Covalent Bonding and Nomenclature Flashcards

Fundamentals of Chemical and Covalent Bonding

• Chemical bonding is the fundamental process through which two or more elements combine to attain a stable, full valence energy level, achieving either an octet ($8$ valence electrons) or a duet ($2$ valence electrons).

• Covalent bonding occurs specifically when two or more nonmetal atoms share valence electrons to fulfill their octet or duet requirements.

• The mechanism of electron sharing in covalent bonds can be conceptualized as a mutual attraction or competition over shared electrons between two bonded nuclei:

  • Each participating nonmetal atom exerts an attractive pull on the shared valence electrons to fill its outer valence shell.

  • Due to the opposing attractive pulls from both positively charged nuclei, the shared bonding electrons remain concentrated in the region between the two nuclei.

• Distinction between atoms and molecules:

  • An atom is the smallest structural unit of an element that retains the full chemical identity and properties of that element. Elements consist of only one specific type of atom.

  • A molecule is defined as two or more atoms bound together by covalent chemical bonds. It represents the smallest single unit of a chemical compound that retains all the physical and chemical properties of that compound.

  • Molecules can be composed of multiple atoms of the same element (homonuclear structures) or different elements (heteronuclear compounds), provided all bonds between the constituent atoms are covalent.

Comparison of Ionic and Covalent Bonds

• Types of Participating Atoms:

  • Ionic Bonds: Formed between a metal atom and a nonmetal atom.

  • Covalent Bonds: Formed between two or more nonmetal atoms.

• Relative Positions on the Periodic Table:

  • Ionic Bonds: Formed by elements located on opposite sides of the Periodic Table.

  • Covalent Bonds: Formed by elements located on the same side of the Periodic Table (nonmetals on the upper right, along with hydrogen).

• Mechanism of Valence Electron Interaction:

  • Ionic Bonds: Involve the complete transfer of valence electrons from a metal atom to a nonmetal atom, resulting in the formation of positive cations and negative anions.

  • Covalent Bonds: Involve the mutual sharing of valence electron pairs between nonmetal atoms.

Classification of Chemical Compounds

• A chemical formula is a symbolic notation combining element symbols and numerical subscripts to represent the exact number and identity of each atom present in a compound.

• Classification of representative compounds:

  • $NaCl$ (Sodium Chloride): Ionic compound, formed between a metal ($Na$) and a nonmetal ($Cl$).

  • $CH_3COOH$ (Acetic Acid): Covalent compound, composed entirely of nonmetal elements ($C$, $H$, $O$).

  • $MgF_2$ (Magnesium Fluoride): Ionic compound, formed between a metal ($Mg$) and a nonmetal ($F$).

  • $CO_2$ (Carbon Dioxide): Covalent compound, formed between nonmetal elements ($C$ and $O$).

  • $KNO_3$ (Potassium Nitrate): Ionic compound, formed between a metal ($K$) and a polyatomic nonmetal group ($NO_3$).

  • $H_2O$ (Water): Covalent compound, formed between nonmetal elements ($H$ and $O$).

Covalent Bond Types, Bond Length, and Potential Energy

• Structural classification of covalent bonds:

  • Single Covalent Bond:

  • Shared Electron Pairs: $1$ pair of electrons.

  • Total Shared Electrons: $2$ total electrons.

  • Symbolic Notation: Represented as a single line between atoms ($H-H$) or a single pair of Lewis dots ($H:H$).

  • Relative Bond Length: Longest of the three bond types.

  • Relative Bond Strength: Weakest of the three bond types.

  • Representative Example: Molecular hydrogen ($H-H$).

  • Double Covalent Bond:

  • Shared Electron Pairs: $2$ pairs of electrons.

  • Total Shared Electrons: $4$ total electrons.

  • Symbolic Notation: Represented as two parallel lines between atoms ($O=O$) or two pairs of Lewis dots ($O::O$).

  • Relative Bond Length: Intermediate (middle) length.

  • Relative Bond Strength: Intermediate (middle) strength.

  • Representative Example: Molecular oxygen ($O=O$).

  • Triple Covalent Bond:

  • Shared Electron Pairs: $3$ pairs of electrons.

  • Total Shared Electrons: $6$ total electrons.

  • Symbolic Notation: Represented as three parallel lines between atoms ($N \equiv N$) or three pairs of Lewis dots ($N:::N$).

  • Relative Bond Length: Shortest of the three bond types.

  • Relative Bond Strength: Strongest of the three bond types.

  • Representative Example: Molecular nitrogen ($N \equiv N$).

• Definitions of Bond Length:

  • Physical Distance Definition: The physical separation distance measured between the centers of two bonded atomic nuclei.

  • Energy State Definition: The specific internuclear distance at which the total potential energy of the bonded atomic system reaches its absolute minimum value.

• System Potential Energy dynamics relative to internuclear distance:

  • Optimal Distance ("Just Right Distance"): At the equilibrium bond length, two covalently bonded atoms exist at an optimal separation distance. Each nucleus exerts an attractive electrostatic pull on the shared pair of valence electrons without the positive nuclei repelling one another. System potential energy is at its lowest possible point.

