Unit 1 - Annotations

In this unit we will establish a basic understanding of a molecule. In organic and biological chemistry, most matter is made up of molecules. Not all matter, however, is molecular, therefore, we need to be able to distinguish between matter that is molecular and matter that is not molecular. This understanding will be important for describing and predicting the types of interactions that take place between atoms when they form molecules and for describing and predicting the types of interactions that take place between molecules once they are formed. Both types of interactions are important to understanding both organic and biological chemistry.

This unit will draw material from Chapters 3 and 4 of Raymond's text. You should be familiar with all of the material that is covered in Chapters 1 through 3 Raymond, however, we will not cover it in this class. This is material that you should have seen before in your general chemistry course. You might want to review this material on your own. The objectives that we will focus on include

Concepts: Electronic Structure
(Section 3.2)

This section looks at a model for atomic structure that will be useful for predicting the types of ions that an atom of a particular element will form. Read through Section 3.1, if you need to review what an ion is. Later, we will see that these models will also help us predict how non-metal atoms combine to form organic molecules.

The Bohr model of an atom is presented, which places the electrons in different energy levels. It is possible for more than one electron to occupy the same energy level, but there are rules that govern how many electrons can do this.

See the Elaboration - Atomic Structure page for a review of are working model for the strcuture of an atom. Atoms are the basic building blocks for all matter, including organic and biological molecules.

When electrons move from one level to another, they absorb and release energy. This energy can be released or absorbed as light energy. The energy of light is related to its color; light at the violet end of the rainbow has a higher energy than light at the red end of the rainbow.

Of all of the electrons that an atom contains, the most import ones for determining an atom’s chemical and physical properties are the electrons located in the highest occupied energy level. These electrons are called the valence electrons.

Because the valence electrons are so important in chemistry, the American chemistry Gilbert Lewis developed a symbolic way of indicating the number of valence electrons that an atoms has when writing chemical structures. His symbolism is called the Lewis electron dot structures.

Concepts: The Octet Rule
(Section 3.3)

The octet rule is useful for predicting much of what happens when elements react with one another, basically, atoms react with one another in an effort to gain, lose, or share valence electrons in order to end up with eight valence electrons. In the process, the atoms form monoatomic ions and become isoelectronic with one of the inert gases.

See the Elaboration - The Octet Rule page for a description on using the octet rule to predict the charges on monoatomic ions formed from the representenative elements.

Interactions: Ionic Compounds
(Section 3.4)

This section starts off by describing what compounds are. Compounds are made up of more than one type of atom that combine in a defined ratio and produce substances with characteristic chemical and physical properties. Most matter is made of compounds. Atoms form compounds as a way of obeying the octet rule. There are different types of compounds. This section focuses on ionic compounds, which are formed from ions. Binary ionic compounds are combinations of positively charged metal ions and negatively charged non-metal ions. They are held together by an electrical attraction called an ionic bond and combine in a ratio that produces no net charge. By applying this rule you should be able to predict the formulas for binary ionic compounds given the elements that combine to form an ionic compound.

Other compounds are formed by atoms that combine using covalent bonds. This is the subject of the next section. These compounds are formed from non-metals and represent the major form of compounds that we will encounter in organic chemistry and biochemistry. There is a class of ions that are made from more than one non-metal atom than combine using covalent bonds, but which, as a group, has a net charged. These are called polyatomic ions. An example is the carbonate ion, CO₃^2-, which comprises one carbon atom plus three oxygen atoms held together by covalent bonds, and has a negative 2 charge. Polyatomic ions can form ionic compounds by combining with positive ions, usually metal ions. Sodium carbonate (Na₂CO₃) and (CaCO₃) are examples of ionic compounds that contain the carbonate polyatomic ion.

This section also describes how ionic compounds are named. Given its formula, you should be able to name an ionic compound.

While the metals that are representative elements form only one type of ion, which can be predicted by their group number, transition metals often form more than one type of ion that is not readily predictable. When naming ionic compounds that contain transition metals, the charge is indicated using a Roman numeral, for example, CuCl₂ (one Cu²⁺ ion plus two Cl^- ions) is copper(II) chloride, while CuCl (one Cu⁺ ion plus one Cl^- ions) is copper(I) chloride.

Interactions: Covalent Bonds
(Section 3.5)

In the last two sections we saw that metals are looking for ways to lose electrons as a way of obeying the octet rule, whereas, non-metals are looking for ways to gain electrons as a way of obeying the octet rule. When metals and non-metals combine, they readily accommodates each other's needs: the metals give up their electrons to the nonmetals, resulting the the metals forming positive ions (cations), while the non-metals form negative ions (anions). The oppositely charged ions then combine to form ionic compounds, which are held together by ionic bonds, which arises from the electrostatic attraction that oppositely charged ions have for one one another. In this section we look at an alternative strategy that is used when non-metals, by themselves, combine. Since none of the non-metal ions wish to give up their electrons, they instead share their valence electrons, in pairs, and each atom that participates in the sharing counts the shared electrons as part of their octet. This sharing of electrons is called a covalent bond. This section also introduces using line-bond structures to represent the structures of molecules.

