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ORGANIC CHEMISTRY NOTES
What is it about carbon that makes it the element that nature has chosen for living organisms? Carbon atoms can form strong bonds with other carbon atoms and with other elements like hydrogen, nitrogen, oxygen and sulfur. Because of these bond-forming properties, carbon can be the basis for the huge diversity of compounds necessary for the emergence of living organisms.
Atomic structure
Compounds are made up of elements combined in different proportions Elements are made up of atoms. An atom consists of a dense, positively charged nucleus containing protons and neutrons and a surrounding cloud of electrons. Most of the mass of an atom comes from the nucleus where the protons and neutrons are (they are 1800 times as heavy as electrons). The atomic mass contributed by electrons is almost negligible. However, most of the volume of an atom comes from the electrons; the volume of an atom occupied by the electrons is about 10,000 times larger than that of the nucleus.
Elements commonly found in organic molecules:
CHONPS + fluorine, chlorine, bromine and iodine
Each element is distinguished by its atomic number (Z), a number equal to the number of protons in its nucleus. Because an atom is electrically neutral, the atomic number also equals the number of electrons surrounding the nucleus. Atomic number (Z) = # of protons = # of electrons
Isotopes
Isotopes are atoms of the same element that have different masses. Although all the nuclei of all atoms of the same element will have the same number of protons, some atoms of the same element may have different masses because they have different numbers of neutrons.
Valence electrons
The electrons that surround the nucleus exist in shells of increasing energy and at increasing distances from the nucleus. The valence shell is the outermost shell and the electrons use this shell to make chemical bonds with other elements. The number of electrons in the valence shell (valence electrons) is equal to the group number of the atom. For example, carbon is in the group 4A meaning that it has 4 valence electrons; oxygen is in the group 6A, therefore it has 6 valence electrons. The halogens of the group 7A have 7 valence atoms.
Chemical bonds: the octet rule
Major types of chemical bonds: Ionic: bonds formed by the transfer of one or more electrons from one atom to another to create ions. Covalent: bonds result when atoms share electrons The central idea is that atoms without the electronic configuration of a noble gas generally react to produce such a configuration because these configurations are known to be highly stable. For all of the noble gases, except helium, this means achieving an octet of electrons in the valence shell.
Ionic bonds
Atoms may gain or lose electrons to form charged particles called ions. An ionic bond is the attraction between oppositely charged ions
Covalent bonds and Lewis structures
When two or more atoms of the same or similar electronegativities react, a complete transfer of electrons does not occur. In these instances, the atoms achieve noble gas configurations by sharing electrons. Covalent bonds form by sharing of electrons between atoms of similar electronegativities to achieve the configuration of a noble gas Molecules are composed of atoms joined exclusively or predominantly by covalent bonds Atoms can share two or more pairs of electrons to form multiple covalent bonds. Ions themselves may contain covalent bonds. For example, the ammonium ion Exceptions of the octet rule Atoms share electrons, not just to obtain the configuration of an inert gas, but because sharing electrons produces increased electron density between the positive nuclei. The resulting attractive forces of nuclei for electrons is the “glue” that holds the atoms together Elements of the second period of the periodic table can have a maximum of 4 bonds because these elements have only one 2s and three 2p orbitals available for bonding
Each orbital can contain two electrons and a total of eight electrons fills these orbitals. The octet rule, therefore, only applies to these elements and even here, as we shall see in compounds of beryllium and boron, fewer than eight electrons are possible. Elements of the third period and beyond have d orbitals that can be used for bonding. These elements can accommodate more than eight electrons in their balance shells and therefore can form more than 8 covalent bonds. Examples are compound like PCl5 and SF
Formal charges
First, we examine each atom and, using the periodic table, we determine how many valence electrons it would have if it were an isolated atom. This is equal to the group number of the atom in the periodic table. Next, we examine the atom in the Lewis structure and we assign the valence electrons in the following way: We assign to each atom half of the electrons it is sharing with another atom and all of its unshared (lone) electron pairs. Calculation for the atom
Formal charge = number of valence electrons −
number of shared electrons − number of unshared electrons
Condensed structural formulas
In condensed formulas all of the hydrogen atoms that are attached to a particular carbon are usually written immediately after the carbon. In fully condensed formulas, all of the atoms that are attached to the carbon are usually written immediately after that carbon, listing hydrogens first.
Bond-line formulas
Each line represents a bond. Each bend in a line or terminus of a line represents a carbon atom, unless another group is shown explicitly. No Cs are written for carbon atoms, except optionally for CH3 groups at the end of a chain or branch. No Hs are shown for hydrogen atoms, unless they are needed to give a three-dimensional perspective, in which case we use dashed or solid wedges (as explained in the next section). The number of hydrogen atoms bonded to each carbon is inferred by assuming that as many hydrogen atoms are present as needed to fill the valence shell of the carbon, unless a charge is indicated. When an atom other than carbon or hydrogen is present, the symbol for that element is written at the appropriate location (i.e., in place of a bend or at the terminus of the line leading to the atom). Hydrogen atoms bonded to atoms other than carbon (e.g., oxygen or nitrogen) are written explicitly.
