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Chemistry 2131:
Organic Chemistry for the Life Sciences (3)

Review of Basic Chemistry

1. Electronic Structures of Atoms:

• an atom consists of a nucleus made of protons (positively charged) and neutrons (uncharged) surrounded by a cloud of electrons (negatively charged). Most of the mass of the atom is in the nucleus, but most of the volume or space occupied by the atom is in the cloud of electrons. So, where the diameter of the nucleus is on the order of 10^-14 to 10^-15 metres, the diameter of the atom is on the order of 10^-10 metres.
• given the compact nature of the nucleus and the relatively large size of the protons and neutrons, it is easy to imagine how they fill the nuclear space, but what about the much larger electron cloud, filled with very small particles? The electrons do not move freely around the extranuclear space, rather they are confined to regions of space called principal energy levels or shells.
• the shells of are given numbers, with those having smaller numbers being found closer to the nucleus. So, shell number one is the closest to the nucleus, then 2 ... This number has another importance. The number of electrons that can occupy that shell is 2n². So, shell 1 can hold up to 2 electrons, shell 2 can hold up to 8, and so on.
• the electrons in the first shell are closer to the nucleus, so they are held more strongly to it. One way of looking at this is to say that these electrons are of lowest energy. Another way to look at it is that these electrons require more energy to bump them up to another shell. Electrons occupying the higher numbered shells, being farther from the nucleus are held less tightly, and are more easily bumped up to a higher energy level.
• for all but the simplest shell, the electrons are further restricted in space to orbitals. These are designated s, p, d, and f. What are these orbitals? They are spaces where an electron will spend 90 - 95% of its time. Each orbital can hold a pair of electrons. So, you can easily calculate how many orbitals there must be in a given shell. Shell 1 has 2 electrons, therefore only 1 orbital, which is called 1s. In this course we will only encounter atoms with relatively small numbers of electrons, so we will be concerned mainly with the s and p orbitals.
• the s orbital is spherical in shape. The 1s, being the closest to the nucleus, is the smallest. 2s is a larger sphere.
• what happens in the second shell, when we need to accomadate not 2 but 8 electrons? The extra electrons occupy 2p orbitals. Note, there are no 1p orbitals. The 3 p orbitals are calledp_x, p_y, p_z. The x, y and z refer to the axes of 3-space. The orbitals have a lobed shape, 2 lobes on either side of the nucleus (where the chance on finding an electron is obviously very small). The three different orbitals are perpendicular to each other.
• now that you know which orbitals exist, and what they look like, how do you know which orbitals are filled with electrons in different atoms. Numerous possibilities exist, but the most stable or most frequently encountered state is the one referred to as the ground state electronic configuration. This is the lowest energy configuration, and obviously other configurations occur when the atom absorbs energy.
• there are three rules to follow when determining ground state electronic configurations for atoms. Before starting you have to figure out how many electrons an atom has. I'm hoping that this is perfectly obvious for you.
  
  1. Fill orbitals in order of increasing energy from the lowest energy to the highest. The order is 1s, 2s, 2p, 3s, 3p, 4s. Note that 4s is of lower energy than 3d.
  2. Each orbital can hold up to 2 electrons.
  3. When orbitals of equivalent energy are available but there are not enough electrons to fill them all, one electron is added to each before adding a second one.
• have a look at Table 1.3 in the text for ground state configurations for a number of relevant elements.

