Chapter 3

Question 1

What is the octet rule?

Answer 1

Atoms gain/lose/share e⁻ to obtain 8 valence electrons (noble‑gas configuration)

Question 2

What are incomplete-octet elements?

Answer 2

H 2 e⁻; He 2; Li 2; Be 4; B 6

Question 3

What are expanded-octet elements?

Answer 3

Period 3 + (P 10 e⁻, S 12, Cl 14, etc.)

Question 4

What are odd-electron molecules?

Answer 4

Species with an odd total of valence e⁻, e.g. NO, NO₂

Question 5

What does ΔEN 0‑0.5 indicate?

Answer 5

Non‑polar covalent bond (equal sharing)

Question 6

What does ΔEN 0.5‑1.7 indicate?

Answer 6

Polar covalent bond (unequal sharing, dipole)

Question 7

What does ΔEN > 1.7 indicate?

Answer 7

Ionic bond (electron transfer)

Question 8

What are the characteristics of ionic compounds?

Answer 8

Metal + non‑metal; crystalline lattice; very high melting/boiling points; conduct electricity when molten/aqueous; often soluble in polar solvents

Question 9

What are the characteristics of covalent compounds?

Answer 9

Non‑metals; discrete molecules; lower melting/boiling points; poor conductors in any phase

Question 10

What is a coordinate covalent bond?

Answer 10

Lone pair from one atom (Lewis base) donated to empty orbital of Lewis acid; after formation identical to any covalent bond

Question 11

What happens to metals and non-metals in terms of electron transfer?

Answer 11

Metals lose e⁻ become cations; non‑metals gain e⁻ become anions

Mnemonic “MeTals → caTions (+); NoN‑metals → aNions (–)”

Question 12

How does bond order relate to bond length and energy?

Answer 12

Higher bond order → shorter length → greater bond energy; triple > double > single

Question 13

What is the dipole moment equation?

Answer 13

μ = q × d (charge magnitude × distance)

Question 14

What are the steps to draw a Lewis structure?

Answer 14

1) Arrange skeleton; 2) Sum valence e⁻ (+ for ‑charge, – for +charge); 3) Single bonds; 4) Complete octets on outer atoms then central; 5) Make multiple bonds if central lacks octet; 6) Check formal charges

Question 15

What is the formal charge formula?

Answer 15

FC = valence – non‑bonding e⁻ – ½ (bonding e⁻)

Question 16

What are the rules for the most stable resonance structure?

Answer 16

Minimize formal charges; minimize charge separation; place negative charge on most electronegative atom

Question 17

What is a sigma bond?

Answer 17

Head‑on orbital overlap; free rotation

Question 18

What is a pi bond?

Answer 18

Parallel p‑orbital overlap; no rotation; accompanies sigma in double/triple bonds

Question 19

What is the result of molecular orbital overlap?

Answer 19

Same sign → bonding MO (stable); opposite sign → antibonding MO (unstable)

Question 20

What is the VSEPR shape for 2 regions?

Answer 20

Linear, 180°

Question 21

What is the VSEPR shape for 3 regions?

Answer 21

Trigonal planar, 120°

Question 22

What is the VSEPR shape for 4 regions?

Answer 22

Tetrahedral, 109.5°

Question 23

What is the VSEPR shape for 5 regions?

Answer 23

Trigonal bipyramidal, 90°/120°/180°

Question 24

What is the VSEPR shape for 6 regions?

Answer 24

Octahedral, 90°/180°

Question 25

How do lone pairs affect bond angles?

Answer 25

Lone pairs repel more than bonding pairs; compress angles (e.g. NH₃ ≈ 107°, H₂O ≈ 104.5°)

Question 26

What is the definition of coordination number?

Answer 26

Number of atoms bonded to central atom (used for molecular geometry determination)

Question 27

What is the molecular polarity rule?

Answer 27

Non‑polar bonds ⇒ non‑polar molecule; polar bonds + symmetric geometry ⇒ non‑polar; polar bonds + asymmetric geometry ⇒ polar

Question 28

Give an example of a polar molecule.

Answer 28

H₂O (bent, net dipole)

Question 29

Give an example of a non-polar molecule with polar bonds.

Answer 29

CCl₄ (tetrahedral, dipoles cancel)

Question 30

What are London dispersion forces?

Answer 30

Instantaneous/induced dipoles; present in all atoms/molecules; strength ↑ with size and polarizability; weakest IMF

Question 31

What are dipole–dipole interactions?

Answer 31

Permanent molecular dipoles align; stronger than London; negligible in gas phase at high T

Question 32

What is a hydrogen bond?

Answer 32

Strong dipole–dipole between H attached to F, O, or N and a lone pair on F/O/N in another molecule; strength ≈ 10 % of covalent bond

Question 33

What is the hydrogen-bond mnemonic?

Answer 33

“Pick up the FON (phone)” – F O N required

Question 34

What is the order of intermolecular force strength?

Answer 34

Hydrogen bonding > Dipole–dipole > London dispersion

Question 35

How does IMF affect physical properties?

Answer 35

Higher IMF → higher boiling/melting points, higher viscosity, lower vapor pressure

Question 36

What is the diatomic element mnemonic?

Answer 36

BrINClHOF – Br₂ I₂ N₂ Cl₂ H₂ O₂ F₂

Question 37

What is the bond classification quick line?

Answer 37

“0‑0.5 same, 0.5‑1.7 share, >1.7 tear” (non‑polar, polar, ionic)

Question 38

What is the formal charge on N in NH₄⁺?

Answer 38

+1

Question 39

How many σ and π bonds are in CH₃‑C≡N?

Answer 39

6 sigma, 2 pi

Question 40

What is the geometry and angle of SF₆?

Answer 40

Octahedral, 90°

Question 41

Rank the IMF of CH₃OH vs HCl vs CO₂.

Answer 41

CH₃OH (H‑bond) > HCl (dipole) > CO₂ (London only)

Question 42

Does BrF₃ have a net dipole?

Answer 42

Yes, T‑shaped, vectors do not cancel

Question 43

Rank the bond length order of C‑C, C=C, C≡C.

Answer 43

C≡C shortest; C‑C longest

Question 44

Is water a polar solvent?

Answer 44

Yes – bent geometry, large net dipole