Chapter 1:

ranking resonance structures

PRIORITY #1: structures with filled valence shells contribute more than structures with unfilled valence shells PRIORITY #2: structures with more covalent bonds contribute more than structures with fewer covalent bonds PRIORITY #3: structures with less charge separation contribute more than structures with more charge situation (opposite charges attract so having them be separate identities is such close proximity requires energy and is unfavorable) PRIORITY #4: structures with negative charge on the more electronegative atom contributes more than structures with negative charges on a less electronegative atom

polar vs nonpolar covalent bonds

*Polar covalent*: Unequal electron sharing between 2 oppositely charged ends vs *Nonpolar covalent*: equal electron sharing and no noticeable charge in ANY covalent bond, the electrons are shared, the AMOUNT of sharing just differs depending on if it's polar or nonpolar in *polar covalent bonds*, the MORE ELECTRONEGATIVE atom gets the more electron density, which leads to the formation of partial charges. *Partial charges* are indicated by a plus/minus arrow or δ+/δ-

ionic vs covalent bonds

*ionic*- formed when valence electrons are *transferred* from one atom to another, creating a *cation *(when an atom loses an electron, positive charge) and *anion* (when an atom gains an electron, negative charge) ~~~ ions have charges because the # of protons and # of electrons differ *covalent*- formed when two atoms *share* valence electrons; when sharing is equal

Rules of Drawing Resonance Structures

1) do NOT change the connectivity of the atoms 2) do NOT change the number of valence electrons (must have the same overall formal charge between all resonance structures) 3) do NOT exceed an octet of valence electrons

Drawing chair conformations

1) draw two parallel slanted lines 2) connect the tops and bottoms with shallow "V's", Make sure the center of the V's don't bump into each other (no bowties allowed! we want lots of space in the middle of the chair) 3) draw the axial substituents using vertical lines (if the carbon points up, the sub. points up and vice versa) 4) draw the equatorial substituents using horizontal lines (if the axial bond is up then the eq. bond points down and vice versa) Note: make sure to never draw the substituents INSIDE the chair!!! there are two different ways of drawing chair conformations, the only difference between them is the direction the first set of parallel lines is facing

Assigning R and S

1) the 4 groups of a chiral carbon are prioritized by atomic number -directly bonded atom with the largest atomic number is highest priority -if the same atom is directly bonded to the chiral carbon, then keep going out one atom at a time until you come to a point of difference -double/triple bond? treat this as if each bond is an individual bond 2) the molecule is positioned so that the lowest priority group is placed in the back on a WEDGE 3) draw an arrow showing rotation from priority 1 to 2 to 3 >>> clockwise == R >>> counter-clockwise == S 4) add the R or S configuration in parentheses at the beginning of the name of the molecule in question. If there are multiple chiral centers, the position number is included in the parentheses as well {ex. (1S, 3S)- 1-(tert-butyl)-3-methylcyclohexane} note: if the higher priority substituent is on the wedge rather than the dash, assign the R/S configuration then flip is so that R → S or S → R

Drawing Newman Projections

1. notice what bond the question is asking you to look down 2. draw the substituent on the front carbon, starting from the center of the circle, which represents the front carbon 3. draw the substituents on the back carbon starting from the edge of the circle. The back substituents do NOT connect to the center note: for eclipsed conformations, slightly offset the back substituents from the front ones by moving them slightly to the side. Make sure the eclipsed substituents are still parallel though

Naming Cycloalkanes

1. Substituent placed in front of cycloalkane 2. Number not needed if only one branch 3. If 2 or more substituents are attached, ring is assigned according to which substituent is first alphabetically, making sure to assign lowest numbers if possible basically follow the same rules as naming alkanes!!! just as cyclo

Drawing Resonance Structures

1. start with the Lewis structure including formal charges 2. Do one or more of the following using curved arrows; ~~~ convert a lone pair to an adjacent pi bond ~~~ convert a pi bond to an adjacent lone pair

sp hybridization

2 regions of e- density → 2 orbitals to hybridize (1s+ p)→ *sp2* --> 180 degrees --> linear geometry A type of hybridization that results from the combination of the s orbital and only one of the three p orbitals in the second energy level of carbon, resulting in two hybrid sp orbitals and two unhybridized p orbital. This occurs when a carbon atom is bonded to two other atoms. The geometric arrangement of those two hybrid orbitals is called linear the 2s the one 2p atomic orbitals from two C atoms coming together all mix to form two sp hybridized molecular orbitals. Each of the two sp hybrid orbitals are 180 degrees apart, yielding a linear arrangement *TWO UNHYBRIDIZED P ORBITALS REMAIN*

