Review of First-Semester Organic Chemistry
and
Functional Groups
(20 Points)
During this lab period, we will review some of the important principals of organic chemistry. The functional groups that were introduced during the first semester will also be reviewed and some new ones introduced. Your grade on this lab will be your score on a 20-point diagnostic evaluation of your knowledge of bonding and organic families, as presented in the following pages. After you complete the two self-quizzes, see the instructor who will give you the 20-point evaluation. Complete the 20-point evaluation and turn it in to the instructor.
Organic Compounds
Organic compounds contain carbon. The word compound means an electrically neutral aggregate or collection of molecules or ions. As a simplification, there are three kinds of organic compounds, covalent compounds. salts and organometallic compounds. The most common type of organic compound contains only carbon and other nonmetals such as hydrogen, oxygen or nitrogen. Atoms of nonmetals are joined together by covalent bonds to form molecules. Thus, the compounds that contain only nonmetals are called covalent compounds. If a metal is part of the compound, the compound is an ionic compound, because metals form ions. When the metal is bonded to an atom other than carbon, the ionic compound is a salt. When the metal is bonded to carbon, the ionic compound is an organometallic compound.
Covalent Compounds
Carbon is a nonmetal and forms covalent bonds with other nonmetals. A covalent bond is two electrons or a pair of electrons that hold two atoms together. Thus, when carbon and another nonmetal form a bond, that bond is a covalent bond. When only carbon and hydrogen form a molecule, that molecule is a hydrocarbon. The simplest compound between carbon and hydrogen contains only one carbon. Carbon has a normal covalence of four and hydrogen one. Therefore, the simplest molecule is methane, CH4. The compound methane contains an aggregate or collection of methane molecules. Depending on the circumstances, the word methane can refer to one molecule of methane or to a collection of methane molecules. The context of its usage determines the meaning.
Ionic Compounds—Salts and Organometallics
When a compound contains a metal, it is either a salt or an organometallic compound. Organic acids are readily converted into salts by sodium hydroxide. The product is a salt, because the sodium is bonded to oxygen, a heteroatom. When butyl bromide reacts with lithium metal to form lithium bromide and butyl lithium, the butyl lithium is an organometallic compound because lithium is bonded to carbon.
Problem 1. Classify each of the following compounds as a covalent compound, a salt, or an organometallic compound.
acetone ethyl alcohol butane sodium acetate sodium acetylide
Solution 1. Step 1. Draw the structure of each compound, showing the bonding.
Step 2. Look at each structure. If the structure has no metal atoms, it is a covalent compound. If the structure has one or more metal atoms, it is either a salt or an organometallic compound. Thus, acetone, ethyl alcohol and butane are covalent compounds, and butane is a hydrocarbon.
Step 3. If the structure contains a metal, determine whether the metal is bonded to carbon or to a heteroatom. In sodium acetate, the metal sodium is bonded to oxygen. Therefore, sodium acetate is a salt. In sodium acetylide, the metal sodium is bonded to carbon. Therefore, sodium acetylide is an organometallic compound.
Bonding in Organic Compounds
One property of carbon is that it forms single bonds, double bonds and triple bonds. Two carbon atoms can join by single, double or triple bonds. In a complete molecule, each carbon atom will have four bonds. Let us consider two carbon atoms joined in turn by a single bond, a double bond and a triple bond. Carbon must have four bonds. Thus, hydrogen is necessary to make sure carbon always has four bonds in stable organic molecules. We classify organic compounds into families. Hydrocarbons with only single bonds are alkanes, those with double bonds are alkenes, and those with triple bonds are alkynes.
The hydrocarbon families are very useful in the study of covalent bonding. A covalent bond is a pair of electrons that join two atoms. Another way of looking at a covalent bond is that the two electrons exist in a space between the two atoms. In most cases, one electron comes from each atom. In an ideal or hypothetical way, we can consider the individual atoms joining to make a molecule. Thus, we need two carbon and six hydrogen atoms to make ethane. When the atoms join, only the valence electrons will be involved in bonding. Valence electrons are found in the outermost shell of an atom. Each main shell of an atom has subshells that consist of orbitals. Any given orbital can have zero, one or two electrons in it. For hydrogen, there is only one main shell (1) and one orbital (s). Thus, we can describe the one electron of a given hydrogen atom as a 1sor simply selectron. Every hydrogen atom contains an s electron in an s orbital. The shape or geometry of an s orbital is a sphere. The one electron gets its name from its orbital. Remember that the orbital is the three-dimensional space where the electron is found. A general principal is that an orbital can hold a maximum of two electrons. The lone s electron of hydrogen is a valence electron, because it is found in the outermost main shell. A carbon atom contains six electrons and has the electron configuration 1s22s22p2. The outermost main shell of carbon is shell 2, which has four electrons. Thus, carbon has four valence electrons. The second main shell of carbon has two subshells s and p. Like the first main shell, this s subshell is simply one spherical s orbital that can hold two electrons as a max. The p subshell contains three p orbitals. These three orbitals are labeled px, py, and pz just so we can tell them apart. Each of these orbitals can hold a maximum of two electrons. From the electron configuration, we see that carbon has two p electrons. One of these electrons is found in the px orbital and the other in the py orbital. Each p orbital gets one electron before any p orbital gets two electrons (Hund’s rule). The shape of p orbitals is like a dumbbell, with a node in the middle. The orbital representations of hydrogen and carbon atoms are shown on the next page. In our hypothetical model, we will have these atoms join to make molecules.
