Source: CHEM 237 Practice Exam 1, University of Washington | Chapters 1, 2
Tags: electronegativity, polar bonds, partial charges, functional groups, hybridization, sp3, sp2, sp, molecular geometry, VSEPR, cortisol, CHEM 237
Difficulty: Intermediate
Prerequisites: Nomenclature and Lewis structures (see Part 1 notes). Familiarity with electronegativity trends from general chemistry.
Once you can draw a molecule, the next step is understanding its electronic character. Bond polarity tells you where electron density sits unevenly, which governs how molecules interact and where reactions happen. Functional groups are the reactive "hot spots" of organic molecules, and recognising them on sight is essential for predicting chemical behaviour. Hybridization and geometry link the Lewis structure to the three-dimensional shape of a molecule, which matters for everything from reaction mechanisms to drug design.
This material builds directly on Lewis structures and nomenclature. If those are shaky, revisit Part 1 before continuing.
A bond is polar when the two atoms have different electronegativities; the more electronegative atom pulls electron density toward itself and carries a partial negative charge (δ-). Functional groups (alcohols, ketones, alkenes, etc.) are specific atom arrangements that determine a molecule's reactivity, and you need to spot them by sight. Hybridization (sp3, sp2, sp) describes how an atom's orbitals blend to form bonds, and it dictates molecular geometry: tetrahedral, trigonal planar, or linear.
Electronegativity
A measure of an atom's ability to attract bonding electrons toward itself. On the Pauling scale, common values for organic chemistry: H 2.2, B 2.0, C 2.6, N 3.0, O 3.5, F 4.0.
In simple terms, the higher the number, the more that atom hogs the electrons.
Polar bond
A covalent bond between two atoms with different electronegativities. The electron density shifts toward the more electronegative atom.
Think of it as a tug-of-war where one side is stronger. The electrons spend more time near the winner.
Partial charge (δ+ and δ-)
The symbols used to indicate the uneven distribution of electron density in a polar bond. δ- sits on the more electronegative atom; δ+ on the less electronegative one.
Functional group
A specific arrangement of atoms within a molecule that determines its chemical reactivity and properties. Examples: hydroxyl (-OH, alcohol), carbonyl (C=O, ketone or aldehyde), carboxyl (-COOH, carboxylic acid), amino (-NH2, amine), alkene (C=C).
Think of functional groups as the parts of the molecule where things happen. The carbon skeleton is relatively inert; the functional groups are where reactions occur.
Hybridization
The mixing of atomic orbitals (s and p) to form new, equivalent hybrid orbitals used for bonding.
sp3: one s + three p orbitals mix to form four equivalent hybrid orbitals. Tetrahedral geometry, ~109.5° bond angles. Found on carbons with four single bonds (e.g. methane, ethane).
sp2: one s + two p orbitals mix to form three equivalent hybrid orbitals, with one unhybridised p orbital remaining for a π bond. Trigonal planar geometry, ~120° bond angles. Found on carbons with one double bond (e.g. ethene, carbonyl carbons).
sp: one s + one p orbital mix to form two equivalent hybrid orbitals, with two unhybridised p orbitals for two π bonds. Linear geometry, 180° bond angles. Found on carbons with a triple bond or two double bonds (e.g. acetylene, CO2).
Molecular geometry (VSEPR)
The three-dimensional arrangement of atoms around a central atom, determined by the number of electron groups (bonding pairs + lone pairs). Lone pairs take up more space and compress bond angles slightly.
A bond is polar when the electronegativity difference between the two atoms is roughly 0.4 or greater (though any nonzero difference creates some polarity). To identify all polar bonds in a molecule:
Compare electronegativity values for each pair of bonded atoms.
Label the more electronegative atom δ- and the less electronegative atom δ+.
C-H bonds (difference of 0.4) are often treated as essentially nonpolar in organic chemistry, though some courses count them. Check your instructor's convention.
Working through the practice exam molecule (Q4)
Given: H 2.2, B 2.0, C 2.6, N 3.0, O 3.5.
C-N bonds: N is more electronegative, so C is δ+ and N is δ-.
C-O bonds: O is more electronegative, so C is δ+ and O is δ-.
O-H bonds: O is more electronegative, so H is δ+ and O is δ-.
N-H bonds: N is more electronegative, so H is δ+ and N is δ-.
B-H bonds: H is more electronegative than B, so B is δ+ and H is δ-.
B-C bonds: C is more electronegative than B, so B is δ+ and C is δ-.
C-H bonds: small difference (0.4), borderline polar.
C-C bonds: nonpolar (same atom).
The practice exam asks you to identify functional groups on cortisol. Common functional groups to look for:
Alcohol (hydroxyl, -OH): an oxygen bonded to a carbon and a hydrogen. Cortisol has multiple -OH groups.
Ketone (C=O within a chain): a carbonyl group flanked by two carbons.
Alkene (C=C): a carbon-carbon double bond.
Primary, secondary, tertiary alcohol: classified by how many carbons are attached to the carbon bearing the -OH.
When counting hybridised carbons on a complex molecule like cortisol:
sp3 carbons: four single bonds (or single bonds plus bonds to heteroatoms). These are the majority in a steroid skeleton.
sp2 carbons: involved in a double bond (C=C or C=O).
sp carbons: involved in a triple bond. Cortisol has none.
