Exhaustive Study Notes on Molecules, Compounds, and Chemical Equations

Elements, Mixtures, and Compounds

Elements combine with one another to form compounds, yielding the vast diversity of physical substances observed in nature.
When elements combine chemically to form a compound, an entirely new substance is produced with chemical and physical properties distinct from those of its constituent elements.

Selected Properties of Hydrogen, Oxygen, and Water

Comparison of Physical and Chemical Properties:
- Hydrogen ( $H_2$ ): Boiling point of $-253\,^{\circ}\text{C}$ , gas state at room temperature, highly explosive.
- Oxygen ( $O_2$ ): Boiling point of $-183\,^{\circ}\text{C}$ , gas state at room temperature, necessary for combustion.
- Water ( $H_2O$ ): Boiling point of $100\,^{\circ}\text{C}$ , liquid state at room temperature, used to extinguish flames.

Distinction between Mixtures and Compounds:
- Mixtures: Elements can mix in any arbitrary ratio. A sample containing hydrogen gas ( $H_2$ ) and oxygen gas ( $O_2$ ) can exist in any proportion.
- Compounds: Elements combine in fixed, definite proportions. Water ( $H_2O$ ) molecules always maintain a fixed atomic ratio of 2 hydrogen atoms to 1 oxygen atom.

Chemical Bonds and Energy

Compounds are held together by chemical bonds, which arise from electrostatic attractions between positively charged nuclei (protons) and negatively charged electrons.
Chemical bonds are broadly classified into two primary types: ionic bonds and covalent bonds.

Ionic Bonds

Occur between metals and nonmetals.
Involve the transfer of one or more electrons from the metal atom to the nonmetal atom.
- The metal atom loses electrons to become a positively charged ion (cation).
- The nonmetal atom gains electrons to become a negatively charged ion (anion).
Oppositely charged ions attract each other via electrostatic forces, forming an ionic bond.
In the solid phase, ionic compounds assemble into a lattice, which is a regular three-dimensional array of alternating cations and anions.

Example of Ionic Bond Formation (Sodium Chloride, $NaCl$ ):
- Neutral sodium atom ( $\text{Na}$ , $11e^-$ ) loses an electron to form a sodium cation ( $\text{Na}^+$ , $10e^-$ ).
- Neutral chlorine atom ( $\text{Cl}$ , $17e^-$ ) gains an electron to form a chloride anion ( $\text{Cl}^-$ , $18e^-$ ).
- Electrostatic attraction binds $\text{Na}^+$ and $\text{Cl}^-$ into a crystalline lattice.

Covalent Bonds

Occur between two or more nonmetals.
Involve the sharing of electron pairs between nonmetal atoms to form discrete molecules.
Covalently bonded compounds are referred to as molecular compounds.
Potential Energy in Covalent Bonding:
- Electrostatic potential energy is lowest (most stable state) when a shared negative charge (electron density) is situated directly between two positively charged nuclei, reducing inter-nuclear repulsion and maximizing attraction.

Representing Compounds: Chemical Formulas and Molecular Models

A chemical formula indicates the specific elements present in a compound and the relative number of atoms or ions of each.
- Water: $H_2O$
- Sodium Chloride: $NaCl$
- Carbon Dioxide: $CO_2$
- Carbon Tetrachloride: $CCl_4$

Types of Chemical Formulas

Empirical Formula: Gives the relative, simplest whole-number ratio of atoms of each element in a compound.
- For Hydrogen Peroxide ( $H_2O_2$ ), the greatest common factor is 2; the empirical formula is $HO$ .
- For Diborane ( $B_2H_6$ ), the greatest common factor is 2; the empirical formula is $BH_3$ .
- For Carbon Tetrachloride ( $CCl_4$ ), the common factor is 1; the empirical formula and molecular formula are identical ( $CCl_4$ ).
Molecular Formula: Gives the actual number of atoms of each element present in a single molecule of a compound.
- $\text{Molecular Formula} = n \times (\text{Empirical Formula})$ , where $n$ is a positive integer.
Structural Formula: Uses structural lines to represent covalent bonds and depicts the connectivity and spatial arrangement of atoms within a molecule.
- Structural formulas communicate the highest amount of information regarding a compound; empirical formulas communicate the least.

