Chapter 2 - Atoms, molecules and atoms

This chapter, "Atoms, Molecules, and Ions," is from the book "Chemistry: The Central Science" by Dr. Brown. By the end of this module, you'll be able to list the basic postulates of Dalton's atomic theory. You will describe the key experiments that led to the discovery of electrons and the nuclear model of the atom. You will also describe the structure of the atom in terms of protons, neutrons, and electrons. Additionally, you'll be able to describe the electrical charge and relative masses of protons, neutrons, and electrons. You will learn how to use chemical symbols, together with atomic number and mass number, to express the atomic composition of isotopes. You will know how to calculate the atomic weight of an element from the masses of individual atoms and knowledge of natural abundance. You will describe how elements are organized in the periodic table by atomic number and similarities in chemical behavior that give rise to periods and groups. You will identify the locations of metals and nonmetals in the periodic table. You'll be able to distinguish between molecular substances and ionic substances in terms of their composition. You will also be able to distinguish between an empirical formula and a molecular formula. You will describe how molecular formulas and structural formulas are used to represent the composition of molecules. You will explain how ions are formed by gaining or losing electrons, and you'll be able to use the periodic table to predict the charge of common ions. You will learn how to write empirical formulas of ionic compounds given the charges of their component ions. You will also be able to write the name of an ionic compound given its chemical formula or write a chemical formula given its name. The same applies to binary inorganic compounds. Additionally, you will be able to identify organic compounds and name simple alkanes and alcohols. So, let's start with the atomic theory of matter. This one is known as Dalton's atomic theory, the theory that atoms are the fundamental building blocks of matter. It emerged in the early 19th century by John Dalton. He mentioned that the elements are composed of extremely small particles called atoms. An atom is basically the smallest particle you can find in matter with specific characteristic(s). All atoms of a given element are identical to one another in mass and other properties, but the atoms of one element are different from the atoms of all other elements. So, here we have, for example, the representation of oxygen in the red spheres. All the atoms of oxygen are the same. Here, you have the atoms of nitrogen represented by blue spheres. All of these atoms are the same, but the atom of oxygen is different from the atom of nitrogen. When you're talking about an element, all of the atoms of that specific element are the same. But when you're comparing atoms from different elements, those atoms are different. Atoms of an element do not change into atoms of a different element by chemical reaction. Atoms are neither created nor destroyed in a chemical reaction. So, if you have a hydrogen atom, it's going to be a hydrogen atom always. You can't change that to another element, same way with oxygen, nitrogen, etc. You can't change the element to another one by chemical reaction. Atoms of more than one element combine to form compounds. A given compound always has the same relative numbers and kinds of atoms. Subsection 2.1 talks about the law of conservation of mass. It says that the total mass of the substances present at the end of a chemical process is the same as the mass of the substances present before the process took place. So, the total mass of all substances at the beginning of a reaction has to be the same at the end, even though they may be different kinds of compounds because of the chemical reaction. So, you start with A and B, and you're going to produce C and D, different compounds, but the total mass has to be equal at the end as well as at the beginning. This law was one of the laws on which Dalton's atomic theory was based. The law of multiple proportions states that if two elements A and B form more than one compound, the masses of B that combine with a given mass of A are in the ratio of small whole numbers. Dalton also predicted and observed this law while developing his atomic theory. When two or more compounds exist from the same element, they cannot have the same relative numbers of A and B. So, we can have, for example, H2O, but also we can have H2O2. H2O is water, and H2O2 is hydrogen peroxide. Even though they are basically the same elements, but the ratio that holds them in those whole numbers are different. In H2O, it is 2:1 while in H2O2 it is 1:1. That's basically the law of multiple proportions. Now, the atom has a few subatomic particles. In that sense, the atom was not the smallest particle possible. The discoveries led to the fact that the atom itself was made up of smaller particles. You can see that in Section 2.2, where you can find electrons and cathode rays. Also, back to the activity, the nucleus, protons, and neutrons. I've also provided