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molecular formula

shows the number of atoms in each part of the compound, but doesn’t indicate structure.

Example: C₅H₁₂ = pentane

condensed formula

Specifies order of attachment and distinguishes compound structure.

Example: CH₃CH₂CH₂CH₂CH₃ = pentane

structural formula

• shows bonds between atoms (Lewis Structure)

  •  image is Pentane structure

bond line notation

Better depiction, easier to compare large molecule structures

• Vertex = carbon atom

• Hydrogens = implied when bonded to carbon

• Charged carbons form 3 bonds

• Wedge = coming out of page, dash = going into page

• Functional groups - non carbon atoms indicated by element symbol

Heteroatom - any atom in an organic molecule that isn’t C or H

structural / constitutional isomers

same formula, different connectivity

degrees of unsaturation

sum of all pi bonds in an organic molecule (double/triple bonds, rings), which tells you how many hydrogens are absent compared to a structure with the same number of carbons and no pi bonds.

s, p, d, f

order of orbitals

2

# of electrons in each orbital

periodic table row

corresponds with orbital number (Ex: row 2 = 2s and 2p)

filling shells rule

all orbitals receive one electron before any receive two

hybridization example

1 s and 3 p orbitals are combined to form 4 sp³ orbitals

electron domains

a lone pair or bond (single, double, triple = one domain). Number of domains = steric number.

sp³ orbitals

four electron domains, no unhybridized p orbitals, 25% “s” character

sp² orbitals

three electron domains, one unhybridized p orbital (allows for pi bond), 33.3% “s” character

sp orbitals

two electron domains, two unhybridized p orbitals (allows for 2 pi bonds), 50% “s” character

covalent bonds

• atomic orbitals overlap and two atoms share a pair of electrons

• Sigma and pi bonds

Sigma and pi bonds

• Sigma bonds - formed by direct overlap of s orbitals and/or hybridized orbitals.

• Pi bonds - Formed by lateral overlap of UNHYBRIDYZED p orbitals

• All single bonds are sigma

• Double bonds are one sigma and one pi

• Triple bonds are one sigma and two pi

nonpolar bonds

Atoms share electrons equally, and occurs when bonded atoms share similar electronegativity.

polar bonds

Atoms share electrons unequally, and occurs when bonded atoms have different electronegativity (more electronegative attracts electrons, arrow points toward more electronegative atom)

ionic bonds

One atom transfers an electron to another atom

bond strength

Single < Double < Triple

bond length

Single > Double > Triple

molecular geometry factors

central atom hybridization

central atom electron domain

tetrahedral geometry

sp³, 4 bonds, 0 lone pairs, 109.5^o angles, methane

trigonal pyramidal geometry

sp³, 3 bonds, 1 lone pair, < 109.5^o angles, ammonia

sp3 bent geometry

sp³, 2 bonds, 2 lone pairs, < 109.5^o angles, water

trigonal planar geometry

sp², 3 bonds, 0 lone pairs, 120^o angles, formaldehyde

sp2 bent geometry

sp², 2 bonds, 1 lone pair, < 120^o angles, pyridine

linear geometry

sp, 2 bonds, 0 lone pairs, 180^o angles, acetylene

cations and radicals

do not contribute to molecular geometry; still trigonal planar.

molecular geometry summary table

resonance

• The delocalization of electrons across p orbitals.

• For an atom to participate in __________, it must contain overlapping p orbitals.

• __________ increases stability.

• Charge of the molecule must be preserved.

Lone pair ←→ pi bond resonance

pi bond ←→ pi bond resonance

conjugated ring resonance

resonance RULES

• Maintain the same skeleton (no breaking sigma bonds).

• All structures must have valid Lewis structures (periodic table rows 1&2 can not have expanded octets).

• All structures must have the same net charge.

resonance hybrids

• combination of all possible resonance structures.

• Represents the distribution of electron density

• Structures with more stability contribute more to resonance hybrids.

• If an atom's lone pair is delocalized through resonance, it does not contribute to total electron domains.

factors that increase resonance contribution

• Maximized # of complete octets

• Minimized # of formal charges

• Negative charges on more electronegative atoms

• Positive charges on less electronegative atoms

  • Electron withdrawing halogens stabilize negative charge.

partial positives and negatives

represented with δ- and δ+

resonance hybrid images

resonance structures example

conjugation

more conjugated = more stable molecule

3 definitions of acids and bases

Arrhenius, Bronsted Lowry, and Lewis

Arrhenius acid and base

• Acid - Produces H⁺ in aqueous solutions.

