Comprehensive Chemical Bonding and Inorganic Nomenclature Study Guide

Administrative Details and Course Schedule

• Homework Deadlines:

  • Aktiv HW: Chapter 4 Part 1 (Intro & Naming) – Due Tuesday, 09/29 at 11:59 PM.
  • Aktiv HW: Chapter 4 Part 2 (Lewis Structure) – Due Saturday, 10/03 at 11:59 PM.

• In-Class Quiz #2:

  • Administered during the last 10 minutes of class on Wednesday, 09/30 or Thursday, 10/01.
  • Accommodation setup requires students to stay until everyone has completed the exam.

• Aktiv Quiz #3 (Chapter 4):

  • Covers Chapter 4 contents up to Thursday, 10/01 or Friday, 10/02.
  • Window of Availability: Opens Saturday, 10/03 at 12:00 AM; Closes Sunday, 10/04 at 11:59 PM.
  • Format & Rules: Closed-book format. Students are permitted to print and use the official reference sheet. A strict time limit is enforced. Proctoring protocols are active: receiving two warning emails will result in an immediate academic dishonesty report.

• Extra Credit Opportunity:

  • Chemical Compound Naming Exercise package is accessible via the Canvas Module.
  • Completion Deadline: Wednesday, Sep 30th.
  • Submission Procedure: Must be submitted in person during Student Instructor (SI) or Learning Assistant (LA) office hours.
  • Grading Methodology: A single random page will be selected from the submitted package for grading, followed by a live Naming Quiz.
  • Requirement: Answering 3 out of 3 questions correctly on the naming quiz grants 1 full point added directly to the upcoming quiz score.

Chemical Fundamentals: Compounds vs. Molecules and Types of Chemical Formulas

• Definitions:

  • Compound: A substance composed of two or more different kinds of elements chemically bonded together in fixed proportions.
  • Molecule: An electrically neutral group of two or more atoms bonded together.
  • Fundamental Relationship: All compounds are molecules, but not all molecules are compounds (e.g., $\text{O}_2$ is a molecular element, not a compound).

• Chemical Formulas:

  • Express the composition of a substance using elemental chemical symbols.
  • Subscripts positioned to the right of each element's symbol specify the exact number of atoms of that element in the formula unit or molecule.

• Three Major Types of Chemical Formulas:

  • Empirical Formula: Gives the simplest, lowest whole-number ratio of atoms of each element present in a chemical compound.
  • Molecular Formula: Gives the actual, exact number of atoms of each element present in a single molecule of a compound.
  • Structural Formula: A two-dimensional representation showing how atoms in a molecule are covalently bonded to one another, using solid lines to represent single, double, or triple covalent bonds.

• Comparative Examples of Chemical Formulas:

  • Water: Molecular Formula = $\text{H}_2\text{O}$, Empirical Formula = $\text{H}_2\text{O}$.
  • Glucose: Molecular Formula = $\text{C}_6\text{H}_{12}\text{O}_6$, Empirical Formula = $\text{CH}_2\text{O}$.
  • Ozone: Molecular Formula = $\text{O}_3$, Empirical Formula = $\text{O}$.
  • Ethane: Structural Formula shows two central carbon atoms bonded together, each attached to three hydrogen atoms ($\text{H}_3\text{C}-\text{CH}_3$). Molecular Formula = $\text{C}_2\text{H}_6$, Empirical Formula = $\text{CH}_3$. Higher-order compounds sharing the same empirical ratio include butane ($\text{C}_4\text{H}_{12}$ theoretical structure or related hydrocarbon ratios).

Molecular Models and Visual Representations

• Ball-and-Stick Models:

  • Represent atoms as solid spheres and covalent bonds as rigid rods or sticks.
  • Ideal for visualizing 3D molecular geometry, arrangement, and determining exact bond angles.

• Space-Filling Models:

  • Represent atoms as overlapping spheres scaled to their van der Waals radii.
  • Provide the most accurate representation of the actual space occupied by electrons and the realistic physical contour of the molecule.

