Organic Chemistry Chapter 1

What forces keep the bond at optimal length?

(1) the force of repulsion between the two negatively charged electrons, (2) the force of repulsion between the two positively charged nuclei, and (3) the forces of attraction between the positively charged nuclei and the negatively charged electrons.

Polar Covalent bonds

- Electrons tend to shift away from lower electronegativity atoms to higher electronegativity atoms. - The greater the difference in electronegativity, the more polar the bond. - Some bonds are acceptable to write as a covalent bond or an ionic bond. - Example: Picture> The electronegativity difference is 1.5, so it is on the cusp of polar covalent and ionic, according to just one method used for determining electronegativity values. So, the absolute difference in electronegativity is to be taken with a grain of salt.

Boiling point branching vs no branching

A branched hydrocarbon generally has a smaller surface area than its corresponding straight-chain isomer, and therefore, branching causes a decrease in boiling point.

Formal Charges

A formal charge is associated with any atom that does not exhibit the appropriate number of valence electrons. 1. Determine the appropriate number of valence electrons for an atom. 2. Determine whether the atom exhibits the appropriate number of electrons.

ETHYLENE

Each carbon atom in ethylene has three sp2-hybridized orbitals available to form σ bonds. - One σ bond forms between the two carbon atoms, and then each carbon atom also forms a σ bond with each of its neighboring hydrogen atoms we have seen that the carbon atoms of ethylene are connected via a σ bond and a π bond. The σ bond results from the overlap of sp2-hybridized atomic orbitals, while the π bond results from the overlap of p orbitals. These two separate bonding interactions (σ and π) comprise the double bond of ethylene.

Ethene

Each carbon in ethene must bond to three other atoms, so only three hybridized atomic orbitals are needed

C=O

It has considerable ionic character, rendering it extremely reactive.

London dispersion forces

London dispersion forces are stronger for higher molecular weight hydrocarbons because these compounds have larger surface areas that can accommodate more interactions. As a result, compounds of higher molecular weight will generally boil at higher temperatures.

Degenerate Orbitals

Orbitals with the same energy level are called degenerate orbitals. The 2p orbitals are of equal energy, and thus are degenerate orbitals

debye

Partial charges (δ+ and δ−) are generally on the order of 10−10 esu (electrostatic units) and the distances are generally on the order of 10−8 cm. Therefore, for a polar compound, the dipole moment (μ) will generally have an order of magnitude of around 10−18 esu · cm Iit is more convenient to report dipole moments with a new unit, called a debye (D), where

Dry Cleaning

Rather than surrounding the nonpolar compound with a micelle so that it will be water soluble, it is actually conceptually simpler to use a nonpolar solvent. This is just another application of the principle of "like dissolves like." Dry cleaning utilizes a nonpolar solvent, such as tetrachloroethylene, to dissolve the nonpolar compounds. This compound is nonflammable, making it an ideal choice as a solvent. Dry cleaning allows clothes to be cleaned without coming into contact with water or soap.

Soap

Soaps are compounds that have a polar group on one end of the molecule and a nonpolar group on the other end

Fleeting Dipole-Dipole Interactions

Some compounds have no permanent dipole moments, and yet analysis of boiling points indicates that they must have fairly strong intermolecular attractions

intermolecular forces .

The attractive forces between molecules. All intermolecular forces are electrostatic—that is, these forces occur as a result of the attraction between opposite charges. The electrostatic interactions for neutral molecules (with no formal charges) are often classified as (1) dipole-dipole interactions, (2) hydrogen bonding, and (3) fleeting dipole-dipole interactions.

Bond for a H2 molecule

The bond for a H2 molecule results from constructive interference. The bond is formed from the overlap of the 1s orbitals of each hydrogen atom. The electron density of this bond is primarily located on the bond axis. This type of bond is called a sigma (σ) bond and is characterized by circular symmetry with respect to the bond axis. To visualize what this means, imagine a plane that is drawn perpendicular to the bond axis. This plane will carve out a circle The bonded electrons spend most of their time in the overlapping atomic orbital space... which is called a sigma bond. all single bonds are σ bonds.

MO Theory example with 2 hydrogen atoms

The bond formed is the result of the overlap of two atomic orbitals (s orbitals), each of which is occupied by one electron. According to MO theory, when two atomic orbitals overlap, they cease to exist. Instead, they are replaced by two molecular orbitals, each of which is associated with the entire molecule Atomic orbitals are combined mathematically (using the LCAO method) to produce two molecular orbitals. the antibonding MO has one node, which explains why it is higher in energy. Both electrons occupy the bonding MO in order to achieve a lower energy state.

