Bonds and Compounds

Introduction to Compounds

Compounds are substances formed from two or more elements in a fixed ratio, held together by chemical bonds. The properties of a compound differ from those of its constituent elements, and compounds cannot be separated by physical methods.

• Molecular formula: Shows the exact number and type of atoms in a molecule (e.g., H2O).

• Structural formula: Indicates the spatial arrangement of atoms (e.g., H-O-H).

• Ball-and-stick model: Visualizes atoms as spheres and bonds as sticks.

• Space-filling model: Represents the relative sizes and spatial relationships of atoms.

Chemical Reactions and Formulas

Chemical Equations

Chemical equations use chemical formulas to express the identities and quantities of substances involved in a reaction.

• Reactants: Substances consumed during the reaction.

• Products: Substances formed as a result of the reaction.

• Example:

Law of Multiple Proportions

Combining Elements in Different Ratios

The same elements can combine in different ratios to form different compounds. The law of multiple proportions states that if two elements form more than one compound, the masses of one element that combine with a fixed mass of the other are in the ratio of small whole numbers.

• Example: Carbon and oxygen form CO and CO2 in different ratios.

Types of Compounds and Bonding

Classification of Compounds

Compounds are classified based on the types of bonds holding their atoms together.

• Molecular (Covalent) Compounds: Atoms held together by covalent bonds (sharing electrons).

• Ionic Compounds: Composed of positively charged cations and negatively charged anions held together by electrostatic attraction.

• Metallic Bonds: Attraction between positive ions and freely mobile electrons ("electron sea").

Properties Related to Bond Type

Bonding and Physical Properties

The type of bonding in a compound determines its physical properties, such as malleability, conductivity, and solubility.

• Metals: Malleable due to the absence of electrostatic repulsion; metallic bonds allow deformation.

• Ionic Crystals: Brittle due to repulsion between like charges when stress is applied.

Nomenclature of Compounds

Molecular and Empirical Formulas

• Molecular formula: Exact number and type of atoms in a molecule (e.g., C6H6).

• Empirical formula: Simplest whole-number ratio of elements (e.g., CH for benzene).

• Formula unit: Smallest electrically neutral unit in an ionic compound.

Binary Molecular/Covalent Compounds

Binary molecular compounds consist of two nonmetals.

• First element is named first.

• Second element is named with the suffix "-ide".

• Prefixes indicate the number of atoms: mono-, di-, tri-, tetra-, penta-, hexa-, hepta-, octa-, nona-, deca-.

• Never use "mono-" for the first element.

• Example: P2O5 is diphosphorus pentoxide.

Examples of Binary Molecular Compounds

• NF3: nitrogen trifluoride

• NO: nitrogen monoxide

• NO2: nitrogen dioxide

• N2O: dinitrogen monoxide (laughing gas)

• N2O4: dinitrogen tetraoxide

• PCl5: phosphorous pentachloride

• S2F10: disulfur decafluoride

• H2O: dihydrogen monoxide (water)

• NH3: nitrogen trihydride (ammonia)

• PH3: phosphorous trihydride (phosphine)

Binary Ionic Compounds

Binary ionic compounds consist of cations (usually metals) and anions (usually nonmetals).

• Formulas must be electrically neutral.

• Cation is named first using the elemental name.

• Anion is named with the "-ide" ending.

• No numerals for the number of atoms in the name.

• Example: Na2O is sodium oxide, not disodium monoxide.

Naming Cations

• Cations can be single atom ions or polyatomic ions (e.g., NH4+ is ammonium).

• Group I metals (charge 1+), Group II metals (charge 2+), Al3+ (aluminium), Zn2+ (zinc), Ag+ (silver).

• Metals with multiple charges: name followed by Roman numerals or suffixes (-ous for lower, -ic for higher valency).

• Examples: Cu+ (copper(I), cuprous), Cu2+ (copper(II), cupric), Fe2+ (iron(II), ferrous), Fe3+ (iron(III), ferric).

Naming Anions

• Negatively charged elements have the suffix "-ide" (e.g., O2- is oxide, S2- is sulfide).

• Polyatomic ions (usually oxoanions) have suffixes "-ate" or "-ite". "-ate" indicates more oxygen than "-ite".

• Examples: SO42- (sulfate), SO32- (sulfite), NO3- (nitrate), NO2- (nitrite).

Common Oxoanions

Other Common Polyatomic Ions

• Acetate: C2H3O2-

• Carbonate: CO32-

• Perchlorate: ClO4-

• Nitrate: NO3-

• Sulfate: SO42-

• Chromate: CrO42-

• Hydrogen carbonate: HCO3-

Naming Ionic Compounds: Examples

• NaOH: sodium hydroxide

• CaCO3: calcium carbonate

• FeSO4: iron(II) sulfate or ferrous sulfate

• FeSO3: iron(II) sulfite or ferrous sulfite

• (NH4)3PO4: ammonium phosphate

Naming Acids and Acid Salts

Binary Acids

Binary acids contain hydrogen and a nonmetal (usually a halogen).

• Name: "hydro" + base name + "ic" + "acid"

• Examples: HBr(aq): hydrobromic acid, HCl(aq): hydrochloric acid

Oxoacids

Oxoacids have the general formula HaXbOc, where X is an element other than hydrogen or oxygen. The name is derived from the polyatomic oxoanion.

• If the anion ends in "-ate", the acid name ends in "-ic".

• If the anion ends in "-ite", the acid name ends in "-ous".

• Examples: SO32- (sulfite) → H2SO3 (sulfurous acid); SO42- (sulfate) → H2SO4 (sulfuric acid)

Naming Acid Salts

Acid salts contain extra hydrogens on the anion, which are acidic and can be released as H+ in water.

• Name: [cation name] [anion name]

• Examples: calcium hydrogen carbonate, magnesium dihydrogen phosphite

Nomenclature Summary Table

Conclusion

Understanding the types of compounds, their bonding, and the rules for naming them is fundamental in general chemistry. Mastery of these concepts enables students to interpret chemical formulas, write chemical equations, and predict compound properties. Additional info: Some images and tables were inferred to clarify the academic context and reinforce the explanations.