Organic Chemistry Solomon and Fryhel: Ch1 Carbon Compounds and Chemical Bonding

• Atomic Orbitals
• Electron probability density - the square of a wave for a particular x,y,z location expresses the probability of finding an electron there.
• Orbitals - a place where finding an electron is high
• Atomic orbital - plots of wave functions in three dimension, this is where we get the shapes for the S, P, and D orbitals.
• Molecular orbital - shows the region in space where one or two electrons of a molecule are likely to be found.
• Energy levels
• As we get further away from the nucleus of an atom, the electrons gain energy. Ie 1s has less energy than 2s and even less energy than 2p. But each electron in an electron pair have the same amount of energy.
• Aufbau Principle - orbitals will fill lowest energy to highest energy
• Pauli exclusion principle - a max of two electrons can be places in each orbital as long as the electrons spin in opposite directions.
• Hund's rule - when it comes to degenerate orbitals, we fill each orbital with one electron first then go back and add the second electron to each orbital. This allows the electrons to be further apart is they can ( two negatives will repel)
• Chemical bonding.
• Ionic bonding - (electrovalent) electrons are taken from one atom and given to another.
• This occurs due to the attractive force between two oppositely charged ions.
• Covalent bonding - two atoms will share an electron.
• This occurs when two atoms with similar electronegativity interact and end up sharing an electron, this is the most common and ALL molecules are formed by covalent bonds.
• Electronegativity is an atoms ability to attract electrons.

 

• Octet rule - the tendency of an atom to try to achieve 8 electrons in its valence shell
• There are exceptions too this rule. All elements in the third period and beyond have "d" orbital and can make more than the normal 4 bonds (ie 8 electrons) these are usually P, S, CL
• And some highly reactive molecules or ions may have less than 8 electrons.

• Lewis structure
• Lewis structures show the connection between atoms in a molecule or ion using only the valence electrons of the atom involved.
• Main group elements have the same number of valence electrons as their group number on the periodic table.
• Each atom will attempt to obtain eight electrons
• Hydrogen wants to have 2 valence electrons, meaning it will be like helium and can make one covalent bond.
• Isomers - different substances that share the same molecular formula (the atomic connectivity is different) these molecules will have different physical and chemical properties.

• Formal charge - the number of valence electrons minus half of the number of shared electron minus the number of unpaired electrons.
• F=Z-(1/2)s-u
• The sum of all the charges in a molecule or ion will equal the overall charge of the ion or molecule.
• Resonance  = the state attributed to certain molecules having a structure that con not adequately be represented by a single structure or formula, but is a composite of two or more structures of higher energy. (only the electrons are moving around)
• None of the structures are an accurate depiction of the molecule or ion, and will be better represented by a hybrid structure.

 

• It is important to understand that a single barred arrow is sign for resonance and that the double barred arrow is the sign for equilibrium.
• The energy of the resonance hybrid is lower than the energy of any contributing structure; but if the resonance structures are equivalent, then its energy is large.  (the more stable a contributing factor is by itself, the greater its contribution to the hybrid.
• Stability of structures in resonance.
• The more covalent bonds the more stable
• Charge separation decreases stability  (Ϩ+ and Ϩ-)
• Structures that have an octet are more stable.
• Molecular orbitals
• Molecular orbitals are formed when atomic orbitals combine, the number of molecular orbitals always equals the number of atomic orbital that combined, and they will combine either in or out of phase.

 

• Bonding molecular orbital occurs when the two orbitals are in the same phase
• Anti-bonding molecular orbitals occur when the two orbitals are of an opposite phase. This state occurs when the molecule in the ground state absorbs a photon of light of the proper energy level.
• Bond length - the inter-nuclear distance between the two atoms.
• Heisenberg theory of uncertainty - there is no way to know the position AND momentum of an electron at a given time.
• There are four types of atomic orbitals

 

• S orbitals
• Each S orbital can hold two electrons
• S block elements, elements that have an S orbital as their valence group, include group 1 (alkali metals) and group 2 (alkaline earth metals)
• P orbitals
• There are three p orbitals
• P block elements , elements in groups 13-18, all have a fully filled S orbital and a full or partially full P orbital.
• D orbitals
• There are 5 D orbitals with a possible 10 electrons
• D block elements are in groups 3-12 on the periodic table (transition metals)
• F orbitals
• There are 7 F orbital with a possible 14 electrons
• F block elements are in the lanthanide and actinide families at the very bottom of the periodic table
• These orbitals combine to form molecules

• HYBRIDIZATION
• Hybridization - the concept of "mixing" atomic orbitals into new hybrid orbitals that are more suitable for the pairing of electrons to form chemical bonds in valence bond theory.
• Percent S and percent P character.
• The amount of s- and p- character in a molecular orbital relates to the amount of energy and stability of the bond or orbital. This is important in carbocations.
• p- orbitals have more energy and stability than s-orbitals.
• To find the % s or p divide the number of s or p orbitals by the total number of orbitals.
• Sp3 hybridization
• This is the simplest of the three carbon hybridizations, where one s-orbital combines with three p-orbitals

 

•  ex. Methane - the hydrogen's s-orbital overlaps with one of the carbons half filled sp3 orbitals forming a covalent bond. 
• This bond is also an example of a single bond or sigma σ, which is cylindrically symmetrical allowing for a lot of movement.
• The 2 cylinders will overlap, this will allow for rotation without changing the overall shape of the.
• Sp2 hybridization

 

• Uses two p-orbitals and  one s-orbital to hybridize, when this happens three sp2 orbitals and one p orbital is left un-hybridized which tends to lead to a double bond with the sp2 hybridized carbon, oxygen, or nitrogen. Resulting in one π bond and one σ bond.
• The p2 orbitals are responsible for the π bond
• This type of bond is more ridged than the sp3 bond but is still partially cylindrically symmetrical. The σ bond is more stable than the π bond and must be formed first, in a reaction the π bond will be the first to break.
• Sp hybridization

 

• One s-orbital and one p-orbital combine to for 2sp orbitals, this bond is responsible for the extremely inflexible triple bond.
• This will form either one triple bond with a similarly hybridized carbon, nitrogen, or oxygen… or it can form two double bonds.
• MOLECULAR GEOMETRY (VSEPR Theory)