Just a reminder with the narrated lectures, after you have viewed the lecture, make sure you go back through and examine some of the slides more closely, especially some of the slides that we might skim over more quickly during this abbreviated version. So starting in our second chapter, we're going to move into our first level of the biological hierarchy, and that's going to be with our chemical structures. So as we move through this chapter, one of our main focuses is going to be looking at the structure of the atom, as well as how atoms combine together to form things like molecules.
So all throughout this chapter, we're going to be seeing that initial theme of the emergent properties, that as you move up from one level to the next one, we start to produce some different functionalities. And we're going to see that right off the bat here, looking at some of our structures with the elements and the matter here, showing you that as things combine into larger forms, they do in fact change what their function can be. So begin at the most basic part with our chemistry we can see that everything around us is composed of matter and essentially we know there is matter because that material is taking up space and we can give it a weight we can take its mass for it even the air that is currently surrounding you we can say is composed of matter because we know that if we put that air into a jar and try to compress it we cannot compress it fully.
That's because there is material, these are matter there in that air, whether it's things like carbon dioxide or nitrogen or oxygen or whatever else, there is material making up its composition. Now, that matter itself is not just a bunch of nothing. It is composed of different types of elements. So we talked about things like carbon dioxide, that's carbon and oxygen. If we're looking at oxygen or we're looking at something like nitrogen in terms of other materials.
These are all different elements. And the basic idea with an element is this is the smallest that we can break matter down into and not actually go through and lose its basic properties. So it's the most basic we can get in terms of how we go about the process of moving up the biological hierarchy by using the elemental form.
So if we have an element like sodium here, you're going to find that based on its properties, sodium actually comes out to be a heavy metal. and it's also very toxic. Same thing with chlorine gas. This is an element in the most basic form.
It has a great toxicity. When you combine these two together though, so sodium and chlorine combine to make a reaction and produce what is called sodium chloride. Sodium chloride now we're going to find has no toxicity.
So this is one of our first ideas of the emergent properties. Combining two of the lowest low levels of biological hierarchy, we now can show this new form. which has a change in functionality. We're going from toxicity at the basic elemental form.
Now in what we call the compound form, we have the functionality of a seasoning for our food, but not having that same toxicity. Now, technically, as we get into these chemical structures later on, we're going to see that sodium chloride is in fact a compound, but also classified as what is called an ionic compound. And typically, as we look at materials, you're going to find that if we are designating something that's a compound, compound many times, not always, we will see that it's based on using what's called the ionic bond or ionic connection. As you look at things like water later on, we're going to see this is what's called a molecule. And molecules are based on the idea that's held together by a covalent bond.
But compound, again, is one of these things that as you're reading literature, as you're looking through the text, keep in mind it may not always illustrate exactly what's considered an ionic compound. A compound's most basic form is saying that we have at least two or more elements, and there's a fixed ratio. So water, which is H2O, technically fits that definition. We've got two or more elements.
We've got hydrogen and oxygen, and there's a fixed ratio. There's two hydrogens and one oxygen. But the idea is that as you're looking at this combination, whether it's compound, ionic compound, or molecule, we are moving up in that bilateral hierarchy, and we're seeing these emergence. new properties or this output of new properties or new functionalities for that material. Now in terms of the elements that make up our living systems, since we were in biology, we want to think about those primarily as the major constituents.
In our chart here on the right, you're going to see there's a whole series of different elements. The four most important ones for us though as a living organism are made up of oxygen, carbon, hydrogen, and nitrogen. And you can see here that in terms of the human body, That's about 96% of our overall body mass, and that does include water in terms of that makeup. This is part why oxygen here has such a great percentage, around 65% of the overall body mass.
Now, later on, as we get into Chapter 4, we will see that carbon really is the major element here, especially for living systems, because carbon is going to make up the backbone for things like carbohydrates and proteins, lipids. nucleic gases and then we're going to find that things like oxygen or hydrogen and nitrogen or other elements here will branch off from a carbon skeleton so these four elements the oxygen the carbon the hydrogen the nitrogen are actually four elements we're going to see the most throughout this semester one of the most abundant ones we find in most living organisms in our second set here about 4% of the overall body mass it's not that these are less important but they're in smaller quantities. And we'll talk later on about a process where we use our potassium and we use our sodium and what is called an active transport process. We'll see, of course, the sodium and our chlorine in terms of forming salts.
We'll talk about how phosphorus is a major part of things like our membranes of our cells or even how it's part of our DNA structure. Same idea with sulfur. Sulfur we're going to look at later on in terms of proteins.
how it forms a strong linkage in those compositions. So these percentages, even though they are smaller in their numbers, they're still equally as important. Same thing with our bottom one, these trace elements.
Trace elements mean that we're going to have these small or minute quantities. In this case, we can see less than 0.01% of the mass, but they're just as important. So something like iodine is a vital trace element in your diet. This is a lot of times we're going to get from seafood or other materials.
And if you're not getting that from that intended intake, you can also get it through things like salt. In fact, if you look at the condiment aisle in the grocery store, you're going to see salt oftentimes in two forms. Regular salt, and then you have iodized salt.
And it's just one way to get that extra iodine that's needed to help our bodies function. Same idea with things like iron. Take a box of cereal.
Look on the back in terms of the... levels of the calories and looking at your caloric input. In that list, you're going to see different levels also of some of these trace elements, one of which would be iron.
If you crush up that cereal and you run a magnet through it, you're going to pull out small little black flecks. Those little flecks are in fact the iron that your body's consuming and then using that again in certain metabolic processes. So this list of all these elements that we're seeing here, these are essentially the elements that we're going to find as the primary components of most living organisms.
