Chapter 3 - Compounds

Course Overview, Assessment Objectives, and Required Knowledge

• Assessments and Curricular Structure:

  • Quiz Component: Evaluates core chemical vocabulary, main-group ionic charge predictability, and specific elemental symbols and spellings.

  • Test Component: Comprehensive evaluation covering periodic trends, atomic interactions, bond types, ionic and covalent compound properties, chemical nomenclature, and Lewis structures.

  • Test Extra Credit Verse: Romans 6:13 (NIV) — "Do not offer any part of yourself to sin as an instrument of wickedness, but rather offer yourselves to God as those who have been brought from death to life; and offer every part of yourself to him as an instrument of righteousness."

• Core Vocabulary and Formal Definitions:

  • Atomic Radius: Half the distance between the nuclei of identical atoms that are chemically bonded together.

  • Electronegativity ($\chi$): A measure of the attraction an element has for the electrons participating in a chemical bond.

  • Chemical Bond: A mutual electrostatic attraction between the nuclei and valence electrons of different atoms that binds the atoms together and stabilizes them.

  • Ionic Bond: A chemical bond resulting from the electrostatic attraction between oppositely charged ions (cations and anions), formed when valence electrons are transferred from one atom to another due to a large electronegativity difference ($\Delta\chi$).

  • Covalent Bond: A chemical bond formed when atoms share one or more pairs of valence electrons due to a low electronegativity difference ($\Delta\chi$).

  • Metallic Bond: A chemical bond formed by the electrostatic attraction between positive metal ions and a surrounding delocalized, mobile "sea of electrons."

  • Ion: An atom or bonded group of atoms that bears a net positive or negative electrical charge due to the loss or gain of electrons.

  • Cation: A positively charged ion formed when an atom loses one or more valence electrons.

  • Anion: A negatively charged ion formed when an atom gains one or more valence electrons.

  • Polyatomic Ion: A group of two or more covalently bonded atoms that carries an overall net electrical charge.

  • Octet Rule: The chemical principle stating that atoms tend to gain, lose, or share electrons in order to acquire a stable electron configuration containing eight valence electrons (a full outer shell), with the exception of the first energy level ($H$ and $He$), which is filled by two electrons.

  • Formula Unit: The simplest representative collection of cations and anions that preserves overall electrical neutrality in an ionic compound.

  • Molecule: A neutral group of atoms held together exclusively by covalent bonds; the representative unit of a covalent compound.

  • Electrolyte: A chemical substance that dissociates into freely mobile ions when dissolved in water (or molten), thereby conducting an electric current.

  • Chemical Formula: A symbolic representation indicating the simplest whole-number ratio or exact number of atoms of each element comprising a unique chemical compound.

  • Single Bond: A covalent bond in which two atoms share one pair of valence electrons ($2\,e^-$).

  • Double Bond: A covalent bond in which two atoms share two pairs of valence electrons ($4\,e^-$).

  • Triple Bond: A covalent bond in which two atoms share three pairs of valence electrons ($6\,e^-$).

  • Nonpolar Covalent Bond: A covalent bond characterized by electrical symmetry and an even distribution of electron density due to an identical or very low electronegativity difference ($\Delta\chi$) between the bonded atoms.

  • Polar Covalent Bond: A covalent bond characterized by electrical asymmetry, in which electrons are shared unequally due to a significant electronegativity difference ($\Delta\chi$), producing a bond dipole with a partial positive pole ($\delta^+$) and a partial negative pole ($\delta^-$).

• Required Elemental Memorization:

  • Group 1 (Alkali Metals and Hydrogen):

  • Hydrogen ($H$)

  • Lithium ($Li$)

  • Sodium ($Na$)

  • Potassium ($K$)

  • Rubidium ($Rb$)

  • Cesium ($Cs$)

  • Francium ($Fr$)

  • Group 2 (Alkaline Earth Metals):

  • Beryllium ($Be$)

  • Magnesium ($Mg$)

  • Calcium ($Ca$)

  • Strontium ($Sr$)

  • Barium ($Ba$)

  • Radium ($Ra$)

  • Group 17 (Halogens):

  • Fluorine ($F$)

  • Chlorine ($Cl$)

  • Bromine ($Br$)

  • Iodine ($I$)

  • Astatine ($At$)

  • Elements 46 through 52:

  • Element 46: Palladium ($Pd$)

  • Element 47: Silver ($Ag$)

  • Element 48: Cadmium ($Cd$)

  • Element 49: Indium ($In$)

  • Element 50: Tin ($Sn$)

  • Element 51: Antimony ($Sb$)

  • Element 52: Tellurium ($Te$)

  • Elements 78 through 84:

  • Element 78: Platinum ($Pt$)

  • Element 79: Gold ($Au$)

  • Element 80: Mercury ($Hg$)

  • Element 81: Thallium ($Tl$)

  • Element 82: Lead ($Pb$)

  • Element 83: Bismuth ($Bi$)

  • Element 84: Polonium ($Po$)

Periodic Trends (Section 3.1)

• Nature of Periodic Trends:

  • Many elemental properties change in a predictable, systematic manner across a period (horizontal row) or down a group (vertical column) of the periodic table.

