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Flashcards · AP® Chemistry · Unit 2

Unit 2: Compound Structure and Properties flashcards

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2.1 Types of Chemical Bonds

What is electronegativity, and how does it vary in the periodic table?
How strongly an atom in a bond attracts the shared electrons (Pauling scale, no unit; fluorine is highest at 3.98). It increases across a period (larger effective nuclear charge, same shell) and decreases down a group (extra shell, bonding electrons farther away and more shielded).
What is the difference between a nonpolar and a polar covalent bond?
Nonpolar: electrons shared between atoms of equal or similar electronegativity (H₂, Cl₂; C–H counts as nonpolar). Polar: unequal sharing; the more electronegative atom is δ− and the other δ+. For single bonds, a bigger electronegativity difference means a bigger bond dipole.
Is there a fixed electronegativity cut-off that makes a bond ionic?
No. Ionic and covalent are ends of a continuum; every polar bond has some ionic character. A fixed cut-off such as 1.7 fails: HF (difference 1.78) is covalent. Use the elements involved and, above all, the compound's properties.
How do you decide which atom gets δ− in a polar bond?
The more electronegative atom is δ−, even if both are nonmetals. In O–F, F is δ−; in Si–H (Si 1.90, H 2.20), H is δ−. Use δ+ and δ−, never full charges like Cl⁻, for a polar covalent bond.
What property evidence points to ionic, metallic or covalent molecular bonding?
Conducts as a solid and malleable → metallic. Conducts only when molten or dissolved, brittle, high melting point → ionic. Low melting point and does not conduct → covalent molecules.
How is metallic bonding described?
Metal cations with delocalized valence electrons that belong to no single atom. Do not describe metals as sharing fixed electron pairs between particular atoms.

2.2 Intramolecular Force and Potential Energy

What does a potential energy curve for two bonding atoms show?
Energy (kJ mol⁻¹) against internuclear distance (pm). Zero is the atoms far apart and not interacting. As they approach, attraction lowers the energy; too close, nucleus–nucleus and electron–electron repulsion raise it steeply.
How do you read bond length and bond energy from a potential energy curve?
The bond length is the distance at the minimum (not at the zero crossing). The bond energy is the depth of the minimum below zero, quoted as a positive value. A minimum at 120 pm and −500 kJ mol⁻¹ means 120 pm and 500 kJ mol⁻¹.
Does breaking a bond release or absorb energy?
Bond breaking always absorbs energy. The minimum of the curve is negative, but the bond energy (the energy needed to separate the atoms) is positive.
How do atom size and bond order affect bond length and bond energy?
Larger atoms (down a group) give longer, weaker bonds: C–F 134 pm, 485 kJ mol⁻¹; C–I 213 pm, 240 kJ mol⁻¹. Higher bond order gives shorter, stronger bonds: C–C 154 pm, 347 kJ mol⁻¹; C≡C 120 pm, 839 kJ mol⁻¹.
Is a double bond twice as strong as a single bond?
No. Bond energy rises with bond order, but not in proportion: C–C is 347 kJ mol⁻¹ and C=C is 614 kJ mol⁻¹, less than double.
How do you rank the strength of attraction between pairs of ions?
Coulomb's law, F ∝ q₁q₂ / r², with r ≈ sum of the ionic radii. Larger charges give stronger attraction (Mg²⁺O²⁻ > Na⁺F⁻); for the same charges, smaller ions do (Na⁺F⁻ > Na⁺Cl⁻; K⁺F⁻ > K⁺I⁻). Use both charge and size.
What features must a sketched potential energy curve have?
Labelled axes with units (energy in kJ mol⁻¹, distance in pm), a steep rise at short distance, a single minimum below zero at the bond length, and a curve that levels off at zero at large distance.