Again, a majority of this will come up here from the first four, about 96% of overall mass coming from the oxygen, a carbon-based skeleton, additions of hydrogen, and additions of nitrogen. So moving up with our hierarchy, if we know everything around us is made of matter, and we know that matter consists of elements, we also want to see that the most basic form of an element, away from looking at, say, a gram of gold or looking at that... say two kilograms of sodium whatever it might be the basic unit of all that matter is in fact the atom so the different forms of elements out there they're going to have a certain atomic makeup which gives them their set of properties which also means as you recombine those into a compound or a molecule they're going to be affecting how the atoms themselves are going to interact now with the atom we're going to see we could break it down further into smaller pieces But if we did so, we would lose all properties of the intended atom of that element. So breaking down an atom of sodium or an atom of chlorine or an atom of carbon or oxygen down into these subatomic particles, you would not have that given property of being toxic or being skeletal makeup for most of our we call them macromolecules.
These subatomic particles, though, given their certain ratios, are what provide the properties for that elemental structure. So we've got ones that are called neutrons, we've got some protons, and we've got electrons. Now in neutrons, these are going to carry what's considered a neutral charge, or no electrical charge. Protons, we're going to see as a positive charge.
Electrons, we're going to get a negative charge. So just be careful when trying to remember something like the neutrons, protons, and electrons, because one of the first things you might do is say, oh gee, protons, P on the proton, P on the positive. Then you get the neutrons, and you go, ooh. N. That means negative.
Be careful though, because in the electrons here... These are really the negative charge ones. When you're looking at this in terms of trying to do word association, you can do positive with protons.
When you get to neutrons, though, instead of thinking negative, think about neutral. The N-E-U-T, in terms of neutrons here, could also go into the N-E-U-T for the neutral. In fact, you can go to the R if you want.
It's got the whole first part saying that we've got no charge or we've got a neutral association. The electrons then have a negative part. Maybe you think about the fact that electricity or electrons kind of sound the same here.
Electricity can be bad if you touched it. You have a net of interaction on yourself. And looking at these in terms of their associated charges or lack thereof, and using word association does help in some cases, but just be careful you don't mess up the neutrons are neutral and that they're not negative.
Now, the two structures of the proton, the neutrons, are actually larger in size than electrons. And we're going to find these first two structures. held within what's called the atomic nucleus.
Now we talked about the nucleus in the first chapter in terms of cell configuration. This is not the same in terms of the atom. It's a similar idea in the fact that the nucleus is kind of a central structure in terms of the cell and the atom, but they're composed of entirely different kind of materials for us. So with the atom's makeup, the inner structure is in fact the atomic nucleus, and then surrounding that would be are set in what we call the electron shells.
And that's where we're going to find these little clouds of electrons circling the actual atomic nucleus. So two ways, or actually multiple ways, we can represent the atom. The one on the left side, which is our part A, is more kind of a theoretical view, where you've got the atomic nucleus in the middle with your positive protons and your neutral neutrons. And then surrounding it in kind of a three-dimensional cloud, we have these two associated electrons.
In the Part B version, then, we see that we do have the same exact configuration in the middle with two positive protons, two neutral neutrons. And then surrounding that, we have the electrons in the fellow electron shell. The B version is probably the one you're going to see most frequently throughout the chapter, because this gives us a way to designate and show how many exact electrons are there.
In the one on the Part A, you're going to see that the electrons... would be identified in this cloud, but they're not put there in the exact form we're showing with that negative charge. We have to imagine that these little electrons kind of are floating around that nucleus at any one given point around the atomic structure here with helium.
Now in terms of actually designating how these atoms are going to look, you're going to find that every single atom is kind of unique in terms of the number of protons, the neutrons, and the electrons. The most important part out of all those subatomic particles, though, is the number of the protons. This is the unique identity for that element. So this number of protons is actually what's designated as the atomic number. And this is one of the items you're going to find usually predominantly displayed on the periodic table.
It's usually the top number above the symbol for the element. The other number you're going to see on the periodic table is typically going to be what's called the atomic mass. And the atomic mass is all of the subatomic particles. And it's the total mass associated with the element.
Now, this atomic mass here is what we're going to find sometimes maybe a little bit different than what we call the mass number. The mass number is only the ratio here of, or the sum of the protons being positive and the neutrons being neutral. does not account for the electrons. Now to get to our example here in just a moment with the periodic table and showing some of these numbers, we're going to look at the mass number and we're going to use this as a way to diagram the actual atoms.
But the number you're going to see on the periodic table is the atomic mass, which is your total mass of all materials. And we're going to have to go through and approximate this to give us this mass number. So we're going to round up or round down to the nearest whole number.
depending on where we're at in those ratios. So first little example here with oxygen. You can see that we've got the atomic number usually at the top.
So this is your number of protons. And then we have the atomic mass, 15.9994. This is all the protons, all the neutrons, all the electrons combined to give us this mass. If we round that to the nearest whole number, in this case being 16, you would find that this would give us then the atomic I'm sorry, the mass number of 16. which tells us there are a total of 16 structures combined with the protons and neutrons. So using these two numbers, we can very easily decode how many neutrons, and also we can infer how many electrons are making up this structure.
So in this little number here being 16, if we do 16 minus the 8, because again the 16 here is the protons plus neutrons, If we subtract away the number of protons, we will be left with the number of neutrons. And this case happens to come out that we end up with eight neutrons. The number of protons and neutrons that will not always be the same for every single element. You're going to find as you get further down on this chart, especially getting beyond the oxygen here, that many times this mass number gets larger in such a way that when you subtract away that atomic number or the number of protons, your neutron number.
much greater. You're going to see as you go further further down the chart the neutron amount gets greater and greater compared to that proton number as you move along in progression. The only thing left for us to figure out is number of electrons. One thing you're going to see on the periodic table is that there is no associated charge listed on this elemental symbol.