Molecular Models

Ball-and-Stick Model: Represents atoms as spheres and chemical bonds as sticks, accurately depicting the three-dimensional geometry of the molecule.
Space-Filling Model: Represents atoms as expanded spheres filling interatomic space, providing an accurate estimate of a molecule's outer electron density surface and overall shape.

Ways of Representing Benzene, Acetone, Glucose, and Ammonia

Standard Color Coding for Molecular Models:
- Hydrogen: White
- Carbon: Black / Dark Gray
- Nitrogen: Blue
- Oxygen: Red
- Fluorine: Light Green
- Phosphorus: Dark Blue
- Sulfur: Yellow
- Chlorine: Green

Classification of Pure Substances

Classification of Elements and Compounds

Pure substances are categorized into elements and compounds:
- Elements:
- Atomic Elements: Exist in nature as single isolated atoms as their basic structural units (e.g., $\text{Ne}$ , $\text{Na}$ , $\text{K}$ , $\text{Mg}$ ).
- Molecular Elements: Exist in nature as multi-atom molecules composed of two or more bound atoms of the same element.
- Compounds:
- Molecular Compounds: Composed of two or more covalently bonded nonmetal atoms; basic repeating units are discrete molecules (e.g., $H_2O$ , $CO_2$ , $C_3H_8$ ).
- Ionic Compounds: Composed of cations and anions held by ionic bonds; basic structural unit is the formula unit (e.g., $NaCl$ ).

Molecular Elements on the Periodic Table

Detailed Listing of Molecular Elements:
- Diatomic Elements (7 elements): Hydrogen ( $H_2$ ), Nitrogen ( $N_2$ ), Oxygen ( $O_2$ ), Fluorine ( $F_2$ ), Chlorine ( $Cl_2$ ), Bromine ( $Br_2$ ), Iodine ( $I_2$ ).
- Polyatomic Elements: Phosphorus ( $P_4$ ), Sulfur ( $S_8$ ), Selenium ( $Se$ ).

Nomenclature of Ionic Compounds

Formula rules for ionic compounds:
- Cations are listed first, followed by anions.
- Total positive charge must equal total negative charge (overall electrical neutrality).
- Subscripts reflect the smallest whole-number ratio of ions.

Categorization of Metals in Ionic Compounds

Type I Metals: Form only one type of cation with a fixed, invariant charge across all compounds.
Type II Metals: Form more than one type of cation with variable charges depending on the compound.

Metals Whose Charge Is Invariant from One Compound to Another

Invariant Charge Metals (Type I):
- Group 1A Alkali Metals: $\text{Li}^+$ , $\text{Na}^+$ , $\text{K}^+$ , $\text{Rb}^+$ , $\text{Cs}^+$ (charge is always $1+$ ).
- Group 2A Alkaline Earth Metals: $\text{Mg}^{2+}$ , $\text{Ca}^{2+}$ , $\text{Sr}^{2+}$ , $\text{Ba}^{2+}$ (charge is always $2+$ ).
- Selected Transition/Main Group Metals: Scandium ( $\text{Sc}^{3+}$ ), Aluminum ( $\text{Al}^{3+}$ ), Zinc ( $\text{Zn}^{2+}$ ), Silver ( $\text{Ag}^+$ ).

Naming Type I Binary Ionic Compounds

Binary compounds contain exactly two different elements.
Naming format: [name of cation (metal)] + [base name of anion (nonmetal) + -ide]
Common Monoatomic Anions:
- Fluorine ( $\text{F}^-$ ): Fluoride
- Chlorine ( $\text{Cl}^-$ ): Chloride
- Bromine ( $\text{Br}^-$ ): Bromide
- Iodine ( $\text{I}^-$ ): Iodide
- Oxygen ( $\text{O}^{2-}$ ): Oxide
- Sulfur ( $\text{S}^{2-}$ ): Sulfide
- Nitrogen ( $\text{N}^{3-}$ ): Nitride
- Phosphorus ( $\text{P}^{3-}$ ): Phosphide
Examples:
- $KCl$ : Potassium chloride
- $CaO$ : Calcium oxide