2.3 Structure of Ionic Solids

What is the structure of an ionic solid?
A repeating three-dimensional array of cations and anions extending to the surface of the crystal. Each ion is surrounded by ions of opposite charge, which maximises attraction and keeps like charges apart. There are no molecules.
What does the formula of an ionic compound like CaF₂ tell you?
Only the ratio of ions. One set of ions in that ratio is a formula unit. CaF₂ needs two F⁻ for each Ca²⁺, so a drawing with 5 Ca²⁺ needs 10 F⁻.
Which matters more for attraction in an ionic solid, ion charge or ion size?
Charge usually dominates. Attraction ∝ |q₊ × q₋| / d with d ≈ r₊ + r₋. A 2+/2− pair has a charge product 4 times that of a 1+/1− pair, so do not rank a 1+/1− solid above a 2+/2− solid because of a small size difference.
How do you compare attraction in two ionic solids with ions of equal charge?
Compare the distance between ion centres. Smaller ions give a shorter d and stronger attraction. LiF has d = 76 + 133 = 209 pm against 296 pm for LiI, so LiF has the stronger attraction.
What rules should a particulate drawing of an ionic solid follow?
Alternate + and − so like charges never touch; draw cations usually smaller than anions (K⁺ 138 pm, Cl⁻ 181 pm); match the formula ratio so the total charge is zero; label charges with signs or symbols, not colour alone. Never draw ion pairs.
Why does an ionic crystal shatter when a layer is forced to slide?
Sliding a layer lines up ions of the same charge next to each other, and they repel strongly, so the crystal cleaves. This is why ionic solids are brittle.
Why does a solid ionic compound not conduct electricity, although it contains ions?
Conduction needs mobile charged particles. In the solid the ions are fixed in the lattice. When molten or dissolved, the ions can move and the substance conducts.

2.4 Structure of Metals and Alloys

What is the model of bonding in a metal?
Each atom releases its valence electrons into a shared sea of delocalized electrons. The remaining positive ion cores sit in a regular, closely packed array. Metallic bonding is the attraction between the cores and the electron sea in every direction. It is not ionic.
How does the electron-sea model explain the properties of metals?
Mobile delocalized electrons conduct electricity and heat. Metals are malleable and ductile because layers of cores can slide while the electron sea adjusts, so there is no like-charge repulsion.
What is the difference between an interstitial and a substitutional alloy?
Interstitial: added atoms much smaller than the host fill gaps, and no host atom is removed (steel: C in Fe). Substitutional: added atoms of comparable size replace host atoms (brass: Zn for Cu).
How do you decide whether an alloy is likely interstitial or substitutional?
Compare radii: percent difference = |r(host) − r(added)| ÷ r(host) × 100. A small difference suggests substitutional (Cu/Zn 4.7%); a large one suggests interstitial (Fe/C 40.5%).
How do alloys differ in properties from the pure metal?
Interstitial alloys have a more rigid lattice: harder, less malleable and ductile (steel is less ductile than iron). Substitutional alloys have uneven rows, so layers slide less easily and the alloy is usually harder. Alloys still conduct.
How many delocalized electrons are in a model of a metal?
Electrons in the sea = number of cores × valence electrons per atom. 16 Mg atoms give 16 Mg²⁺ cores and 32 electrons; 12 Al atoms give 36 electrons.
Why can mass percent not be used as atom percent for an alloy?
The elements have different molar masses. Convert each mass to moles first. For steel (C and Fe) the difference is large; for 70% Cu / 30% Zn brass, 29% of the atoms are Zn.

2.5 Lewis Diagrams

What does a Lewis diagram show, and what does it not show?
All the valence electrons in a molecule or ion: shared pairs as lines and lone pairs as pairs of dots. Core electrons are not drawn. It shows connectivity, not shape.
What are the steps for drawing a Lewis diagram?
Count total valence electrons (add the charge for an anion, subtract for a cation). Put the least electronegative atom in the centre, H and F outside. Add single bonds, complete outer octets with lone pairs, put leftovers on the central atom, then make multiple bonds if the centre is short of 8. Recount; bracket ions.
Which atoms obey, fall short of, or exceed the octet rule?
C, N, O and F almost always end with 8 electrons and never exceed 8. H ends with 2. B and Be can have fewer than 8 (B in BF₃ has 6). Period 3 and beyond can exceed 8 (S in SF₄ has 10).
How do you count valence electrons for a polyatomic ion?
Add the group valence electrons of every atom, then add one per negative charge or subtract one per positive charge. NH₄⁺ has 8; CN⁻ has 10; SO₃²⁻ has 6 + 18 + 2 = 26.
What shortcut finds the number of bonds in an octet molecule?
Bonds = (needed − available) ÷ 2, where needed = 8 per non-H atom + 2 per H. For C₂H₂: needed 20, available 10, so 5 bonds = 2 C–H + 1 C≡C.
What are the usual bonding patterns of neutral H, C, N, O and outer halogens?
H: 1 bond, 0 lone pairs. C: 4 bonds, 0. N: 3 bonds, 1. O: 2 bonds, 2. Outer halogen: 1 bond, 3 lone pairs. PCl₃ has 3 bonds and 1 lone pair on P.
What are common errors when drawing Lewis diagrams?
Ignoring the ion's charge; putting H in the centre; leaving leftover electrons undrawn (SF₄'s last pair goes on S); expanding the octet of a period 2 atom; making a double bond to an outer halogen; missing brackets and charge on an ion.