  • Trends are governed by two physical variables:

  • The number of protons located in the nucleus (effective nuclear charge, which increases across a period).

  • The principal energy level ($n$) and spatial distribution of the valence electrons (electron shielding, which increases down a group).

• Atomic Radius:

  • Definition: Half the internuclear distance ($2r$) between two identical atoms chemically bonded together: $r = \frac{d_{\text{nuclei}}}{2}$.

  • Across Periods (Left to Right):

  • Atomic radii generally decrease.

  • Cause: Protons are added to the nucleus while electrons enter the same principal energy level. The resulting increase in positive nuclear charge exerts a stronger electrostatic pull on the electron cloud, drawing valence electrons closer to the nucleus.

  • Down Groups (Top to Bottom):

  • Atomic radii generally increase.

  • Cause: Outer electrons occupy successively higher principal energy levels ($n = 1, 2, 3, \dots$). The additional core electron shells increase the distance from the nucleus and provide shielding, decreasing the net attraction felt by the valence electrons.

• Calculated Atomic Radii Values (in picometers, $\text{pm}$):

  • Period 1: $H$ ($53\,\text{pm}$), $He$ ($31\,\text{pm}$).

  • Period 2: $Li$ ($167\,\text{pm}$), $Be$ ($112\,\text{pm}$), $B$ ($87\,\text{pm}$), $C$ ($67\,\text{pm}$), $N$ ($56\,\text{pm}$), $O$ ($48\,\text{pm}$), $F$ ($42\,\text{pm}$), $Ne$ ($38\,\text{pm}$).

  • Period 3: $Na$ ($190\,\text{pm}$), $Mg$ ($145\,\text{pm}$), $Al$ ($118\,\text{pm}$), $Si$ ($111\,\text{pm}$), $P$ ($98\,\text{pm}$), $S$ ($88\,\text{pm}$), $Cl$ ($79\,\text{pm}$), $Ar$ ($71\,\text{pm}$).

  • Period 4: $K$ ($243\,\text{pm}$), $Ca$ ($194\,\text{pm}$), $Sc$ ($184\,\text{pm}$), $Ti$ ($176\,\text{pm}$), $V$ ($171\,\text{pm}$), $Cr$ ($166\,\text{pm}$), $Mn$ ($161\,\text{pm}$), $Fe$ ($156\,\text{pm}$), $Co$ ($152\,\text{pm}$), $Ni$ ($149\,\text{pm}$), $Cu$ ($145\,\text{pm}$), $Zn$ ($142\,\text{pm}$), $Ga$ ($136\,\text{pm}$), $Ge$ ($125\,\text{pm}$), $As$ ($114\,\text{pm}$), $Se$ ($103\,\text{pm}$), $Br$ ($94\,\text{pm}$), $Kr$ ($88\,\text{pm}$).

  • Period 5: $Rb$ ($265\,\text{pm}$), $Sr$ ($219\,\text{pm}$), $Y$ ($212\,\text{pm}$), $Zr$ ($206\,\text{pm}$), $Nb$ ($198\,\text{pm}$), $Mo$ ($190\,\text{pm}$), $Tc$ ($183\,\text{pm}$), $Ru$ ($178\,\text{pm}$), $Rh$ ($173\,\text{pm}$), $Pd$ ($169\,\text{pm}$), $Ag$ ($165\,\text{pm}$), $Cd$ ($161\,\text{pm}$), $In$ ($156\,\text{pm}$), $Sn$ ($145\,\text{pm}$), $Sb$ ($133\,\text{pm}$), $Te$ ($123\,\text{pm}$), $I$ ($115\,\text{pm}$), $Xe$ ($108\,\text{pm}$).

  • Period 6: $Cs$ ($298\,\text{pm}$), $Ba$ ($253\,\text{pm}$), $La$ ($217\,\text{pm}$), $Hf$ ($208\,\text{pm}$), $Ta$ ($200\,\text{pm}$), $W$ ($193\,\text{pm}$), $Re$ ($188\,\text{pm}$), $Os$ ($185\,\text{pm}$), $Ir$ ($180\,\text{pm}$), $Pt$ ($177\,\text{pm}$), $Au$ ($174\,\text{pm}$), $Hg$ ($171\,\text{pm}$), $Tl$ ($156\,\text{pm}$), $Pb$ ($154\,\text{pm}$), $Bi$ ($143\,\text{pm}$), $Po$ ($135\,\text{pm}$), $At$ ($127\,\text{pm}$), $Rn$ ($120\,\text{pm}$).