For the practice exam molecule in Q6 (an amino acid derivative):
A carbon with four single bonds (e.g. -CH2-) is sp3, tetrahedral geometry around it.
A carbon with a double bond (e.g. C=O or C=C) is sp2, trigonal planar geometry.
A carbon with a triple bond (e.g. C≡C or C≡N) is sp, linear geometry.
Nitrogen with three bonds and one lone pair (as in an amine) is sp3, trigonal pyramidal geometry (tetrahedral electron geometry, but the shape defined by atoms alone is pyramidal).
Oxygen with two bonds and two lone pairs (as in an alcohol or ether) is sp3, bent geometry.
The nitrogen in an amide or when conjugated with a double bond can be sp2 (trigonal planar), because the lone pair participates in resonance.
The practice exam (Q10) presents structures with subtle errors. Common drawing mistakes to watch for:
A carbon with five bonds (carbon can only form four).
Incorrect number of hydrogens implied by a line-angle drawing. At each vertex, count the lines, then add hydrogens to reach four bonds.
A tetrahedral carbon drawn with impossible geometry in a chair or wedge-dash drawing (e.g. two wedge bonds on the same side when they should alternate).
An sp2 carbon drawn with tetrahedral geometry, or vice versa.
A double bond drawn inside a ring that would violate Bredt's rule (relevant for bridged bicyclic systems, less so at this stage).
Students often assume that a carbon with a double bond is sp3. If a carbon is part of a double bond (C=C or C=O), it is sp2. If it has a triple bond, it is sp. Only carbons with exclusively single bonds are sp3.
Geometry and hybridization are frequently conflated. Hybridization describes orbital mixing. Geometry describes the shape defined by the positions of atoms (or electron groups). A nitrogen with three bonds and one lone pair is sp3 hybridised, but its molecular geometry is trigonal pyramidal (not tetrahedral), because VSEPR counts the lone pair as an electron group but the shape name only describes atom positions.
Students sometimes label C-C bonds as polar. Two identical atoms bonded together always form a nonpolar bond, regardless of what else is attached to them.
When identifying functional groups, students occasionally confuse an alcohol (-OH bonded to a saturated carbon) with a carboxylic acid (-COOH). The presence of a neighbouring carbonyl (C=O) is what makes it an acid, not just the -OH.
⚠️ Expect a question where you are given electronegativity values and must label every polar bond in a molecule with δ+ and δ-. Do not skip any bonds.
⚠️ Functional group identification on a complex molecule (like a steroid) is a classic exam question. Be able to name each group, not just circle it.
⚠️ Hybridization questions often ask you to state the hybridization of every carbon in a molecule and then give the geometry around each non-hydrogen atom. Practise this on molecules with mixed hybridization (some sp3, some sp2, some sp).
⚠️ "Find the mistake" questions test whether you can spot a carbon with too many or too few bonds, or geometry that does not match the hybridization. These are designed to be subtle.
True or false: a C-O bond is polar, with carbon carrying the partial positive charge. (True.)
Fill in the blank: an sp2 hybridised carbon has ______ geometry with bond angles of approximately ______°. (Trigonal planar, 120.)
True or false: a nitrogen atom with three bonds and one lone pair has trigonal planar geometry. (False. It is trigonal pyramidal.)
Fill in the blank: the functional group -OH bonded to a saturated carbon is called an ______. (Alcohol.)
True or false: boron is more electronegative than hydrogen. (False. H is 2.2, B is 2.0.)
Q: Given electronegativity values H 2.2, C 2.6, N 3.0, O 3.5, identify all polar bonds in methanol (CH3OH) and label the partial charges.
A: The C-O bond is polar (C is δ+, O is δ-). The O-H bond is polar (O is δ-, H is δ+). C-H bonds have a small difference (0.4) and are borderline; if your course treats them as polar, C is δ- and H is δ+.
Q: What is the hybridization and geometry of the carbon atom in formaldehyde (H2C=O)?
A: The carbon is sp2 hybridised (one double bond to O, two single bonds to H). The geometry around carbon is trigonal planar, with bond angles of approximately 120°.
Q: A molecule of cortisol has two ketone groups and three hydroxyl groups. How many sp2 hybridised carbons are contributed by the ketone functional groups alone?
A: Two. Each ketone C=O contributes one sp2 carbon (the carbonyl carbon).
Q: A structure is drawn with a carbon that has two single bonds, one double bond, and one lone pair. What is wrong?
A: That carbon would have five items around it (two single bonds, one double bond, one lone pair), which exceeds carbon's maximum of four bonds and zero lone pairs. Carbon never has lone pairs in a neutral, valid structure (except in carbenes, which are not part of introductory organic chemistry).
Bond polarity feeds directly into understanding dipole moments and intermolecular forces (London dispersion, dipole-dipole, hydrogen bonding), which determine physical properties like boiling point and solubility. Functional group recognition is the foundation for the entire second half of organic chemistry, where each reaction chapter is organised by functional group. Hybridization and geometry become critical when you study stereochemistry, because the three-dimensional arrangement of groups around sp3 carbons creates chirality.
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