  • Example - HCl

• Base - Produces OH^- in aqueous solutions.

  • Example - NaOH

Bronsted Lowry acid and base

• Acid - H⁺ donor

  • Examples - HCl, NH₄⁺, HOAc

• Base - H⁺ acceptor

  • Examples - NaOH, NH₃, OAc^-.

Lewis acid and base

• Acid - Electron pair acceptor (electrophile)

• Base - Electron pair donor (nucleophile)

water

can act as a base or an acid (amphoteric)

Ka = [H+][A-] / [HA]

Acid dissociation constant

• [H⁺] = concentration of hydrogen or hydronium.

• [A^-] = concentration of conjugate base.

• [HA] = concentration of remaining undissociated acid.

• Larger K_a = acid dissolves/dissociates more

  • >1 = strong acid

  • <1 = weak acid

pK_a = -log(K_a)

pK_a goes down as K_a goes up.

lower pKa = stronger acid

Ka and pKa image

solution protonation

• pH > pK_a → deprotonated solution (H is removed)

• pH < pK_a → protonated solution (H stays on)

equilibrium

determined by pK_a of acid and conjugate acid. ____________ favors the formation of the weaker acid (higher pK_a)

CARDIO acronym (the order matters)

Used for ranking acid strength

• Charge

• Atom

• Resonance Delocalization

• Induction

• Orbital Hybridization

Charge (CARDIO)

• More positive charge = stronger acid = more likely to let go of protons.

• Strength: H₃O⁺ > H₂O > OH^- > O^2-.

Atom (CARDIO)

• Acidity changes based on periodic table rows and columns. Down and to the right on PT makes a stronger acid and more stable anion. Find the conjugate base and look specifically at the atom carrying the negative charge.

  • Columns - going down the column increases acidity. HI > HBr > HCl > HF.

  • Rows - going to the right within a row increases acidity. HF > H₂O > NH₃ > CH₄.

Resonance delocalization (CARDIO)

Molecules with more resonance structures are more acidic because they can balance out charges (meaning conjugate base is more stable).

Induction (CARDIO)

• Highly electronegative atoms connected to carbons can stabilize the conjugate base due to pulling electrons toward itself, which increases acidity/stability.

  • The closer the electronegative atom is to the negatively charged atom in the conjugate base, the more stable it is.

  • More substituents also increases acidity (3F > 1F)

Orbital hybridization (CARDIO)

Greater “s” character increases acidity. sp > sp² > sp³ because in sp hybridized carbons, the atom connected to the carbon is closer to the nucleus of the carbon (closer = attraction is stronger)

INTRAmolecular forces

Forces between atoms within the SAME molecule. Important for understanding molecule structure.

INTERmolecular forces

Forces between SEPARATE neighboring molecules. Important for studying molecule interactions.

London Dispersion Forces (LDFs)

present in all molecules, made of temporary dipoles that form due to randomized electron movement.

• Molecules with more electrons experience more significant ________________.

• The weakest type of intermolecular force.

• This is the only force present in nonpolar molecules.

Dipole dipole interactions

arise from polar covalent bonds (EN difference between 0.5 and 1.7; higher difference = stronger interaction)

• Partial positive and negative charges will attract molecules together.

Hydrogen bond

Very strong D/D interaction; only occurs if NH, OH, or FH bonds are present.

• Partial + on H, partial - on N, O, or F.

Ion ion interaction

Cation interacts with anion. Ions bear a full + or - charge, resulting in a very strong force. More charge = stronger interaction.

IMF strength ranking

Ion ion > H bond > Dipole dipole > LDF

Solubility

molecules dissolve best in solvents with the same type of IMFs (“like dissolves like”)

• Example: hexane dissolves well in benzene because both are only capable of LDFs.

• Example: Lengthening a hydrocarbon chain attached to an OH DECREASES solubility because C-H is nonpolar while C-O is polar

Factors that increase melting/boiling point

• Stronger IMFs

• Larger molecule size

• Less branching

• More symmetricality*

Arrows

show movement of electrons, sometimes resulting in charges on atoms. They can originate at a bond or atom, and can form either a new bond, a lone pair, or radicals.

double headed arrow

depicts TWO electrons moving.

single headed arrow

depicts ONE electron moving.

poop intermolecular force

when the atoms of your nose are attracted to the smell of poop