PRACTICE PROBLEMS
Three-dimensional formulas
A dashed wedge represents a bond that projects behind the plane of the paper. A solid wedge represents a bond that projects out of the plane of the paper. An ordinary line represents a bond that lies in the plane of the paper. In general, for carbon atoms that have only single bonds
In a similar way we can change structure 2 into structure 3: Structures 1–3, although not identical on paper, are equivalent. None of them alone, however, fits important data about the carbonate ion. The resonance theory states that whenever a molecule or ion can be represented by two or more Lewis structures that differ only in the positions of the electrons, two things will be true:
- None of these structures, which we call resonance structures or resonance contributors, will be a realistic representation for the molecule or ion. None will be in complete accord with the physical or chemical properties of the substance.
- The actual molecule or ion will be better represented by a hybrid (average) of these structures. Resonance structures, then, are not real structures for the actual molecule or ion; they exist only on paper. The actual carbon–oxygen bond, since it is a hybrid, must be something in between a double bond and a single bond. Because the carbon–oxygen bond is a single bond in two of the structures and a double bond in only one, it must be more like a single bond than a double bond. It must be like a one and one- third bond. We could call it a partial double bond. And, of course, what we have just said about any one carbon–oxygen bond will be equally true of the other two. Thus, all of the carbon–oxygen bonds of the carbonate ion are partial double bonds, and all are equivalent. All of them should be the same length, and this is exactly what experiments tell us. The bonds are all 1.28 Å long, a distance which is intermediate between that of a carbon–oxygen single bond (1.43 Å) and that of a carbon–oxygen double bond (1.20 Å). One other important point: by convention, when we draw resonance structures, we connect them by
double-headed arrows ( ⟷ ) to indicate clearly that they are hypothetical, not real. For the carbonate
ion we write them this way:
We should not let these arrows, or the word “resonance,” mislead us into thinking that the carbonate ion fluctuates between one structure and another. These structures individually do not represent reality and exist only on paper; therefore, the carbonate ion cannot fluctuate among them because it is a hybrid of them. Resonance structures do not represent an equilibrium
An equilibrium is indicated by ⇌ and resonance by ⟷
How can we write the structure of the carbonate ion in a way that will indicate its actual structure? The bonds in the structure on the left are indicated by a combination of a solid line and a dashed line. This depiction is to indicate that the bonds are something in between a single bond and a double bond. As a rule, we use a solid line whenever a bond appears in all structures, and a dashed line when a bond exists in one or more but not all. We also place a δ− (read partial minus) beside each oxygen to indicate that something g less than a full negative charge resides on each oxygen atom. In this instance, each oxygen atom has two-thirds of a full negative charge. Calculations from theory show the equal charge density at each oxygen in the carbonate anion. Figure 1.5 shows a calculated electrostatic potential map of the electron density in the carbonate ion. In an electrostatic potential map, regions of relatively more negative charge are red , while more positive regions (i.e., less negative regions) are indicated by colors trending toward blue. Equality of the bond lengths in the carbonate anion (partial double bonds as shown in the resonance hybrid above) is also evident in this model
4. The energy of the resonance hybrid is lower than the energy of any contributing structure. Resonance stabilizes a molecule or ion. This is especially true when the resonance structures are equivalent. Chemists call this stabilization resonance stabilization. If the resonance structures are equivalent, then the resonance stabilization is large. 5. The more stable a structure is (when taken by itself), the greater is its contribution to the hybrid.
How To Decide When One Resonance Structure Contributes More to the Hybrid Than
Another
The following rules will help us:
1. The more covalent bonds a structure has, the more stable it is. Consider the resonance structures for formaldehyde below. (Formaldehyde is a chemical used to preserve biological specimens.) Structure A has more covalent bonds, and therefore makes a larger contribution to the hybrid. In other words, the hybrid is more like structure A than structure B. 2. Charge separation decreases stability. It takes energy to separate opposite charges, and therefore a structure with separated charges is less stable. Structure B for formaldehyde has separated plus and minus charges; therefore, on this basis, too, it is the less stable contributor and makes a smaller contribution to the hybrid. 3. Structures in which all the atoms have a complete valence shell of electrons (i.e., the noble gas structure) are more stable. Look again at structure B. The carbon atom has only six electrons around it, whereas in A it has eight. On this basis we can conclude that A is more stable and makes a larger contribution.