So there's no positive, there's no negative up here in this corner. That means that this essentially is neutral in its formation. And that means that if we have eight positive protons, we have to have also eight negative charged electrons to balance it out. And we'll come back to that in a moment here when we talk about things like the electron shells and give you some more illustrations about configuring out what the atomic mass is and what the atomic number is to see kind of some scenarios for us.
Now, one other thing you will see as you move throughout these chapters. is that a lot of times as we look at a three-dimensional structure of a molecule, we're going to use these little spherical forms of atoms as a designation. And the first four are really the four we want to think about.
The red coloration for oxygen, the light blue for nitrogen, the kind of white or light gray for hydrogen, and the darker gray or maybe black for the carbon. The symbols are on here now, but later on you're going to see that symbols won't be on the molecular structures. This is one way you kind of think about this combination that we see this dark one, we're going to associate that with carbon, the light or the white is hydrogen, the red oxygen, and the light blue is nitrogen. Now, give an idea here with the neutrons.
This is actually the only time we're going to talk about what are called isotopes, because in our view, as we look at biology, we're not going to spend a lot of time going into some more of the practical concepts. But if you're going into further courses, maybe like microbiology, going into anatomy course, maybe even going into something like a genetics course, you might come back and revisit some of the ideas about isotopes. But the essential idea with isotopes is that fact that these are corresponding to the same exact element, but varying in the number of neutrons.
So we had the oxygen previously, and if we go back and look at that number just for a moment, we can see oxygen. has eight protons, and we've got the mass here about 16. If this mass read as 17, that number of protons is not going to change, but your neutron number is going to go up by one. So instead of being just eight protons and eight neutrons, it's now going to be eight protons but nine neutrons to account for that extra mass there in your atomic mass or that mass number.
So having these extra neutrons doesn't do anything in terms of the charge, but does on extra weight to that particular element. And in some scenarios, that additional weight added on to the atom produces what are called the radioactive energy, or the ability to get off energy from that structure. This is very useful in terms of trying to look at the data of a fossil in terms of comparing what the soil is around it, tracing a process in terms of metabolism, seeing where different elements end up.
based on how they're marked with different isotopes, or even looking at things like diagnosis of maybe a cancerous placenta and showing the isotopes fluoresce under the right conditions for it. So another example we can look at here is carbon. Now, just to give an idea again about the proton number, neutron number, electron number on your periodic table, for the symbol for carbon, you would see a six at the top indicating the six protons. and roughly a mass number of about 12. So it's pretty close to atomic mass of 12 here in terms of those numbers. So doing that math again, doing the 6, which is your total mass of protons and neutrons.
sorry 12 total mass of protons and neutrons and subtracting that atomic number which is six that'll give you the six neutrons again because we have six positive charged protons we can infer then that we should have six negative charged electrons here so this first one is your basic structure for carbon and we'll come back to see how this diagram is going to work in the next couple slides here now we're looking at carbon 13 as a different isotope for carbon 14 you're going to see atomic number didn't change. It's the carbon's elemental symbol in terms of being six protons. Your electron number did not change.
We're still balanced out with positive and negative. But your mass number here is going from 13 to 14 showing that we've added on one extra neutron in carbon 13 or two extra neutrons in carbon 14. So we've talked about positive charge protons as the identity of the element and we've talked about neutron potential to give us the form with the neutral charge as isotopes. The last thing to look at now, which is really where we're going to get into the rest of our chemical structure, and that's the electrons. And associated with electrons with negative charge, we do come into the view of energy or its ability for things to cause change. Now in the course of our semester, we're going to look at two forms of energy.
In the chapter today, we're going to focus on potential energy. So the energy based on structure or location. So for molecular forms or the placement of electrons around an atom, this is associated with potential energy.
So every single one of the elements that we look at, given their number of electrons, will have an associated amount of energy or potential energy with them, that as you combine that into a molecule or a compound, we start to increase the energy, but we also change of how it's going to be actually shared among the associated atoms. elements there or associated atoms that make up that molecule. So as we saw in a little diagram, the electrons are found in what is called the electron shell. And the electron shells are actually one way we can start to kind of visualize this potential energy, because based on where the electrons are in a certain structure with the atom, if they happen to move out to further electron shells or further energy levels per se.
This means that they're going to go through and absorb energy to move them in that further direction. If they fall back down levels, though, this can be associated with potential energy that's now being lost as we're going back to a different position. Now, it's not going to make as much sense today with this kind of rise and fall of the electrons where they're absorbing energy and then losing energy.
We're going to come back to this idea, though, as kind of a fundamental principle as we look at photosynthesis. and revisiting this idea that we can make electrons move out to absorb energy and have them drop back down to pass off the energy or lose the energy and terminate that process. One way to associate potential energies today though is kind of using our little stair technique up top. And on the stairs you're going to find that if this ball is at the highest position we would associate that with a very high amount of potential energy. Kind of associated down here with the electron.
further away from the atomic nucleus. As that ball drops down from step to step to floor, you're going to find that every drop down is a reduction in the amount of potential energy. Because now if that ball is on that middle step or that bottom step, it's not going to have as far to fall, which is not going to give off as much energy as it would if it was on the very top step. The same thing that we can look at in this diagram is the fact that if the ball was down here at the base or on the floor... and you want to put it back up to that first, the second, or that top step, you're going to have to go through and absorb energy.