number of stereoisomers

2^n (n = number of chiral carbons)

sp2 hybridization

3 regions of e- density → 3 orbitals to hybridize (1s+ 2p)→ *sp2* --> 120 degrees --> trigonal planar geometry A type of hybridization that results from the combination of the s orbital and only two of the three p orbitals in the second energy level of carbon, resulting in three sp3 hybrid orbitals an one unhybridized p orbital. This occurs when a carbon atom is bonded to three other atoms. The geometric arrangement of those three hybrid orbitals is called trigonal planar the 2s the two 2p atomic orbitals from two C atoms coming together all mix to form 3 sp2 hybridized molecular orbitals. Each of the three sp2 hybrid orbitals are 120 degrees apart, yielding a trigonal planar arrangement *ONE UNHYBRIDIZED P ORBITALS REMAIN*

sp3 hybridization

4 regions of e- density → 4 orbitals to hybridize (1s+ 3p)→ *sp3* --> 109.5 degrees --> tetrahedral geometry A type of hybridization that results from the combination of the s orbital and all three p orbitals in the second energy level of carbon, resulting in four hybrid orbitals and occurs when a carbon atom is bonded to four other atoms. The geometric arrangement of those four hybrid orbitals is called tetrahedral. the 2s the three 2p atomic orbitals from two C atoms coming together all mix to form 4 sp3 hybridized molecular orbitals. Each of the four sp3 hybrid orbitals are 109.5 degrees apart, yielding a tetrahedral arrangement *NO UNHYBRIDIZED P ORBITALS REMAIN*

mesocompound

A molecule that contains a chiral center but is achiral (not chiral) as it has a plane of symmetry. Has chiral centers but isn't chiral due to symmetry!!! mesocompounds don't follow the 2n rule for stereoisomers... they tend to have *one less stereoisomer* than predicted by the 2n rule

achiral

A molecule that either does not contain a chiral center or contains chiral centers and a plane of symmetry; as such, it has a superimposable mirror image. A molecule that is superimposable on its mirror image

Chirality

A property of a compound to exist in both left and right forms; occurs whenever a compound contains an asymmetric carbon (carbon that is bonded to four different substituent groups)

energy level diagram

A schematic drawing used to arrange atomic orbitals in order of increasing energy levels. steps: 1) draw various shells/orbitals needed in INCREASING energy 2) fill out the electrons according to the atomic number of the element you are making the diagram for (ex. carbon has 6 electrons in its diagram). 3) Make sure you are applying Aufbau, Pauli, and Hund's principles 4) write the shorthand notation, also called its *electronic configuration*. ex) carbons electronic configuration is 1s²2s²2p² note: the electrons in the outermost shells are the valence electrons! you can see them in energy level diagrams!

Lewis structures and how to draw them:

A structural formula in which electrons are represented by dots; dot pairs or dashes between two atomic symbols represent pairs in covalent bonds. STEPS: 1) Determine the total number of valence electrons in the molecule by looking at the group # of the atoms in the periodic table. (ex. C has 4) 2) Determine the connectivity 3) Connect the relevant atoms by single bonds first 4) Add the remaining valence electrons as lone pairs 5) Convert lone pairs to double/triple bonds if needed to generate octets for every atom... *recall that each single bond represents TWO electrons* 6) Assign formal charges as needed!

Bonding vs antibonding molecular orbitals

Bonding orbitals created from constructive interference are LOWER in energy than the atomic orbitals from which they derive ~~~ why? because the electrons can spread out more in their sigma MO (for example) than in their original atomic orbitals, therefore significantly reducing electron-electron repulsion Antibonding orbitals created from destructive interference are HIGHER in energy than the atomic orbitals from which they derive ~~~ why? the electron cannot spread out into the space between the atoms due to the nodal plane, which increases their mutual repulsion (unstable, higher energy) than in their original atomic orbitals

carboxylic esters

COOC derivatives of carboxylic acids in which the carboxyl hydrogen is replaced by a carbon containing group (O-C bond rather than OH bond connected to the carbonyl group)

carboxylic acids

COOH carbonyl group + OH at the end of a molecule/substituent note: the OH group in the carboxylic acid is NOT an alcohol!!! Alcohols only exist when OH is bound to an sp3 hybridized carbon where as carboxyl OH's have an OH bound to an sp2 hybridized carbon

carboxylic amides

COON a derivative of a carboxylic acid in which the -OH is replaced by an nitrogen containing group Note: just as the OH in a carboxylic acid is NOT an alcohol, the nitrogen containing group in a carboxylic amide is NOT an amine!!!

