Difficulty: Intermediate | Prerequisites: Periodic table trends, electron configuration basics (Unit 1)
Bond polarity depends on the electronegativity difference between bonded atoms: the larger the gap, the more polar (or ionic) the bond. Ionic compounds form lattice structures held together by coulombic attractions whose strength depends on ion charge and ion size. Potential energy diagrams show you bond length (the minimum) and bond energy (the depth of the well), and comparing curves tells you which bond is shorter or stronger.
Electronegativity
A measure of an atom's ability to attract shared electrons in a chemical bond. In simple terms, the more electronegative an atom is, the harder it pulls the bonding electrons toward itself.
Bond polarity
The unequal distribution of electron density between two bonded atoms caused by a difference in their electronegativities. Think of it as one atom hogging the shared electrons more than the other.
Nonpolar covalent bond
A bond between atoms with the same (or very similar) electronegativity, so the electrons are shared equally. In simple terms, neither atom pulls harder than the other.
Polar covalent bond
A bond in which electrons are shared unequally because of a moderate electronegativity difference. Think of it as a tug of war where one side is slightly stronger.
Ionic bond
A bond formed by the electrostatic attraction between a cation and an anion, resulting from a large electronegativity difference (typically a metal and a nonmetal). In simple terms, one atom takes the electron outright rather than sharing it.
Partial charge (delta plus / delta minus)
A notation (d+ and d-) showing that electron density is shifted toward the more electronegative atom in a polar bond. The more electronegative atom carries the partial negative charge.
Lattice energy
The energy required to completely separate an ionic solid into individual gaseous ions. In simple terms, it tells you how tightly the ions are packed together in the crystal.
Coulomb's law (for ionic interactions)
The electrostatic force between two charged particles is proportional to the product of their charges and inversely proportional to the square of the distance between them. Think of it as: bigger charges and smaller ions mean stronger attractions.
Bond dissociation energy (BDE)
The energy required to break one mole of a specific bond in a gaseous molecule into gaseous atoms. A higher BDE means a stronger bond.
Bond length
The equilibrium distance between the nuclei of two bonded atoms, corresponding to the minimum on the potential energy curve.
Ionic character
The degree to which a bond resembles a full ionic bond. Increases with larger electronegativity differences between the bonded atoms.
The polarity of a bond is determined by the electronegativity difference (Delta EN) between the two bonded atoms. A larger difference means a more polar bond.
Ranking bond polarity in compounds
Calculate Delta EN for each bond in question.
Example: given electronegativities H = 2.1, C = 2.5, S = 2.5, F = 4.0, Cl = 3.0, Si = 1.8:
C-H bond: Delta EN = 0.4
Si-Cl bond: Delta EN = 1.2
S-F bond: Delta EN = 1.5
So in order of increasing bond polarity: CH4 < SiCl4 < SF4
Decreasing polarity series
H-F (Delta EN = 1.9) > N-F (Delta EN = 1.0) > F-F (Delta EN = 0): this is a valid ranking because polarity decreases as the difference shrinks and reaches zero for identical atoms.
HF as a case study
Fluorine (EN = 4.0) is far more electronegative than hydrogen (EN = 2.1).
The bond is highly polar: electron density shifts toward F, giving F a partial negative charge and H a partial positive charge.
Students sometimes pick "partial negative on H" by mistake. The partial negative always sits on the more electronegative atom.
Most polar bond in a set of carbon-halogen bonds
C-Cl has the largest Delta EN among C-Cl, C-I, C-H, C-Br because Cl is the most electronegative halogen in the set (EN = 3.0).
Ionic character spectrum
Nonpolar covalent: Delta EN close to 0 (e.g. Cl2, F-F).
Polar covalent: moderate Delta EN (e.g. H-F, C-Cl).
Ionic: large Delta EN, typically metal + nonmetal (e.g. CaCl2, NaCl).