a link, which you can find in the folder of Chapter 2, to a very interesting video on how the discovery of electrons and neutrons was made. You may take a few minutes to watch that video, as it may be helpful in understanding those experiments. Let's talk about the subatomic particles now. Protons have a charge of +1, electrons have a charge of -1, and neutrons are neutral. So, those are the three most important subatomic particles. Protons, electrons, and neutrons. Protons have a charge of +1, electrons have a charge of -1, and neutrons are neutral. Protons and neutrons have essentially the same mass, the relative mass of one. Whereas the mass of an electron is so small that we ignore it. It is basically considered to be zero when compared to the protons and neutrons. Protons and neutrons are found in the nucleus, while the electrons travel around the nucleus. So, basically, we can envision the atom as having a nucleus, and inside that nucleus, we have the protons and neutrons. All around that nucleus, we're going to find the electrons orbiting the nucleus. Here, we have a table for some characteristics: protons have a charge of +1, neutrons stay neutral, and electrons have a charge of -1. We can also see the mass of each particle. Protons and neutrons basically have the same mass, while the electron is really, really small in comparison. You can say that it is basically zero compared to the protons and neutrons. The atomic mass: atoms have extremely small masses. And as you may think, that mass is going to depend on the number of neutrons and protons. The heaviest known atom has a mass of approximately 4 x 10^-2 grams. So, it has a really, really small mass. A scale on the atomic level is used, where an atomic mass unit (amu) is the base unit. One atomic mass unit (1 amu) is equal to 1.66054 x 10^-24 grams. So, basically, this is a conversion factor for atomic mass. Consolidate or define that as 1 amu. This is basically the atomic mass unit - the atomic and molecular weight can be measured with greater accuracy using a mass spectrometer. The mass that you usually compare all the elements against is the mass of a carbon atom that has six protons and six neutrons and is established as carbon-12. Here we have a basic example of using a mass spectrometer. Okay, so how does it work? We have a sample that goes through here and then it can be detected and separated. For example, a sample of chlorine 35 and chlorine 37 can be separated in that detector. So it has a very high accuracy to determine the atomic mass of elements. Now let's talk about the symbols of elements. Here we have carbon-12 and C6. The C is the symbol of the element. In this case, carbon. The number 6 represents the atomic number. The atomic number also represents the number of protons or electrons for the element. In blue, we have 12. This is the mass number. The mass number is the number of protons plus neutrons. So here we have the mass number, atomic number, and symbol of the element. Elements are represented by one or two-letter symbols. This is the symbol for carbon. As I mentioned, if we have two letters, the first one is going to be a capital while the other one is going to be a lowercase letter. All atoms of the same element have the same number of protons, which is called the atomic number. The atomic number is also the number that represents the periodic table. We're going to see that in a few more slides. So it's going to be organized by that atomic number. The atomic number is also represented as Z and is written as a subscript before the symbol. So most of the time, it's going to be here, the Z that represents the atomic number, also the number of protons. The mass number is the total number of protons and neutrons in the nucleus of an atom. It is written as a superscript before the symbol. So we're always going to put here the mass number, and down here we're going to write the atomic number. That's the number of protons and the mass number represent the number of protons and neutrons in the element. Elements are represented by one or two-letter symbols. If it's one, it's always a capital letter. If it's two, the first one is going to be a capital letter, and the second one is going to be a lowercase letter. Now let's talk a little bit about isotopes. Isotopes are atoms of the same element with different masses. Isotopes have a different number of neutrons but the same number of protons. They need to have the same number of protons because that is a characteristic of the element. So if we are talking about they are of the same element, that means that the number of protons can't change. So the only difference is going to be in the mass number. It's going to be the number of neutrons. For example, carbon has four isotopes: carbon-11, carbon-12, carbon-13, and carbon-14. The atomic number or the number of protons for carbon is 6, so the atomic number for all of these is going to be 6 because it's carbon. The number of electrons is also going to be 6 because we have the same number of protons and electrons. This is neutral, so they can cancel each other out. 6 plus 6 charges and 6 negative charges is equal to 0. So that's why you have the same number of protons and electrons because we're talking about an element that has a neutral charge, 0 charge. But we have a different number of neutrons in each of the isotopes. Here we're going to have 5, 6, 7, and 8 because the mass number here is 11 and the number of protons is 6. That means that 11 minus 6 is equal to 5. That's the number of neutrons. Remember, this number is the sum of protons plus neutrons. So we know the number of neutrons. 