Hydrocarbons

Compounds composed of only carbon and hydrogen alkanes: single bonds only, have the general formula of *C(n)H(2n+2)*, also referred to as *saturated hydrocarbons* alkenes: double bonds alkynes: triple bonds prefixes: meth- eth- pro- but- pent- hex- hept- oct- non- dec-

Stereoisomers

Compounds with the same structural formula but with a different arrangement of the atoms in space. cannot be interconverted because you would have to break bonds to make them identical stereoisomers have... >>> the same molecular formula (therefore some type of isomer) >>> the same connectivity (rules out constitutional isomers) >>> different arrangements in space

filling molecular orbitals

MO populate according to the same three rules as atomic orbitals! 1) Aufbau 2) Pauli 3) Hund's

Hydrocarbon prefixes

Meth- (1) Eth- (2) Prop- (3) But- (4) Pent- (5) Hex- (6) Hept (7) Oct- (8) Non (9) Dec (10)

Aldehydes

R-CHO contains a carbonyl group have one hydrogen and one carbon bonded to the carbonyl carbon at the end of a molecule/substituent

Ketones

R-CO-R contains a carbonyl group have two carbon atoms bonded to the carbonyl carbon in the middle of a molecule/substituent

Alcohols

R-OH (hydroxyl group) can be primary, secondary, or tertiary depending on the *number of carbons bonded to the carbon that bears the OH*

Cycloalkanes

Saturated hydrocarbons with carbon atoms joined in a ring general formula: *CnH2n*

sigma vs pi bonds

Sigma bonds are stronger, more stable, and are shorter than pi bonds...this is because the orbitals are able to overlap much more and thus the electron sharing is much stronger...however, pi bonds are the result of 2 p orbitals that are parallel to each other that overlap side by side...this overlap is not as effective and thus the overlap that results creates a weaker, longer bond both sigma and pi orbitals can mix constructively (leading to sigma and/or pi bond formation) or destructively (no bond formation) sigma bonds are always localized between two atoms where as pi bonds can extend between two OR MORE atoms, therefore establishing the possibility of *resonance*

octet rule

States that atoms lose, gain, or share electrons in order to acquire a full set of eight valence electrons (an especially stable electronic configuration) This need for an octet drive bonding! Atoms can achieve an octet in two ways: 1) They can become *ions* by gaining or losing electrons, therefore resulting in *ionic bond formation* via the attraction between oppositely charged ion 2) they can *share* electrons between atoms to form *covalent bonds*

conformational isomers

Stereoisomers that differ by rotation about one or more single bonds, usually represented using Newman projections. single bonds, unlike double or triple bonds, are free to rotation, therefore allowing different spatial arrangements/conformations of a molecules different conformers have differing stabilities/energies associated with them types of conformers: ~~~ *staggered*, which include *gauche* (60 degrees), *anti* (180 degrees), and *eclipsed* (0 degrees) note: conformations are NOT different molecules... just different spatial arrangements. They are NOT frozen in place either... they interconvert rapidly creating "spinning" around single bonds

Stereocenter (chiral center)

Tetrahedral atom (usually carbon) with four different substituent groups chiral centers can have two different configurations known as R and S to figure out how many possible stereoisomers are for a molecule use 2^n where n is the # of chiral centers

valence electrons

The electrons in the outermost shell (main energy level) of an atom; these are the electrons involved in *forming bonds*. are the *highest energy electrons* most of organic chemistry comes down to *where* these valence electrons are, because *the behavior of organic molecules can be predicted based on where the electrons are*

atomic orbitals vs molecular orbitals:

They are both 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. when atomic orbitals (1s, 2s, 2px, etc) overlap with atomic orbitals on another atom (1s, 2s, 2px, etc), they mix to form molecular orbitals mixing of atomic orbitals can be CONSTRUCTIVE or DESTRUCTIVE since electrons act like waves. ~~~ constructive mixing → bonding MO (ex. sigma) ~~~ destructive mixing → antibonding MO (ex. sigma*)

