that's hydrogen phosphate. When there's two hydrogens there, H2PO4-, that's dihydrogen phosphate here. CO3 2 minus carbonate, that's an important one to know.

HCO3 minus, that's hydrogen carbonate, also known as bicarbonate. CLO minus hypochlorite, CLO2 minus chlorite, it is CLO3 minus chlorate, and I should have said this as ite. Sometimes I can mispronounce that if I'm going too fast.

So again, CLO minus is hypochlorite, and CLO3 minus is chlorate. Then we have ClO4-, perchlorate. This is acetate. Permanganate here, MnO4-.

We have dichromate, chromate, peroxide. So all of these really important to know. So make sure that you spend some time committing these to memory.

Now we can go ahead and move into talking about naming ionic compounds with polyatomic ions. So far we've been able to deal with ionic compounds with monatomic ions. It was easy to look at sodium chloride and recognize that sodium is the cation, chloride is the anion, and work from there.

But now we want to deal with ionic compounds with polyatomic ions. And from this table that we've looked at, there are a couple of things that I want to point out that are considered polyatomic cations, and the rest are going to be polyatomic anions. And being able to recognize that means now we can work with ionic compounds with polyatomic ions and easily recognize which part is the cation, which part is the anion, and then proceed with naming and writing chemical formulas in the same manner we have up to this point.

So from the table, here are the polyatomic cations. I'm gonna write them down here. Polyatomic cations that we should be familiar with.

There's just really three main ones to keep in mind right now. Okay, we have ammonia, ammonium, I should say, ammonium NH4 plus polyatomic cation. Then we have hydronium, hydronium.

I'm not sure it's in this table, but very important to know. Hydronium is H3O plus. And then the next polyatomic cation to know is mercury one.

All right, it's also the first one on this table right here. Okay, so obviously these first two. polyatomic cations mercury one okay also written as h g two with a two plus charge so these are our polyatomic cations all the other things that are listed in this table are polyatomic anions and their names.

Okay, so now that we are able to distinguish polyatomic cations from polyatomic anions, and we've also discussed monatomic cations and monatomic anions, now we can do some more complicated problems. Okay, so again, ionic compounds with polyatomic ions are assigned special names that must be memorized to name the compounds that contain them. And we're going to learn these by doing Example problems.

So this first problem here says, give the systematic name for each of the following compounds. First one here is Na2SO4. Let's figure out what the cation is and what the anion is. So obviously here, this is our cation, Na+, this is sodium.

Our anion is this polyatomic ion, SO4. SO4 and it's going to have a 2-charge hence why when we crisscross the charges the written chemical formula is Na2SO4 So we use this to figure out what the charge is for the SO4 by kind of like working backwards from our crisscross So again, we have sodium plus that's our cation SO4 2-as our anion Now same rules for naming ionic compounds we keep the cation name, the same as the element name. And here for the anion, this is a polyatomic anion.

Let's remember what SO4 2-. We go back to our table. SO4 2-is right here. That is sulfate. So let's write that down, sulfate.

That means the name of this chemical formula is going to be sodium sulfate. All right, the name is sodium sulfate. Awesome, so nothing really changed with the naming protocol that we learned.

It's just a matter of fact that we had to learn the names of several polyatomic ions so that we can be able to do this. Okay, next up, let's do this next one together again. Our cation is obvious here.

It is potassium. Let's write that down in a different color, potassium K+. That is our cation.

Now, our anion is this H2PO4, and it's going to have a... minus charge right here. Okay.

Minus charge right here. We conclude that from looking at the fact that potassium, there's one of them. So we worked backward from our crisscross. So that's our anion. Let's write that down.

H2PO4 minus. Let's go back to our table and see if we can figure out what that is. That is right here. H2PO4 minus dihydrogen phosphate, dihydrogen phosphate. So we're going to write that down.

Di. hydrogen phosphate. So we figured out the cation and the anion and the name of this molecule is potassium dihydrogen phosphate. Go ahead and do C and let me know what you got in the comments below.

Let's work on this next example where we want to write the formula for each. So A says sodium, hydrogen, carbonate. Okay, sodium is the cation. That's Na plus. And then hydrogen carbonate.

This is going to be HCO3. And you can look at the table if you need to when you're first working through these. Hopefully, as you do more and some of these names and anions reappear here and there, it becomes a lot easier to work with.

So sodium, hydrogen carbonate. Sodiums are cation, Na+. And then our hydrogen carbonate is HCO3.

All right. And it's up here. HCO3 minus hydrogen carbonate. So let's not forget the minus. And again, we crisscross the charges and the formula, the chemical formula for sodium hydrogen carbonate is NaHCO3.

Wonderful. Let's do this next one also. Okay. Caesium perchlorate.

Okay. So caesium is our cation. cesium is going to have a plus one charge. And then perchlorate is going to be CLO4. And we can double check the charge here.

It's a minus. Okay. So CLO4 minus, we crisscross the charges to get cesium perchlorate CS CLO4. So there we go.

That is cesium perchlorate. Wonderful. Now we can move on to talking about binary compounds type 3. Now, binary covalent compounds are formed between two nonmetals, and although these compounds do not contain ions, they are named very similarly to binary ionic compounds.

So, in the naming of binary covalent compounds, the following rules apply. 1. The first element in the formula is named first using the full element name. 2. The second element is named as if it were an anion. Prefixes are used to denote the number of atoms present. And four, the prefix mono is never used for naming the first element.

So for example, CO is called carbon monoxide, not monocarbon monoxide. And here we see some examples. So N2O would be named dinitrogen because there's two nitrogens.