• Electronegativity ($\chi$):

  • Definition: A dimensionless chemical property measuring the relative attraction an atom exerts on shared electrons within a chemical bond.

  • Trend Relationships:

  • Increases from left to right across a period as effective nuclear charge rises.

  • Decreases from top to bottom down a group as increased distance and inner-shell shielding weaken nuclear hold on shared electrons.

  • Fluorine possesses the maximum value ($\chi = 3.98$), making it the most reactive and electronegative of all nonmetals.

  • Cesium and Francium possess the minimum values ($\chi = 0.79$ and $\chi = 0.7$).

  • Selected Pauling Electronegativity Values:

  • Group 1: $H$ ($2.20$), $Li$ ($0.98$), $Na$ ($0.93$), $K$ ($0.82$), $Rb$ ($0.82$), $Cs$ ($0.79$), $Fr$ ($0.7$).

  • Group 2: $Be$ ($1.57$), $Mg$ ($1.31$), $Ca$ ($1.00$), $Sr$ ($0.95$), $Ba$ ($0.89$), $Ra$ ($0.9$).

  • Transition Elements: $Sc$ ($1.36$), $Ti$ ($1.54$), $V$ ($1.63$), $Cr$ ($1.66$), $Mn$ ($1.55$), $Fe$ ($1.83$), $Co$ ($1.88$), $Ni$ ($1.91$), $Cu$ ($1.90$), $Zn$ ($1.65$), $Ag$ ($1.93$), $Pt$ ($2.2$), $Au$ ($2.4$).

  • Group 13: $B$ ($2.04$), $Al$ ($1.61$), $Ga$ ($1.81$), $In$ ($1.78$), $Tl$ ($1.8$).

  • Group 14: $C$ ($2.55$), $Si$ ($1.90$), $Ge$ ($2.01$), $Sn$ ($1.96$), $Pb$ ($1.8$).

  • Group 15: $N$ ($3.04$), $P$ ($2.19$), $As$ ($2.18$), $Sb$ ($2.05$), $Bi$ ($1.9$).

  • Group 16: $O$ ($3.44$), $S$ ($2.58$), $Se$ ($2.55$), $Te$ ($2.1$), $Po$ ($2.0$).

  • Group 17: $F$ ($3.98$), $Cl$ ($3.16$), $Br$ ($2.96$), $I$ ($2.66$), $At$ ($2.2$).

  • Group 18: Noble gases lack standard Pauling electronegativity values under normal bonding conditions.

Atomic Interactions and Fundamental Bond Types (Section 3.2)

• Mechanics of Atomic Collisions and Bonding Energetics:

  • Chemistry is a physical science occurring on an atom-to-atom basis; a physical interaction or collision is required for bonding.

  • During atomic approach, competing subatomic electrostatic forces interact:

  • Attractive forces: Positively charged protons in the nucleus of one atom attract the negatively charged electrons of the neighboring atom.

  • Repulsive forces: Protons repel protons, and electrons repel electrons.

  • Energetics of Bond Formation:

  • Chemical bonds stabilize atoms by reducing their potential energy.

  • When atoms are infinitely separated: Internuclear distance is large; zero overlap occurs; electrostatic attraction is negligible; potential energy is approximately zero.

  • When atoms reach the bond distance: Electron clouds overlap; attractive forces balance repulsive forces; potential energy reaches a minimum well, establishing a stable chemical bond.

  • When atoms are forced too close: Positively charged nuclei and dense core electron clouds strongly repel; potential energy spikes sharply.

• Classification of Chemical Bonds by Element Type:

  • Ionic Bond: Formed between a metal and a nonmetal.

  • Covalent Bond: Formed between two or more nonmetals.

  • Metallic Bond: Formed between metal atoms.

• Electronegativity Difference ($\Delta\chi$) and Bond Character:

  • The nature of a chemical bond is a continuum determined by $\Delta\chi = |\chi_A - \chi_B|$:

  • Covalent Character: Low $\Delta\chi$ values (typically $< 1.7$–$2.0$) result in shared valence electrons. Representative examples include diatomic halogens ($Cl_2$, $\Delta\chi = 0$), interhalogens ($IBr$, $HI$, $ICl$), and hydrogen halides ($HBr$, $HCl$).

  • Ionic Character: High $\Delta\chi$ values (typically $> 1.7$–$2.0$) result in complete valence electron transfer. Representative examples include alkali halides ($LiI$, $LiBr$, $LiCl$, $NaCl$, $KCl$, $KBr$, $CsCl$, $KF$, $LiF$).