Ranking by increasing ionic character: Cl2 (zero difference, pure covalent) < PCl3 (nonmetal-nonmetal, moderate difference) < CaCl2 (metal-nonmetal, large difference).
Equal electronegativity case
When two elements share the same EN value (e.g. C and Se, both 2.55), the bond between them is nonpolar covalent. There is no partial charge on either atom.
Least polar bonds
O-F has a very small Delta EN (3.5 vs 4.0 = 0.5), making it less polar than N-H, H-F, I-F, or O-H.
When does an ionic bond form?
An alkali metal reacting with a halogen will form an ionic compound. This is the most valid general claim you can make about such a pairing, since the formula, solubility, and other properties depend on the specific elements.
Compounds like NH4Cl contain both ionic bonds (between NH4+ and Cl-) and covalent bonds (N-H bonds within the ammonium ion).
Coulomb's law and melting/boiling points
The strength of ionic attractions follows Coulomb's law:
Force is proportional to (charge of cation x charge of anion) / (distance between ion centres)^2.
Higher ion charges = stronger attraction = higher melting/boiling points.
Smaller ions = shorter distance = stronger attraction.
MgO vs NaF: the classic comparison
MgO melts at 2852 degrees C; NaF melts at 993 degrees C.
Mg2+ has double the charge of Na+; O2- has double the charge of F-.
Both charge factors contribute, making MgO's lattice far stronger.
The O2- ion is actually slightly larger than F- (140 pm vs 133 pm), so the higher charges, not ion size, are the dominant factor here.
Predicting boiling points with ionic radius data
NaCl vs NaF: both have +1/-1 ions (same charges), but Cl- (181 pm) is larger than F- (133 pm).
Larger anion means greater distance between centres, weaker coulombic attraction.
Prediction: NaCl has a lower boiling point than NaF.
Lattice energy comparisons
NaF lattice energy: 930 kJ/mol; MgF2 lattice energy: 2978 kJ/mol.
The Mg cation carries a +2 charge vs Na's +1 charge. This charge difference is the primary reason MgF2's lattice energy is so much larger.
For separating ion pairs: Mg2+ and O2- require the most energy (among Na+/Cl-, Cs+/Br-, Mg2+/O2-, Ca2+/O2-) because both ions carry a charge of 2 and Mg2+ is smaller than Ca2+, giving a shorter distance.
Dissolution enthalpy
CaF2 dissolving in water has a positive Delta H (endothermic).
The Ca2+ ions have very strong ion-ion interactions with F- in the crystal lattice. Breaking apart that lattice requires more energy than is released by hydration.
Crystal lattice arrangement
Ionic solids form a regular, alternating pattern of cations and anions. Each positive ion is surrounded by negative ions and vice versa.
In a correct diagram of solid KF, the ions alternate in all directions: +, -, +, -, with each ion surrounded by ions of the opposite charge. All ions in KF are similar in size.
A diagram where like charges cluster together (all positives in one row, all negatives in another) does not represent an ionic lattice.
Conductivity of ionic solids
Ionic solids do not conduct electricity because their ions are locked in fixed positions within the lattice and cannot move freely.
Common wrong answer: "it conducts because it is made of ions." Having ions is necessary but not sufficient. The ions must be free to move (as in a molten salt or aqueous solution).
Brittleness and cleavage
When stress is applied to an ionic crystal, planes of ions slide relative to each other.
This brings like charges into alignment (positive next to positive, negative next to negative), generating strong repulsive forces.
The crystal shatters rather than deforming. This is brittleness.
The correct diagram shows an alternating +/- lattice that, after displacement, splits apart with like charges now facing each other across the break.
Density and lattice spacing (KCl vs KF)
KCl has a greater molar mass than KF, yet its density is actually lower.
The Cl- ion is significantly larger than F-, so the ions in KCl are spaced further apart, and the unit cell volume is larger.
A correct particulate diagram for KCl shows the same alternating pattern as KF, but with larger anions, creating more space in the lattice.