Naming Type II Binary Ionic Compounds

Applied to metals (mostly transition metals, along with main group metals like $\text{Pb}$ , $\text{Sn}$ , $\text{Ti}$ ) that exhibit variable charges.
Naming format: [name of cation (metal)] + (charge of cation in Roman numerals) + [base name of anion (nonmetal) + -ide]
Examples:
- $FeS$ : Iron(II) sulfide (since $\text{S}^{2-}$ requires $\text{Fe}^{2+}$ )
- $Fe_2S_3$ : Iron(III) sulfide (since $3 \times \text{S}^{2-} = 6-$ , requiring $2 \times \text{Fe}^{3+} = 6+$ )
- $Cu_2O$ : Copper(I) oxide (since $\text{O}^{2-}$ requires $2 \times \text{Cu}^+$ )
- $CuO$ : Copper(II) oxide (since $\text{O}^{2-}$ requires $\text{Cu}^{2+}$ )
- $CrBr_3$ : Chromium(III) bromide ( $\text{Cr charge} + 3(-1) = 0 \rightarrow \text{Cr} = +3$ )
Systemic versus Older/Common Nomenclature for Variable Metal Cations:
- Chromium: $\text{Cr}^{2+}$ = Chromium(II) / Chromous; $\text{Cr}^{3+}$ = Chromium(III) / Chromic
- Iron: $\text{Fe}^{2+}$ = Iron(II) / Ferrous; $\text{Fe}^{3+}$ = Iron(III) / Ferric
- Cobalt: $\text{Co}^{2+}$ = Cobalt(II) / Cobaltous; $\text{Co}^{3+}$ = Cobalt(III) / Cobaltic
- Copper: $\text{Cu}^+$ = Copper(I) / Cuprous; $\text{Cu}^{2+}$ = Copper(II) / Cupric
- Tin: $\text{Sn}^{2+}$ = Tin(II) / Stannous; $\text{Sn}^{4+}$ = Tin(IV) / Stannic
- Mercury: $\text{Hg}_2^{2+}$ = Mercury(I) / Mercurous; $\text{Hg}^{2+}$ = Mercury(II) / Mercuric
- Lead: $\text{Pb}^{2+}$ = Lead(II) / Plumbous; $\text{Pb}^{4+}$ = Lead(IV) / Plumbic

Ionic Compounds Containing Polyatomic Ions

Polyatomic ions are groups of covalently bonded atoms possessing an overall net ionic charge.
Naming rule: Follow standard ionic nomenclature, substituting the exact name of the polyatomic ion.
- $NaNO_3$ : Sodium nitrate ( $\text{Na}^+$ and $\text{NO}_3^-$ )
- $CaCO_3$ : Calcium carbonate ( $\text{Ca}^{2+}$ and $\text{CO}_3^{2-}$ )
- $Mg(ClO_3)_2$ : Magnesium chlorate ( $\text{Mg}^{2+}$ and $\text{ClO}_3^-$ )
- $NaNO_2$ : Sodium nitrite ( $\text{Na}^+$ and $\text{NO}_2^-$ )
Standard Polyatomic Ions Reference:
- Acetate: $C_2H_3O_2^-$
- Carbonate: $CO_3^{2-}$
- Hydrogen carbonate (bicarbonate): $HCO_3^-$
- Hydroxide: $OH^-$
- Nitrite: $NO_2^-$
- Nitrate: $NO_3^-$
- Chromate: $CrO_4^{2-}$
- Dichromate: $Cr_2O_7^{2-}$
- Phosphate: $PO_4^{3-}$
- Hydrogen phosphate: $HPO_4^{2-}$
- Dihydrogen phosphate: $H_2PO_4^-$
- Ammonium: $NH_4^+$
- Hypochlorite: $ClO^-$
- Chlorite: $ClO_2^-$
- Chlorate: $ClO_3^-$
- Perchlorate: $ClO_4^-$
- Permanganate: $MnO_4^-$
- Sulfite: $SO_3^{2-}$
- Hydrogen sulfite (bisulfite): $HSO_3^-$
- Sulfate: $SO_4^{2-}$
- Hydrogen sulfate (bisulfate): $HSO_4^-$
- Cyanide: $CN^-$
- Peroxide: $O_2^{2-}$