  • 50% Ionic Character Cutoff: A recognized arbitrary cutoff dividing predominantly covalent from predominantly ionic character.

  • Notable Boundary Cases: Hydrogen fluoride ($HF$) exhibits a large difference ($\Delta\chi = 3.98 - 2.20 = 1.78$) but falls just under the $50\%$ ionic character threshold, behaving as an intensely polar covalent molecule.

• Ionic Bonds (Metal with Nonmetal):

  • Driven by large $\Delta\chi$.

  • Metal atoms lose valence electrons to form positively charged cations (e.g., $Na \rightarrow Na^+ + 1\,e^-$).

  • Nonmetal atoms gain valence electrons to form negatively charged anions (e.g., $Cl + 1\,e^- \rightarrow Cl^-$).

  • The bond is the electrostatic attraction binding the resulting cations and anions in a continuous crystal lattice.

• Covalent Bonds (Nonmetal with Nonmetal):

  • Driven by small $\Delta\chi$.

  • Atoms share pairs of electrons to achieve stable, full valence electron shells.

  • Atoms conform to the Octet Rule (acquiring eight valence electrons identical to noble gas configurations), except Hydrogen ($H$) and Helium ($He$), which achieve a stable valence duet of two electrons.

• Metallic Bonds (Metal with Metal):

  • Small $\Delta\chi$ among metal atoms prevents full ionic transfer, yet low electronegativity permits valence electrons to delocalize.

  • Sea of Electrons Model: The valence electrons do not belong to any specific atom pair; they float freely through a rigid lattice of positive metal cation cores.

  • Example: Silver lattice ($Ag^+$ surrounded by delocalized $e^-$); Potassium metal where each neutral atom ($[Ar]4s^1$) sheds its single outer $4s^1$ valence electron into the communal mobile electron pool.

  • Macroscopic Properties Explained by the Model:

  • Malleability and Ductility: Metal cation planes slide past one another under mechanical stress without shattering, as the mobile electron sea continuously buffers positive-positive repulsion.

  • Electrical Conductivity: Delocalized electrons drift freely through the lattice under an applied voltage.

  • Thermal Conductivity: Kinetic energy is rapidly dispersed through collisions of mobile valence electrons.

  • Hardness and Tensile Strength: Multi-directional electrostatic attraction binds the cations to the pervasive electron sea.

Bonding Behavior and Versatility of Hydrogen

• Unique Atomic Structure and Placement:

  • Electron configuration: $1s^1$ (one proton, one electron, zero neutrons in protium).

  • Group 1 Commonalities: Listed atop Group 1 because it possesses one valence electron.

  • Group 17 Commonalities: Only requires one additional electron to achieve a filled shell ($1s^2$ duet matching Helium), resembling halogens.

  • Electronegativity ($\chi = 2.20$): Intermediate between electropositive Group 1 alkali metals ($\chi \approx 0.7$–$1.0$) and electronegative Group 17 halogens ($\chi \approx 2.5$–$4.0$).

• Four Distinct Chemical Behaviors:

  1. Diatomic Nonpolar Covalent Bonding ($H_2$):

  • Two hydrogen atoms share their single electrons equally ($\Delta\chi = 0$), forming a nonpolar single covalent bond ($H-H$).

  1. Cation Formation ($H^+$) with Highly Electronegative Elements:

  • When bonded to high-electronegativity elements (e.g., $Cl$ in $HCl$), hydrogen acts like an alkali metal.

  • It releases its electron upon dissociation in aqueous solution, yielding a free hydrogen ion (proton, $H^+$) and a chloride anion ($Cl^-$):        $HCl_{(aq)} \rightarrow H^+_{(aq)} + Cl^-_{(aq)}$

  1. Anion Formation ($H^-$) with Low Electronegativity Elements:

  • When bonded to strongly electropositive metals (e.g., Sodium, $Na$), hydrogen acts like a halogen.

  • It pulls an electron away from the metal, forming an ionic compound containing the hydride ion ($H^-$):        $Na + H \rightarrow Na^+ H^-\quad (\text{sodium hydride})$

  1. Polar/Nonpolar Covalent Molecular Bonding with Nonmetals:

  • When bonded with standard nonmetals (e.g., Carbon in methane, $CH_4$), the electronegativity difference is small ($\Delta\chi = |2.55 - 2.20| = 0.35$).

  • Electrons are shared covalently without forming discrete ions.

Properties and Classification of Ionic and Covalent Compounds (Section 3.3)

• Electrical Conductivity and Electrolytes:

  • Both ionic compounds (formula units) and covalent compounds (molecules) are macroscopic, electrically neutral species.

  • Ionic Behavior in Solution:

  • Dissolve by dissociating into mobile, hydrated cations and anions.