Hybridized Atomic Orbitals

The carbon must undergo hybridization to form 4 equal atomic orbitals, with symmetrical geometry The atomic orbitals must be equal in energy to form four equal-energy symmetrical C-H bonds the shape of an sp3 orbital results from have 25% s-character, and 75% p-character

sp3-hybridized orbitals

The four sp3-hybridized orbitals are equivalent in energy (degenerate) and will therefore position themselves as far apart from each other as possible, achieving a tetrahedral geometry. Also notice that hybridized atomic orbitals are unsymmetrical. That is, hybridized atomic orbitals have a larger front lobe (shown in red in Figure 1.22) and a smaller back lobe (shown in blue). The larger front lobe enables hybridized atomic orbitals to be more efficient than p orbitals in their ability to form bonds.

What is wrong with this statement? "ψ2 represents the probability of finding an electron in a particular location."

This statement seems to treat an electron as if it were a particle flying around within a specific region of space. But remember that an electron is not purely a particle—it has wavelike properties as well. Therefore, we must construct a mental image that captures both of these properties. That is not easy to do, but the following cloud analogy might help.

Bond angles with electrons and positive charge

bond angle decrease with electrons

sp2- hybridized carbon atom

carbon atom with one p orbital and three sp2-hybridized orbitals they were obtained by averaging one s orbital and two p orbitals.

Bent geometry

due to . lone pairs we have more repulsion

what repels more strongly?

presume that lone pairs repel more strongly than σ bonds (lone pairs are not bound by another nucleus, so they occupy more space than bonding electron pairs

Electron density

term used to refer to probability of finding an electron (the orbital shape is 90-95% of the space where an electron "probably" is). Beyond this region, the remaining 5-10% of the electron density tapers off but never ends. In fact, if we want to consider the region of space that contains 100% of the electron density, we must consider the entire universe.

C-C and C-H bonds

the EN difference is so small! This is why we need functional groups for reactions and such. Since reactions don't happen at C-C or C-H bonds normally/often.

Covalent bond

the result of two atoms sharing a pair of electrons

Molecular Orbital Theory

uses mathematics as a tool to explore the consequences of atomic orbital overlap. The mathematical method is called the linear combination of atomic orbitals (LCAO). According to this theory, atomic orbitals are mathematically combined to produce new orbitals, called molecular orbitals. MOs are a more complete analysis of bonds, because they include both constructive and destructive interference. The number of MOs created must be equal to the number of AOs that were used.

History of electron

was first proposed in 1874 by George Johnstone Stoney (National University of Ireland), who attempted to explain electrochemistry by suggesting the existence of a particle bearing a unit of charge. Stoney coined the term electron to describe this particle. In 1897, J. J. Thomson (Cambridge University) demonstrated evidence supporting the existence of Stoney's mysterious electron and is credited with discovering the electron.

Induction

withdrawal of electrons in polar covalent bonds shown with an arrow Induction causes the formation of partial positive and partial negative charges, symbolized by the Greek symbol delta (δ).

Clouds vs electron clouds

• Clouds in the sky can come in any shape or size. Electron clouds have specific shapes and sizes (as defined by the orbitals). • A cloud in the sky is comprised of billions of individual water molecules. An electron cloud is not comprised of billions of particles. We must think of an electron cloud as a single entity, even though it can be thicker in some places and thinner in other places. This concept is critical and will be used extensively throughout the course in explaining reactions. • A cloud in the sky has edges, and it is possible to define a region of space that contains 100% of the cloud. In contrast, an electron cloud does not have defined edges. We frequently use the term electron density, which is associated with the probability of finding an electron in a particular region of space. The "shape" of an orbital refers to a region of space that contains 90-95% of the electron density. Beyond this region, the remaining 5-10% of the electron density tapers off but never ends. IF we want to consider the region of space that contains 100% of the electron density, we must consider the entire universe.