Figure 1. Atomic Orbitals of Hydrogen and Carbon
Atomic and Molecular Hydrogen
Consider the formation of molecular hydrogen from atomic hydrogen. Molecular hydrogen has the formula H2 and forms from two hydrogen atoms. That is, we make a covalent bond between two hydrogen atoms. The covalent bond is two electrons, which hold the molecule together. The following equation shows the formation of H2 from two hydrogen atoms.
2 H• H2 = H-H
Let’s look at the actual formation of the covalent bond. The one electron of each hydrogen atom is found in an s orbital. The orbitals are where the electrons are found.
The s orbitals are spherical spaces where the electrons are found. The two orbitals that contain valence electrons come together and share the same space. When two orbitals share the same space, they are said to overlap. The two electrons, one from each orbital of each H atom can now be found in the overlapped space. The two electrons in the same space make a covalent bond when they join two atoms together. That is, they are bonding electrons, because they make a bond. When two valence electrons share the same space but do not form a bond between two atoms, they are called nonbonding electrons.
Overlap of Two s Orbitals
Note that we start with two atoms of hydrogen and we form one hydrogen molecule. We also start with two atomic orbitals, and we form two molecular orbitals. Thus, when we overlap two atomic orbitals, we make two new molecular orbitals. The two atomic orbitals are in different atoms. The two molecular orbitals are within the same molecule.
All orbitals can hold a maximum of two electrons, whether they are atomic or molecular orbtials. Thus, the two electrons go into the lowest energy molecular orbital called a sigma () orbital. The higher energy orbital called sigma star (*) is empty, because we have only two electrons. In the valence-bond theory of bonding, we generally ignore the * orbital. The * orbital is an antibonding orbital. Antibonding orbitals are important in the molecular orbital (MO) theory of bonding. The bond between two hydrogen atoms is a single bond that forms by the overlap of ones orbital with anothers orbital. The bond is called a sigma or bond. The bond is called a bond because it is formed by the overlap of two s orbitals. That is s + s = sigma is a Greek s). The two electrons share the space on a straight line between the two hydrogen nuclei. Since a hydrogen atom has only ones orbital, a hydrogen atom can only make bonds. This is true when H bonds with itself to form H2 and when it bonds with carbon or any other nonmetal atom. In a sense, hydrogen is simply necessary in many organic compounds to fill up the bonding sites of carbon. For example, hydrocarbons are made up of carbon backbones and filled in with hydrogen. Every time hydrogen forms a bond with carbon, that bond is called a sigma bond. Let us see how hydrogen bonds with carbon to make methane.
Methane
In order to make methane, four hydrogen atoms must form bonds with one carbon atom. Since hydrogen is involved, we know that every bond will be called a sigma bond. Let us see how these four bonds are formed. To understand how the atomic orbitals overlap, we must go back to Figure 1 and see how the valence electrons are distributed in carbon. We know that the four hydrogen atoms each have one valence electron in an s orbital.
Figure 2. Valence Electrons in a Carbon Atom
The valence electrons in a carbon atom are distributed as shown in Figure 2. The s orbital has two electrons, and two p orbitals each have one electron. Four H atoms cannot bond to one C atom that has this orbital arrangement and give four identical bonds. Linus Pauling recognized this fact and concluded that the four valence orbitals of a carbon atom must change into four identical orbitals. We now call this process hybridization. To form methane, a carbon atom undergoes hybridization and forms a hybridized carbon atom called a sp3-hybridized carbon atom. The four valence electrons are then redistributed so that each of the four identical sp3 orbitals gets one electron. Figure 3 shows the sp3 hybridization of a carbon atom.
Figure 3. sp3 Hybridization of a Carbon Atom
When atomic carbon goes to an sp3 carbon, the name of the carbon atom and the name of the four identical orbitals are sp3. Thus, an sp3-hybridized carbon atom has four sp3-hybrid orbitals. The symbol sp3 means that four atomic orbitals—an s and three p orbitals have changed into four new sp3 orbitals. The symbol is pronounced ess pee three. From the pronunciation, we get one s and three p orbitals. The symbol tells us exactly how many and which orbitals were combined to make an sp3-hybridized carbon atom.
