Atomic Structure - Online Tutor, Practice Problems & Exam Prep

The atom is the fundamental unit of matter, consisting of a nucleus with protons and neutrons, surrounded by electrons in energy levels called shells. The atomic number, unique to each element, equals the number of protons, while the mass number is the sum of protons and neutrons. Isotopes have the same atomic number but different mass numbers due to varying neutrons. Electrons can create ions when their number differs from protons, resulting in cations (positive) or anions (negative). Understanding electron configuration principles—Aufbau, Pauli exclusion, and Hund's rule—is essential for predicting atomic behavior.

The difference between atomic numbers and atomic mass.

Hey everyone. So we're explaining from general chemistry that the atom is the basic unit of matter and a collection of atoms is what helps to make a molecule. Now, talking about an atom, we have an example of an atom in the top right corner here. Remember, at the center of an atom is where we have the majority of the mass of the atom. That's where we have our protons and our neutrons. Here it's not drawn to scale; the nucleus is very small, here we're just zooming in on it so we can see the protons and the neutrons. Also remember that zooming around this are electrons which are found within different orbits. Now, also remember that the atomic number of an atom is equal to the number of what? That's right, the number of protons. Also remember that the atomic number is unique to a given element. So an element has an atomic number and only that element has that atomic number. Besides the atomic number, we have the mass number. The mass number of an atom is equal to the number of protons plus neutrons. So we're talking about the total number of these subatomic particles within the nucleus of a particular atom. With this whole idea of protons, neutrons, and electrons, we also know another term, isotopes. Now isotopes, these are a type of element that have the same number of protons and therefore the same atomic number, but they have differing neutrons. So their mass numbers are different from them. So they all have the same number of protons, but it's just their mass numbers that are different. We'll go into further explanations for what exactly is an isotope and some examples in the following video.

Understanding the hydrogen isotopes.

Hey everyone. So here we're going to take a look at hydrogen isotopes. Now before we begin, first realize that an isotope is not a new type of atom. All it is is a heavier form of the same atom, and we make it heavier by adding neutrons. Here we have hydrogen, deuterium, and tritium. For each of them, they all are just hydrogen in different forms because all of them have an atomic number of 1. Just remember, the bottom number here represents our atomic number, the number of protons. The number up top, we're going to say that represents our mass number, which remember is the number of protons plus neutrons. So if we take a look here at the first one, which is just regular hydrogen, it has an atomic number of 1. So it has 1 proton, and its mass number is also 1. If we were to subtract those two numbers, that will give us the number of neutrons. So we'd say here that hydrogen, regular hydrogen, has 1 proton and 0 neutrons. Next, we add 1 neutron and when we do that we create deuterium. Now deuterium is going to be an isotope of hydrogen you're going to see later on throughout organic chemistry. You won't find it on the periodic table, you'll just see hydrogen there. Remember deuterium is just another form of the hydrogen atom. It's gotten heavier by adding an additional neutron. Here we abbreviate it as D. Here we'd say that since its atomic number is still 1, it still has 1 proton, so it's still 1 p+, and then if we subtract those two numbers from each other, 2 minus 1, we get 1, which is the number of neutrons. Then what we can do is, we can add another neutron, and that would help to create tritium. Tritium is abbreviated as just T. Again, you wouldn't find this on the periodic table, this is just another form of the hydrogen atom. Here we have 1 proton still and then 2 neutrons. So you can see that all of them are hydrogen because all of them have 1 proton, but they're all different forms of hydrogen because they have different numbers of neutrons. Also realize that adding neutrons comes at a cost. If you're adding more neutrons to create different isotopes, you're going to affect the stability of your atom. It's going to become less stable. So here, tritium is incredibly unstable. It doesn't last but for a few moments because we're adding too many neutrons. Also remember that besides stability, we also talk about scarcity. So here, hydrogen before we've added any neutrons is the most stable of these 3, so the majority of hydrogen is found in this form. As we go towards tritium, the percentage of hydrogen existing in that form goes lower and lower. Alright, so just remember that isotopes are just different forms of the same atom, they're just heavier forms of that atom and we do that by adding neutrons.