• Standard Element Color-Coding System (CPK Colors):

  • Hydrogen ($\text{H}$): White sphere
  • Carbon ($\text{C}$): Black / Dark Gray sphere
  • Nitrogen ($\text{N}$): Blue sphere
  • Oxygen ($\text{O}$): Red sphere
  • Fluorine ($\text{F}$): Light Green sphere
  • Phosphorus ($\text{P}$): Blue-Violet / Purple sphere
  • Sulfur ($\text{S}$): Yellow sphere
  • Chlorine ($\text{Cl}$): Green sphere

• Comparison Table of Structural Formulas and Models:
  • Benzene: Empirical Formula = $\text{CH}$, Molecular Formula = $\text{C}_6\text{H}_6$. Structural formula displays a six-carbon ring with alternating single and double bonds, each carbon bound to one hydrogen atom.
  • Acetylene: Empirical Formula = $\text{CH}$, Molecular Formula = $\text{C}_2\text{H}_2$. Structural formula displays a linear carbon-carbon triple bond (\text{H}-\text{C}\t\text{E}\text{C}-\text{H}).
  • Glucose: Empirical Formula = $\text{CH}_2\text{O}$, Molecular Formula = $\text{C}_6\text{H}_{12}\text{O}_6$. Structural formula displays an open-chain aldehyde sugar with hydroxyl groups.
  • Ammonia: Empirical Formula = $\text{NH}_3$, Molecular Formula = $\text{NH}_3$. Structural formula displays central nitrogen with three hydrogen atoms in a trigonal pyramidal arrangement.

Types of Chemical Bonding

• Nature of Chemical Bonds:

  • Chemical bonds represent distinct methods by which atoms interact and distribute valence electrons to achieve stable electronic configurations.

• Summary Table of Chemical Bond Types:

  • Ionic Bonding:
  • Atomic Combination: Metal bound to Nonmetal (or containing polyatomic ions).
  • Valence Electron Behavior: Complete transfer of valence electrons from the metal atom to the nonmetal atom.
  • Resulting Structure: Metal becomes a positively charged cation; nonmetal becomes a negatively charged anion. Opposites attract to form an extended 3D crystalline lattice.
  • Covalent Bonding:
  • Atomic Combination: Nonmetal bound to Nonmetal.
  • Valence Electron Behavior: Valence electrons are shared mutually between nonmetal atoms.
  • Resulting Structure: Distinct, discrete units termed molecular compounds.
  • Metallic Bonding:
  • Atomic Combination: Metal bound to Metal.
  • Valence Electron Behavior: Valence electrons are pooled into a delocalized "sea of electrons" ($\text{e}^-$ sea) surrounding fixed metal cation cores.

Classification of Pure Substances

• Pure substances are categorized into two primary divisions: Elements and Compounds.

• Elements:

  • Atomic Elements: Exist naturally with single individual atoms as their basic structural unit (e.g., Neon, $\text{Ne}$).
  • Molecular Elements: Do not normally exist as single isolated atoms, but as molecules containing two or more identical atoms bound together (e.g., Oxygen gas, $\text{O}_2$).

• Compounds:

  • Molecular Compounds: Composed of two or more different nonmetal elements covalently bonded into discrete molecules (e.g., Water, $\text{H}_2\text{O}$).
  • Ionic Compounds: Composed of cations and anions bound by electrostatic forces in an extended array defined by the simplest formula unit (e.g., Sodium chloride, $\text{NaCl}$ formula unit).

Stoichiometric Relationships Within Chemical Formulas

• Subscripts within a chemical formula represent exact mole ratios between the entire compound and its constituent atoms.

• Quantitative Example Calculation:

  • Problem: Determine the number of moles of oxygen atoms and the total number of oxygen atoms present in a $1.25\,\text{mol}$ sample of glucose ($\text{C}_6\text{H}_{12}\text{O}_6$).
  • Calculation 1 (Moles of Oxygen):     $\text{Moles of O} = 1.25\,\text{mol C}_6\text{H}_{12}\text{O}_6 \times \frac{6\,\text{mol O}}{1\,\text{mol C}_6\text{H}_{12}\text{O}_6} = 7.50\,\text{mol O}$
  • Calculation 2 (Total Oxygen Atoms):     $\text{Oxygen Atoms} = 7.50\,\text{mol O} \times 6.022 \times 10^{23}\,\text{atoms/mol} = 4.52 \times 10^{24}\,\text{oxygen atoms}$