Reading a PE vs internuclear distance diagram
The x-axis is the distance between nuclei. The y-axis is potential energy.
At very large distances, the atoms do not interact, and PE is approximately zero.
As atoms approach, attractive forces lower the PE (curve dips below zero).
At very short distances, nuclear repulsion dominates and PE rises steeply.
The minimum of the curve corresponds to the bond length (equilibrium distance).
The depth of the well (from zero down to the minimum) equals the bond dissociation energy.
Correct shape of the H-H potential energy curve
Starts at zero for large separation.
Dips to a minimum below zero (at the bond length, about 74 pm for H2).
Rises steeply as distance approaches zero.
The curve does not cross zero again on the way down, and it does not have a minimum above zero. The correct graph shows a smooth well shape with a single minimum below the zero line.
At the minimum (bond length), forces are balanced
At distance x corresponding to the PE minimum, attractive and repulsive forces are exactly balanced.
The atoms maintain this average internuclear distance. This is the bond length.
A net attractive force exists at distances greater than the bond length (atoms are pulled closer).
A net repulsive force exists at distances shorter than the bond length (atoms are pushed apart).
Comparing diatomic PE curves (H2, N2, O2)
H2 has the shortest bond length (smallest x at the minimum) and moderate bond energy.
N2 has the deepest well (largest bond dissociation energy, approximately 945 kJ/mol) because of its triple bond.
O2 has intermediate bond energy (approximately 498 kJ/mol, double bond).
On a graph, the curve that sits furthest left with a shallow well is H2 (short bond, moderate energy). The deepest well is N2.
Comparing Br2 and I2
Br2 BDE = 193 kJ/mol; I2 BDE = 151 kJ/mol.
Br2 has a deeper well (stronger bond) and a shorter bond length than I2.
On the diagram, the Br2 curve dips lower and its minimum sits to the left of the I2 curve.
HCl vs HBr on a PE diagram
HCl has a shorter bond length and higher bond energy than HBr.
The dashed HBr curve should have its minimum further to the right and not as deep as HCl's solid curve.
Bond enthalpies of F2, O2, N2
F2 has only a single bond (BDE approximately 159 kJ/mol).
O2 has a double bond (BDE approximately 498 kJ/mol).
N2 has a triple bond (BDE approximately 945 kJ/mol).
Order from least to greatest: F2 < O2 < N2.
Internuclear distances: H2, HCl, Cl2
H2 is the shortest (two small atoms, triple/single notwithstanding, H is tiny).
Cl2 is the longest (two large chlorine atoms).
HCl falls in between.
From shortest to longest: H2 < HCl < Cl2.
Bond length and bond order
N2 (triple bond) has the shortest bond length among common diatomics (N2, O2, Cl2, Br2, I2).
Higher bond order = shorter bond and higher bond energy.
Students often think ionic solids conduct electricity because they contain ions. They do not. The ions must be free to move, which only happens when the solid is melted or dissolved.
Students sometimes assign the partial negative charge to the less electronegative atom. The partial negative charge always goes on the atom with the higher electronegativity.
Confusing bond polarity with molecular polarity. A molecule can have polar bonds yet be nonpolar overall if the bond dipoles cancel due to symmetry (e.g. CCl4, CO2).
Assuming that a higher molar mass automatically means a higher density. Density also depends on how much space the particles take up. KCl has a higher molar mass than KF but a lower density because the larger Cl- ions spread the lattice out.
Thinking that bond length is where the PE curve crosses zero. Bond length is at the minimum of the curve, not where it crosses zero.
Electronegativity difference questions appear on nearly every AP Chemistry exam. Be ready to calculate Delta EN and rank bonds.
Coulomb's law reasoning (charge and distance) is the go-to explanation for melting point, boiling point, and lattice energy comparisons of ionic compounds.
PE curve interpretation is a recurring free-response and MCQ topic. Know how to identify bond length, bond energy, and the correct curve shape.
