Organic Chemistry for Dummies 2nd Edition by Arthur Winter, PhD
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Protons: number of protons cannot be changed without changing the identity of the atom itself; if the number of protons change, it becomes a different element.
Ion: an atom that has more or fewer electrons than the amount of protons. The atom is also electrically charged.
Electrons: they are located in the shells surrounding the nucleus, not in the nucleus itself.
First Shell: closest to the nucleus of the atom, has the lowest energy, and can hold up to 2 electrons.
Second Shell: higher in energy, farther away from the nucleus, and can hold up to eight electrons.
Third Shell: higher in energy than the first and second shells, even farther away from the nucleus, and can hold up to 18 electrons.
%%Note: there are other higher shells, but they are not dealt with in organic chemistry.%%
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Orbitals: electron shells are further subdivided into these; they are the actual location in which an electron can be found.
Difference between Shells and Orbitals
Analogy to Better Understand Electrons:
Heisenberg uncertainty principle: the uncertainty in knowing the locations of electrons at a given moment.
Orbitals (Used in Organic Chemistry)
Each orbital can hold up to 2 electrons, but if there are 2 in one orbital, they have opposite spins.
s orbital: spherical in shape.
p orbital: shaped like a dumbbell.
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1s orbital: spherically symmetric, holds 2 electrons, only orbital in the first shell.
@@Second Shell:@@ contains both s and p orbitals, holds up to 8 electrons.
2s orbital: spherical shape like 1s orbital, but larger and higher in energy.
2p level: consists of 3 individual p orbitals:
Px: an orbital that points in the x direction.
Py: an orbital that points in the y direction.
Pz: an orbital that points in the z direction.
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Ground-state electron configuration: list of orbitals occupied by electrons in a particular atom.
You start by placing electrons into lower energy orbitals and build up from there.
The lowest-energy orbital is 1s, followed by 2s, 2p, 3s, 3p, 4s, and so on.
The Aufbau chart is helpful for remembering which orbitals to fill first.
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Filling Orbitals
Fill two electrons per orbital, starting with lowest-energy orbitals and work up until you run out of electrons.
However, the last electrons in an orbital must be placed according to the Hund’s rule.
Hund’s rule: the electrons should go into different orbitals with the same spin, instead of pairing up into a single orbital with the opposite spin.
Electrons repel each other and want to be as far away from each other as possible.
%%Example:%% Carbon’s electron configuration is 1s^2 2s^2 2px^1 2py^1 2pz^0 (not 1s^2 2s^2 2px^2 2py^0 2pz^0, which violates Hund’s rule).
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Example of Ionic Bonding reaction:
Sodium (Na) and Chlorine (Cl) to make (NaCl), also known as table salt.
Sodium is an atom in the first column of the periodic table and has 1 valence electron.
Chlorine is an atom in the second-to-last column of the periodic table and has 7 valence electrons.
In order to achieve valence octet, sodium could either gain 7 electrons or lose 1; chlorine could either lose 7 electrons or gain 1.
However, atoms generally do not gain or give up more than 3 electrons, so sodium gives up 1 of its valence electron to chlorine. Chlorine, then, has a full octet.
Sodium becomes a cation as it lost 1 electron, and chlorine becomes an anion because it gained 1 electron. Hence, the positive and negative signs. The dots symbolize the valence electrons.
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When giving up an electron, sodium imitates the electron configuration of noble gas neon (Ne), which has a full electron configuration.
When gaining 7 electrons, chlorine imitates the noble gas argon (Ar).
Attraction between sodium cation and chlorine anion in sodium chlorine is an ionic bond.
Electrons are NOT shared. They are taken away from one atom by the other atom.
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A hydrogen atom has 1 electron, so it needs another 1 to fill its shell.
Because both hydrogen atoms need 1 electron to fill its shell, they share their electrons equally (instead of grabbing an electron from each other).
Now, they both achieved the electron configuration of noble gas helium (He).
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How to know whether a bond is ionic or covalent?