Naming Ionic Compounds (Inorganic Nomenclature)

• General Principles of Ionic Formulas:
  • Ionic compounds are electrically neutral overall.
  • Criss-Cross Rule: The charge magnitude on the cation becomes the subscript on the anion, and the charge magnitude on the anion becomes the subscript on the cation. Subscripts must be reduced to the lowest whole-number integer ratio.
  • Example: Formation of Magnesium Nitride from $\text{Mg}^{2+}$ and $\text{N}^{3-}$.     $\text{Cation: Mg}^{2+}, \text{ Anion: N}^{3-} \rightarrow \text{Mg}_3\text{N}_2$

• Type I Ionic Compounds (Invariant Cations):

  • Metals that possess the exact same electrical charge in every chemical compound they form.
  • Includes Group 1A metals ($+1$), Group 2A metals ($+2$), Aluminum ($\text{Al}^{3+}$), and three transition elements with invariant charges: Zinc ($\text{Zn}^{2+}$), Scandium ($\text{Sc}^{3+}$), and Silver ($\text{Ag}^+$).
  • Naming Pattern: [Name of Cation Metal] + [Base Name of Anion Nonmetal + -ide]

• Table 5.2: Invariant Metals:

  • Lithium ($\text{Li}^+$) — Group 1A
  • Sodium ($\text{Na}^+$) — Group 1A
  • Potassium ($\text{K}^+$) — Group 1A
  • Rubidium ($\text{Rb}^+$) — Group 1A
  • Cesium ($\text{Cs}^+$) — Group 1A
  • Beryllium ($\text{Be}^{2+}$) — Group 2A
  • Magnesium ($\text{Mg}^{2+}$) — Group 2A
  • Calcium ($\text{Ca}^{2+}$) — Group 2A
  • Strontium ($\text{Sr}^{2+}$) — Group 2A
  • Barium ($\text{Ba}^{2+}$) — Group 2A
  • Aluminum ($\text{Al}^{3+}$) — Group 3A
  • Zinc ($\text{Zn}^{2+}$) — Transition Metal (Invariant)
  • Scandium ($\text{Sc}^{3+}$) — Transition Metal (Invariant)
  • Silver ($\text{Ag}^+$) — Transition Metal (Invariant; rare exceptions exist)

• Table 4.2: Common Monoatomic Anions:
  • Fluorine ($\text{F}^-$) — Base: Fluor — Anion Name: Fluoride
  • Chlorine ($\text{Cl}^-$) — Base: Chlor — Anion Name: Chloride
  • Bromine ($\text{Br}^-$) — Base: Brom — Anion Name: Bromide
  • Iodine ($\text{I}^-$) — Base: Iod — Anion Name: Iodide
  • Oxygen ($\text{O}^{2-}$) — Base: Ox — Anion Name: Oxide
  • Sulfur ($\text{S}^{2-}$) — Base: Sulf — Anion Name: Sulfide
  • Nitrogen ($\text{N}^{3-}$) — Base: Nitr — Anion Name: Nitride
  • Phosphorus ($\text{P}^{3-}$) — Base: Phosph — Anion Name: Phosphide

• Type II Ionic Compounds (Multivalent Cations):

  • Metals that form cations with variable charges depending on the compound.
  • Includes transition metals, inner transition metals, and select p-block metals (e.g., Lead, $\text{Pb}$; Tin, $\text{Sn}$; Thallium, $\text{Tl}$).
  • Naming Pattern: [Name of Cation Metal] + (Charge in Roman Numerals) + [Base Name of Anion Nonmetal + -ide]

• Table 3.3: Multivalent Cations and Classical System:

  • Chromium ($\text{Cr}$):
  • $\text{Cr}^{2+}$: Chromium(II) / Chromous
  • $\text{Cr}^{3+}$: Chromium(III) / Chromic
  • Iron ($\text{Fe}$):
  • $\text{Fe}^{2+}$: Iron(II) / Ferrous
  • $\text{Fe}^{3+}$: Iron(III) / Ferric
  • Cobalt ($\text{Co}$):
  • $\text{Co}^{2+}$: Cobalt(II) / Cobaltous
  • $\text{Co}^{3+}$: Cobalt(III) / Cobaltic
  • Copper ($\text{Cu}$):
  • $\text{Cu}^+$: Copper(I) / Cuprous
  • $\text{Cu}^{2+}$: Copper(II) / Cupric
  • Tin ($\text{Sn}$):
  • $\text{Sn}^{2+}$: Tin(II) / Stannous
  • $\text{Sn}^{4+}$: Tin(IV) / Stannic
  • Mercury ($\text{Hg}$):
  • $\text{Hg}_2^{2+}$: Mercury(I) / Mercurous
  • $\text{Hg}^{2+}$: Mercury(II) / Mercuric
  • Lead ($\text{Pb}$):
  • $\text{Pb}^{2+}$: Lead(II) / Plumbous
  • $\text{Pb}^{4+}$: Lead(IV) / Plumbic
  • Classical Rules Note: The suffix -ous indicates the ion with the lesser positive charge; -ic indicates the ion with the greater positive charge.

Polyatomic Ions and Oxyanions

• Definition: A polyatomic ion is a group of covalently bonded atoms that carries an overall net electrical charge. Ionic compounds containing polyatomic ions display both covalent bonding (within the polyatomic unit) and ionic bonding (electrostatic attractions throughout the crystal lattice).

• Comprehensive Table of Polyatomic Ions:

  • $\text{NH}_4^+$: Ammonium
  • $\text{C}_2\text{H}_3\text{O}_2^-$: Acetate
  • $\text{CN}^-$: Cyanide
  • $\text{SCN}^-$: Thiocyanate
  • $\text{HCO}_3^-$: Bicarbonate (or Hydrogen carbonate)
  • $\text{CO}_3^{2-}$: Carbonate
  • $\text{ClO}^-$: Hypochlorite
  • $\text{ClO}_2^-$: Chlorite
  • $\text{ClO}_3^-$: Chlorate
  • $\text{ClO}_4^-$: Perchlorate
  • $\text{CrO}_4^{2-}$: Chromate
  • $\text{Cr}_2\text{O}_7^{2-}$: Dichromate
  • $\text{MnO}_4^-$: Permanganate
  • $\text{NO}_2^-$: Nitrite
  • $\text{NO}_3^-$: Nitrate
  • $\text{OH}^-$: Hydroxide
  • $\text{O}_2^{2-}$: Peroxide
  • $\text{H}_2\text{PO}_4^-$: Dihydrogen phosphate
  • $\text{HPO}_4^{2-}$: Hydrogen phosphate
  • $\text{PO}_4^{3-}$: Phosphate
  • $\text{HSO}_3^-$: Hydrogen sulfite
  • $\text{SO}_3^{2-}$: Sulfite
  • $\text{HSO}_4^-$: Hydrogen sulfate
  • $\text{SO}_4^{2-}$: Sulfate

• Oxyanion Nomenclature System:

  • Oxyanions are polyatomic anions containing oxygen bonded to a central element.
  • Two-member Series:
  • Member with more oxygen atoms ends in -ate (e.g., $\text{NO}_3^-$ Nitrate; $\text{SO}_4^{2-}$ Sulfate).
  • Member with fewer oxygen atoms ends in -ite (e.g., $\text{NO}_2^-$ Nitrite; $\text{SO}_3^{2-}$ Sulfite).
  • Four-member Series (Halogen Oxyanions):
  • Most oxygen atoms (+1 O relative to -ate): Prefix per- + Base + Suffix -ate (e.g., $\text{ClO}_4^-$ Perchlorate).
  • Standard base oxygen count: Base + Suffix -ate (e.g., $\text{ClO}_3^-$ Chlorate).
  • One less oxygen atom (-1 O relative to -ate): Base + Suffix -ite (e.g., $\text{ClO}_2^-$ Chlorite).
  • Least oxygen atoms (-2 O relative to -ate): Prefix hypo- + Base + Suffix -ite (e.g., $\text{ClO}^-$ Hypochlorite).