  • Function as strong or weak electrolytes, completing electrical circuits to illuminate testing lightbulbs brightly (e.g., aqueous $NaCl_{(aq)}$).

  • Covalent Behavior in Solution:

  • Remain intact as neutral, uncharged molecules (e.g., aqueous sucrose, $C_{12}H_{22}O_{11(aq)}$).

  • Function as nonelectrolytes, incapable of conducting current (testing lightbulb remains unlit).

  • Select polar organic substances (e.g., ethanol, $CH_3CH_2OH_{(aq)}$) can act as very weak electrolytes or nonelectrolytes depending on ionization extent.

• Comparative Properties: Ionic vs. Covalent Compounds:

  • Constituent Elements:

  • Ionic: Metal bonded to nonmetal.

  • Covalent: Nonmetals bonded exclusively to nonmetals.

  • Bonding Mechanism:

  • Ionic: Electrostatic attraction between positive cations and negative anions.

  • Covalent: Sharing of electron pairs between localized atomic nuclei.

  • Representative Unit:

  • Ionic: Formula unit (simplest empirical ratio in an extended lattice).

  • Covalent: Molecule (discrete, individual molecular entity).

  • Formula Significance:

  • Ionic: Simplest whole-number ratio of cations to anions.

  • Covalent: Actual number and identity of bonded atoms.

  • Physical State at Room Temperature ($25^\circ\text{C}$):

  • Ionic: Crystalline solid.

  • Covalent: Gas, liquid, or low-melting solid.

  • Water Solubility:

  • Ionic: Usually high.

  • Covalent: Variable (governed by molecular polarity).

  • Melting and Boiling Temperatures:

  • Ionic: Generally high (requires breaking strong ionic lattice bonds).

  • Covalent: Generally low (governed by weak intermolecular forces).

  • Electrical Conductivity:

  • Ionic: High when molten (liquid) or dissolved in aqueous solution; nonconductive as solid.

  • Covalent: Poor/nonconductive in all states.

  • Aqueous Dissolution State:

  • Ionic: Dissociates fully or partially into discrete mobile ions.

  • Covalent: Remains intact as discrete neutral molecules.

• Practical Applications of Everyday Ionic Compounds:

  • De-icing / Road Salt: Magnesium chloride ($MgCl_2$), lowers freezing point of water.

  • Dental Fluoride Treatment: Sodium fluoride ($NaF$), provides fluoride ions to strengthen tooth enamel.

  • Radiological Thyroid Protection: Potassium iodide ($KI$), saturates the thyroid with stable iodine to block radioactive iodine absorption during emergencies.

Chemical Nomenclature and Formula Writing

• Main-Group Ion Charge Predictability:

  • Group 1 (Alkali Metals): Form $+1$ cations ($Li^+$, $Na^+$, $K^+$).

  • Group 2 (Alkaline Earth Metals): Form $+2$ cations ($Mg^{2+}$, $Ca^{2+}$, $Ba^{2+}$).

  • Group 13: Form $+3$ cations ($Al^{3+}$, $Ga^{3+}$).

  • Group 14: Form $\pm 4$ ions (rarely form simple monatomic ions; commonly share electrons).

  • Group 15 (Pnictogens): Form $-3$ anions ($N^{3-}$, $P^{3-}$).

  • Group 16 (Chalcogens): Form $-2$ anions ($O^{2-}$, $S^{2-}$, $Se^{2-}$).

  • Group 17 (Halogens): Form $-1$ anions ($F^-$, $Cl^-$, $Br^-$, $I^-$).

  • Group 18 (Noble Gases): Charge is $0$; chemically unreactive with full valence octets.

• Transition Metal Complexity (Groups 3–12 and Post-Transition Metals):

  • Monatomic transition metal cations exhibit variable, unpredictable oxidation states because they can lose varying numbers of outer $s$ and inner $d$ electrons.

  • Cation charges are distinguished by Roman numerals in parentheses:

  • Titanium: $Ti^{3+}$

  • Vanadium: $V^{2+}$, $V^{3+}$, $V^{4+}$, $V^{5+}$

  • Chromium: $Cr^{2+}$, $Cr^{3+}$, $Cr^{6+}$

  • Manganese: $Mn^{2+}$, $Mn^{7+}$

  • Iron: $Fe^{2+}$ (Iron (II)), $Fe^{3+}$ (Iron (III))

  • Cobalt: $Co^{2+}$

  • Nickel: $Ni^{2+}$

  • Copper: $Cu^{2+}$

  • These variable oxidation states often yield distinct, vivid colors in aqueous solutions.

• Polyatomic Ions:

  • Groups of atoms covalently bonded together carrying a net electrical charge; participate in ionic bonding as modular units.