12 minus 6 is equal to 6. That's the number of neutrons in carbon-12. In carbon-13, we have that the number of protons is going to be 6 because we're talking about carbon. So 13 minus 6 is going to be 7, and that is the number of neutrons. The same thing with carbon-14. Carbon-14, 14 minus 6 is going to give us the number of neutrons, which is 8. Okay, so that's how we can determine the number of neutrons. And now we understand that the difference in the different isotopes is going to be the number of neutrons. The number of protons needs to be the same because we're talking about different atoms of the same element. Now, let's talk about the atomic weight. In the real world, we use large amounts of atoms and molecules. We use average masses in calculations. An average mass is found using all isotopes of an element weighted by their relative abundance. This is the element's atomic weight. In other words, the atomic weight is equal to the sum of the isotope mass times the fractional natural abundance. This is the sum of all the different isotopes. So, for example, for carbon, we have four isotopes. So we're going to look for the isotope mass of all four carbons and the fractional abundance of each of them. And we're going to sum all of them to give us the atomic weight. Let's see an example of this to clarify this concept a little bit more. The atomic weight of naturally occurring chlorine is 75.78 for the isotope chlorine-35, and the atomic mass for this isotope is 34.969. 24.22% of chlorine-37. So here we have that the element chlorine has two isotopes: chlorine-35 and chlorine-37. So we can calculate the atomic weight by multiplying the abundance of each isotope by its atomic mass and then summing these products. But first, we need to change those percentages to decimals. So we divide 75.78 by 100, and it's going to be equal to 0.7578. While 24.22 divided by 100 is going to be equal to 0.2422. And now, these abundances will be multiplied by their respective atomic masses. So this one is going to be multiplied by this, and the 0.2422 by the atomic mass of chlorine-37. And that way, we have here the abundance in decimal, expressing decimal. We have the atomic mass of that isotope 35, and then we have the abundance of the second isotope 37. We multiply each of them. We're going to get 26.550 atomic mass units plus the product of this multiplication, which is 8.953. Now, remember, here we have five significant figures, four significant figures. So this is going to be represented in four as well. Here we have four and you have five. So we represent it with four significant figures. Okay, we can follow the rules that we mentioned in Chapter 1. And now we need to add. So that our answer needs to be with two decimal positions because we now look for the decimal precision. We have here, we have two, we have three. So the result must be 35.45 atomic mass units. This is basically the atomic weight for chlorine. It's due to the presence of the two different isotopes. This answer makes sense. Atomic weight, which is actually the average atomic mass, is weighted to the mass of the two isotopes. And it's closer to the value of chlorine-35. What do you think? It's closer to chlorine-35 because that isotope has the highest abundance than the other ones. So the average is going to be closer to the 35 than to the 37. And this is how you calculate the atomic weight of an element by using the isotopic abundance. Now let's talk about the periodic table. The periodic table is a systematic organization of elements. The elements, as I mentioned before, are arranged in order of atomic number. Here we have the element hydrogen. Symbol H and the atomic number is on top of the symbol. We have here the symbol for silver, which is AG, and the atomic number is on top of the AG. Remember that the atomic number also represents the number of protons and electrons. So, silver has 47 protons and 47 electrons as an element. As you can see, all the elements that have a symbol with two letters, the first one is capitalized and the second one is lowercase. The atomic weight of an element appears most of the time at the bottom of the box, but here we are only showing the atomic number and symbol. In the other elements, you will also see the atomic weight most of the time at the bottom of the box. Now, the rows on the periodic table are called periods. Period 1 here is composed of hydrogen and helium. Period 2 goes from magnesium to aluminum, and so on until period 7. From left to right, you will find the periods. The columns are called the groups. We have group 1A, 2A, 3A, 4A, 5A, 6A, 7A, 8A, and the B group, which is known as transition