Quantum Mechanics

- In the 1920s, Quantum Mechanics was established as a theory to explain the wave properties of electrons - The solution to wave equations are wave functions; The 3D plot of a (wave function)^2 gives an image of an atomic orbital • An equation is constructed to describe the total energy of a hydrogen atom (i.e., one proton plus one electron). This equation, called the wave equation, takes into account the wavelike behavior of an electron that is in the electric field of a proton. • The wave equation is then solved to give a series of solutions called wavefunctions. The Greek symbol psi (ψ) is used to denote each wavefunction (ψ1, ψ2, ψ3, etc.). Each of these wavefunctions corresponds to an allowed energy level for the electron. This result is incredibly important because it suggests that an electron, when contained in an atom, can only exist at discrete energy levels (ψ1, ψ2, ψ3, etc.). In other words, the energy of the electron is quantized. • Each wavefunction is a function of spatial location. It provides information that allows us to assign a numerical value for each location in three-dimensional space relative to the nucleus. The square of that value (ψ2 for any particular location) has a special meaning. It indicates the probability of finding the electron in that location. Therefore, a three-dimensional plot of ψ2 will generate an image of an atomic orbital (Figure 1.5).

New definition of Organic and inorganic compounds

- Organic: Carbon atoms - Inorganic Lacking Carbon atoms

How were compounds organized in the early 19th century?

- Organic: living (plants and animals) - Inorganic: non-living (minerals and gases) Why? Saw that organic compounds had different properties, where difficult to isolate and purify. When heated they decompose easier than inorganic. Because of this they thought organic had a "vital force" - Vitalism

Valence Bond Theory

- a covalent bond is formed from the overlap of atomic orbitals. - Remember we treat electrons as waves there for we know we get interference when waves interact. - So Valence bond theory says that A bond occurs when atomic orbitals overlap. A bond is simply the sharing of electron density between two atoms as a result of the constructive interference of their atomic orbitals

Double Bonds and sp2 Hybridization

- compounds with double bonds - look at ethylene - Ethylene exhibits a planar geometry - A satisfactory model for explaining this geometry can be achieved by the mathematical maneuver of hybridizing the s and p orbitals of the carbon atom to obtain hybridized atomic orbitals. In the case of ethylene, each carbon atom forms bonds with three atoms, not four. Therefore, each carbon atom only needs three hybridized orbitals. So in this case we will mathematically average the s orbital with only two of the three p orbitals (Figure 1.26). The remaining p orbital will remain unaffected by our mathematical procedure.

Atomic Orbital

- just space, has no energy. - if you put an electron in it then have energy. - We think of an atomic orbital as a cloud of electric density

Constitutional isomers

- same molecular formula different constitution - have different physical properties and different names.

Linus Pauling

- solved the problem! - suggested that the electronic configuration of the carbon atom in methane does not necessarily have to be the same as the electronic configuration of a free carbon atom. - Specifically, Pauling mathematically averaged, or hybridized, the 2s orbital and the three 2p orbitals, giving four degenerate hybridized atomic orbitals (Figure 1.21). The hybridization process in Figure 1.21 does not represent a real physical process that the orbitals undergo. Rather, it is a mathematical procedure that is used to arrive at a satisfactory description of the observed bonding. This is called sp3-hybridized orbitals. Using valence bond theory, each of the four bonds in methane is represented by the overlap between an sp3-hybridized atomic orbital from the carbon atom and an s orbital from a hydrogen atom.

Trigonal Planar

- sp2 - with three equivalent sp2-- hybridized orbitals and one empty p orbital - angles are 120°

The Structural Theory of Matter

- suggest substances are defined by a specific arrangement of atoms. - According to the structural theory of matter, each element will generally form a predictable number of bonds. - Example: Note have different properties but same molecular formula.

Dipole Moment

- the center of negative charge and the center of positive charge are separated from one another by a certain distance. - used as an indicator of polarity, where μ is defined as the amount of partial charge (δ) on either end of the dipole multiplied by the distance of separation

Methane and Sp3

- the ground state electron configuration for carbon can't explain how carbon makes four bonds - Only two orbitals have unpaired electrons to be shared in the ground state. So consider the excited state. Now carbon has 4 AO capable of forming 4 bonds. - However problem: The geometry of the 2s and three 2p orbitals does not explain the observed three-dimensional geometry of methane when considering the excited state, it doesn't explain how carbon makes 4 equivalent bonds, like the 4 bonds to H in a methane molecule. In reality all bond angles are 109.5°, and the four bonds point away from each other in a perfect tetrahedron. This geometry cannot be explained by an excited state of carbon because the s orbital and the three p orbitals do not occupy a tetrahedral geometry. The p orbitals are separated from each other by only 90° rather than 109.5°. How do we solve this problem?