Ionic bonds are usually found in inorganic compounds
LiF, NaCl, KBr, and MgBr2.
Covalent bonds are found in organic compounds.
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Dipole moment: separation of charge in the bond because the more electronegative atom “bullies” most of the bonding electrons away from the less electronegative atom.
For example: in hydrochloric acid (HCl), chlorine is the more electronegative atom of the two, so the electrons between hydrogen and chlorine are “hogged” mostly by chlorine.
Because electrons spend most of the time around chlorine, chlorine gets a partially negative charge.
Because electrons spend less time around hydrogen, hydrogen gets a partially positive charge.
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Dipole vector: a distinct arrow that is used to show the direction of the dipole moment, or separation of charge.
Head of the arrow points in the direction of the partial negative charge, and the tail (which looks like a plus sign) points in the direction of the partial positive charge.
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How to predict dipole moment for a molecule?
Find the dipole vectors of each of the individual bonds.
Add up each of the individual bond vectors.
Line up the vectors from head to tail (the order doesn't matter).
Example: Chloroform: The dipole moment points to the right.
However because individual bonds have dipole moments, it doesn't mean that those molecules have dipole moments.
Example: Carbon Dioxide: Oxygen is more electronegative than carbon, so two dipole vectors point out.
The net dipole moment is zero because the oxygens are pulling in equal and opposite directions; therefore, they cancel each other out.
VSEPR theory: stands for valence shell electron pair repulsion; predicts the approximate geometry of bonds around an atom.
Main Geometries in Organic Chemistry
Linear: 180°
Trigonal Planar: 120°
Tetrahedron: 109.5°
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What is the solution, then?
Carbon atom promotes an electron from the filled 2s orbital into the last empty p orbital.
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But why would carbon promote the electron?
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Naming Hybridized Orbitals
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sp3 orbital
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%%Note:%% The numbers of orbitals that are mixed must equal the number of hybridized orbitals that come out at the end. If four atomic orbitals are mixed, four hybridized orbitals must come out.
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sp3 orbital: four sp3 hybridized orbitals (three 2p orbitals and 2s orbital), bond angles are at 109.5°, tetrahedron formation.
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sp2 orbital: three sp2 hybridized orbitals (two 2p orbitals and 2s orbital); since one of the p orbitals are not mixed, it remains in its original unhybridized form, bond angles are at 120°, trigonal planar formation.
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sp orbital: two sp hybridized orbitals (one 2p orbital and 2s orbital); since two of the p orbitals are not mixed, they remain in their original unhybridized form, bond angles are at 180°, linear formation.
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Count the number of substituents (or number of atoms bonded to that particular atom) and lone pairs of electrons around that atom.
For BeH2, the beryllium (Be) has two substituents (two identical H atoms), so it’s sp hybridized.
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For BH3 (also refer to Figure 2-15), the boron (B) has three substituents (three H atoms), so it’s sp2 hybridized.
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In methane, CH4, which has four substituents, the carbon is sp3 hybridized.
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Covalent bonds occur when the orbitals of bonding atoms overlap each other.
Two kinds of covalent bonds can be formed in organic molecules:
Sigma bonds: bonds in which orbital overlap occurs between the two bonding nuclei.
Pi bonds: bonds where orbital overlap occurs above and below the nuclei, and not directly between them.
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Draw the Orbital Diagram of a Molecule
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Example of drawing orbital picture for ethylene:
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Determine the hybridization of each of the atoms
Draw each of the atoms with its valence electrons.
Determine which orbitals overlap to make the bonds.
For each of the C-H bonds (sigma bond), the bond will result from the overlap between an sp2 orbital on the carbon and the 1s orbital on the hydrogen
For the double bond (between the C-C bond), one of the bonds comes from the two sp2 hybridized orbitals overlapping between the carbon nuclei to make a sigma bond, while the other bond comes from the two p orbitals overlapping sideways to make a pi bond above and below the carbon nuclei.
Final Step: orbitals that overlap to make each bond in the molecule are accounted for.
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