VSEPR theory

Valence-shell electron-pair repulsion theory; Because electron pairs repel, molecules adjust their shapes so that valence electron pairs are as far apart as possible (*maximizes distance between electron density*) describes the three-dimensionality of molecules based on the # of regions of electron density around the central atom 4 regions of e- density → tetrahedral/ trigonal pyrimidal → 109.5° 3 regions of e- density → trigonal planar → 120° 2 regions of e- density → linear → 180°

tert-butyl

aka 1,1-dimethylethyl a very common substituent name!

isopropyl

aka 1-methylethyl a very common substituent name!

neopentyl

aka 2,2-dimethylpropyl a very common substituent name!

amide vs amine

amide: N is bound to a carbonyl carbon/carbonyl group amine: N bound to a non carbonyl carbon/group

line-angle formula

an abbreviated way to draw structural formulas in which vertices and line endings represent carbons hydrogens implied

diaxial strain

axial substituents can collide with each other creating strain, which is unfavorable since areas of electron density are CLOSER together in this position because of this, equilibrium favors larger substituents in the equatorial position because it relieves diaxial strain stability: bulky sub. on equatorial bonds are far more stable than bulky sub. on axial due to the increased electron density present from diaxial strain

constitutional isomers

compounds with the same molecular formula but different connections among their atoms

Amines

contain an amino group, which is a nitrogen atom bounded to one, two, or three carbon atoms by single bonds can be primary, secondary, or tertiary based on the *number of carbons bounded to the nitrogen atom directly*

cyclohexane conformations

cyclohexane is special because it can form an *unstrained conformation*, formally known as a *chair conformation* chair conformations have: >>> NO angle strain >>> NO steric strain >>> NO torsional strain (no eclipsing interactions)

primary, secondary, tertiary, quaternary carbons

depends upon how many carbons are bonded to the carbon under consideration - 1° (primary) = C attached to 1 other C - 2° (secondary) = C attached to 2 other C's - 3° (tertiary) = C attached to 3 other C's - 4° (quaternary) = C attached to 4 other C's primary H is attached to a primary carbon, a secondary H is attached to a secondary carbon, etc...

torsional strain

energy required to enforce the eclipsed conformation, relative to the staggered conformations that are unstrained molecules relief this strain by adopting staggered gauche or staggered anti conformations

formal charges

exist when an atom is surrounded by more or fewer valence electrons than it has in its neutral state *Formal charge== (# of valence electrons in the neutral, unbonded atom) - ( all unshared electrons + 1/2 of all shared electrons)* formal charges are useful because they indicate if the atom is electron-rich or electron-poor, which affects the molecule's behavior note: The sum of all formal charges must equal the total charge of the molecule

Naming Alkanes

for unbranched alkanes: 1. identify prefix via indicating the number of carbon atoms in the chain 2. add the sufix "-ane" to the end of the molecule name for branched alkanes: 1. Find longest parent chain in the compound 2. Number the chain (ex. propane) 3. Name the substituents, giving them the suffix "-yl". (ex. methyl). The substituent goes in front of parent chain name . (ex. methylpropane) 4. Assign a number to each substituent so that they get the lowest number possible. (ex. 2-methylpropane). If two of more substituents exist Prioritize alphabetical order (ex. 3-ethyl-2-methylpropane), making sure to not include di-, tri- tetra- prefixes in alphabetization (ex.3-ethyl-2,-dimethylpropane NOT 2,2-dimethyl-3-ethylpropane ) 5. Complete the name situational circumstances: 1) *identical substituents:* if two or more identical substituents exist, use the prefixes di-, tr-, tetra- and commas to separate position numbers (ex. 2,4- dimethylhexane). 2) *two parent chains having the same length*: chose the parent chain that has MORE substituents 3) *branched substituents*: if a substituent itself is branched, first name is as you would if it were a separate molecule, drop the -ane suffix for it and replace it with -yl, and then iniclude it in the molecule name in parenthesis alphabetized appropriately

hybridization and bonding

hybrid orbitals also mix in phase and out of phase to allow formation of bonds and/or the formation of nodal planes thus far we know that orbitals can mix to create sigma bonds (aka single bonds), but no more than two electrons can occupy the same space (pauli exclusion principle). So in order to form double/triple bonds we must utilize the unhybridized p orbitals ! sigma bonds only --> sp3, since it has no remaining unhybridized orbitals sigma + 1 pi bond --> allows UP TO double bond formation --> sp2, since it has one unhybridized orbital sigma + 2 pi bond --> allows UP TO triple bond formation --> sp, since it has two unhybridized orbitals

detecting chirality

is HARD to do but we can detect whether we have one enantiomer or both present in a mixture