Drawing Lewis Structures

- treating electrons as particles! not waves.. duh

Solubility

As you learned in general chemistry, like-dissolves-like Polar compounds generally mix well with other polar compounds If the compounds mixing are all capable of H-bonding and/or strong dipole-dipole, then there is no reason why they shouldn't mix Nonpolar compounds generally mix well with other nonpolar compounds If none of the compounds are capable of forming strong attractions, then no strong attractions would have to be broken to allow them to mix

Dipole-Dipole Interactions

Compounds with net dipole moments can either attract each other or repel each other, depending on how they approach each other in space. In the solid phase, the molecules align so as to attract each other. In the liquid phase, the molecules are free to tumble in space, but they do tend to move in such a way so as to attract each other more often than they repel each other. The resulting net attraction between the molecules results in an elevated melting point and boiling point

Phases of Atomic Orbitals

Because they are generated mathematically from wavefunctions, orbital regions can also be (-), (+), or ZERO In this p-orbital, there is a nodal plane. The sign of the wavefunction will be important when we look at orbital overlapping in bonds The sign of the wave function has nothing to do with electrical charge. The value of ψ (+ or −) is a mathematical convention that refers to the phase of the wave . ψ2 (which describes the electron density as a function of location) will always be a positive number. At a node, where ψ = 0, the electron density (ψ2) will also be zero. This means that there is no electron density located at a node.

Atomic orbitals vs molecular orbitals

Both types of orbitals are used to accommodate electrons, but an atomic orbital is a region of space associated with an individual atom, while a molecular orbital is associated with an entire molecule. That is, the molecule is considered to be a single entity held together by many electron clouds, some of which can actually span the entire length of the molecule. These molecular orbitals are filled with electrons in a particular order in much the same way that atomic orbitals are filled. Specifically, electrons first occupy the lowest energy orbitals, with a maximum of two electrons per orbital.

Three types of Bonds

COVALENT BOND: electrons shared between two atoms, where electronegativity difference is less than 0.5 POLAR COVALENT BOND: electrons shared between two atoms with electronegativity difference between 0.5 and 1.7 IONIC BOND: the electrons are not really shared, the two atoms differ in electronegativity by more than 1.7, and so the more electronegative atom owns the electrons.

Why is there differences/scale for EN's

Different methods to determine this shit!

Covalent bonds H2

Each hydrogen atom has one electron. When these electrons are shared to form a bond, there is a decrease in energy, indicated by the negative value of ΔH. The energy of the two hydrogen atoms as a function of the distance between them. Focus on the right side of the diagram, which represents the hydrogen atoms separated by a large distance. Moving toward the left on the diagram, the hydrogen atoms approach each other, and there are several forces that must be taken into account. As the hydrogen atoms get closer to each other, all of these forces get stronger. Under these circumstances, the electrons are capable of moving in such a way so as to minimize the repulsive forces between them while maximizing their attractive forces with the nuclei. This provides for a net force of attraction, which lowers the energy of the system. As the hydrogen atoms move still closer together, the energy continues to be lowered until the nuclei achieve a separation (internuclear distance) of 0.74 angstroms (Å). At that point, the force of repulsion between the nuclei begins to overwhelm the forces of attraction, causing the energy of the system to increase if the atoms are brought any closer together. The lowest point on the curve represents the lowest energy (most stable) state. This state determines both the bond length (0.74 Å) and the bond strength (436 kJ/mol).

Pauli exclusion principle.

Each orbital can accommodate a maximum of two electrons that have opposite spin. In order for the orbital to accommodate two electrons, the electrons must have opposite spin states.

1s orbital

Electrons are most stable (lowest in energy) if they are in the 1s orbital? The 1s orbital, like every atomic orbital, can have up to 2 electrons in it. If there are more electrons in the atom they fill up the 2s the 2p orbitals the 1s orbital is closest to the nucleus and it has no nodes (the more nodes that an orbital has, the greater its energy)

Electron as waves

Electrons behave as both particles and waves. How can they be BOTH? Maybe the theory is not yet complete - The theory does match experimental data, and it has predictive capability. - Like a wave on a lake, an electron's wavefunction can have a positive (+) value, a negative (-) value, or zero (a node).

HOMO and LUMO

For every molecule, two of its molecular orbitals will be of particular interest: (1) the highest energy orbital from among the occupied orbitals is called the highest occupied molecular orbital, or HOMO, and (2) the lowest energy orbital from among the unoccupied orbitals is called the lowest unoccupied molecular orbital, or LUMO. These are the MO's in play when undergoing a chemical rxn

Rejection of Vitalism

Friedrich Wohler showed that ammonium cyanate can be turned into Urea!. More examples came after

Anti bonding and Bonding MO

The lower energy molecular orbital, or bonding MO, is the result of constructive interference of the original two atomic orbitals. The higher energy molecular orbital, or antibonding MO, is the result of destructive interference. The antibonding MO has one node, which explains why it is higher in energy. Both electrons occupy the bonding MO in order to achieve a lower energy state.