  • Key Polyatomic Ions:

  • Ammonium: $NH_4^+$ (charge: $+1$)

  • Hydroxide: $OH^-$ (charge: $-1$)

  • Nitrate: $NO_3^-$ (charge: $-1$)

  • Phosphate: $PO_4^{3-}$ (charge: $-3$)

  • Sulfate: $SO_4^{2-}$ (charge: $-2$)

  • Carbonate: $CO_3^{2-}$ (charge: $-2$)

• Systematic Ionic Nomenclature:

  • Formula to Name:

  1. Name the cation first; name the anion second.

  2. For monatomic cations, use the unaltered element name ($Mg^{2+} = \text{magnesium}$, $Al^{3+} = \text{aluminum}$).

  3. For transition metal cations with variable charges, append Roman numerals in parentheses indicating the positive oxidation state ($FeO \rightarrow \text{iron (II) oxide}$; $Fe_2O_3 \rightarrow \text{iron (III) oxide}$).

  4. For monatomic anions, take the root of the nonmetal name and append the suffix -ide ($O^{2-} = \text{oxide}$, $S^{2-} = \text{sulfide}$).

  5. For polyatomic ions, state their specific standard chemical name without alteration ($(NH_4)_2O \rightarrow \text{ammonium oxide}$; $Al_2(CO_3)_3 \rightarrow \text{aluminum carbonate}$).

  • Name to Formula:

  1. Write the elemental or polyatomic symbols for the cation and anion along with their ionic charges.

  2. Determine the lowest common multiple of the charges to establish how many of each ion are needed to achieve net zero charge.

  3. Assign these quantities as subscripts in the chemical formula.

  4. If more than one polyatomic ion is required to balance the charge, enclose the entire polyatomic formula in parentheses before adding the subscript:

     • Magnesium oxide: $Mg^{2+} + O^{2-} \rightarrow 1\,Mg^{2+} + 1\,O^{2-} \rightarrow MgO$

     • Aluminum sulfide: $Al^{3+} + S^{2-} \rightarrow 2\,Al^{3+} + 3\,S^{2-} \rightarrow Al_2S_3$

     • Ammonium oxide: $NH_4^+ + O^{2-} \rightarrow 2\,NH_4^+ + 1\,O^{2-} \rightarrow (NH_4)_2O$

     • Aluminum carbonate: $Al^{3+} + CO_3^{2-} \rightarrow 2\,Al^{3+} + 3\,CO_3^{2-} \rightarrow Al_2(CO_3)_3$

• Covalent Bonding Rules by Nonmetal Group:

  • Group 18: Possess 8 valence electrons; chemically inert; form 0 bonds.

  • Group 17 (and Hydrogen): Possess 7 valence electrons (H has 1); require 1 electron; form 1 single bond.

  • Group 16: Possess 6 valence electrons; require 2 electrons; form 2 covalent bonds total (either 2 single bonds or 1 double bond).

  • Group 15: Possess 5 valence electrons; require 3 electrons; form 3 covalent bonds total (3 single bonds, or 1 single and 1 double bond, or 1 triple bond).

  • Group 14: Possess 4 valence electrons; require 4 electrons; form 4 covalent bonds total (4 single bonds, or 2 single and 1 double bond, or 2 double bonds, or 1 single and 1 triple bond).

• Greek Numerical Prefixes for Covalent Compounds:

  • $1 = \text{mono-}$

  • $2 = \text{di-}$

  • $3 = \text{tri-}$

  • $4 = \text{tetra-}$

  • $5 = \text{penta-}$

  • $6 = \text{hexa-}$

  • $7 = \text{hepta-}$

  • $8 = \text{octa-}$

  • $9 = \text{nona-}$

  • $10 = \text{deca-}$

• Systematic Binary Covalent Nomenclature:

  • Formula to Name:

  1. Name the first element using its full, unmodified name.

  2. Name the second element using the root of its name plus the suffix -ide.

  3. Apply numerical Greek prefixes to denote the number of atoms of each element present, matching the subscripts.

  4. Prefix Rules:

     • Never use the prefix "mono-" on the first element.

     • If adding a prefix creates two consecutive vowels (e.g., "tetraoxide" or "pentaoxide"), drop the trailing vowel of the prefix to improve pronunciation (e.g., "tetroxide", "pentoxide").

  • Examples:

    • $CO_2 \rightarrow \text{carbon dioxide}$

    • $P_2O_5 \rightarrow \text{diphosphorus pentoxide}$

    • $NF_3 \rightarrow \text{nitrogen trifluoride}$

  • Name to Formula:

  1. Write the symbol of the first element, adding a subscript corresponding to its prefix (if no prefix, 1 atom is implied; write no subscript).