Three general principals about hybrid orbitals are:
1. Hybrid orbital all have the same shape—a teardrop shape.
2. Hybrid orbitals always repel each other and orient as far away as possible from other hybrid orbitals.
3. Hybrid orbitals always form sigma () bonds.
How will the four sp3 orbitals of an sp3-hybridized carbon orient?
Answer: The four hybrid orbitals repel each other and orient to be as far away from each other as possible. That gives a tetrahedral orientation of the four orbitals. Figure 4 shows how the four hybrid orbitals orient in an sp3-hybridized carbon atom to make methane.
Figure 4. Orientation of sp3-hybrid orbitals of a sp3-hybrid carbon atom.
When carbon is part of a molecule, it is always a hybridized carbon atom. Hybrid orbitals always make sigma bonds. Therefore, when we form the four bonds of methane, they will all be sigma bonds. Covalent bonds are made by the overlap of two orbitals, making a three dimensional space for two electrons. The sp3-hybridized carbon atom has four sp3–hybrid orbitals that must overlap with four s orbitals from four hydrogen atoms. Figure 5 shows how these orbitals overlap.
Figure 5. Orbital Overlap in Methane
Methane has four bonds. Each bond is made by the overlap of an sp3 orbital with an s orbital. Each bond is a sigma bond. An sp3 + s = sigma bond (Note: a sigma bond forms whenever two orbitals that start with the letter s overlap). Now, let’s start with methane, and determine its bonding. Methane has the formula CH4. We know that hydrogen can form only sigma bonds. So the four bonds must all be sigma bonds. We also know that the carbon atom of methane has four identical sp3-hybrid orbitals and that all hybrid orbitals only make sigma bonds. So all four bonds must be sigma bonds. Sigma bonds from one atom always orient as far away from each other as possible. How do we know this? Because these bonds are made up of hybrid orbitals, and the hybrid orbitals, where the electrons are found, are as far away from each other as possible. Therefore, electrons, which are found in these orbitals, are also as far away from each other as possible. The structure of methane is shown below.
We can tell from the structure whether hybrid or unhybridized orbitals are involved. From the structure we see that the carbon atom has four sigma bonds. That means that all four of the carbon atom’s bonding orbitals are hybrid orbitals, because hybrid orbitals always make sigma bonds. To get four hybrid orbitals we need to blend or hybridize four atomic orbitals--one s and three p atomic orbitals, and we get four sp3-hybrid orbitals. The carbon atom must be sp3 hybridized, because it has four sp3 orbitals. The name of the hybridized carbon atom is the same as the name of its hybrid orbitals. An sp3-hybridized carbon atom has four sp3 orbitals.
Ethane
Ethane has the structure shown below.
All of the bonds are bonds. What kinds of orbitals overlap to make the bond between the two carbon atoms? Answer: An sp3 hybrid overlaps with an sp3 hybrid. An sp3 + sp3 makes a sigma bond. A general rule is that anytime an orbital with an s in its symbol overlaps with another orbital with an s in its symbol, we get a sigma bond. Ethane is made up of two carbon and six hydrogen atoms. Each hydrogen has one electron to donate to its bond with carbon, and each carbon has four electrons to donate, one to each bond. Hydrogen can only make sigma bonds, so the s orbital of each H atom overlaps with an sp3 orbital of a carbon atom to make a bond. Therefore, a sigma bond can arise from the overlap of a spherical s orbital with a teardrop hybrid orbital. A sigma bond can also arise from the overlap of two hybrid orbitals. Hybrid orbitals always make sigma bonds. Figure 6 shows the formation of sigma bonds.
Figure 6. Three Ways to Make a Sigma Bond
Ethene
Ethene has the structure shown below.
Ethene has a double bond. The double bond is between two carbon atoms, because hydrogen (H) never has a double bond. H always has a single bond or bond. The double bond is not two sigma bonds, because only two electrons can be in the same space. A double bond contains four electrons in two bonds. The first bond between the two carbon atoms is a sigma bond formed by the overlap of two hybrid orbitals. The second bond between the two carbon atoms is called a pi () bond. The bond is made by the overlap of unhybridized p orbitals. We only make a bond after we have made a bond. The bond is made by the end-to-end overlap of two hybrid orbitals, and the bond is made by the side-by-side overlap of two unhybridizedp orbitals. Just as carbon always has four bonds in stable molecules, carbon always has four valence orbitals. Valence orbitals are orbitals that hold valence electrons. Valence electrons are the electrons in the outermost main shell of an atom. A carbon atom has four valence orbitals, three of which hybridize into three sp2-hybrid orbitals when a double bond forms. Figure 7 shows the formation of three sp2 orbitals from a carbon atom.
Figure 7. sp2 Hybridization of a Carbon Atom
As always, the hybrid orbitals repel each other and orient as far from each other as possible, in a plane and 120o apart. The remaining unhybridized p orbital is perpendicular to the three hybrid orbitals. Figure 8 shows how a sp2-hybridized carbon forms.