metals. The transitional elements or transition metals are ordered here differently because they have their own characteristic periodicity. That's why you see a jump from group 8 to 3A. Elements in the same group have similar chemical properties. They are also known as families. Group 1A has similar chemical properties, and the same goes for group 7A and 8A. The periodic table is called periodic because of the periodicity of its characteristics. When one looks at the chemical properties of elements, one can observe a repeating pattern of activity. By putting the elements in an orderly line like this, scientists observed this characteristic that from 2 to 10 and 12 to 18, they were basically non-reactive gases. The green ones were soft reactive metals, and they started to look for characteristics that repeated in a certain position. That's how they came up with this table. All of these elements have the same characteristic chemical properties. The groups also have some names. Group 1A is known as the alkali metals, group 2A is the alkaline earth metals, group 6A is the chalcogens, group 7A is halogens, and group 8A is noble gases. These five groups are known by their names. Another important information that we can obtain from the periodic table is that at the left side, we find elements considered as metals. All of these are colored in beige boxes. Hydrogen is an exception, as it is at the left side but not considered a metal. Properties of metals include being shiny, conductive to heat and electricity, and solid (except for mercury). Nonmetals are found at the right side of the periodic table, colored in green and also hydrogen. These are considered nonmetals, even though hydrogen is an exception. Nonmetals can be solid (like carbon), liquid (like bromine), or gas (like neon) at room temperature. Metalloids are the elements that are part of this step-like line. These elements have properties that are sometimes like metals and sometimes like nonmetals. Now, let's talk about chemical formulas. We can combine the elements and create chemical formulas. The subscript to the right of the symbol of an element tells us the number of atoms of that element in one molecule of the compound. For example, H2 means that in a molecule of hydrogen, we have two atoms of hydrogen. O2 means that in a molecule of oxygen, we find two atoms of oxygen. H2O (water) means that in each molecule of water, we have two atoms of hydrogen and one atom of oxygen. In ethylene (C2H4), we have two atoms of carbon and four atoms of hydrogen in each molecule. Molecular compounds are composed of molecules and almost always contain only nonmetals. When we combine nonmetals, we are talking about molecular compounds. Diatomic molecules are a group of molecules that naturally exist as two atoms in each molecule. We can find diatomic molecules like hydrogen, nitrogen, oxygen, fluorine, chlorine, bromine, and iodine. These molecules exist as H2, N2, O2, F2, Cl2, Br2, and I2. There are different kinds of formulas. The first one is the empirical formula, which gives the lowest whole number ratio of atoms of each element in a compound. The molecular formula, on the other hand, gives the exact number of atoms of each element in a compound. If we know the molecular formula of a compound, we can determine its empirical formula, but if we have the empirical formula, we can't know the exact molecular formula. Let's take glucose as an example. Glucose has a molecular formula of C6H12O6. Nitrous oxide, an anesthetic commonly known as laughing gas, has a molecular formula of N2O. Oh, so this is the molecular formula of glucose, okay. And we need to determine the empirical formula. To determine the empirical formula, we need to look for the largest common factor of each of the numbers that we find in the molecular formula. So, we have here the six, we have the two, we have the one. One possible common factor is going to be the two, but we can also have the three and the six. So those are three common factors, but we're going to use the largest common factor. We're going to divide all those numbers by that largest common factor. So, we can see that the largest common factor is going to be 6. 6 divided by 6 is 1, 12 divided by 6 is 2, and 6 divided by 6 is 1. That means that our empirical formula is going to be C H2O. This is the empirical formula for glucose. We just divide it by the largest common factor, which in this case is 6. So 6 divided by 6 is 1, and most of the time, when we're not going to write 1 because there's a presence of the symbol of the element, it means one. So, we're not going to write C1 H2O1. We just write C H2O. Now, what about the nitrous oxide N2O? The largest common factor between the 2 and the 1 is going to be 1. That's basically the lowest integral number that we can find. So in this case, the empirical formula is going to be equal to the molecular formula. Sometimes we can find that, as we saw here. So we can sometimes determine the empirical formula by using the largest common factor, or sometimes the molecular formula is also the empirical formula because we don't have a largest common factor other than 1. Let's