  2. Write the symbol of the second element, adding a subscript corresponding to its prefix.

  • Examples:

    • Dihydrogen monoxide: $H_2O$

    • Disulfur decafluoride: $S_2F_{10}$

    • Nitrogen trihydride: $NH_3$

    • Sulfur dioxide: $SO_2$

    • Carbon tetrachloride: $CCl_4$

Advanced Compound Complexity and Octet Exceptions

• Exceptions to the Octet Rule (Expanded Octets):

  • Nonmetals located in Period 3 or higher on the periodic table (e.g., Phosphorus, Sulfur, Chlorine) have valence electrons in the third principal energy level ($n = 3$) or above.

  • These elements possess accessible, empty $d$-orbitals ($3d$) that can accommodate additional electron pairs beyond the typical 8-electron limit when acting as central atoms.

  • Example: Sulfur hexafluoride ($SF_6$):

  • Sulfur forms 6 single bonds to 6 surrounding Fluorine atoms.

  • Sulfur holds a total of 12 shared valence electrons, forming a stable expanded octet.

• Polarity of Covalent Bonds:

  • Nonpolar Covalent Bonds:

  • Formed between identical nonmetal atoms ($\Delta\chi = 0$) with completely symmetric electron sharing.

  • Includes all diatomic elements: Hydrogen ($H_2$), Nitrogen ($N_2$), Oxygen ($O_2$), Fluorine ($F_2$), Chlorine ($Cl_2$), Bromine ($Br_2$), Iodine ($I_2$).

  • Bonds between Carbon and Hydrogen ($C-H$) are also classified as nonpolar due to their very small electronegativity difference ($\Delta\chi = 0.35$).

• Polar Covalent Bonds:

  • Formed between different nonmetal elements (excluding $C-H$) with significant electronegativity differences ($0.4 \le \Delta\chi \le 1.9$).

  • Electron density is shifted toward the more electronegative atom.

  • Establishes a permanent bond dipole with partial charges:

  • $\delta^-$ (partial negative): Resides on the more electronegative atom.

  • $\delta^+$ (partial positive): Resides on the less electronegative atom.

  • Example: Water ($H_2O$):

  • Oxygen ($\chi = 3.44$) pulls electron density away from both Hydrogen atoms ($\chi = 2.20$).

  • Oxygen acquires a $\delta^-$ charge while each Hydrogen acquires a $\delta^+$ charge.

Lewis Structures and Electron-Dot Modeling (Section 3.4)

• Lewis Structures for Single Atoms (Electron-Dot Diagrams):

  • Electron-dot diagrams show the valence electrons of individual neutral atoms.

  • Procedure:

  1. Write the atomic element symbol.

  2. Determine total valence electrons from group number.

  3. Place dots individually on each of the four sides (top, right, bottom, left) before pairing any dots.

  • Valence Electron Counts across Periods and Groups:

  • Group 1 ($H, Li, Na, K$): 1 valence electron ($1$ unpaired dot).

  • Group 2 ($Be, Mg, Ca$): 2 valence electrons ($2$ single dots on different sides).

  • Group 13 ($B, Al, Ga$): 3 valence electrons ($3$ single dots).

  • Group 14 ($C, Si, Ge$): 4 valence electrons ($4$ single dots).

  • Group 15 ($N, P, As$): 5 valence electrons ($1$ pair and $3$ single dots).

  • Group 16 ($O, S, Se$): 6 valence electrons ($2$ pairs and $2$ single dots).

  • Group 17 ($F, Cl, Br$): 7 valence electrons ($3$ pairs and $1$ single dot).

  • Group 18 ($Ne, Ar, Kr$): 8 valence electrons ($4$ complete pairs; Helium is a duet with $2$ dots).

• Lewis Structures for Ionic Compounds (Formula Units):

  • Step-by-Step Procedure:

  1. Draw individual electron-dot structures for all participating neutral atoms.

  2. Use curved arrows to demonstrate electron transfer from the metal atom(s) to the nonmetal atom(s).

  3. Draw the final Lewis structure representing newly formed ions:

     • Enclose each ion within square brackets [...].

     • Indicate the net ionic charge outside the upper right bracket as a superscript.

     • Metal Cations: Show zero dots inside the brackets with a positive charge (denoting loss of the outer valence shell).

     • Nonmetal Anions: Show a full octet of eight dots inside the brackets with a negative charge.