You're going to have to put energy in to move that ball back up. And that's essentially what we're going to think about with our electrons and this potential energy as we progress through this chapter. That as we change their placement, whether bringing them closer into a nucleus or sending them farther out, or even how we share them among multiple atoms, you're going to see a change in potential energy. So when it comes to diagramming these different atoms, we've talked about the atomic nucleus in terms of the protons and neutrons, where the proton number is the atomic number and your number of neutrons and protons is that mass number.
We can do that very basic math of the mass number minus the atomic number to give us the number of neutrons. And we associated in the fact that if we add a set number of positive protons, we have to balance it with a set number of negative charged electrons. So the number of electrons that we see with every element is actually associated with the chemical connections we can make and how we form molecules or compounds or other associations based on their distribution.
So this kind of modified periodic table here only consists of the first three rows. This is actually where most of the elements that we will encounter are found. In fact, you can see the ones highlighted in red here as the primary building blocks that made up about 96% of all of our body's mass. Now, do not go home and try or do not sit there and try to memorize every single one of these distributions.
I do want you to be familiar though with the ones highlighted in red because this is going to help you when you try to think about structures like our macromolecules or carbohydrates or proteins or lipids in terms of how conductions are made and how a molecule might be stabilized. based on the given number of connections. But you will not be asked ever to go through and actually have to draw these, but you may have to go through and use the given material, like the atomic number or the mass number, to figure out is this distribution of electrons correct in terms of how matter in the first shell, the second shell, and the third shell. Now, before we go back into these four elements of kind of interest here, a couple of rules to look at in terms of distribution. In each one of these yellow lines around the atomic nucleus, these are your electron shells.
These are the different energy levels associated with these atoms. In your very first electron shell, you can only have a maximum, as we see here in helium, of two electrons. In your second shell, you're going to see the very first shell is already filled up with two electrons at its maximum, but the outer shell can only have a maximum of eight. Same idea with the third shell.
First and second shell are already filled, and now your third shell is at capacity, or trying to reach capacity, of the eight electrons. Now, you're going to notice here that the only ones that have reached full capacity is helium, neon, and argon. And when an element is fully complete, in terms of the first shell, or the first and second, or the first, second, and third, we would say that these elements are non-reactive. or we could say also as inert, which basically means they do not want to share, they don't want to donate, they don't want to receive any electrons. They don't want to go through and work a chemical reaction.
But everything else, though, you will notice is missing at least one electron, as in case of hydrogen or chlorine, or maybe missing seven electrons, in case of lithium and sodium. So these, everything except the last three in that final row there, everything else is missing. not complete. It's not stable in its form. And the goal to make them stable is the formation of chemical bonds.
These could be covalent bonds, could be ionic bonds to make them complete. So first shell, maximum of two electrons. Second shell, maximum of eight. Third shell, maximum of eight.
Now one thing we can designate here in terms of distribution, even though every line is an electron shell, the outermost one that we're trying to complete is what is called the valence shell. So this is a valence shell for hydrogen, valence shell for helium. The second shell here for each one of these is the valence shell.
The third shell for each one of these is the valence shell. The valence shell usually is what is considered the reactive portion of the atom in order for us to go through and make that chemical connection. So focusing on the four main elements for us, hydrogen you can see has only one electron in its valence shell. which means it needs one more to be complete. Carbon has four electrons in its valence shell.
It needs four more to be complete. Nitrogen has a total of five electrons in its valence shell. It needs three more to be complete. And oxygen has a total of six electrons in the valence shell.
It needs two more to be complete. Now, this missing number of electrons, the two for oxygen, three for nitrogen, four for carbon, or the one for hydrogen, is also what's considered the valence for this atom. And we'll see later on that the valence corresponds to the actual number of bonds that have to be made to make that element stable, or make the atom stable in terms of the electrons coming in from other structures with that form.
So again, if we give you a number like the atomic number and the mass number, you should be able to go through and figure out how many protons and neutrons are in that nucleus. And then given that number of protons that we've got, You can also infer how many electrons should be around that atomic structure and where they're placed. So in this case, two in the first shell, four in that second shell based on the fact of having six positive protons and now six negative charge electrons. So again, like we said, that very outer shell, this is what's called a valence shell. This also gives us what's called a valence capacity.
In the case of neon here where that valence shell is filled, the valence number is zero. and that corresponds to the fact that this element is chemically inert. If that valence shell is filled, it does not want to make any other interactions. So every other element in terms of carbon or nitrogen or oxygen or chlorine, which is not filled, has the potential or the capability to go off and form some kind of chemical interaction.
Now, one thing to kind of give us a little different vision here just for a moment is the fact that even though we show those electrons either as a cloud or now as an electron shell, in reality, they are actually found in what are called orbitals. And the orbital, like that cloud, is a three-dimensional area where the electrons themselves are found about 90% of their time. So your very first electron shell actually corresponds to more of a spherical-type structure.
In your second shell, we have a sphere. We also have what are called these little p-orbitals. almost like little infinity symbols or little dumbbell shapes where you find the electrons to find along this x the y and that z range in terms of placement now eventually we're going to see towards the end of this chapter that if you go through and superimpose all the available orbitals on top of themselves you're going to find that this particular atomic structure actually goes to what's called a hybridization which means that some of these little p orbitals here actually going to bend a little bit and things are going to fuse together in a certain way to give a distinct shape to the molecules that we are forming based on elements that are there so we know proton number we know neutron number we know electron number given that electron number now we're going to move forward with looking at chemical bonds and associating how that functionality or the missing number of electrons will configure into larger and larger forms or movements up that biological hierarchy.