Hydrated Ionic Compounds

• Definition: Hydrates are ionic compounds that possess a specific, fixed number of water molecules ($\text{H}_2\text{O}$) bound within their solid crystalline structure per formula unit.

• Thermal Removal: Heating drives off the waters of hydration, transforming the hydrated solid into an anhydrous compound.

• Greek Numerical Prefixes Table:

  • $1/2$: Hemi-
  • $1$: Mono-
  • $2$: Di-
  • $3$: Tri-
  • $4$: Tetra-
  • $5$: Penta-
  • $6$: Hexa-
  • $7$: Hepta-
  • $8$: Octa-
  • $9$: Nona-
  • $10$: Deca-

• Naming Rules & Examples:
  • Formula Pattern: [Name of Ionic Compound] + [Greek Prefix]-hydrate
  • Magnesium sulfate heptahydrate (Epsom Salt): $\text{MgSO}_4 \cdot 7\text{H}_2\text{O}$
  • Cobalt(II) chloride hexahydrate: $\text{CoCl}_2 \cdot 6\text{H}_2\text{O}$ (Pink magenta powder)
  • Cobalt(II) chloride anhydrate: $\text{CoCl}_2$ (Purple-blue powder)

Binary Molecular Compounds

• Composed of two distinct nonmetal elements covalently bonded together.

• Naming Template:   [Prefix] + [Name of 1st Element] + [Prefix] + [Base Name of 2nd Element + -ide]

• First Element Exception: If the formula contains only one atom of the first element, omit the prefix mono- (e.g., $\text{CO}_2$ is named carbon dioxide, not monocarbon dioxide).

• Examples:
  • $\text{NI}_3$: Nitrogen triiodide
  • $\text{P}_4\text{S}_{10}$: Tetraphosphorus decasulfide
  • Dinitrogen monoxide: $\text{N}_2\text{O}$
  • Xenon difluoride: $\text{XeF}_2$

Acids: Classification and Nomenclature

• General Properties: Acid names derive from Latin acidus (meaning sour or tart). Acids dissolve many metals and neutralize basic solutions.

• Structural Definition: Chemical formulas for acids are written with hydrogen ($\text{H}$) listed as the first element, representing a releasable $\text{H}^+$ cation (proton). Acids are written to be overall electrically neutral.

• Classification of Acids:

  • Binary Acids: Contain an $\text{H}^+$ proton combined with a single nonmetal monoatomic anion.
  • Oxyacids: Contain an $\text{H}^+$ proton combined with a polyatomic oxyanion.

• Binary Acid Naming Rule:   hydro- + [Base Name of Nonmetal + -ic] + acid
  • $\text{HF}$: Hydrofluoric acid
  • $\text{HCl}$: Hydrochloric acid
  • $\text{HBr}$: Hydrobromic acid
  • $\text{HI}$: Hydroiodic acid
  • $\text{H}_2\text{S}$: Hydrosulfuric acid
  • $\text{H}_2\text{Se}$: Hydroselenic acid
  • $\text{H}_3\text{P}$: Hydrophosphoric acid
  • $\text{H}_2\text{O}$: Hydroxic acid