Oxyanions Nomenclature

Oxyanions are polyatomic anions containing oxygen and another nonmetal element.
When a series of oxyanions contains two species:
- The ion with more oxygen atoms receives the suffix -ate ( $\text{NO}_3^-$ : nitrate; $\text{SO}_4^{2-}$ : sulfate).
- The ion with fewer oxygen atoms receives the suffix -ite ( $\text{NO}_2^-$ : nitrite; $\text{SO}_3^{2-}$ : sulfite).
When an oxyanion series contains four species, prefixes are added:
- hypo- = fewest oxygen atoms
- per- = most oxygen atoms
- Halogen Oxyanion Series Example:
- $ClO^-$ : Hypochlorite
- $ClO_2^-$ : Chlorite
- $ClO_3^-$ : Chlorate
- $ClO_4^-$ : Perchlorate
- $BrO^-$ : Hypobromite
- $BrO_2^-$ : Bromite
- $BrO_3^-$ : Bromate
- $BrO_4^-$ : Perbromate

Hydrated Ionic Compounds

Hydrates are ionic compounds containing a specific stoichiometric ratio of water molecules bound within their crystal structure.
Naming format: [Ionic Compound Name] + [Prefix]hydrate
Hydrate Numerical Prefixes:
- $\frac{1}{2}$ = hemi
- $1$ = mono
- $2$ = di
- $3$ = tri
- $4$ = tetra
- $5$ = penta
- $6$ = hexa
- $7$ = hepta
- $8$ = octa
Examples:
- $MgSO_4 \cdot 7H_2O$ : Magnesium sulfate heptahydrate
- $CaSO_4 \cdot \frac{1}{2}H_2O$ : Calcium sulfate hemihydrate
- $BaCl_2 \cdot 6H_2O$ : Barium chloride hexahydrate
- $CuSO_4 \cdot 6H_2O$ : Copper(II) sulfate hexahydrate

Nomenclature of Molecular Compounds

Molecular compounds consist of two or more covalently bonded nonmetals.
Ordering elements in formula/name:
- Element with the smaller group number is listed first.
- If both elements belong to the same group, the element with the greater row (period) number is listed first.
Naming format: [Prefix][name of 1st element] + [Prefix][base name of 2nd element + -ide]
Numerical Prefixes for Molecular Compounds:
- 1 = mono- (Note: mono- is omitted for the first element)
- 2 = di-
- 3 = tri-
- 4 = tetra-
- 5 = penta-
- 6 = hexa-
- 7 = hepta-
- 8 = octa-
- 9 = nona-
- 10 = deca-
Examples:
- $NO$ : Nitrogen monoxide
- $NO_2$ : Nitrogen dioxide
- $N_2O$ : Dinitrogen monoxide
- $N_2O_3$ : Dinitrogen trioxide
- $N_2O_4$ : Dinitrogen tetroxide
- $N_2O_5$ : Dinitrogen pentoxide
- $P_2O_5$ : Diphosphorus pentoxide

Acids: Classification and Nomenclature

Acids are molecular compounds that ionize in water to form hydrogen cations ( $H^+$ ) and an accompanying anion.
Characteristics of Acids:
- Sour taste.
- Ability to dissolve reactive metals (e.g., $\text{Zn}$ , $\text{Fe}$ , $\text{Mg}$ ; does not dissolve unreactive noble metals like $\text{Au}$ , $\text{Ag}$ , $\text{Pt}$ ).
- Chemical formulas begin with $H$ (e.g., $HCl$ , $H_2SO_4$ ).
- Dissolution behavior: $HCl(g) \xrightarrow{H_2O} H^+(aq) + Cl^-(aq)$ .

Binary Acids

Binary acids consist of a hydrogen cation ( $H^+$ ) paired with a monoatomic nonmetal anion.
Naming format: hydro- + [base name of nonmetal + -ic] + acid
Examples:
- $HCl(aq)$ : Hydrochloric acid
- $H_2S(aq)$ : Hydrosulfuric acid
- $HF(aq)$ : Hydrofluoric acid
- $HBr(aq)$ : Hydrobromic acid