2. Lewis Structures, Bonding and Electronegativity:

• now that you can easily tell how many electrons a given atom has in its elemental form we can move on to the concept of valence electrons or the valence shell. These are the electrons that are involved in chemical bonds and reactions. A quick look at the periodic table and you will see that the number of electrons in the valence shell will be different for different elements. For hydrogen and helium, the shell only contains 2 electrons, whereas for the next series has 8 valence electrons.
• since these electrons are important for reactivity and bond formation we often write these atoms in with their valence electrons depicted by a dot. This is called a Lewis Structure. Remember that only valence shell electrons are represented.
• one of the first things that you should notice when you are trying these is that the noble gases all have full valence shells. For helium this is 2 electrons, for neon and argon it is 8 electrons. The other to notice is that none of the other simples elements have full valence shells.
• the very observations that we have just talked about lead Gilbert Lewis to a hypothesis, that the chemical inertness of noble gases indicates a degree of stability in the electronic configurations of these elements. He proposed that atoms react in such a way that they achieve an outer or valence shell that is full. For most simple elements (those in groups IA to VIIA, the main group elements) an outer shell of 8 electrons is full.
• this is the basis of the octet rule. Basically, it means that atoms will react to lose or gain electrons to be left with a complete valence shell octet.
• this leads us to the concept of bonding, that atoms should "bond" together in a way that each atom participating acquires a completed outer shell. There are two ways that atoms can gain or lose electrons to obtain a full outer shell.
• the first way is to become an ion or to ionize. If the atom gains an electron, it is no longer neutral in charge, but negatively charged and we call it an anion. Similarly, and atom that loses one or more electrons becomes positively charged and is called a cation.
• we call the bond between a cation and an anion an ionic bond
• the other way that an atom can fill its outer shell is to share electrons with other atoms to complete its outer shell. In this case, the bond is called a covalent bond
• these two kinds of bonds are very different. The ionic bond relies on the attractive forces between oppositely charged ions. In essence, electrons are transferred from one element to the other. For example the salt LiCl is made from Li with one valence electron and Cl with 7 valence electrons. Li loses its valence electron to achieve a full shell, and Cl obtains an electron to fill its shell.
• how do we know which atom will lose electrons and which will gain them? This can be predicted by looking at the property of the element known as the electronegativity. This is a very important concept in organic chemistry, and it will come up again and again later in the course.
• Linus Pauling (the vitamin C guru, now dead, didn't do him that mush good did it?) devised a scale of electronegativity, assigning a value to each of the atoms. What is this electronegativity? It is a measure of how strongly a given element is attracted to electrons from other atoms that it shares in chemical bonds.
• in the Pauling scale, fluorine is assigned the highest value, 4.0. Everything else is rated relative to fluorine. There are a number of thisngs to notice on this scale. First of all, if you look from left to right along a given row, the electronegativity increases. The other direction to look at is top to bottom. Values increase from bottom to top, or put in a different way, the smaller the member of the group the more electronegative the element.
• to understand these trends you have to think about what the differences are. As you proceed along a row, from left to right, the number of protons increases, so the positive charge on the nucleus increases, increasing the attraction for electrons. Obviously this only extends aalong a row, once a shell is full, at the end of the row, you start again on the next row.
• the other tendency has to do with the shell number the valence electrons occupy. As you go down the table the number increases, this means that the electrons are farther from the nucleus, thus less strongly attracted.
• these values are estimations only though, and the environment in which an atom is found affects its actual electronegativity.
• so, we can predict that a more electronegative atom will gain electrons and become an anion and a less electronegative atom will lose electrons and become a cation.
• generally we say that a chemical bond is ionic if the difference between the electronegativities of the atoms is 1.9 or greater. So, you look up the electronegativity values on the chart and determine the difference. In the caseof LiCl, the difference is 2.0, a clear case for an ionic bond.
• before leaving ionic bonds, it should be noted that in the solid state, or the crystal, the ionic salts don't interact with only one other ion, they form a kind of lattice, each interacting with a number of surrounding ions.
• the ionic bond is one that we have seen before, but what about the covalent bond?. Before it was noted that electrons must be shared between the atoms in order to get a covalent bond. To use the guideline that we used for ionic bonds, a covalent bond exists if the difference in electronegativites of the atoms is less than 1.9. In essence it means that neither atom is strong enough (electronegative enough) to rip an electron off of the other, so they share the electron instead.