So very first one we're going to count in terms of chemical bonds is the covalent bond. And for our considerations this semester, this is your strongest bond. This is the one that if you want to hold a molecule together, you want to use a covalent bond. And that's because with covalent bonds, this is considered a sharing of electrons.
So in the case of our little hydrogen molecule over here, you're going to find that this association with the two clouds kind of melding together is the sharing of electrons. There's one electron around the first hydrogen atom, one electron around the second one. As they come closer together, those electrons are now being shared in the orbiting areas around each one of these.
At any given point in time, you could have two electrons around this one and none over here. We could have two around this one and none over here. We're going to go back and forth in terms of the sharing.
We now made a chemical connection in the form of a covalent bond. So one pair of electrons being shared among the outermost electron shell or that valence shell is making up this covalent bond. So if, in fact, we are making covalent bonds, this is what's going to give us the building of what are considered molecules.
So in the very beginning, we said compound, two or more elements with a fixed ratio. And molecules... We're going to use that kind of same idea that we have some kind of ratio, and we've got two or more atoms now, but you notice we don't have to have two or more elements.
So hydrogen is a molecule, but hydrogen molecule would not also be a compound. Even though it's got two atoms, it's not two different elements. But water, we could say is a molecule because of the covalent bond we're going to find with it, which also might consider a compound because it is a fixed ratio with two hydrogen and one oxygen. suffices those two forms with the number of elements. Now for most of our bonding capabilities, you're going to see that the equivalent bond here is a single bond with one pair of electrons being shared or a double bond with two pairs of electrons being shared.
In most scenarios, you're going to see single bonds, especially between things used like nitrogen and hydrogen and a lot of times a carbon, but you will also see double bonds a lot of times. between oxygen and carbon, oxygen and oxygen, and sometimes maybe even between like a phosphorus and a carbon-based structure. So given a couple ideas of how these molecules start to look, especially using single bonds or potentially double bonds, you're going to find just like the way we draw the atom has multiple ways, there are many ways to illustrate the molecular form of these molecules. So first thing, of course, is the naming. We talked about the hydrogen molecule first off.
The second part here is what's called the molecular formula. It tells you how many of each element is there. there's two hydrogens or O2 there's two oxygens H2O two hydrogens one oxygen CH4 one carbon four hydrogens the second part looks at the distribution of electrons so illustrating where the valence shell starts to overlap and how we're sharing either a pair of electrons here or two pairs of electrons or pair here and a pair here or pairs all the way around with the carbon structure Now the electron distribution is probably one of the more suitable ways to think about this if we want to be able to show every single electron and all the different electron shells.
But it's very time consuming. So the Lewis dot and the structural formula are ways to kind of abbreviate that formation. Your Lewis dot illustrates the pair of electrons or the two pairs that are being shared where the structural formula focuses on taking that pair and putting it as a line formation. So a single line here corresponds just to one single bond.
Two lines correspond now to a double bond. So there's one pair of electrons being shared here. There's two pairs of electrons being shared in this one.
The last part here of the space-filling model is more of a three-dimensional shape. This is where a lot of our form and the functionality will actually be evident as we start to go into larger and larger molecular structures. Now, in most cases, if you're looking at drawings in the text, you're probably going to be seeing things like the structural formula for a lot of the space-flowing models.
And if you were to draw one, usually that structural formula is the fastest way to draw that composition to illustrate how things are connected. So again, looking back at that periodic table, all of these outside valence shells that are not complete, everything from hydrogen, from lithium to fluorine, sodium to chlorine, their valence shell is not at complete. capacity.
Again, that number of electrons that are missing is what is called the atom's valence. So fluorine has one missing. That's a valence of one.
Oxygen has two missing, valence of two. Nitrogen has a valence of three. Carbon, valence of four. Boron, valence of five.
Beryllium, valence of six. Lithium, valence of seven. Now, one of the things we will see with this valence capacity is that typically a valence of four.
the greatest amount in which the elements themselves can kind of connect with. You're not going to find things like boron, beryllium, lithium, sodium, magnesium, and aluminum tending to form these kind of covalent bonds. Typically the carbon through the fluorine or the silicon through the chlorine are better candidates based on this little lower number of the valence capacity. We also see that this kind of unique number of four bonds here with carbon also comes back into play with the orbitals and how they hybridize together for that three-dimensional shape.
So as we get into our compound here, you're going to see that again, compound is being used in a very simple form. Combination of two or more different elements. In a molecule, there was two atoms coming together, but a compound strictly is now more combination with two or more different elements.
And a lot of times we can have a covalent bond there. But typically if we say there's a molecule of some structure, that's one way to illustrate that we're dealing with this covalent bond connection as opposed to, say, an ionic bond connection. Now one last thing we can look at with our covalent bonds is that sometimes when the electrons are being shared between the two valence shells, they don't get shared equally.
And a lot of times that comes down to the fact that the atom itself is what is called electronegative. Now in terms of all the molecules that we look at, the only two atoms you have to think about as being electronegative are oxygen and nitrogen. And with each one of these atoms, you're going to find that they tend to pull the electrons in the outer valence shell more closely towards their atomic nucleus than say something like carbon or silicon or fluorine.
That stronger pull results in a change in where the electrons are being shared. among the molecule that's being formed. So if we have an atom or a set of atoms that are all electronegative or are not electronegative, you're going to find that we produce what is called a polar covalent bond.