2.6 Resonance and Formal Charge

What are resonance structures?
Lewis diagrams with the same skeleton and electron count that differ only in where lone pairs and multiple-bond pairs are placed. They are joined with ↔. Atoms never move between resonance structures.
What is a resonance hybrid?
When resonance structures are equivalent (formate, nitrite, nitrate, carbonate), the real species is one structure, a hybrid, not a molecule that flips between forms. The resonance bonds are identical in length and energy, and any charge is shared.
How do you calculate formal charge?
FC = V − N − ½B, where V is valence electrons, N nonbonding electrons and B bonding electrons. Shortcut: FC = V − dots − lines. The formal charges must add to the overall charge. Example: N in NH₄⁺ is 5 − 0 − 4 = +1.
How do you choose the best Lewis diagram using formal charge?
Octets for C, N, O and F (never more than 8); formal charges as close to zero as possible; any negative formal charge on the most electronegative atom; no like charges on neighbouring atoms. If octet rule and formal charge disagree (SO₄²⁻), name the criterion you use.
How do you find the average bond order in a resonance hybrid?
Average bond order = bonds to the equivalent atoms ÷ number of positions. CO₃²⁻ and NO₃⁻: 4 ÷ 3 = 1.33, not 1.5. Higher bond order means a shorter, stronger bond.
What is a limitation of the Lewis model shown by NO₂ or NO?
Species with an odd number of electrons (NO has 11 valence electrons; NO₂, ClO₂) cannot give every atom an octet. The model also assumes electrons sit in fixed pairs; resonance patches over electrons that are really spread out.
What formal charges go with each bonding pattern for C, N and O?
C: 4 bonds = 0; 3 bonds, 0 lone pairs = +1; 3 bonds, 1 lone pair = −1. N: 3 bonds + 1 lone pair = 0; 4 bonds = +1; 2 bonds + 2 lone pairs = −1. O: 2 bonds + 2 lone pairs = 0; 3 bonds + 1 lone pair = +1; 1 bond + 3 lone pairs = −1.

2.7 VSEPR and Hybridization

What is the basis of VSEPR theory, and how do you count electron domains?
Electron domains around an atom repel (Coulombic repulsion) and get as far apart as possible. Each lone pair counts as 1 domain; each bonded atom counts as 1, whether the bond is single, double or triple.
What shapes and angles go with 2, 3 and 4 electron domains?
2 domains: linear, 180°. 3: trigonal planar, 120° (1 lone pair: bent, just under 120°). 4: tetrahedral, 109.5°; 1 lone pair: trigonal pyramidal, about 107° (NH₃); 2 lone pairs: bent, 104.5° (H₂O).
What shapes go with 5 and 6 electron domains?
5 domains (lone pairs go equatorial): trigonal bipyramidal, seesaw (1 lone pair), T-shaped (2), linear (3). 6 domains: octahedral, square pyramidal (1), square planar (2, the second lone pair opposite the first).
Why do lone pairs reduce bond angles?
Repulsion order is lone pair–lone pair > lone pair–bond > bond–bond, so lone pairs push bonded atoms closer together. Multiple bonds also push a little harder than single bonds. The shape name uses atom positions only.
How do you assign hybridization to an atom?
Count its electron domains, including lone pairs: 2 → sp, 3 → sp², 4 → sp³. Hybridization labels are used only for 2–4 domains in this course.
What is the difference between σ and π bonds?
σ: head-on overlap on the internuclear axis. π: side-on overlap above and below the axis. A single bond is 1 σ; a double is 1 σ + 1 π; a triple is 1 σ + 2 π. π is weaker than σ (C=C 614 vs C–C 347 kJ mol⁻¹) and prevents rotation, allowing cis/trans isomers.
When does a molecule have a dipole moment?
When it has polar bonds and a shape that does not cancel them. SF₆ is octahedral, so its six bond dipoles cancel: no dipole moment. Polar bonds alone do not guarantee a polar molecule.
How many σ and π bonds are in HCN, and what is carbon's hybridization?
H–C≡N has 2 σ bonds and 2 π bonds. Carbon has 2 domains, so it is sp and the molecule is linear, 180°.

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