Chem Unit 5 - Bond Types & Lewis Structures

Unit 5 Bonding quiz prep (Covalt Honors Chem)

Ionic bond

metal + nonmetal; electrons TRANSFERRED; + and − ions attract

Covalent bond

nonmetal + nonmetal; electrons SHARED

Metallic bond

metal + metal; 'sea' of delocalized electrons around + metal ions

ΔEN under 0.5

nonpolar covalent (Covalt treats C–H as nonpolar)

ΔEN 0.5 to 1.7

polar covalent (unequal sharing)

ΔEN 1.7 or more

ionic

Why do atoms bond?

to reach lower (more stable) energy; bond length = distance where energy is lowest

Properties of ionic compounds

high melting/boiling points, hard, brittle, conduct only when melted or dissolved

Properties of metals

conduct electricity and heat as solids, malleable, ductile, shiny

Properties of molecular (covalent) compounds

low melting/boiling points, often gases/liquids, don't conduct

Number of valence electrons

last digit of the group number (Group 1 → 1, Group 13 → 3, Group 17 → 7, Group 18 → 8)

Lewis dot symbol

element symbol with valence electrons as dots around it (N has 5 dots)

Lewis symbol for a cation (Na⁺)

dots removed, charge written: Na⁺

Lewis symbol for an anion (Cl⁻)

8 dots, in [brackets], charge outside: [:Cl:]⁻

Octet rule

atoms bond to get 8 valence electrons (H only needs 2)

Lewis step 1

count total valence e⁻; ADD 1 per − charge, SUBTRACT 1 per + charge

Lewis step 2

central atom = least electronegative (never H; usually C)

Lewis step 3

connect atoms with single bonds (2 e⁻ each)

Lewis step 4

fill outer atoms to 8 (H gets 2)

Lewis step 5

put leftover e⁻ on the central atom as lone pairs

Lewis step 6

if central atom has fewer than 8, move an outer lone pair in to make a double/triple bond

Electron exceptions: H, Be, B

H = 2, Be = 4, B = 6 (BF₃ is fine with 6 on B)

Expanded octet

period 3+ central atoms (P, S, Cl, Xe) can hold more than 8: PCl₅, SF₆

Covalt Honors rule on formal charge

do NOT add extra double bonds to minimize formal charge (that's AP). Use plain octet structures.

Brackets on ions in Lewis structures

polyatomic anions go in [brackets] with charge outside (Covalt's video drew NH₄⁺ without brackets)

Bonding pair vs lone pair

bonding pair = shared between 2 atoms (a line); lone pair = 2 dots on one atom, not shared

Single / double / triple bond electrons

2 / 4 / 6 electrons

Bond strength and length trend

triple > double > single in strength; triple is SHORTEST

Resonance

when a double bond could be in more than one equivalent spot (NO₃⁻, CO₃²⁻, O₃); real molecule is an average

Total valence e⁻ in CO₂

16 → O=C=O

Total valence e⁻ in H₂O

8 → 2 bonds + 2 lone pairs on O

Total valence e⁻ in NH₃

8 → 3 bonds + 1 lone pair on N

Total valence e⁻ in CH₄

8 → 4 single bonds, no lone pairs

Total valence e⁻ in NH₄⁺

8 (5 + 4 − 1)

Total valence e⁻ in SO₄²⁻

32 (6 + 24 + 2)

Total valence e⁻ in NO₃⁻

24 (5 + 18 + 1) → one N=O double bond

Total valence e⁻ in HCN

10 → H–C≡N