Monoxide because there's one oxygen and the oxygen is the second element. And we name second elements as if they were anion. If we look at no here, there's only one nitrogen, but we don't say mononitrogen monoxide.

We just say nitrogen monoxide. And again, here in this table, there are the prefixes that are associated with each number. So if you have one, use the prefix mono.

Two is di, three is tri, four is tetra, five is penta, six is hexa, seven is hepta, eight is octa, nine is nona, and ten is deca. So that is binary compounds type three. And here in this next page, we kind of put it all together because we did talk about a lot.

And so are you dealing with a binary compound? If the answer is no, then the next thing you want to follow up with is, well, is there polyatomic ions present? If the answer is no, then this topic we have not covered yet.

We don't know how to name those kinds of compounds just yet. But if the answer is yes, then we not name compounds using the procedure that is very similar to binary compounds. We identify our cation, we identify our anion, and then we name it as we would if we were dealing with monoatomic cations and anions, but just making sure that we can recognize the nomenclature that goes with the present either polyatomic cation or polyatomic anion or both.

So that is for polyatomic cation. ions. Now, if you are dealing with a binary compound, then the next question to ask is, is the metal present? If no, then you are dealing with binary compounds type three. So binary covalent compounds, which is what we just talked about.

And you want to follow the rules we just covered there. But if there are metals present, then the next question to ask is, well, does the metal form more than one cation? If the answer is no, then what you have is a binary compound type one.

So follow those rules, identify your cation, identify your anion. Your cation, you use the element name, and for the anion, you take the element root name, and then you add ide at the end. If the metal forms more than one cation, then what you're dealing with is a binary compound type 2. Here we use Roman numerals to indicate the charge that's associated with the cation.

So this is the workflow, the flow chart that you want to fall on. whenever you are trying to work through naming compounds. So with that being said, let's do a mixed problem set.

Here in the first example, we're going to attempt to give the systematic name for the following compounds. A is P4O10. So here we have two non-metals, phosphorus and oxygen.

So this is a binary covalent compound type three. So prefixes are used. The first element is phosphorus and there's four of them. So the first part of the name is tetraphosphorus.

And then we have oxygen. But remember in the rules, we want to name the second element as if it were an anion. So oxide. And there are 10. So we want to use the prefix deca. But here sometimes the A is dropped because then you're going to have like two vowels right next to each other.

Typically, it's written like decoxide like this. Then for B, we have NB2O5. Now, what we have here is a monatomic cation and anion. But that cation is niobium. It's a transition metal, and it can have different charges.

So it's going to require a Roman numeral. We're going to have to figure out what that is. But let's write. We have our cation, niobium.

We have oxygen. This has a 2-charge. Here, what's really cool is that, yeah, we do this crisscross to figure out how much of each cation and anion is.

This O, this 2 from the O goes to the niobium. And then if niobium had a 5 plus charge, then it would give us back exactly this. And this was a fast way of figuring that out here.

So that means that we have for naming, we'll name the cation first. Niobium charges 5. And then the anion is oxide. I'll let you figure out C on your own and let me know down below. Actually, what I'll probably do is for all the ones that I've asked you to attempt, I'll...

probably put the answers, I'll try to, hopefully I don't forget, in the description box, but later on, so then you can like look at it after you've attempted these problems to double check. Now, let's also do this next example where we write the formula for each. So we have vanadium 5 fluoride.

So the compound contains vanadium 5 plus ions, and it requires 5 fluoride ions for charge balance here, right? Because vanadium is going to have that 5 plus charge, fluoride that minus 1, then we crisscross those charges and what we get is VF5. So that's the chemical formula for vanadium 5 fluoride.

Now for this next one, we have dioxygen difluoride. So here the prefixes really help us out. We have two oxygens and two fluoride, fluorine, so O2F2.

And again here, I want you to attempt to see on your own and let me know what you got or double check your answer in the description box. But with that, now we can move on to our last and final category for nomenclature, and that is acids. An acid can be viewed as a molecule with one or more hydrogen ions attached to an anion.

The rules for naming acids depends on whether the anion contains oxygen. So If the anion does not contain oxygen, the acid is named with the prefix hydro and the suffix ik. When the anion contains oxygen, then the acidic name is formed from the root name of the anion with a suffix of ik or ois depending on the name of the anion. So here we get into those subrules.

If the anion's name ends in ik, then the suffix ik is used. And if the anion has an i ending, then that is replaced with ois. So here...

This first category, acids that don't contain oxygens. We can look at a couple of examples. HF, this is hydrofluoric acid.

So notice the hydro prefix, the ik suffix, and then we end with acid. And again, if we look at HCl, hydrochloric acid. Here we have a list of acids that do contain oxygen. So looking at the first one, HNO3.

Now, NO3, that's the anion. This is nitrate, so it ends in A-T-E. That means that we want to use the suffix"-ic", and add that to the root name.

So the name of this acid is actually nitric acid. HNO2 though, however, this is called the anion, NO2 is nitrite with an I-T-E ending. And so that's going to be replaced with O-S.

So the name of this acid, HNO2, is nitrous acid. And the same goes with... Other examples, H2SO4.

SO4 is sulfate with A-T-E. And so the name of this acid is sulfuric acid. We use the IC suffix. But if we look at H2SO3 instead of 4, this ends in"-ite".

And so the name of this acid is sulfurous acid. With that, we've completed all the topics for this chapter. In the next video... We'll work on problems that relate to all the topics we covered here in chapter two. I really hope this was helpful.

Let me know if you have any questions, comments, concerns down below. Other than that, good luck, happy studying, and have a beautiful, beautiful day.