Enantiomers

isomers that are non-superimposable mirror images of each other you know you have an enantiomer if... >>> The mirror image molecules only have one chiral center each, and one is "R" while the other is "S" >>> The mirror image molecules differ at ALL the chiral centers in question. (Ex. a molecule with 3 chiral centers R,S,R and its mirror image molecules having the configuration S,R,S)

shells vs atomic orbitals

key difference: shells house orbitals *orbitals*: regions around the nucleus in which a given electron or electron pair is likely to be found. Various exist, each with corresponding energies. Each subshell within an orbital can only have TWO electrons each. ex) *1s, 2s, 2px, 2py, 2pz, 3s, 3px, 3py, 3pz, plus five 3d orbitals*, where energy generally increase as shells increase *shells*: each electron shell can have different combinations of electronic orbitals ex) shell 1,2, and 3 shell 1 houses the 1s orbital, shell 2 houses 2s, 2px, 2py, 2pz orbitals, and shell 3 houses 3s, 3px, 3py, 3pz, plus five 3d orbitals *the last shell houses the valence electrons!!!*

biggest mistakes when naming alkanes:

not finding the longest parent chain forgetting to alphabetize numbering from the wrong side of the chain

electronegativity table trends for a periodic table:

on a periodic table, *electronegativity generally increases from left to right and down to up*. Following this trend, F is the most electronegative atom Why do we see an increase in electronegativity from bottom to top? ~~~ because the outer shell of the electrons is closer to the nucleus, which holds the positive protons Why do we see an increase in electronegativity from left to right? ~~~ because the number of protons in the nucleus is increasing the need for negatively charged electrons (balance)

showing stereochemistry

we show stereoisomers by indicating the relative orientation of their substituents with dashes and wedges dashes show substituents that are BEHIND the plane of the page wedges show substituents that are IN FRONT of the plane of the page isomers can be cis or trans; >>> cis; substituents on the same side of the ring >>> trans; substituents on opposite sides of the ring

chair flip

one chair conformation is converted to the other briefly passes through the boat conformation RULES: >>>all axial groups become equatorial >>>all equatorial groups become axial >>> substituents remain pointing in the same general direction (ex. if the substituent was "up" and axial then it will be "up" and equatorial as well) >>>dashed lines remain dashed >>>wedges remain wedges >>> keep in mind cis, trans isomerism conversion is slowed if there is bulky groups attached bulkiest group will favor the equatorial position interconversion between chairs happens rapidly so each cyclohexane mixture consists of BOTH chair conformations at all times

3) Hund's Rule

orbitals of equal energy are each *occupied by one electron* before any orbital is occupied by a second electron, and all electrons in *singly occupied orbitals must have the same spin* ex) 2px, 2py, and 2pz are all equal in energy. They will all get one electron first before they get a second electron

Conformations of Cycloalkanes

planar and puckered/envelope >>> puckering of cycloalkanes increases stability because it adjusts the bond angles so that they are closer to those predicted by VSEPR >>>planar conformation is ALWAYS less stable/less favorable than puckered conformation because of the greater angle strain and torsional strain that exists in planar conformation puckering reduces angle strain (allows more ideal bond angles) and reduces torsional strain because it reduces eclipsing interactions *all cycloalkanes AFTER cyclopropane will pucker*

angle strain

results when bond angles deviate from their ideal values by being stretched or compressed - As *ring sizes ^* * angle/ring strain (energy) decreases* in addition to torsional strain (due to eclipsing interactions), small cycloalkanes suffer from angle strain (also called small-ring strain) when the actual angles differ from the predicted angles given via VSEPR generally, small ring strain molecules are less stable than linear counterparts... linear alkanes are always more stable examples: cyclopropane and cyclobutane are good examples of small ring strain

shapes of atomic orbitals (s, p, d)

s orbital: Sphere p orbital: dumbbell d orbital: clover

2) Pauli Exclusion Principle

states that a *maximum of two electrons* can occupy a single atomic orbital but only if the electrons have *opposite spins* this means that the 1 p orbital can only have at MOST two electrons, for example

Diastereomers

stereoisomers that are not mirror images of each other. Differ in R/S configuration at some but not all chiral centers you know you have a diastereomer if... >>> There's more than one chiral center and they have the same R/S configuration at at least one chiral center (Ex. a molecule with 3 chiral centers R,S,R and its mirror image molecules having the configuration S,S,R)

torsional vs. steric strain

steric strain is when the atoms physically collide and is due the types of atoms present on molecule torsional strain is present in eclipsing interactions due to rotation around a bond