If we go back and look at the structure here with our methane, you're going to find that carbon here is not electronegative. Neither is hydrogen. That means these bonds being shared between the carbon and the hydrogens all the way around are nice equal sharings. In fact, if we modify this little structural formula and put a pair of dots here and a pair of dots here and a pair here and a pair over here, you would see that those pairs are equally shared. In this oxygen, though, you're going to find it's actually electronegative, which means it's not going to share the electrons equally.
So the little pair of the electrons in the line here or the pair being shown here, you might find those electrons are closer towards the oxygen in terms of placement. And that's going to result in forming that covalent bond still, but now in what's called a polar covalent bond. So most of our molecules, especially with carbon as the kind of skeletal structure, these will form non-polar covalent bonds, where the electrons are being equally shared between the two different elements.
With a polar covalent bond, though, this is where, again, you've got one atom, which is more electronegative, in this case the oxygen, and the electrons are not being shared equally. That unequal sharing results in a scenario where now the electrons are further away from this hydrogen. or further away from that hydrogen, resulted in a very slight positive charge on each one of those atoms of hydrogen. And if there's slightly more electrons now around the atomic nucleus of the oxygen, it picks up a slightly more negative charge there. Now, we have to be careful here, because this slight negative versus slight positive is not the same as we're going to see with ions, or as a full negative or a full positive.
The association here is that this creates this kind of imbalance in where the electrons are. Now you might be thinking, well, gee, I only have one negative but two positives. How does that kind of balance out? One thing you'll see if we go back to our distribution with water is that technically there are two pairs of electrons here in this autovalent shell.
And usually we'll find an associated negative polarity on this pair of electrons, negative polarity here. and then slight positive here and slight positive there. So actually there are four locations of polarity around a water molecule. And we'll come back and revisit that in Chapter 3, looking at water and its properties that are based on this polar covalent bond connection and this polarity now as a structure.
So if the atom of oxygen is there or if the atom of nitrogen is there, one thing you might see, especially with the availability of hydrogen around it, is that it's going to form these polar covalent bonds. which means some part of that oxygen atom or some part of the nitrogen atom might be slightly negative, and some of the hydrons might be slightly positive. And that will change how we form other interactions in terms of chemical bonds later on. So if we have carbon, which can form kind of the ultimate number of covalent bonds, being 4, what happens with boron or magnesium or aluminum, which have that higher valence number, so valences of 5, 6, or 7?
how are they going to go about the process of fulfilling that chemical bond connection that's needed to balance out those electron shells? This is where what are called ionic bonds come in. This is also where we're going to form what are called ionic compounds.
So if we don't have enough space to share electrons, or maybe it's just not feasibly possible at that point in time to share electrons, we have the chance for many atoms to go through and receive or to donate the electrons from their valence shell to reach that full capacity. So transferring electrons out or bringing electrons back in. This will now change the charge of the atom into what is called an ion. So the first example to look at here is sodium and chlorine. In sodium, you can see the first valence electron shell is filled.
We've got eight in a second electron shell. You've got one in the outer valence shell. Now, in a sodium atom, in order for it to make covalent bonds, you would have to form seven more pairs of electrons being shared in terms of covalent bond capacity.
Now, maybe in a really strange lab scenario, you might find that take place. But under natural conditions, conditions... you're not going to really see the feasibility of taking seven other electrons and sharing it with that sodium because typically that comes from maybe seven other elements or seven other atoms surrounding that. There's not just enough space to actually make those connections.
So the one kind of alternative for sodium, instead of trying to fulfill that valence capacity, is for it to actually give one up. And as it gives up that one electron, it drops down to that second shell now, which now becomes that valence shell, and it's already filled. So that sodium atom is now stable in what's considered a sodium ion. In this case a cation, which shows this positive charge. So now when you look at the number of protons in sodium and the number of electrons, they're no longer equal.
So if we're looking back over in the first example, you can see the two in the first shell, the eight in the second shell, that's 10, and the last one, there's 11 electrons here, and that's associated with 11 protons in that nucleus. Over here, you have 11 protons, but now only 10 electrons. So this is where that single positive charge came from.
If sodium had lost two electrons for some reason, you would see it as a 2+, because you have lost now two negative charges associated with the number of positive charges. Now, chlorine, on the other hand, has two ways of going about forming chemical connections. It can go about the receiving of electrons that we're seeing here, forming this chloride ion, or it could share.
With only having a valence of one in the outer shell, so only one electron missing, it's got options in terms of what kind of chemical bond it wants to make. whether covalent bond or ionic bonds. So because sodium's giving up the electrons here to reach instability, chlorine's gonna just take those in.
And when that chlorine picks up one extra electron, you're now gonna find there is one extra negative charge around that atomic structure compared to the number of positive charges within that nucleus. This is where that negative charge comes from and formation now of what's called this chloride ion or this anion. Now covalent bond is your strongest bond.
Your ionic bond is a little bit weaker. And it's weaker in that sense because we're now just attracting two opposite charges together to form this interaction. And think about it as like two magnets.
Magnets are kind of that same idea where the opposite poles attract and we get that connection. Yes, some magnets are very strong, very hard to pull apart, but you can get them apart given the right scenario. The same thing is the case for ionic bonds.
Covalent bonds, we can break them, but they're very resilient because they're sharing electrons. Ionic bonds, though, they're not sharing anything. It's just a positive associated with a negative.
So there is no actual change there in terms of that formation. So ionic bonds, by far, are usually a more weaker bond compared to, say, covalent bonds. So when you take a whole bunch of ionic bonds and combine them together, we now form into what we call the ionic compounds. And ionic compounds are associated in this format that, again, we have a set number of elements and we have a fixed ratio.