Shells, orbitals and types of ions

Now we're going to talk a little bit about just atomic structure. Okay? So remember that electrons orbit regions around a nucleus. Okay? And they orbit it based on energy level. Okay? So that region of space that has a certain energy is called a shell. Okay? A shell is that space that has a certain energy where the electrons can move. Alright? Now shells, remember shells can hold a lot of electrons. Some shells can hold up to 18 electrons, something like that. But there's a smaller subset of space within a shell that holds exactly enough, that has enough room for exactly one pair of electrons. Do you guys remember what that is called? That would be called an orbital. Okay? An orbital is that region of space where only 2 electrons, an up spin and a down spin, can exist in. Okay? And we're going to talk a lot about orbitals later. Alright? Then what happens when we already talked about what happens when atoms have different amounts of neutrons. Alright? But when atoms possess a different number of electrons than protons, so now we're talking about what happens if they have different amounts of electrons. What's going to happen is that instead of being heavier or lighter, electrons don't really contribute to mass very much. Remember that electrons are tiny, tiny, tiny. They do not they're not very heavy. Instead, they're going to contribute towards the charge. So these are called ions. Okay? Remember the word ion just describes something that is a charged atom. Alright? So remember there are 2 kinds of charges you could have. You could have a positive or a negative. You guys should be able to fill this in. What is it called when it's a positively charged atom? That's a cation. Okay? Please don't say cation. Okay? That is not a cation. That is a cation. Alright? A negatively charged atom would be called an anion. Okay? So it's important that you guys know this distinction. Basically, different amounts of neutrons affect the weight. So you get a different atomic mass. Different amounts of electrons don't affect the weight that much, but it does affect the charge. Why? Because if you think about it, electrons have a negative charge. Protons have a positive charge. In atoms, these are supposed to perfectly balance out. You're supposed to have exactly as many protons as you have electrons. If you have a difference, then you're going to have a net charge for that atom. Alright?

Understanding the hydrogen ions.

So here's an example of some really good ions right here. These are just the simplest form of ions you can make, which should be the positive hydrogen and the negative ion of hydrogen. So let's talk about the positive one. Normally, a hydrogen atom consists of a proton and an electron. If I take away one of those electrons, what's going to happen is that I'm going to get a positive charge. Okay? I'm going to get a positively charged atom. And the reason is because I have nothing to counteract the positive charge of the proton. Alright? Now, when you have one of these H+ ions, it's actually just called a proton. Why? Because there's nothing else. There are no neutrons. There are no electrons. So literally, we just call it a proton. So when I say, hey, there's a proton whizzing around, that means it's a hydrogen that does not have an electron on it. Okay? And that means it has a positive charge. Then a hydride is the name of a hydrogen that has a negative charge. Now why would hydrogen ever have a negative charge? Well, if it has more than one electron. For example, if it has two electrons, then the electrons are going to win. There's going to be more electrons than protons, so you'd have a net negative. So H− is hydride. H+ is a proton. And you guys need to know that that's actually super important for organic chemistry.

Three rules about orbitals you need to know.