Oxyacids

Oxyacids consist of a hydrogen cation ( $H^+$ ) combined with a polyatomic oxyanion.
Naming rules based on oxyanion suffix:
- If the oxyanion name ends in -ate, change ending to -ic and add acid.
- If the oxyanion name ends in -ite, change ending to -ous and add acid.
Examples:
- $HNO_3(aq)$ : Nitric acid (oxyanion: nitrate, $NO_3^-$ )
- $HNO_2(aq)$ : Nitrous acid (oxyanion: nitrite, $NO_2^-$ )
- $H_2SO_4(aq)$ : Sulfuric acid (oxyanion: sulfate, $SO_4^{2-}$ )
- $H_2SO_3(aq)$ : Sulfurous acid (oxyanion: sulfite, $SO_3^{2-}$ )
- $HClO_3(aq)$ : Chloric acid (oxyanion: chlorate, $ClO_3^-$ )
- $HC_2H_3O_2(aq)$ : Acetic acid (oxyanion: acetate, $C_2H_3O_2^-$ )
- $H_3PO_4(aq)$ : Phosphoric acid (oxyanion: phosphate, $PO_4^{3-}$ )

Environmental Significance: Acid Rain

Industrial pollutants such as gaseous nitrogen oxides ( $NO$ , $NO_2$ ) and sulfur oxides ( $SO_2$ , $SO_3$ ) react with atmospheric water vapor to yield corrosive nitric and sulfuric acids, causing acid rain that lowers the pH of natural aquatic environments and damages forestry.

Summary Flowchart for Inorganic Nomenclature

Flowchart Decision Logic:
1. Ionic (Metal + Nonmetal):
- Metal forms one type of ion (Type I): [Name of cation] + [Base name of anion + -ide] (e.g., $CaI_2$ = calcium iodide).
- Metal forms multiple ions (Type II): [Name of cation](Roman numeral charge) + [Base name of anion + -ide] (e.g., $FeCl_3$ = iron(III) chloride).
1. Molecular (Nonmetals only):
- [Prefix][1st element] + [Prefix][Base name of 2nd element + -ide] (e.g., $P_2O_5$ = diphosphorus pentoxide).
1. Acids ( $H$ + Nonmetals):
- Binary Acids (two elements): hydro- + [Base name of nonmetal + -ic] + acid (e.g., $HCl$ = hydrochloric acid).
- Oxyacids (contains oxygen):
  - Oxyanion ends in -ate: [Base name of oxyanion + -ic] + acid (e.g., $H_3PO_4$ = phosphoric acid).
  - Oxyanion ends in -ite: [Base name of oxyanion + -ous] + acid (e.g., $H_2SO_3$ = sulfurous acid).

Formula Mass, Molar Mass, and Stoichiometric Conversions

Formula Mass

Formula mass (also called molecular weight or molecular mass) is the average mass of an individual molecule or formula unit of a compound.
Equation: $\text{Formula Mass} = \sum_{i} \left( \text{Number of atoms of element } i \text{ in formula} \times \text{Atomic mass of element } i \right)$
Calculation Example for Water ( $H_2O$ ): $\text{Formula Mass of } H_2O = 2(1.008\,\text{amu}) + 1(16.00\,\text{amu}) = 18.02\,\text{amu}$

Molar Mass of Compounds

The molar mass of a compound is defined as the mass in grams of 1 mole ( $6.022 \times 10^{23}$ units) of its molecules or formula units.
Molar mass in $\text{g\,mol}^{-1}$ is numerically equal to formula mass in $\text{amu}$ . $\text{Molar mass of } H_2O = 18.02\,\text{g\,mol}^{-1}$
1 mole of $H_2O$ molecules contains 2 moles of $H$ atoms ( $2 \times 1.008\,\text{g} = 2.016\,\text{g}$ ) and 1 mole of $O$ atoms ( $1 \times 16.00\,\text{g} = 16.00\,\text{g}$ ).

Counting Molecules by Weighing

To convert from mass of a substance to number of molecules: $\text{Number of molecules of } A = \text{g } A \times \frac{1\,\text{mol } A}{\text{Molar mass of } A\,\text{g}} \times \frac{6.022 \times 10^{23}\,\text{molecules } A}{1\,\text{mol } A}$

Percent Composition and Conversion Factors from Formulas

The mass percentage of an element in a compound expresses the element's mass relative to the total mass of the compound: $\text{Mass percent of element } X = \frac{\text{Mass of element } X \text{ in 1 mol of compound}}{\text{Mass of 1 mol of compound}} \times 100\%$