So it does meet that compound kind of definition. But now we're going to see that it's kind of an alternating pattern. Here we have the positive sodium in green and that negative chlorine in that yellow. So in terms of forming this larger structure, it's no longer a molecule we're looking at, but this ionic compound. Now one thing we will see with ionic compounds is because the ionic bonds are slightly weaker than the valenic bonds, if we associate this ionic compound in a white condition, say a water-based solution, that water can be used to very easily separate out the ions into individual structures, just because we have this positive and negative association.
Now, as we move on to our chemical bonds here, we go into our weakest ones. And one thing we want to illustrate with the weaker bonds is that even though their strength is not as great as, say, ionic or that greatest covalent bond, these weak interactions are just as important. And we'll see that as part of the case, especially with hydrogen bonds, as we move into things like DNA or we look at things like proteins.
Because even though the bond itself is weak, the interactions we get and the stability that we form... is a much better kind of reinforcement based on using these weak chemical bonds. The advantage, though, with a weak chemical bond is you have the ability to break it more easily if needed to.
And we'll kind of illustrate that later on as we get into some of these properties, not just with water, but also some of our larger macromolecules. So hydrogen bond is our third bond we're looking at for chemical interactions. But unlike the ionic bond and the covalent bond, which forms some kind of structure, hydrogen bonds are not going to form necessarily the combination of atoms into a structure, but it's there to attract different molecules or different ions to each other.
So it's not really there to build a structure, but it's there to reinforce or to add into the formation. in combination with ionic bonds or covalent bonds. Now, the one way we form hydrogen bonds is by going back and have that polarity. So hydrogen bonds are always going to be associated with some type of polar molecule, which means you're going to have some form of a polar covalent bond, usually with that connection.
In the case here of our water molecule, remember oxygen is electronegative. In the case of our ammonia molecule, the nitrogen is electronegative. So in each one of these molecules being formed, we do produce this polarity. We have the three positive or slightly positive poles here and one negative pole.
We're only showing two positive ones here and only one negative, but really there are two negative poles on that oxygen within that water molecule. Now, when you see the hydrogen bond, you'll see it take place between one of these. partially positive charged hydrogens and then a negative charged other atom using nitrogen or oxygen in terms of the makeup.
In fact, in chapter three, as we go into our water molecule properties, we will see that hydrogen bonds are formed between hydrogen and oxygen or the oxygen of a different molecule again and again across all these small interactions with that structure. Now, the hydrogen bond is recognized here as a dotted line to illustrate the fact that it's not a permanent connection or not a strong connection. This is a weak interaction between these two associated parts of these molecules as a way of keeping them closer together for us.
Now even weaker than hydrogen bonds are what are called van der Waals interactions. And like hydrogen bonds where you've got an association of kind of a positive and negative, in a van der Waals interaction this is actually the result of what are called little hot spots. The hotspots are where a molecule, maybe even a compound, are going through and changing just a little bit in their structure.
As they fold in or open up a little bit, that puts atoms closer or farther apart. That closeness or that distance creates these little positive or negative areas that are ever so momentarily present that it creates some of these interactions. Again, like the hydrogen bond, even though the Van der Waals interactions are very weak, when they combine in the force of the millions upon millions, it creates a really, really strong effect within that interconnection.
And I'll show you just how collectively strong these little hydrogen bonds are. If you look at the toe of a gecko, where you've got these little ridges, the ridges have these little projections, and the projections are even then flattened out further. and like a little paintbrush that's been squashed out. These little connecting points with those toe pads are what allows that gecko to adhere to a surface, but also to very easily break that connection.
So just imagine yourself trying to walk up a wall. And if you're forming covalent bonds, kind of equate that to putting a superglue on your hand and trying to move your hand up the wall by supergluing again and again. More than likely, it's going to be very hard to...
dislodge that hand. They're going to be pulling skin away from your hand as they're trying to break that connection between the wall and your hand. With a gecko though using hydrogen bonds, because the hydrogen bonds are so weak, by a little added effort they can break that bond, but then quickly restore it.
As you're moving that toe or that foot or that leg from one position to the next one. So covalent bonds are the strongest bonds. Ionic bonds are a little bit weaker.
hydrogen bonds even weaker, van der Waals one of our weakest. But even with our weak interactions, you're going to find that they are there in many of our larger molecules as a way to reinforce a shape or to hold it together momentarily and allow it to be broken easy under the right conditions. Now going back to our shape here just for a moment and looking back at those orbitals, we did say that as those orbitals line up, they undergo what's called this hybridization. The hybridization here essentially shows that as those spheres and p-orbitals line up, it produces what's considered kind of a tetrahedral shape.
And this is one of the things that as we start to build into our larger macromolecules. we will see as evident in the three-dimensional structure. So looking at the water molecule here, you're going to see that we always draw it in kind of this angled portion.
And that's because when you look at the orbitals, two orbitals are always occupied by these other two pairs of electrons that aren't bound up with a chemical bond. That leaves then the other orbitals here that are hybridized to make the connection between the hydrodome. you will always see that this water molecule creates this 104.5 degree angle between the hydrogen and the oxygen here in terms of that connection.
It's not going to be a straight line. It's more evident even for the tetrahedral shape with methane. With the carbon at the middle in these four hydrodons, you're going to find that it produces this more pyramid-based shape based on where those orbitals end up. This is what you're going to see a lot, especially as you go into a larger form.
carbohydrate or protein even nucleic acids but they're not flat structures they in fact are three-dimensional in how they interact now in our very first chapter we talked about the fact that form and function or shape and function are related and if we change the shape you're going to change the function in this example we're seeing two molecules that look pretty different you've got the naturally made endorphin quite a large molecule and it's got sulfur in it, it's got a lot of nitrogen, it's got some carbon, it's got oxygen, it's got hydrogen. And then you see the man-made or the lab-made morphine, which is a lot smaller in size. At first glance, just due to the shape and structure, you might think that these do not do anything near the same effect in terms of interactions with the body.