• Oxyacid Naming Rules:
  • If the polyatomic oxyanion ends in -ate:     Change -ate to -ic and add the word acid.     Template: [Base Name of Oxyanion + -ic] + acid
  • $\text{ClO}_3^-$ (Chlorate) $+ \text{H}^+ \rightarrow \text{HClO}_3$ (Chloric acid)
  • $\text{ClO}_4^-$ (Perchlorate) $+ \text{H}^+ \rightarrow \text{HClO}_4$ (Perchloric acid)
  • $\text{SO}_4^{2-}$ (Sulfate) $+ 2\text{H}^+ \rightarrow \text{H}_2\text{SO}_4$ (Sulfuric acid)
  • $\text{PO}_4^{3-}$ (Phosphate) $+ 3\text{H}^+ \rightarrow \text{H}_3\text{PO}_4$ (Phosphoric acid)
  • $\text{CrO}_4^{2-}$ (Chromate) $+ 2\text{H}^+ \rightarrow \text{H}_2\text{CrO}_4$ (Chromic acid)
  • $\text{C}_2\text{H}_3\text{O}_2^-$ (Acetate) $+ \text{H}^+ \rightarrow \text{HC}_2\text{H}_3\text{O}_2$ (Acetic acid)
  • If the polyatomic oxyanion ends in -ite:     Change -ite to -ous and add the word acid.     Template: [Base Name of Oxyanion + -ous] + acid
  • $\text{ClO}_2^-$ (Chlorite) $+ \text{H}^+ \rightarrow \text{HClO}_2$ (Chlorous acid)
  • $\text{ClO}^-$ (Hypochlorite) $+ \text{H}^+ \rightarrow \text{HClO}$ (Hypochlorous acid)
  • $\text{SO}_3^{2-}$ (Sulfite) $+ 2\text{H}^+ \rightarrow \text{H}_2\text{SO}_3$ (Sulfurous acid)
  • $\text{PO}_3^{3-}$ (Phosphite) $+ 3\text{H}^+ \rightarrow \text{H}_3\text{PO}_3$ (Phosphorous acid)
  • $\text{NO}_2^-$ (Nitrite) $+ \text{H}^+ \rightarrow \text{HNO}_2$ (Nitrous acid)

• Common Acids and Their Practical Uses:
  • Hydrochloric acid ($\text{HCl}$): Metal cleaning; food preparation; ore refining; primary component of stomach acid.
  • Sulfuric acid ($\text{H}_2\text{SO}_4$): Fertilizer and explosives manufacturing; dye and glue production; automobile batteries; electroplating of copper.
  • Nitric acid ($\text{HNO}_3$): Fertilizer and explosives manufacturing; dye and glue production.
  • Acetic acid ($\text{HC}_2\text{H}_3\text{O}_2$): Plastic and rubber manufacturing; food preservation; active component of vinegar.
  • Citric acid ($\text{H}_3\text{C}_6\text{H}_5\text{O}_7$): Found in citrus fruits such as lemons and limes; used to adjust pH in foods and beverages.

Inorganic Nomenclature Decision Flowchart Strategy

• Sequential Decision Tree for Naming:

  1. Does the formula begin with Hydrogen ($\text{H}$)?
  • YES: It is an Acid.
    • Does it contain Oxygen?
      • NO: Binary Acid \t\text{R} Use hydro- + base name of nonmetal + -ic acid.
      • YES: Oxyacid \t\text{R} Check the polyatomic oxyanion ending:
      • Anion ends in -ate \t\text{R} Change to -ic acid.
      • Anion ends in -ite \t\text{R} Change to -ous acid.
  • NO: Proceed to Step 2.
  1. Is the substance Ionic or Molecular?
  • MOLECULAR (All Nonmetals):
    • Use Greek numerical prefixes to specify atom counts.
    • Format: [Prefix] + 1st Element + [Prefix] + 2nd Element Base + -ide (Omit mono- for first element).
  • IONIC (Metal + Nonmetal, or contains Polyatomic Ion):
    • Check metal cation type:
      • Type I (Invariant metal charge: Groups 1A, 2A, 3A, $\text{Zn}^{2+}$, $\text{Sc}^{3+}$, $\text{Ag}^+$): [Cation Name] + [Anion Base + -ide / Polyatomic Ion Name].
      • Type II (Multivalent metal charge: Transition metals, $\text{Pb}$, $\text{Sn}$, $\text{Tl}$): [Cation Name] + (Roman Numeral Charge) + [Anion Base + -ide / Polyatomic Ion Name].
  1. Is it a Hydrate ($\cdot n\text{H}_2\text{O}$)?
  • Append [Greek Prefix]-hydrate to the complete ionic compound name.

Organic Compounds: Hydrocarbons

• Definition: Organic compounds composed primarily of Carbon ($\text{C}$) and Hydrogen ($\text{H}$).

• Carbon Bonding Capacity: Carbon always forms four covalent bonds to achieve an octet.

• Simplest Hydrocarbon: Methane ($\text{CH}_4$).

