Unit 2 Practice: Hard

Answer: Shifting the lattice can place ions with the same charge next to each other, creating strong repulsion.

Explanation: In the normal crystal, positive and negative ions alternate.

When force shifts one layer, positive ions may become aligned with positive ions and negative ions with negative ions.

Like charges repel strongly.

The crystal can then fracture instead of bending.

Answer: LiF

Explanation: Both contain ions with charge magnitudes of 1.

The major difference is ionic size.

Li⁺ is much smaller than Cs⁺, allowing Li⁺ and F⁻ to approach more closely.

Coulombic attraction increases as distance decreases.

Answer: CaO

Explanation: CaO contains:

Ca²⁺ and O²⁻

KBr contains:

K⁺ and Br⁻

The larger charge product in CaO produces much stronger electrostatic attraction. Its ions are also relatively compact.

Answer: Graphite contains delocalized electrons that can move through its layers.

Explanation: Each carbon in graphite forms three major covalent bonds, leaving electrons that can become delocalized across the sheet.

Diamond uses each carbon in a four-bond three-dimensional network.

Its electrons are localized in covalent bonds and are not free to carry current.

Answer: The strong bonds are mainly within each sheet, while attractions between sheets are much weaker.

Explanation: The carbon layers can slide over one another because the forces holding different layers together are much weaker than the covalent bonds inside each layer.

Answer: 1.5

Explanation: Each resonance structure has:

one O=O double bond one O—O single bond.

Because the two positions are equivalent in the resonance hybrid:

(2 + 1)/2 = 1.5

Both real O—O bonds therefore have intermediate character.

Answer: The actual electron distribution is delocalized and is represented by the resonance hybrid.

Explanation: Resonance structures are drawing tools.

They represent different valid ways of placing electrons on paper.

The molecule does not need to physically switch between the drawings.

Answer: It helps identify structures with more reasonable electron distributions.

Explanation: When several structures satisfy basic electron-counting rules, chemists generally favor structures with:

smaller formal-charge magnitudes

less unnecessary charge separation

negative formal charge on more electronegative atoms when appropriate

Formal charge is a bookkeeping model, but it helps compare Lewis structures.

Answer: Seesaw and polar

Explanation: SF₄ has five electron groups:

4 bonding 1 lone pair

The electron geometry is trigonal bipyramidal.

The lone pair prefers an equatorial position, producing a seesaw molecular shape.

The bond dipoles do not cancel completely, so SF₄ is polar.

Answer: T-shaped and polar

Explanation: Cl has five electron groups:

3 bonding 2 lone pairs

The electron geometry is trigonal bipyramidal.

The two lone pairs prefer equatorial positions, producing a T-shaped molecule.

Its bond dipoles do not cancel, so it is polar.

Answer: Square planar and nonpolar

Explanation: Xe has:

4 bonding groups 2 lone pairs

The electron geometry is octahedral.

The lone pairs occupy opposite positions.

The four identical Xe—F bonds form a symmetric square, allowing their dipoles to cancel.

Answer: Lone pairs repel neighboring electron groups more strongly than bonding pairs do.

Explanation: A bonding pair is attracted by two nuclei and is more concentrated between them.

A lone pair is localized mainly around one central atom and occupies more space around that atom.

Its stronger repulsion pushes bonding groups closer together.

Answer:

LP–LP > LP–BP > BP–BP

Explanation: Lone pairs occupy more space around the central atom than bonding electron pairs.

Therefore interactions involving lone pairs tend to create greater repulsion.

Answer: sp hybridized

The C≡C bond contains:

1 sigma + 2 pi bonds

Explanation: Each carbon has two electron groups:

one C—H region one C≡C region

Two electron groups correspond to sp hybridization.

The sp orbitals form the sigma framework while two remaining unhybridized p orbitals form the two pi bonds.

Answer: 5 sigma bonds and 1 pi bond

Explanation: Ethene contains four C—H single bonds:

4 sigma

The C=C bond contains:

1 sigma + 1 pi

Total:

5 sigma

1 pi

Answer: All of the bonding electron density connects the same two atoms and occupies the same general region around the central atom.

Explanation: VSEPR counts regions of electron density, not individual bond lines.

Therefore:

single bond = one group double bond = one group triple bond = one group

Answer: Electron geometry: tetrahedral

Molecular geometry: bent

Explanation: Four total electron groups always give tetrahedral electron geometry.

But molecular geometry considers only the positions of atoms.

With two bonded atoms and two lone pairs, the visible shape is bent.

H₂O is the classic example.

Answer: Its geometry can arrange the bond dipoles so they cancel.

Explanation: Bond polarity does not automatically mean molecular polarity.

For example:

CO₂ is linear.

Its two C=O dipoles point in opposite directions and cancel.

BF₃ is trigonal planar.

Its three equal B—F dipoles cancel because of the molecule's symmetry.

Answer: Electron sharing can occur with different amounts of charge separation depending largely on electronegativity difference.

Explanation: When atoms have similar electronegativities, electrons are shared relatively evenly.

As the electronegativity difference grows, electron density becomes increasingly uneven.

The bond develops greater ionic character.

So real bonds range from relatively nonpolar covalent to highly polar/ionic rather than always fitting perfectly into two boxes.

Answer: Higher bond order generally means a shorter and stronger bond.

Explanation: Compare:

single → double → triple

As bond order increases, more electron density exists between the two nuclei.

That increases attraction between the nuclei and shared electrons.

The atoms are pulled closer together, shortening the bond, and more energy is required to separate them.

Therefore:

bond order ↑ → bond length ↓ → bond strength ↑

For example:

C—C is longer and weaker than C=C, while C≡C is generally the shortest and strongest of the three.