Stoichiometric Ratios from Chemical Formulas

Chemical formulas give mole relationships between a compound and its constituent elements.
For water ( $H_2O$ ):
- $1\,\text{mol } H_2O : 2\,\text{mol } H$
- $1\,\text{mol } H_2O : 1\,\text{mol } O$
- Derived ratio conversion factor: $\frac{2\,\text{mol } H}{1\,\text{mol } H_2O}$

Determining Chemical Formulas from Experimental Data

Empirical Formula Determination Procedure

Convert Mass/Percentage to Grams: If given mass percentages, assume a total sample mass of $100\,\text{g}$ so that percentage values equal mass values in grams.
Convert Grams to Moles: Divide the mass of each element by its atomic molar mass ( $\text{g\,mol}^{-1}$ ).
Write Pseudoformula: Express subscripts using the calculated mole values.
Divide by Smallest Mole Value: Divide all mole values by the lowest mole quantity obtained to simplify subscripts.
Obtain Whole-Number Subscripts: If fractional subscripts persist after step 4, multiply all subscripts by a common integer:
- Subscript decimal $\approx 0.5$ : multiply all by 2.
- Subscript decimal $\approx 0.33$ or $0.67$ : multiply all by 3.
- Subscript decimal $\approx 0.25$ or $0.75$ : multiply all by 4.

Calculating Molecular Formula from Empirical Formula

The molecular formula is calculated by finding the integer factor $n$ : $n = \frac{\text{Molar Mass of Compound}}{\text{Empirical Formula Mass}}$
$\text{Molecular Formula} = n \times (\text{Empirical Formula subscript values})$

Combustion Analysis

Analytical technique used for organic compounds containing carbon, hydrogen, and oxygen.
Process:
1. A known mass of the compound is combusted completely in the presence of excess $O_2$ gas in a furnace.
2. All carbon in the sample is converted into $CO_2$ gas and trapped in a $CO_2$ absorber.
3. All hydrogen in the sample is converted into $H_2O$ vapor and trapped in an $H_2O$ absorber.
4. Both absorbers are weighed before and after combustion to determine the masses of $CO_2$ and $H_2O$ produced.
5. Elemental masses are calculated: $\text{Mass C} = \text{Mass } CO_2 \times \frac{12.011\,\text{g C}}{44.01\,\text{g } CO_2}$ $\text{Mass H} = \text{Mass } H_2O \times \frac{2.016\,\text{g H}}{18.02\,\text{g } H_2O}$
6. The mass of oxygen in the original sample is determined by subtraction: $\text{Mass O} = \text{Total sample mass} - (\text{Mass C} + \text{Mass H})$
7. Empirical formula is calculated using elemental mole ratios.

Organic Compounds

Organic compounds are carbon-based compounds primarily composed of carbon ( $\text{C}$ ) and hydrogen ( $\text{H}$ ), often incorporating oxygen ( $\text{O}$ ), nitrogen ( $\text{N}$ ), phosphorus ( $\text{P}$ ), sulfur ( $\text{S}$ ), and halogens.

Carbon Bonding Characteristics

Carbon forms four covalent bonds (tetravalent).
Carbon atoms form single ( $C-C$ ), double ( $C=C$ ), and triple ( $C\equiv C$ ) bonds.
Carbon possesses the ability to undergo catenation—linking to itself to form linear chains, branched structures, and aromatic/aliphatic rings.

Classification of Organic Compounds

Hydrocarbons: Compounds consisting exclusively of carbon and hydrogen.
Functionalized Hydrocarbons: Hydrocarbons incorporating functional groups (heteroatoms or double/triple bonds) that impart distinct chemical properties.

Hydrocarbons Nomenclature

Alkanes: Hydrocarbons containing only single bonds (suffix -ane).
Alkenes: Hydrocarbons containing one or more double bonds (suffix -ene).
Alkynes: Hydrocarbons containing one or more triple bonds (suffix -yne).
Base Prefixes for Carbon Chain Length:
- 1 Carbon = meth-
- 2 Carbons = eth-
- 3 Carbons = prop-
- 4 Carbons = but-
- 5 Carbons = pent-
- 6 Carbons = hex-
- 7 Carbons = hept-
- 8 Carbons = oct-
- 9 Carbons = non-
- 10 Carbons = dec-
Examples of Common Hydrocarbons:
- Methane ( $CH_4$ ): Primary component of natural gas.
- Propane ( $C_3H_8$ ): Fuel for outdoor grills.
- n-Butane ( $C_4H_{10}$ ): Fuel for lighters (where "n-" stands for straight-chain normal alkane).
- n-Pentane ( $C_5H_{12}$ ): Component of gasoline.
- Ethene ( $C_2H_4$ ): Ripening agent in fruit.
- Ethyne ( $C_2H_2$ ): Fuel for welding torches.
- Isobutane ( $C_4H_{10}$ ): Branched alkane isomer.
- Cyclohexane ( $C_6H_{12}$ ): Ring-shaped alkane hydrocarbon.