Quantum Mechanics and Atomic Structure
A theory of atomic and molecular structure was advanced independently and almost simultaneously by three people in 1926: Erwin Schrödinger, Werner Heisenberg, and Paul Dirac. This theory, called wave mechanics by Schrödinger and quantum mechanics by Heisenberg, has become the basis from which we derive our modern understanding of bonding in molecules. At the heart of quantum mechanics are equations called wave functions (denoted by the Greek letter psi, ψ). Each wave function (ψ) corresponds to a different energy state for an electron Each energy state is a sublevel where one or two electrons can reside. The energy associated with the state of an electron can be calculated from the wave function. The relative probability of finding an electron in a given region of space can be calculated from the wave function The solution to a wave function can be positive, negative, or zero The phase sign of a wave equation indicates whether the solution is positive or negative when calculated for a given point in space relative to the nucleus. Wave functions, whether they are for sound waves, lake waves, or the energy of an electron, have the possibility of constructive interference and destructive interference. Constructive interference occurs when wave functions with the same phase sign interact. There is a reinforcing effect and the amplitude of the wave function increases. Destructive interference occurs when wave functions with opposite phase signs interact. There is a subtractive effect and the amplitude of the wave function goes to zero or changes sign. An orbital is a region of space where the probability of finding an electron is high. Atomic orbitals are plots of ψ2 in three dimensions. These plots generate the familiar s, p, and d orbital shapes. ORBITALS
All s orbitals are spheres. A 1s orbital is a simple sphere. A 2s orbital is a sphere with an inner nodal surface (ψ2 = 0). The inner portion of the 2s orbital, ψ2, has a negative phase sign The shape of a p orbital is like that of almost-touching spheres or lobes. The phase sign of a 2p wave function, ψ2p is positive in one lobe and negative in the other. A nodal plane separates the two lobes of a p orbital, and the three p orbitals of a given energy level are arranged in space along the x, y, and z axes in a Cartesian coordinate system. The + and − signs of wave functions do not imply positive or negative charge or greater or lesser probability of finding an electron ψ2 (the probability of finding an electron) is always positive , because squaring either a positive or negative solution to ψ leads to a positive value. Thus, the probability of finding an electron in either lobe of a p orbital is the same.
Electron configurations:
In the case of our hydrogen model above, the shaded spheres represent the 1s orbital of each hydrogen atom. As the two hydrogen atoms approach each other their 1s orbitals begin to overlap until their atomic orbitals combine to form molecular orbitals. A molecular orbital (MO) represents the region of space where one or two electrons of a molecule are likely to be found. An orbital (atomic or molecular) can contain a maximum of two spin-paired electrons (Pauli exclusion principle). When atomic orbitals combine to form molecular orbitals , the number of molecular orbitals that result always equals the number of atomic orbitals that combine. Thus, in the formation of a hydrogen molecule the two ψ1s atomic orbitals combine to produce two molecular orbitals. Two orbitals result because the mathematical properties of wave functions permit them to be combined by either addition or subtraction. That is, they can combine either in or out of phase. A bonding molecular orbital (ψmolec) results when two orbitals of the same phase overlap An antibonding molecular orbital (ψmolec ) results when two orbitals of opposite phase overlap Calculations show that the relative energy of an electron in the bonding molecular orbital of the hydrogen molecule is substantially less than its energy in a ψ1s atomic orbital. These calculations also show that the energy of an electron in the antibonding molecular orbital is substantially greater than its energy in a ψ1s atomic orbital.
Orbital hybridization Orbital hybridization, in its simplest terms, is nothing more than a mathematical approach that involves the combining of individual wave functions for s and p orbitals to obtain wave functions for new orbitals. The new orbitals have, in varying proportions, the properties of the original orbitals taken separately. These new orbitals are called hybrid atomic orbitals. According to quantum mechanics, the electronic configuration of a carbon atom in its lowest energy state—called the ground state—is that given here: The valence electrons of a carbon atom (those used in bonding) are those of the outer level, that is, the 2s and 2p electrons. Restricted Rotation and the Double Bond The σ–π model for the carbon–carbon double bond also accounts for an important property of the double bond: There is a large energy barrier to rotation associated with groups joined by a double bond. Maximum overlap between the p orbitals of a π bond occurs when the axes of the p orbitals are exactly parallel. Rotating one carbon of the double bond 90° (Fig. 1.27) breaks the π bond, for then the axes of the p orbitals are perpendicular and there is no net overlap between them. Estimates based on thermochemical calculations indicate that the strength of the π bond is 264 kJ mol−1. This, then, is the barrier to rotation of the double bond. It is markedly higher than the rotational barrier of groups joined by carbon–carbon single bonds (13–26 kJ mol−1). While groups joined by single bonds rotate relatively freely at room temperature, those joined by double bonds do not.
