Covalent Bonding, Nomenclature, and Molar Mass

  • Covalent bonding vs ionic bonding

    • Ionic bonds arise from attraction between positive and negative charges forming ionic compounds.
    • Covalent bonds arise when atoms share electrons to satisfy the octet rule.
    • In covalent bonding, electrons are shared as bonding pairs between atoms; nonbonding electrons are lone pairs.

  • Key concepts in covalent bonding

    • Bonding pair: the shared electrons between two atoms that holds them together.
    • Lone pair (nonbonding electrons): electrons that are not involved in bonding.
    • Octet rule for nonmetals: most nonmetals form enough covalent bonds so that each atom achieves 8 valence electrons.
    • Duet rule for hydrogen: hydrogen needs 2 electrons to achieve a full outer shell.
    • General idea: the number of covalent bonds a nonmetal forms equals the number of electrons it needs to gain to reach an octet (or a duet for H).

  • Example: fluorine molecule (F2) via covalent bonding

    • Fluorine (group 7A) has 7 valence electrons.
    • Two fluorine atoms share a pair of electrons to complete each atom’s octet.
    • The shared pair forms the covalent bond between the two F atoms.
    • Result: each F atom has a complete octet.
    • Note: covalent bonds are formed between nonmetals by sharing of electron pairs.

  • Electron dot structures vs line bond structures

    • HF as example:
    • H has 1 valence electron; F has 7 valence electrons.
    • They share a pair of electrons (bond) to form HF.
    • Bonding electrons: the shared pair; called the bonding electron pair.
    • Fluorine’s remaining electrons appear as lone pairs on F; H achieves a duet.
    • Electron dot formula vs line bond structure:
    • Electron dot formula shows bonding and lone pairs with dots.
    • Line bond structure replaces a bonding pair with a line (–) to indicate a covalent bond.
    • NH (in HF, H2O, NH3, CH4 examples) shows how lone pairs are arranged.

  • Water (H2O) as a covalent molecule

    • Oxygen (group 6A) has 6 valence electrons.
    • Two hydrogen atoms each contribute 1 valence electron.
    • Shared pairs: one pair between H and O on the left, and one pair between H and O on the right.
    • Electron dot structure shows two bonding pairs between H–O and two lone pairs on O.
    • Line bond structure can depict bonding pairs as lines; water is often drawn in a bent shape.
    • Bent geometry arises because the lone pairs repel bonding pairs, giving a V-shaped molecule.

  • Covalent bonding in other simple molecules

    • Fluorine (F) + Hydrogen (H) forms HF: 1 bond; F needs 1 more electron to complete octet, H needs 2 for duet; sharing gives both full shells.
    • Oxygen (O) forms 2 bonds in many compounds to reach octet (e.g., H2O, CO2).
    • Nitrogen (N) forms 3 bonds in ammonia (NH3) to reach octet; may have a lone pair on N.
    • Carbon (C) forms 4 bonds in methane (CH4) to reach octet; this is foundational for organic chemistry.
    • The number of covalent bonds equals the number of electrons the atom needs to gain to reach an octet (or duet for H).

  • Common covalent molecules and their bonding patterns

    • Ammonia (NH3): nitrogen (group 5A) has 5 valence electrons; forms 3 covalent bonds with H; has one lone pair; can give a trigonal pyramidal arrangement.
    • Methane (CH4): carbon (group 4A) has 4 valence electrons; forms 4 single bonds with 4 H atoms; carbon achieves an octet.
    • Acetylene (C2H2, HC≡CH): carbon–carbon triple bond; each carbon also forms a single bond to a hydrogen; the triple bond shares three pairs of electrons.
    • Hydrogen cyanide (HCN): H–C single bond; C≡N triple bond; N may have a lone pair; carbon and nitrogen reach octets.
    • Carbon dioxide (CO2): central carbon with two double bonds to O atoms (O=C=O); each O and C reach octets via two bonding pairs per O and two double bonds total for C.