But when you're looking at the interaction, one thing to keep in mind is that the entire molecule doesn't always have to react with a certain structure. In fact, in the little outlined area here in endorphins, in the outlined area in the morphine, you're going to see that this little substructure of that larger molecule is the only part that binds to these little proteins as part of the interaction. That's the only part that's going to go through and actually make this change, where in the brain, these two molecules can block that pain-recepting signal to keep it from reaching the brain, keep you from feeling that pain given, say, a certain injury. But when you're looking at even those two little areas, you can see that in morphine, there's another little oxygen right off that side.
So it's not even the whole entire carbon structure that we're looking at that interact. It's the little hydrogen and the oxygen and part of that carbon skeleton, which is binding to these little endorphin receptors as a way of making that connection. So keep that in mind as you look at other chemical structures. They might look very different, but if they share some commonality in terms of small bits and pieces being the same, they might react in a similar way due to that small little section being a part that interacts with something in that chemical interactions.
All right, very last thing to look at in this chapter is taking us kind of the full circle. So we talked about the atoms and the matter in the beginning, so the smaller pieces, and we've talked about how we can form different kind of bonds like covalent bonds or ionic bonds or hydrogen bonds. The last step then is to see kind of the next flow of the hierarchy. So we're going from the basic atom or element into the molecule. How do we go about that?
So when you're forming any kind of chemical bond, you are carrying out what's considered a chemical reaction. It's the making and the breaking of these chemical bonds. Now, for any chemical reaction, there's a few rules you have to follow. The first thing is that with these reactions, you have to start with some kind of molecule or even a compound, and you have to end. with a molecule or a compound.
You cannot start with an individual atom or an individual ion. You want to be able to go through and show that everything is kind of built as it should. In this case here with a water molecule, we could start with a single hydrogen molecule.
We could not, however, start with a single oxygen atom though. You can't find those just off the shelf. We have to start with a water oxygen molecule in order for us to go through and form that water molecule.
The other part that's required for chemical reactions is they have to be balanced, which means whatever goes in, so if you've got four hydrogens here and two oxygens there, you have to produce them as well. So here's four hydrogens, here's the two oxygens. Whatever you begin with, you have to end with. You can't lose anything, you can't gain anything extra. Now typically with chemical reactions, you'll read them in kind of the same way you're reading the text across the screen here, in a left-to-right fashion.
Typically, reactants will fall on the left-hand side of the reaction, and your products will fall on the right-hand side. One way to verify you're looking at the correct products though is also to look for the arrow. The arrow will always point to your products. So the hydrogen molecules and the oxygen molecule here, these are your reactants that we're going to break apart.
And then as you're breaking them apart and reforming into our products of the water molecules, that is the complete reaction. So reactants, the starting material, products, the end material. Now this is one thing you want to kind of remember because as we get into our metabolism with respiration and photosynthesis, you will be asked to know what the reactants and products are for both those equations. So it's essentially, you kind of get the idea down now.
So later on, if you're asked to pick out the reactants and products for that reaction, you know where to find them and to use the arrows in a way to figure out where they're located in that reaction. Now one other thing you will see with chemical reactions is that even though the arrow goes one direction, technically we can go in reverse. And that's because with chemical reactions, a forward reaction can also be reversed to produce the opposite or the return. So combining hydrogen and oxygen to form a water molecule as the one reaction, if we flip the arrow and then going from right to left, we could start with water as a reactant and produce. oxygen and hydrogen as the product.
Again, because we have balanced the equation there, the forward and reverse could be happening at the exact same rate. And if they are, we have reached what is called chemical equilibrium. So your forward reaction and reverse reaction, so reversing with the products and reactants are, is at the exact same speed.
Now, the last equation we're looking at here is photosynthesis. where on the reactant side you've got carbon dioxide and you've got water the product side you've got what's considered a sugar molecule or glucose and you've got oxygen now we will see later on that in respiration it's a reversal of this arrow you would start with glucose and the oxygen as your reactants and you would produce water and carbon dioxide as your product Now, even though we can go through and show both those reactions later on and potentially kind of say this is reaching equilibrium, one thing we will see with most biological reactions, so ones that happen within the human or a living system, when we go off and produce this sugar molecule and this oxygen molecule, you're going to find that they will not always stay with that reaction. Instead, you're going to find that sometimes that element actually removed. It's actually taken away and used someplace else.
So you're looking at our little diagram here with this plant. This is what's actually called Elodea. These little bubbles are the oxygen leaving that reaction.
So yes, we could reverse the arrow and show it going back to produce water and carbon dioxide, but a lot of times these products, once they're made, are actually released from that system or removed from that reaction. So we're no longer there to reverse. go the other direction. Now the only thing missing in this reaction is actually the energy.
And we'll see later on that in something like photosynthesis, we would have sunlight on the reactant side, where in respiration we would show that as glucose and oxygen are being broken apart and rearranged to form the product of carbon dioxide and water, we would show ATP as part of a production. So chemical reactions, again, we can go forward, we can go in reverse. The one thing to keep in mind, especially in a living system, a lot of times the products do not stay in that reaction. They will actually be lost or they will leave the reaction as a way of carrying out some other process. So that's going to be completing the end of our Chapter 2 lecture.
We're moving next into Chapter 3, continuing our view with chemistry. But our primary focus is going to be just on the water molecule as a way of how water itself is a life-giving property. for us on our planet.