do another example. This is the molecular formula for decaborane, so it's B10 H14. What will be the empirical formula for this? The largest common factor will be 2. So we'll divide 10 by 2, which is 5, and 14 divided by 2 is 7. So it will be B5 H7. That will be basically the empirical formula for decaborane. We also have different ways in which we can represent molecules. One is the structural formula. The structural formula shows the order in which atoms are attached and the three-dimensional shape of molecules. So in the structural formulas, the only thing that we know is how they are attached to each other. We have one hydrogen, one carbon, and one hydrogen. We see that one carbon is attached to four hydrogens as well. Here, one carbon, four hydrogens. Okay, now we also have the perspective drawings. The perspective drawings also show the three-dimensional arrangement of the atoms in a compound. These are also represented using models. We're going to use the left representation. So these three representations are the perspective. In this case, the wedge is basically out of the page, so it's in front of you. The dashed line is in the back side of the screen, and the solid line is in the plane of the screen. That's why this representation is different from this. This has some perspective because this hydrogen is in front, and this hydrogen is in the back part, while these hydrogens are in the plane of the screen. Also, this ball-and-stick model is basically the same as this. It's the same perspective. We have this hydrogen and this hydrogen in the plane, and this one is in front of you, and this one is backward of the plane of the screen. And this other one is a space-filling model. And here, the difference between this and this is that we can see here the bond between the hydrogen and the carbon. Here it's basically occupied. The space-filling model occupies that bond and creates that sense of the area where those electrons are found. These are the different ways that we can draw or represent molecules. The structural formulas just tell us the order, and the perspective drawing gives us a three-dimensional perspective of the molecules. Let's talk now about ions. When an atom or a group of atoms loses or gains electrons, it becomes an ion. So remember this. Let's review what we have learned so far about electrons, neutrons, and protons. The number of protons, if we change the number of protons, we're going to have different elements because the atomic number represents the characteristic of each element. So hydrogen is always going to be one, helium two. So if you have a different number of protons, you're talking about different elements. If you have a different number of neutrons, you're going to have different isotopes of the same element. Now, if you change the number of electrons, if you add or remove electrons, now you're going to have ions. You haven't changed the number of protons, so that means you're going to have the same element with a charge. If you gain electrons, it's going to be a negative charge. If you lose electrons, it's going to be a positive charge because the element is neutral. We saw before, for example, carbon that has six electrons and six protons. If you remove one electron, it's going to be positive one. Before, it was zero, but because you remove a negative, now you have a positive. On the other hand, if you add one electron, now you're going to have seven electrons. So you have one more electron than protons. So that carbon will be negative. In the end, you have created an ion. That's how you create ions when you lose or gain electrons. Cations are formed when at least one electron is lost. Monatomic cations are formed by metals. So when you lose electrons, at least one, you are forming a cation, which has a positive charge, and it's known as a cation. And all the metals produce cations. On the other hand, if you gain electrons, now you're going to create an anion, an ion with a negative charge. And most of the time, when you have monatomic anions, they're going to be formed by the nonmetals. So the nonmetals that we can find on the right side of the periodic table are going to produce the anions, while the metals on the left side are the ones that are going to produce the cations. So, for example, how did you need to be here? Because in this case, how did you gain one electron? I mean, you lost an electron, and this is positive. So it's similar to the group 1A. That's why hydroxide is part of the 1A. Now, let's write some chemical symbols for ions. If the chemical symbol includes a superscript indicating the mass number, for the ion with 22 protons, 26 neutrons, and 19 electrons, and also the ion for sulfur that has 16 neutrons and 18 electrons. Let's work with the first one. The atomic number is equal to the number of protons. We have 22 protons. So we go now to the periodic table and look for number 22. We're going to find titanium. Now that we have the symbol of the ion, we can look for the charge of that ion. We're going to look for the number of protons, which is 22, minus the number of electrons, which is 19. So 22 minus 19 is going to be...