Questions about which compounds contain both ionic and covalent bonds (e.g. NH4Cl) are a favourite.
Ionic solid conductivity: expect a question distinguishing solid-state behaviour (no conduction) from molten/dissolved (conduction).
True or False: A bond between two atoms with the same electronegativity is ionic. (False, it is nonpolar covalent.)
Fill in the blank: The bond length on a PE diagram is found at the ______ of the curve. (minimum)
True or False: MgO has a higher melting point than NaF primarily because Mg and O atoms are heavier. (False, it is because Mg2+ and O2- have greater charges.)
True or False: Solid NaCl conducts electricity. (False, the ions are fixed in place.)
Fill in the blank: Among F2, O2, and N2, the molecule with the highest bond dissociation energy is ______. (N2)
Q: Given electronegativities of H (2.1), C (2.5), Si (1.8), and Cl (3.0), rank the bonds C-H, Si-Cl, and C-Cl in order of increasing polarity.
A: C-H (Delta EN = 0.4) < Si-Cl (Delta EN = 1.2) < C-Cl... wait, Si-Cl = 1.2 and C-Cl = 0.5. So: C-H (0.4) < C-Cl (0.5) < Si-Cl (1.2). Calculate each Delta EN, then sort.
Q: Why is the melting point of MgO (2852 degrees C) so much higher than that of NaF (993 degrees C)?
A: The charges on Mg2+ and O2- (both magnitude 2) are greater than the charges on Na+ and F- (both magnitude 1). By Coulomb's law, greater charges produce stronger electrostatic attractions, requiring more energy to separate the ions.
Q: On a PE diagram, two curves are shown. Curve A has a deeper minimum and its minimum is at a shorter internuclear distance than Curve B. Which molecule has the stronger bond? Which has the longer bond length?
A: Curve A has the stronger bond (deeper well = higher BDE) and the shorter bond length (minimum further to the left). Curve B has the weaker, longer bond.
Q: Explain why solid KF does not conduct electricity but molten KF does.
A: In solid KF, the K+ and F- ions are held in fixed lattice positions and cannot migrate. In molten KF, the lattice breaks apart and the ions are free to move, allowing them to carry electric charge.
Q: Which ion pair requires the most energy to separate: Na+ and Cl-, Cs+ and Br-, Mg2+ and O2-, or Ca2+ and O2-?
A: Mg2+ and O2-. Both ions carry a charge of magnitude 2 (larger product of charges than the +1/-1 pairs), and Mg2+ is smaller than Ca2+, so the interionic distance is shorter. Both factors increase the coulombic attraction.
This connects to intermolecular forces (Unit 3) because the type of bonding within a molecule (polar vs nonpolar) determines what kind of IMFs act between molecules. A polar bond in a symmetric molecule (like CO2) results in no net dipole and only weak London dispersion forces, while the same polar bonds in an asymmetric molecule (like H2O) give a strong dipole and hydrogen bonding.
Lattice energy connects to thermodynamics (Unit 5) through Hess's law and Born-Haber cycles, where lattice energy is one step in calculating the enthalpy of formation of an ionic compound.
Potential energy curves link to kinetics (Unit 5) and equilibrium: the concept of energy minima and the balance of attractive and repulsive forces recurs throughout the course.
electronegativity, bond polarity, polar covalent, nonpolar covalent, ionic bond, ionic character, electronegativity difference, delta EN, partial charge, delta plus, delta minus, lattice energy, Coulomb's law, coulombic attraction, melting point ionic compounds, boiling point ionic compounds, ion charge, ion size, ionic radius, crystal lattice, ionic solid conductivity, brittleness ionic crystal, potential energy curve, PE diagram, bond length, bond dissociation energy, BDE, bond enthalpy, bond order, diatomic molecules, H2, N2, O2, F2, Br2, I2, KF, NaCl, MgO, NaF, MgF2, CaF2, AP Chemistry Unit 2