• in the simplest case, atoms with identical electronegativities, such as the 2 hydrogen atoms in H₂, are covalently bonded together. Start by looking at the Lewis structure of hydrogen, with its single valence shell electron. To fill its valence shell it "wants" one more electron. One way to get that electron is to share it with another hydrogen atom. So, each hydrogen atom of the molecule has a full valence shell. In this case both atoms are of equal electronegativity, so the electrons are shared evenly.
• the electrons that are shared form the covalent bond. The electron pair of this bond occupies the region between the two nuclei. Clearly, bringing two positively charged nuclei that close together results in a repulsion between them. This repulsion is shielded or overcome by the electron pair occupying the space between them. The electrons also attract both of the nuclei.
• the distance between two nuclei in a covalent bond is an important parameter called the bond length. Every covalent bond has a characteristic length. What determines this length? A combination of factors is resposible. A plot of potential energy versus distance between nuclei shows this best. Starting out at large internuclear distances, there is no overlap between electron clouds and essentially no attraction between the atoms, thus the energy is essentially zero. As the distance is decreased, the atoms move closer together, some overlap occurs, some sharing of electrons occurs and this becomes favourable, and the energy decreases (things are becoming more stable). There is a distance at which the atoms are ideally spaced. There is a maximum of attraction, sharing of electrons to fill valence shells, but repulsion has not become significant yet. This distance is the bond length (0.074 nm for the H-H bond). This corresponds to the lowest energy or most stable part of the curve. If the distance between atoms decreases further, then repulsion takes over and the energy increases aymptopically.
• what about bonds between atoms that have different electronegativities, but not enough different to give classical ionic bonds? In the previous example, of a bond between atoms of identical size and electronegativity, the electrons are shared evenly. Other covalent bonds are not between atoms of identical electronegativity. For instance the bond between oxygen and hydrogen in an alcohol functional group has an electronegativity difference of 1.4. This is not enough for it to be called an ionic bond, but we think of it as polar covalent.
• the polarization of a bond is said to occur when the difference in electronegativity between the two atoms is between 0.5 and 1.8.
• when the difference in electronegativity is 0 to 0.4 the bond is said to be nonpolar covalent. A very important example is the C-H bond. The difference in electronegativity is 0.4. Thus, this bond is siad to be nonpolar, and the electrons are essentially evenly shared between carbon and hydrogen.
• for the O-H bond, with a difference in electronegativity of 1.4, there is a polarization of the electrons. The more electronegative atom, the one with the greater attraction for electrons has a greater fraction of the shared electron pair than the less electronegative.
• thus we say that the more electronegative atom bears a partial negative charge and the other atom bears a partial positive charge. These are denoted with the greek symbol for a small delta with a plus or minus sign.
• before leaving the topic of covalent bonds and ions, some points on the depiction of such molecules so that others understand what you mean. Read over the rules on page 10 and 11 of the textbook.
• in a nut shell, you should figure out how many electrons the various atoms contribute, subtracting an electron for every positive charge and adding one for every negative charge, you figure out how the atoms will be arranged (you should know this from another source). Then, you connect the atoms with single bonds, each single bond represents a pair of electrons, and add remaining electrons to fill every atom's valence shell. You can have single, double and triple bonds (representing 2,4 and 6 electrons). Unshared electron pairs are depicted as a pair of dots.
• please get in the habit of writing proper Lewis structures for molecules. That means that when you draw molecules include unshared electron pairs.
• there are some numbers to keep in mind when writing Lewis structures. Hydrogen should always have one single bond, that gives it a full valence shell of 2 electrons. Carbon usually has 4 bonds to give it a full valence shell, nitrogen usually has 3 bonds and a lone pair to fill its shell, oxygen has 2 bonds and 2 lone pairs, and halogens have one bond and 3 lone pairs of electrons. These guidelines will help to figure out Lewis structures.
• as you try a number of Lewis structures you will come across the notion of formal charge. Often ions are made up of several atoms. It is important to know which atom of a molecule bears the charge or charges. The charge on a given atom in a molecule is the formal charge.
• to derive a formal charge, you should first draw the correct Lewis structure for the molecule. Then you assign to each atom all of its unshared electrons and one half of its shared electrons (the ones participating in bonds). This number is then compared with the number of electrons in the unbonded atom. If the number of electrons in the atom in a molecule is less than in the unbonded atom, then the atom has a positive formal charge for each electron less. The opposite is also true, if the number of electrons in the atom in a molecule is greater than in the unbonded atom then a negative formal charge is assigned.
• the sum of formal charges in a molecule is the actual charge on the molecule as a whole. A good example of this is given in the text, nitric acid, which is HNO₃. To draw this properly, one of the oxygens has a negative charge and the nitrogen has a positive charge.