  • Distances Too Short (Atoms Too Close): Forcing atoms closer than their equilibrium bond length causes potential energy to rise rapidly due to electrostatic repulsion between the positively charged protons in both nuclei.

  • Distances Too Long (Atoms Too Far Apart): Separating atoms beyond their equilibrium bond length causes potential energy to increase because the separation distance prevents either nucleus from attracting the valence electrons of the opposing atom.

Physical and Chemical Properties of Ionic vs. Covalent Compounds

• Comparative Bond Strength and Electrostatic Attraction:

  • Ionic compounds feature strong electrostatic forces of attraction between oppositely charged positive cations and negative anions.

  • Covalent compounds feature comparatively weaker attractions holding discrete molecules together compared to ionic lattice forces.

• Melting Points and Boiling Points:

  • Ionic Compounds: Require significant thermal energy to overcome strong electrostatic ionic bonds, resulting in high melting points and high boiling points.

  • Covalent Compounds: Require substantially less energy to overcome intermolecular attractions between molecules, resulting in lower melting points and lower boiling points.

• Electrical Conductivity:

  • Ionic Compounds: Conduct electricity when melted (molten) or dissolved in water (aqueous solution) because charged ions are free to move.

  • Covalent Compounds: Do not conduct electricity in any state because they contain no free ions or charged moving particles.

• Physical Appearance:

  • Ionic Compounds: Typically exhibit a highly ordered, crystalline lattice appearance.

  • Covalent Compounds: Generally do not display a crystalline appearance (commonly existing as liquids, gases, or non-crystalline solids at room temperature), though rare crystalline covalent network structures exist.

• Water Solubility:

  • Ionic Compounds: Typically display high solubility in water.

  • Covalent Compounds: Generally display lower solubility in water compared to ionic compounds.

Covalent Nomenclature Rules and Prefixes

• Standard Rules for Naming Binary Covalent Compounds:

  • Order of Elements: Name elements in the order they appear from left to right across the Periodic Table. The element situated furthest to the left is named first.

  • Second Element Suffix: The second element is named using its root name modified with the suffix -ide.

  • Irregular element suffix forms:

    • Oxygen becomes oxide

    • Nitrogen becomes nitride

    • Hydrogen becomes hydride

  • Use of Greek Prefixes: Attach numerical Greek prefixes to element names to specify the exact number of atoms of each element present in a molecule.

  • Exception for First Element (mono- rule): Never start the name of a covalent compound with the prefix mono-. Omit the prefix entirely if only one atom of the first element is present (e.g., $CO$ is Carbon monoxide, not Monocarbon monoxide).

  • Subscripts in Chemical Formulas: Use numerical subscripts following an element's chemical symbol to show how many atoms of that type are in a molecule. Do not write a subscript for $1$.

• Greek Prefixes for Covalent Compounds:

  • $1$ = mono-

  • $2$ = di-

  • $3$ = tri-

  • $4$ = tetra-

  • $5$ = penta-

  • $6$ = hexa-

  • $7$ = hepta-

  • $8$ = octa-

  • $9$ = nona-

Diatomic Elements

• Concept of Diatomic Molecules: Certain nonmetal elements are chemically too reactive to exist as single, unbonded atoms in nature. When not bonded to other elements, two identical atoms of these elements bond covalently to form stable diatomic molecules.

• The "Hydrogen and the 7" Rule: The seven diatomic elements include Hydrogen along with seven elements that form a numerical "7" shape on the Periodic Table starting at element number $7$ (Nitrogen).

• List of Diatomic Formulas:

  • Hydrogen: $H_2$

  • Nitrogen: $N_2$

  • Oxygen: $O_2$

  • Fluorine: $F_2$

  • Chlorine: $Cl_2$

  • Bromine: $Br_2$

  • Iodine: $I_2$

• Naming Diatomic Molecules: When naming an uncombined diatomic element, state the element's standard name without adding numerical prefixes (e.g., $H_2$ is named Hydrogen).

Comprehensive Compound Naming and Formula Practice

• Practice Binary Covalent Compound Names and Formulas:

  • $BF_3$: Boron trifluoride

  • Nitrogen monoxide: $NO$

  • $N_2O_3$: Dinitrogen trioxide

  • Carbon tetraiodide: $CI_4$

  • $IF_5$: Iodine pentafluoride

  • Phosphorus trichloride: $PCl_3$

  • Diphosphorus pentoxide: $P_2O_5$

  • Dinitrogen hexasulfide: $N_2S_6$

  • $CH_4$: Carbon tetrahydride

  • Sulfur hexafluoride: $SF_6$

  • $CO$: Carbon monoxide

  • Hydrogen chloride: $HCl$

  • $NO_2$: Nitrogen dioxide

  • Dihydrogen monoxide: $H_2O$

  • $SCl_6$: Sulfur hexachloride