Concepts: Molecules
(Section 3.6)

This section describes how non-metals can form molecules as a way of obeying the octet rule; the atoms form covalent bonds in which they share pairs of electrons with other non-metals atoms. The number of covalent bonds that an atom makes when forming molecules is generally equal to the number of electrons that it needs to obtain 8 valence electrons (or 2 in the case of hydrogen). For example, oxygen, which is in Group VIA and needs 2 electrons to obtain 8 electrons in it valence shell, typically forms 2 covalent bonds when forming molecules molecules.

See the Elaboration - Compounds page for a discussion on some of differences that exist between ionic and molecular compounds.

Some atoms form multiple covalent bonds with other atoms, i.e., they form double and triple bonds. In a double bond the two atoms are sharing 4 electrons, and in triple bonds, the two atoms are sharing 6 electrons.

There is also a discussion of how binary molecular compounds are named. Because the number of each atom in a binary molecular compound can vary and is not readily predictable, prefixes are used to explicitly designate the number of atoms of each type in the molecule.

Concepts: Structural Formulas and Formal Charge
(Section 4.1)

This section describes different ways of representing a molecule. The molecular formula, which was introduced earlier, simply lists the elements that participate in the formation of a molecule and the number of each atoms of each atom that is present in the molecule; this number is placed as a subscript to the right of the corresponding symbol for the element. Molecular formulas do not describe how these atoms are connected. It is possible to have more than one type of molecule share the same molecular formula. We will learn that different molecules that share the same molecular formula are called isomers. Structural formulas are another type of formula that explicitly show how the atoms in a molecule are connected to one another. There are several different types of structural formulas, including, electron dot structures, line-bond structures, condensed structural formulas, and skeletal structures.

See the Elaboration - Structural Formulas page for a step-by-step comparison of the different types of structural formulas.

This section also introduces a method for determining how the charges are distributed in a molecule. There are different ways of doing this, the concept of formal charge is the one introduced in this section. Knowing how the charges are distributed in a molecule will be important for predicting a molecule's physical properties. The formal charge on an atom in a molecule compares the number of electrons that the atom contributes to a molecule in the form of its valence electrons, to the number of electrons that are in the vicinity of that atom once the molecule is formed. Sample Problem 4.1 (p.86) in Raymond goes step-by-step through the procedure of calculating a formal charges on the atoms in a molecule. To calculate the electrons that are in the vicinity of an atom, count each of the non-bonding electrons and add to this one electron for each covalent bond.

Concepts: Shape and Polarity
(Section 4.2)

We will see that organic and biological chemistry deal in large part with molecular compounds. To understand the behavior of molecular compounds we need to understand how they interact with one another. The subject of the next section is noncovalent interactions, a term which groups together a variety weak intermolecular interactions. A property of a molecule that heavily influences these interactions is polarity. Even though molecules have no net charge, the electrons in a molcules are often not evenly distributed in the molecule. This leads to some regions of the molecule having a partial negative charge, while other regions have an offsetting partial positive charge. This can make the molecule as a whole polar and thereby significally influence how it interacts with other molecules. This section describes how to predict the polarity of a molecule, which involves indentifying polar covalent bonds, determining the shape of the molecule, and looking to see if the combination of shape and polar convalent bonds leads to an overal polarity for the molecule.

See the Elaboration - Polarity for a discussion on determining the polarity of a molecule.

Interactions: Noncovalent Interactions
(Section 4.3)

This section discusses a very important group of interacations called noncovalent interactions. These interactions are much weaker than the covalent bonds that hold molecules together, and are easily distrupted by such factors as increased temperature. Noncovalent interactions are the interactions that hold molecule compounds together in the solid and liquid states and it is these interactions which are disrupted when a solid molecular compound melts to form a liquid, and evaporates to form a gas. In biology, these are the interactions that cause membranes to form, proteins to fold and nucleic acids to form double helices. Your textbook describes how these interactions are primarily electrical in nature and arise from either permanent of transitory charges that are present on molecules. These include the hydrogen bond, which is particularly important in biology, salt bridges, dipole-dipole interactions and London dispersion forces. The weakest of these is the London dispersion force, and all molecules have this interaction. This interaction gets stronger as molecules get larger. The other interactions depend on the presence of dipoles, charges and hydrogen bond donors and acceptors. The more of these interactions a molecule has, the stronger a molecule will interact with its neighbors. This is reflected in higher melting points and boiling points for molecules that can participate in these interactions. We will refer often to these interactions during the semester, so it is worthwhile to spend some time reading the descriptions of these interactions and studying the figures in this section.