How do we determine which atomic orbitals will be occupied for a given atom?

we use THREE principles: 1) Aufbau: 2) Pauli exclusion 3) Hund's

dihedral angle

the angle between two specified groups in a Newman projection aka the angle difference between a front atom of a Newman projection and a back atom on a Newman projection

TWO important principles in organic chemistry:

the behavior of complex molecules can be predicted... 1) based on the behavior of their *component chemical groups* (ex. functional groups, acids/bases) 2) based on *where the electrons are*

functional groups

the components of organic molecules that are most commonly involved in chemical reactions have a characteristic set of physical and chemical properties

steric strain

the interference between two bulky groups that are so close together that their electron clouds experience a repulsion

hybridization

the mixing of several atomic orbitals to form the same total number of equivalent hybrid orbitals *which atomic orbitals hybridize depends on how many areas of electron density surround the atom*! 4 regions of e- density → 4 orbitals to hybridize (1s+ 3p)→ *sp3* 3 regions of e- density → 3 orbitals to hybridize (1s+ 2p)→ *sp2* 2 regions of e- density → 2 orbitals to hybridize (1s+ 1p)→ *sp*

polar molecule vs polar bond

the presence of polar bonds within a molecule DOES NOT mean the molecule itself is polar, you must consider the molecule geometry via VSEPR first to see if the molecular dipoles cancel out After you apply VSEPR to a molecule you can then determine polarity; ~~~ Polar molecule: net dipole moment (dipoles DONT cancel) ~~~ Nonpolar molecule: no not dipole moment (dipoles cancel) use tail-head vector connection to determine net dipole moment

1) Aufbau Principle

the rule that electrons occupy the orbitals of *lowest energy* first

Energy of Conformers

the staggered conformations (gauche and anti) are more stable and LOWER in energy because they spread out electron density moreso than eclipsed conformations because of this staggered conformations are called "unstrained conformations" because they require no energy to adopt unlike eclipsed conformations, which require energy. *the eclipsed conformation is higher in energy* and more strained because it takes energy to keep the substituent's electron density in close proximity *between gauche staggered and anti staggered, gauche is higher in energy due to steric strain* *more gauche interactions== higher in energy than molecules with no gauche interactions*

partial charge

the unequal sharing of electrons which results in a slight negative or positive charge stem from polar covalent bonds *Partial charges* are indicated by a plus/minus arrow or δ+/δ-

fisher projections

they depict molecules with multiple chiral centers RULES: >>> Vertical lines are assumed to be oriented behind the page on a DASH >>> Horizontal lines are assumed to be oriented out of the page on a WEDGE >>> Do NOT rotate molecules by 90 degrees when making fisher projections

Predicting bonds types: Ionic? Covalent (polar or non polar)?

to determine the bond type, we need to *subtract the electronegativity values* (from a electronegativity table) for the two atoms involved in the bond, and note the difference! difference *greater than 1.9 == ionic bond* difference of *0.5 to 1.9 == polar covalent* difference *less than 0.5 == nonpolar covalent*

branched vs unbranched

unbranched structures consist of single carbon chains whereas branched structures have substituents such as methyl groups that is attached to carbon chains

When atoms have multiple atomic orbitals, which ones participate in bonding?

with atoms like H that only have 1 atomic orbital (the 1s orbital) it is easy to see that when two H's form a bond, in both H's we know that the 1s orbital is the orbital that allows bonding... this is not the same with other atoms like C that have multiple different atomic orbitals!!! C has multiple s and p orbitals in various shells! the problem: atoms with multiple orbitals need the space between the orbitals to be as far apart as possible. P orbitals are 90 degrees apart (VSEPR), which is too close for a stable arrangement the solution: hybridization of orbital!!! they allow us to spread out the electron density in atomic orbitals

constructive vs destructive interference

• Constructive Interference - waves added together leading to *INCREASED* electron density between the two atoms, which *allow bond formation* • Destructive Interference - the waves subtract from each other as they overlap leading to DECREASED electron density between the atoms generating a *nodal plane* (area with no electron density).. *no electron density leads to no bonds and the atoms actually repel* mixing of atomic orbitals can be CONSTRUCTIVE or DESTRUCTIVE since electrons act like waves. ~~~ constructive mixing → bonding MO (ex. sigma) ~~~ destructive mixing → antibonding MO (ex. sigma*) note: *Whenever TWO atomic orbitals mix, they always generate TWO molecular orbitals*