Aufbau principle.

The lowest energy orbital is filled first

Steric number

The total number of electron pairs.

Dipole moment and lone pairs

The two electrons of a lone pair are balanced by two positive charges in the nucleus, but the lone pair is separated from the nucleus by some distance. There is, therefore, a dipole moment associated with every lone pair.

molecular dipole moment,

The vector sum is called the molecular dipole moment, and it takes into account both the magnitude and the direction of each individual dipole moment

the pi bond

These p orbitals actually overlap with each other as well, which is a separate bonding interaction called a pi (π) bond (Figure 1.29). Do not be confused by the nature of this type of bond. It is true that the π overlap occurs in two places—above the plane of the molecule (in red) and below the plane (in blue). Nevertheless, these two regions of overlap represent only one interaction called a π bond.

VSPER and water

This analysis demonstrates that the VSEPR model correctly predicts the bent geometry of water, and the observed bond angle is even justified. However, VSEPR theory also predicts that the two lone pairs of H2O should be degenerate (the same energy), and this has proven to be false. Experiments conducted over 30 years ago have revealed that the lone pairs of H2O are indeed different (one lone pair is significantly higher in energy than the other). These observations strongly suggest that at least one lone pair occupies a p orbital, while the other lone pair occupies a lower-energy, hybridized orbital. Since one lone pair occupies a p orbital, the oxygen atom cannot be sp3 hybridized, as we might expect from a classical interpretation of valence bond theory and VSEPR theory. In this case, the VSEPR model fails to explain why the lone pairs are different in hybridization, energy, and orientation. The VSEPR model assumes that steric repulsion of electrons is the only factor that determines electronic and molecular structures, but there are often additional relevant factors. In the case of H2O, the lone pairs do NOT occupy the two corners of a tetrahedron, as predicted. Although VSEPR theory has correctly predicted the bent geometry of H2O, it appears to have done so for the wrong reasons..

Ammonia

Trigonal Pyramidal Geometry the bond angles for ammonia are observed to be 107°, rather than 109.5°. This shorter bond angle can be justified by VSEPR theory if we presume that lone pairs repel more strongly than σ bonds (lone pairs are not bound by another nucleus, so they occupy more space than bonding electron pairs).

Valence bond theory vs MO theory view of bonds

Valence bond theory continues to view each bond separately, with each bond being formed from two overlapping atomic orbitals. In contrast, MO theory treats the bonding electrons as being associated with the entire molecule. The molecule has many molecular orbitals, each of which can be occupied by two electrons.

Hydrogen bonding

When a hydrogen atom is connected to an electronegative atom (usually O or N), the hydrogen atom will bear a partial positive charge (δ+) as a result of induction. This δ+ can then interact with a lone pair from an electronegative atom of another molecule. This type of interaction is quite strong because hydrogen is a relatively small atom, and as a result, the partial charges can get very close to each other.

Hund's rule

When dealing with degenerate orbitals, such as p orbitals, one electron is placed in each degenerate orbital first, before electrons are paired up.

Where do the electrons go when the AO's over lap?

When the AOs overlap the electrons go into the bonding MO rather than the antibonding MO in order to achieve a lower energy state

Percent Ionized Characteristic

Why is the C=O double bond so much more polar than the C-O single bond?

C-Br and C-I

are treated as polar covalent for this course

Tripple bond

formed by sp-hybridized carbon atoms. To achieve sp hybridization, one s orbital is mathematically averaged with only one p orbital/ This leaves two p orbitals unaffected by the mathematical operation. As a result, an sp-hybridized carbon atom has two sp orbitals and two p orbitals. The two sp-hybridized orbitals are available to form σ bonds (one on either side), and the two p orbitals are available to form π bonds, giving the bonding structure for acetylene Triple bond is due to one σ bond and two π bonds. The σ bond results from the overlap of sp orbitals, while each of the two π bonds result from overlapping p orbitals.

2s orbital

has one node and is farther away from the nucleus; it is therefore higher in energy than the 1s orbital

Linear Geometry

has two valence electrons, each of which is used to form a σ bond - sp2 bond angle of 180°.

Electronegativity

how strongly an atom attracts shared electrons.

Vitalism

idea that it is impossible to convert inorganic compounds into organic compounds without introduction of an outside vital force