Families of Functionalized Hydrocarbons

Alcohols:
- Suffix: -ol
- General Formula: $R-OH$
- Example: Ethanol ( $CH_3CH_2OH$ )
- Use/Occurrence: Alcohol in fermented beverages; Methanol ( $CH_3OH$ ); Isopropanol / 2-propanol ( $CH_3CHOHCH_3$ ).
Ethers:
- Suffix: ether
- General Formula: $R-O-R'$
- Example: Diethyl ether ( $CH_3CH_2-O-CH_2CH_3$ )
- Use/Occurrence: Anesthetic, laboratory solvent.
Aldehydes:
- Suffix: -al
- General Formula: $R-CHO$
- Example: Ethanal / acetaldehyde ( $H_3C-CHO$ )
- Use/Occurrence: Perfumes, flavors.
Ketones:
- Suffix: -one
- General Formula: $R-CO-R'$
- Example: Propanone / acetone ( $H_3C-CO-CH_3$ )
- Use/Occurrence: Fingernail polish remover.
Carboxylic Acids:
- Suffix: acid
- General Formula: $R-COOH$
- Example: Acetic acid ( $H_3C-COOH$ )
- Use/Occurrence: Vinegar.
Esters:
- Suffix: -ate
- General Formula: $R-COO-R'$
- Example: Methyl acetate ( $H_3C-COO-OCH_3$ )
- Use/Occurrence: Laboratory solvent.
Amines:
- Suffix: amine
- General Formula: $R-NH_2$
- Example: Ethyl amine ( $CH_3CH_2NH_2$ )
- Use/Occurrence: Smell of rotten fish.

Conceptual Connections & Discussion Questions

Conceptual Connection 3.1

Question: What type of bond—ionic or covalent—forms between nitrogen and oxygen?
Options:
- a. Ionic
- b. Covalent
Correct Answer: b. Covalent
Explanation: Nitrogen and oxygen are both nonmetals. Bonding between nonmetals involves shared electron pairs, forming covalent bonds.

Conceptual Connection 3.2

Question: Select the structural formula for water.
Options:
- a. $H-O$
- b. $H-H$
- c. $H-O-H$
- d. $H_2O$
Correct Answer: c. $H-O-H$
Explanation: A structural formula must show atom connectivity using lines for covalent bonds. Option c correctly shows oxygen bonded to two individual hydrogen atoms.

Conceptual Connection 3.3

Question: What part of the atom do the spheres in the molecular space-filling models represent? If you were to superimpose a nucleus on one of these spheres, how big would you draw it?
Options:
- a. Each sphere represents the hard outer shell of an atom. The nucleus would be too small to see on the same scale.
- b. Each sphere represents the electron cloud of the atom. The nucleus would be too small to see on the same scale.
- c. Each sphere represents the nucleus of an atom. The nucleus is the same size as the sphere.
Correct Answer: b. Each sphere represents the electron cloud of the atom. The nucleus would be too small to see on the same scale.

Conceptual Connection 3.4

Question: Classify the substance represented by a molecular view showing bound identical pairs of spheres.
Options:
- a. Atomic element
- b. Molecular element
- c. Molecular compound
- d. Ionic compound
Correct Answer: b. Molecular element
Explanation: The sample contains discrete molecules made of two identical bonded atoms of a single element (a diatomic molecule).

Conceptual Connection 3.5

Question: Which statement best summarizes the difference between ionic and molecular compounds?
Options:
- a. Molecular compounds contain highly directional covalent bonds, which result in the formation of molecules. Ionic compounds contain nondirectional ionic bonds, which result (in the solid state) in the formation of ionic lattices.
- b. Molecular compounds and ionic compounds both contain molecules as their smallest identifiable unit, but in ionic compounds the molecules are smaller.
- c. A molecular compound is composed of covalently bonded molecules. An ionic compound is composed of ionically bonded molecules (in the solid phase).
Correct Answer: a. Molecular compounds contain highly directional covalent bonds, which result in the formation of molecules. Ionic compounds contain nondirectional ionic bonds, which result (in the solid state) in the formation of ionic lattices.