• Table 4.5: Common Hydrocarbons:

  • Methane ($\text{CH}_4$):
  • Structural Formula: Central C bound to 4 single H atoms.
  • Space-Filling Model: Tetrahedral geometry.
  • Common Use: Primary component of natural gas.
  • Propane ($\text{C}_3\text{H}_8$):
  • Structural Formula: $\text{H}_3\text{C}-\text{CH}_2-\text{CH}_3$
  • Common Use: LP gas for grills and outdoor stoves.
  • n-Butane ($\text{C}_4\text{H}_{10}$):
  • Structural Formula: $\text{H}_3\text{C}-\text{CH}_2-\text{CH}_2-\text{CH}_3$ ("n-" indicates "normal" straight-chain structure).
  • Common Use: Fuel for pocket lighters.
  • n-Pentane ($\text{C}_5\text{H}_{12}$):
  • Structural Formula: $\text{H}_3\text{C}-\text{CH}_2-\text{CH}_2-\text{CH}_2-\text{CH}_3$
  • Common Use: Component of commercial gasoline.
  • Ethene ($\text{C}_2\text{H}_4$):
  • Structural Formula: $\text{H}_2\text{C}=\text{CH}_2$ (Contains a carbon-carbon double bond).
  • Common Use: Fruit ripening agent.
  • Ethyne ($\text{C}_2\text{H}_2$):
  • Structural Formula: \text{H}-\text{C}\t\text{E}\text{C}-\text{H} (Contains a carbon-carbon triple bond).
  • Common Use: High-temperature fuel for welding torches.

Comprehensive Practice Problems & Worked Solutions

• Classification as Ionic or Molecular:

  • $\text{CBr}_4$: Molecular compound (composed exclusively of nonmetals).
  • $\text{FeS}$: Ionic compound (metal cation bound to nonmetal anion).
  • $\text{P}_4\text{O}_6$: Molecular compound (composed exclusively of nonmetals).
  • $\text{PbF}_2$: Ionic compound (metal cation bound to nonmetal anion).
  • $\text{Na}_2\text{O}$: Ionic compound (metal cation bound to nonmetal anion).

• Binary Ionic Compound Naming & Formula Solutions:

  • $\text{CaBr}_2$: Calcium bromide
  • $\text{Ag}_3\text{N}$: Silver nitride
  • $\text{PbCl}_4$: Lead(IV) chloride
  • $\text{FeS}$: Iron(II) sulfide
  • Rubidium sulfide: $\text{Rb}_2\text{S}$
  • Magnesium iodide: $\text{MgI}_2$
  • Ruthenium(IV) oxide: $\text{RuO}_2$
  • Chromium(III) nitride: $\text{CrN}$

• Polyatomic & Hydrated Ionic Solutions:

  • $\text{NH}_4\text{Br}$: Ammonium bromide
  • $\text{Co}(\text{NO}_3)_2$: Cobalt(II) nitrate
  • $\text{NaC}_2\text{H}_3\text{O}_2 \cdot 3\text{H}_2\text{O}$: Sodium acetate trihydrate
  • $\text{NH}_4\text{HSO}_4$: Ammonium hydrogen sulfate
  • Nickel(II) perchlorate: $\text{Ni}(\text{ClO}_4)_2$
  • Calcium carbonate: $\text{CaCO}_3$
  • Zinc dichromate anhydrate: $\text{ZnCr}_2\text{O}_7$
  • Copper(II) sulfate pentahydrate: $\text{CuSO}_4 \cdot 5\text{H}_2\text{O}$

• Acid Naming & Formula Solutions:

  • $\text{H}_2\text{SO}_4$: Sulfuric acid
  • $\text{H}_2\text{CO}_3$: Carbonic acid
  • $\text{H}_2\text{S}$: Hydrosulfuric acid
  • $\text{HNO}_2$: Nitrous acid
  • Chromic acid: $\text{H}_2\text{CrO}_4$
  • Acetic acid: $\text{HC}_2\text{H}_3\text{O}_2$
  • Phosphorous acid: $\text{H}_3\text{PO}_3$
  • Hydroiodic acid: $\text{HI}$