  • Worked Examples:

  • Sodium Sulfide:       $2\,Na\cdot + \cdot\ddot{S}\cdot \longrightarrow [Na]^+ \; [:\ddot{S}:]^{2-} \; [Na]^+ \quad (\text{or } 2\,[Na]^+ [:\ddot{S}:]^{2-})$

  • Calcium Sulfide ($CaS$):       $Ca: + \cdot\ddot{S}\cdot \longrightarrow [Ca]^{2+} [:\ddot{S}:]^{2-}$

  • Magnesium Fluoride ($MgF_2$):       $: \ddot{F} \cdot + \cdot Mg \cdot + \cdot \ddot{F} : \longrightarrow [Mg]^{2+} \; 2\,[:\ddot{F}:]^-$

  • Aluminum Sulfide ($Al_2S_3$):       $2\,Al (3\,e^- \text{ each}) + 3\,S (6\,e^- \text{ each}) \longrightarrow 2\,[Al]^{3+} \; 3\,[:\ddot{S}:]^{2-}$

  • Barium Nitride ($Ba_3N_2$):       $3\,Ba (2\,e^- \text{ each}) + 2\,N (5\,e^- \text{ each}) \longrightarrow 3\,[Ba]^{2+} \; 2\,[:\ddot{N}:]^{3-}$

• Lewis Structures for Molecular Covalent Compounds:

  • Step-by-Step Procedure:

  1. Calculate Total Valence Electrons: Sum valence electrons from all participating atoms. The final Lewis structure must contain this exact number of electrons.

  2. Construct the Molecular Skeleton:

     • Select the central atom: The element with the lowest electronegativity is placed in the center.

     • Rule: Hydrogen can never be the central atom because it only forms one bond.

     • Connect outer atoms to the central atom using single lines. Each single line represents a single covalent bond ($2\,e^-$).

  3. Subtract Bonding Electrons: For each single bond drawn, subtract $2\,e^-$ from the total valence count calculated in Step 1.

  4. Distribute Remaining Electrons as Lone Pairs:

     • Distribute remaining electrons as lone pair dots to complete octets ($8\,e^-$) on all outer terminal atoms first (duet of $2\,e^-$ for Hydrogen).

     • Place any remaining electrons on the central atom last.

  5. Verify Octets and Electron Counts:

     • Recount all shared and unshared electrons to verify that the total matches Step 1.

     • Check that every atom satisfies the Octet Rule (or duet for H).

  6. Form Multiple Bonds if Needed:

     • If the central atom lacks an octet after distributing all electrons, convert one or more lone pairs from outer terminal atoms into double or triple bonds to the central atom.

     • Bond Limitations: Single, double, and triple bonds are physically possible. Quadruple bonds do not form.

  7. Repeat verification until every atom satisfies bonding rules.

• Worked Molecular Lewis Structure Examples:

  • Nitrosyl Chloride ($NOCl$):

  • Total valence electrons: $O (6) + N (5) + Cl (7) = 18\,e^-$.

  • Central atom: Nitrogen (less electronegative than Oxygen and Chlorine).

  • Initial skeleton: Connect $O-N-Cl$ (uses $4\,e^-$, leaving $14\,e^-$).

  • Distribute lone pairs to outer atoms ($6\,e^-$ on Oxygen, $6\,e^-$ on Chlorine) and remaining $2\,e^-$ on Nitrogen.

  • Nitrogen only has $6\,e^-$; convert one lone pair from Oxygen to form an $O=N$ double bond.

  • Final structure: $\ddot{O} = \ddot{N} - \ddot{Cl}:$ (Oxygen has 2 lone pairs, Nitrogen has 1 lone pair, Chlorine has 3 lone pairs; total $18\,e^-$).

  • Nitrogen Trichloride ($NCl_3$):

  • Total valence electrons: $N (5) + 3 \times Cl (7) = 26\,e^-$.

  • Central atom: Nitrogen; skeleton connects three Chlorine atoms ($3 \times 2 = 6\,e^-$ used, $20\,e^-$ remaining).

  • Distribute 3 lone pairs ($6\,e^-$) to each of the three Chlorine atoms ($18\,e^-$ used).

  • Place the final 2 electrons on Nitrogen as a lone pair.

  • Final structure: A central Nitrogen atom with 1 lone pair single-bonded to three Chlorine atoms, each carrying 3 lone pairs (total $26\,e^-$).

  • Carbon Monoxide ($CO$):

  • Total valence electrons: $C (4) + O (6) = 10\,e^-$.

  • Single bond $C-O$ leaves $8\,e^-$.

  • Distributing as lone pairs leaves both atoms deficient.

  • Sharing three electron pairs forms a triple bond ($C \equiv O$), leaving each atom with one lone pair.

  • Final structure: $:\!C \equiv O\!:$ (total $10\,e^-$, both satisfy octet rule).

  • Carbon Dioxide ($CO_2$):

  • Total valence electrons: $C (4) + 2 \times O (6) = 16\,e^-$.

  • Central atom: Carbon; form double bonds to both Oxygen atoms.

  • Final structure: $\ddot{O} = C = \ddot{O}$ (Carbon has 4 shared bonds and 0 lone pairs; each Oxygen has 2 double bonds and 2 lone pairs; total $16\,e^-$).