Conceptual Connection 3.6

Question: Which metal has the same charge in all of its compounds?
Options:
- a. Fe
- b. Mo
- c. Pb
- d. Sr
Correct Answer: d. Sr
Explanation: Strontium ( $\text{Sr}$ ) is an alkaline earth metal (Group 2A) and invariably forms a $2+$ charge. $\text{Fe}$ , $\text{Mo}$ , and $\text{Pb}$ are variable-charge metals.

Conceptual Connection 3.7

Question: Identify the polyatomic ion and its charge in each compound: $KNO_2$ , $CaSO_4$ , Mg(NO_3)_2$.\n- **Options:**\n - a. NO_2^- $,$ SO_4^{2-} $, and$ NO_3^-\n - b. K^+ $,$ Ca^{2+} $, and$ Mg^{2+}\n - c. K^+ $,$ Ca^{2+} $,$ Mg^{2+} $,$ NO_2^- $,$ SO_4^{2-} $, and$ NO_3^-\n - d. NO_2^{2-} $,$ SO_4^- $, and$ NO_3^{2-}\n- **Correct Answer:** a. NO_2^- $,$ SO_4^{2-} $, and$ NO_3^-\n- **Explanation:** The polyatomic anions present are nitrite (NO_2^- $), sulfate ($ SO_4^{2-} $), and nitrate ($ NO_3^-).\n\n## Conceptual Connection 3.8\n- **Question:** The compound NCl_3 $is nitrogen trichloride, but$ AlCl_3 is simply aluminum chloride. Why?\n- **Options:**\n - a. The name forms differ because NCl_3 $is an ionic compound and$ AlCl_3 is a molecular compound.\n - b. The name forms differ because NCl_3 $is a molecular compound and$ AlCl_3 is an ionic compound. Prefixes such as mono-, di-, and tri- are used for molecular compounds but not for ionic compounds.\n- **Correct Answer:** b. The name forms differ because NCl_3 $is a molecular compound and$ AlCl_3 is an ionic compound. Prefixes such as mono-, di-, and tri- are used for molecular compounds but not for ionic compounds.\n\n## Conceptual Connection 3.9\n- **Question:** Which number is the best estimate for the scaling factor used in space-filling molecular models? (By approximately what number would you multiply the radius of an actual oxygen atom to get the radius of the sphere used in a space-filling model?)\n- **Options:**\n - a. 10\n - b. 10^4\n - c. 10^8\n - d. 10^{16}\n- **Correct Answer:** c. 10^8\n- **Explanation:** Atomic radii are roughly 10^{-10}\,\text{m} $, while a model sphere radius is roughly$ 10^{-2}\,\text{m} $. The magnification scaling factor is$ \frac{10^{-2}}{10^{-10}} = 10^8\n\n## Conceptual Connection 3.10\n- **Question:** Without doing any calculations, list the elements in C_6H_6O in order of decreasing mass percent composition.\n- **Options:**\n - a. C > O > H\n - b. O > C > H\n - c. H > O > C\n - d. C > H > O\n- **Correct Answer:** a. C > O > H\n- **Explanation:** Mass of carbon = 6 \times 12.011 = 72.066\,\text{g} $; mass of oxygen =$ 16.00\,\text{g} $; mass of hydrogen =$ 6 \times 1.008 = 6.048\,\text{g} $. Therefore, mass contribution follows$ C > O > H\n\n## Conceptual Connection 3.11\n- **Question:** The molecular formula for water is H_2O. Which ratio can be correctly derived from this formula?\n- **Options:**\n - a. 2\,\text{g } H : 1\,\text{g } H_2O\n - b. 2\,\text{mL } H : 1\,\text{mL } H_2O\n - c. 2\,\text{mol } H : 1\,\text{mol } H_2O\n- **Correct Answer:** c. 2\,\text{mol } H : 1\,\text{mol } H_2O$$
Explanation: Subscripts in chemical formulas express relationships in terms of atom counts or mole counts, not grams or volume units.