  • Number and types of covalent bonds

    • Single bond: 1 bonding pair (2 electrons) shared between two atoms.
    • Double bond: 2 bonding pairs (4 electrons) shared between two atoms.
    • Triple bond: 3 bonding pairs (6 electrons) shared between two atoms.
    • Representation:
    • Single bond: a pair of electrons shown as two dots or a single line.
    • Double bond: two pairs of electrons or two lines.
    • Triple bond: three pairs of electrons or three lines.

  • Binary molecular nomenclature (two elements)

    • Prefixes to indicate number of atoms: Mono-, Di-, Tri-, Tetra-, Penta-, Hexa-, Hepta-, Octa-, Nona-, Deca
    • Rule: First element uses its full name; second element ends with -ide after the root name of the second element; the prefix indicates the number of atoms of that element.
    • First element prefix is not used when there is only one atom of the first element (no mono- on the first element).
    • Examples:
    • CO: carbon monoxide (one carbon, one oxygen) → no prefix on carbon; oxygen gets the -ide suffix as oxide.
    • CO2: carbon dioxide (one carbon, two oxygens) → carbon is named, oxygen uses di- and -ide (oxide).
    • Example with two nitrogens and five oxygens: N2O5 → dinitrogen pentoxide.

  • Notable naming exceptions (as discussed in the transcript)

    • CO is carbon monoxide (not a different name).
    • HCl is hydrogen chloride.
    • HF is hydrogen fluoride.
    • H2O is water.
    • Note: In standard naming, CO = carbon monoxide; HCl = hydrogen chloride; HF = hydrogen fluoride; H2O = water.

  • Atomic weight, formula weight, and molecular weight

    • Atomic weight: weight of a single atom (relative atomic mass) for elements.
    • Formula weight: sum of atomic weights in an ionic formula; applies to ionic compounds.
    • Molecular weight: sum of atomic weights in a molecular formula; applies to covalent (molecular) compounds.
    • For NaCl (sodium chloride): Na = 22.99 amu, Cl = 35.45 amu; formula weight = 22.99+35.45=58.44 extamu22.99 + 35.45 = 58.44 \, ext{amu}.
    • For water, H2O: H = 1.01 amu (each); O = 16.00 amu;
    • Molecular weight = 2imes1.01+16.00=18.02 extamu2 imes 1.01 + 16.00 = 18.02 \, ext{amu}.
    • Distinction:
    • If the compound is ionic (like NaCl), use formula weight.
    • If the compound is molecular (covalent, like H2O), use molecular weight.

  • Mass–mole (molar mass) relationship

    • The molar mass (formula weight or molecular weight) expressed in g/mol is the mass of one mole of that compound.
    • General relation: 1 extmol=Mextformulaormolecular extg1 \, ext{mol} = M_{ ext{formula or molecular}} \, ext{g}
    • Example: Iron(III) oxide formula is Fe₂O₃.
    • Atomic weights: Fe = 55.85 amu, O = 16.00 amu.
    • Formula weight (molar mass) of Fe₂O₃:
      • M<em>extFe</em>2extO3=2imes55.85+3imes16.00=111.70+48.00=159.70 extamuM<em>{ ext{Fe}</em>2 ext{O}_3} = 2 imes 55.85 + 3 imes 16.00 = 111.70 + 48.00 = 159.70 \, ext{amu}
      • Since this is a solid ionic compound, this value is used as the formula weight (molar mass) in g/mol: M<em>extFe</em>2extO3=159.70 extg/molM<em>{ ext{Fe}</em>2 ext{O}_3} = 159.70 \, ext{g/mol}

  • Worked example: mass of five moles of iron oxide (Fe₂O₃)

    • Given: n = 5.0 mol of Fe₂O₃.
    • Step 1: determine molar mass (formula weight) of Fe₂O₃:
    • M=159.70 extg/molM = 159.70 \, ext{g/mol}
    • Step 2: convert moles to mass using the mass–mole relation:
    • m=nimesM=5.0 extmolimes159.70 extg/mol=798.5 extgm = n imes M = 5.0 \, ext{mol} imes 159.70 \, ext{g/mol} = 798.5 \, ext{g}
    • Step 3: report with proper significant figures. Given 5.0 (three significant figures) and 159.70 (five significant figures, though typically reported to three or four), the final answer is
    • $$m ext{ (to 3 s.f.)} = 7.99 imes 10^2 ext{ g} \