So now finally, what I want to talk about are the three principles of electron configuration. We're going to do some practice with this later. But these principles, which you learn in general chemistry, you just need to commit them to memory. Okay? You're still going to need them for organic chemistry. What these do is they describe the way that electrons fill atomic orbitals. So remember that orbitals hold 2 electrons. The Aufbau principle is also called the building up principle. Building up. Okay? And with the building-up principle, what happens is that it just says, hey, if you have orbitals of differing energies, you have to fill the lowest energy first. Okay? So remember, we're going to talk more about this later, but remember that the 1s orbital is your lowest energy orbital. And then it goes up to 2s, and then it goes to 2p, and then they start going into the d's and the f's, and everything. Okay? So that means you wouldn't fill an f orbital or a d orbital before you completely filled all your lower orbitals. Alright? So you always have to start at the bottom and work your way up. Easy. Then we have the Pauli Exclusion Principle. This one you guys already know. All it says is that you can only have 2 electrons per orbital. So I already told you this, but you didn't know that it was Pauli Exclusion. Maybe you forgot. Okay? Or maybe you remembered. Sorry. You could have totally remembered that. I just assume. I teach this class as if you don't remember anything from general chemistry. So the way that I like to think about it is that imagine, like, these 2 electrons are, like, out on a date. Okay? And then, like, someone tries to join in, it's like a third wheel. They're going to be exclusive. They're going to be like, no. Like, go away. Like, I want to be on a date. Right? So Pauli Exclusion, they're like excluding the other electrons. Alright? So that's Pauli Exclusion. Then finally, we have our last one, which is Hund's Rule. Hund's Rule can be compared to a lot of people compare it to, like, you're on a school bus, and it's like the seats of a bus. Alright? So, basically, you have a certain amount of kids. And, you know, well, I don't know. Kids are weird. But if you had a normal bus, a normal, like, bus with adults, what would happen is that people always take the clear seats first. No one doubles up on a seat for no reason. They always take all the clear seats and once all the seats have one person on them, then you start putting your second people. Okay? And basically, that's what Hund's Rule says. It says that you have to evenly fill all of your orbitals of the same energy level before you can start adding second electrons to them. Alright? So in this case, just remember that you're going to equally fill anything that has equal energy with one electron first, and then you can start adding second ones. Alright? So, guys, it's just important that you remember what these rules are about and we're going to apply them in some practice problems so you will get some practice. Alright? So let's go on and try it out.

PRACTICE: Determine the number of protons, neutrons and electrons in the following atoms.

• A.
• B.
• C.
• D.

PRACTICE: Determine the number of protons, neutrons and electrons in the following atoms.

• A.
• B.
• C.
• D.

PRACTICE: Determine which of the three principles of electron configuration is being broken in the electron diagrams below.

• A. Hund's Rule
• B. Aufbau Principle
• C. Pauli Exclusion Principle

PRACTICE: Determine which of the three principles of electron configuration is being broken in the electron diagrams below.

• A. Hund's Rule
• B. Aufbau Principle
• C. Pauli Exclusion Principle

Here’s what students ask on this topic:

What is the atomic number and how is it determined?

The atomic number of an element is the number of protons in the nucleus of an atom of that element. It is unique to each element and determines the element's identity. For example, hydrogen has an atomic number of 1 because it has one proton, while carbon has an atomic number of 6 because it has six protons. The atomic number is denoted by the symbol Z. It is crucial because it defines the chemical properties of the element and its place in the periodic table.

What is the difference between atomic number and mass number?

The atomic number (Z) is the number of protons in the nucleus of an atom and is unique to each element. The mass number (A) is the total number of protons and neutrons in the nucleus. While the atomic number determines the element's identity, the mass number gives the total mass of the nucleus. For example, carbon-12 has 6 protons and 6 neutrons, so its atomic number is 6 and its mass number is 12.

What are isotopes and how do they differ from each other?

Isotopes are variants of a particular chemical element that have the same number of protons (and thus the same atomic number) but different numbers of neutrons, resulting in different mass numbers. For example, carbon-12 and carbon-14 are isotopes of carbon; both have 6 protons, but carbon-12 has 6 neutrons while carbon-14 has 8 neutrons. This difference in neutron number affects the atomic mass but not the chemical properties of the element.

What is the Aufbau principle in electron configuration?

The Aufbau principle, also known as the building-up principle, states that electrons fill atomic orbitals of the lowest available energy levels before occupying higher levels. This means that electrons will first fill the 1s orbital, then the 2s orbital, followed by the 2p orbitals, and so on. This principle helps predict the electron configuration of atoms, ensuring that the lowest energy state is achieved.

What is Hund's rule and how does it apply to electron configuration?

Hund's rule states that electrons will fill degenerate orbitals (orbitals of the same energy) singly before pairing up. This minimizes electron-electron repulsion within an atom. For example, in the case of the 2p orbitals, one electron will occupy each of the three 2p orbitals before any pairing occurs. This rule ensures that the atom remains in the lowest energy state possible.