Chemical Bonding and Molecular Structure

The Octet Rule

• Atoms seek stability by having eight electrons in their outermost shell (valence shell).
• The octet rule helps predict bonds formed by main group elements (carbon, nitrogen, oxygen, halogens, hydrogen, some metals).
• Covalent bonding allows atoms to share electrons to achieve a full octet.
• Noble gases already possess full valence shells, achieving stability.
• Electronegativity: the tendency of atoms to attract electrons.
  • More electronegative elements (hydrogen, halides) are on the outside of a molecule; less electronegative elements (carbon, oxygen, phosphorus) are on the inside.
  • Fluorine is the most electronegative element.

Electron Orbits

• Elements are grouped based on valence electrons:
  • s-block: first two columns (2 valence electrons)
  • d-block: metals in the middle
  • p-block: last six columns (3-8 valence electrons)
  • f-block: actinide and lanthanide series
• Covalent bonds:
  • Single bond: sharing one pair of electrons
  • Double bond: sharing two pairs of electrons
  • Triple bond: sharing three pairs of electrons
  • Quadruple bond: sharing four pairs of electrons
• An s shell holds two electrons, while a p shell can hold up to six electrons.

Lewis Structures

• Lewis dot structures are diagrams showing bonds between atoms and how they share electrons.
• The elemental symbol is surrounded by its valence electrons.
• Shared electrons are represented by two dots or a line between atoms.
• Electrons are not stationary in reality.

Lewis Structures and the Periodic Table

• s-block elements have one or two valence electrons.
• p-block elements have three to eight valence electrons.
• Noble gases have eight valence electrons (full valence shell).

How to Draw Lewis Structures

1. Use the atomic symbol.
2. The least electronegative element goes in the middle; hydrogen is always on the outside. If there are only two atoms, there is no middle.
3. Draw dots to represent electrons around the symbol.
4. Use a line to represent bonded pairs of electrons (up to four bonds).
5. Consider subscripts and ionic charges.
   • Negative charge: add electrons.
   • Positive charge: remove electrons.

Lewis Structures of Atoms

• Column determines the number of valence electrons (1-7).
• Anything under hydrogen has one, under beryllium has two, under boron has three, carbon and under has four, nitrogen and under has five, under oxygen has six, and under fluorine has seven.

Lewis Structures of Molecules

• Account for all valence electrons when drawing Lewis dot structures for compounds.
• Example: Methane (CH4) has single bonds between carbon and four hydrogens.
• Carbon monoxide (CO) and carbon dioxide (CO2) involve carbon and oxygen atoms.
• Carbon monoxide has a triple bond.
• Carbon dioxide has two sets of double bonds.
• Lone pairs: electrons not involved in bonding.
• Bond pairs: electrons shared in bonds (represented by lines).
• The sum of lone and bond pairs should equal the number of valence electrons.

What Are Covalent Bonds?

• Covalent bonds involve the sharing of electrons between two or more atoms to complete their valence shells.
• Example: Oxygen forms covalent bonds with two hydrogen atoms to complete its outer shell.

How Are Covalent Bonds Formed?

• Electrons orbit the nucleus in shells; the outermost shell is the valence shell.
• The octet rule states that the valence shell should contain eight electrons (except for hydrogen and helium).
• Electronegativity is the attraction of atoms to electrons.
• Atoms with higher electronegativity attract shared electrons more strongly, creating a dipole.
  • The more electronegative atom has a negative dipole.
  • The less electronegative atom has a positive dipole.

Types of Covalent Bonds

• Polar covalent bonds: unequal sharing of electrons due to differences in electronegativity, resulting in dipoles.
• Non-polar covalent bonds: equal sharing of electrons due to similar electronegativity; no dipoles.

Multiple Covalent Bonds

• Single covalent bond: one pair of electrons is shared.
• Double covalent bond: two pairs of electrons are shared.
• Triple covalent bond: three pairs of electrons are shared.
• Carbon dioxide is an example of a double covalent bond.

What is a Covalent Compound?

• A covalent compound forms when two or more different atoms are connected by covalent bonds.
• Atoms typically have similar electronegativity and are non-metals (e.g., carbon, hydrogen, oxygen, nitrogen).
• Carbon dioxide: carbon forms two double bonds with each oxygen atom.
• Water: oxygen forms single bonds with two hydrogen atoms.

Properties of Covalent Compounds

• Low boiling and melting points
• Poor conduction of heat and electricity
• Varied colors
• Low solubility
• High vapor pressure
• Lack of brittleness (except for network solids like diamonds)
• Covalent compounds have weak intermolecular forces, making phase changes easier.
• Covalent compounds are often colorless and are good insulators.
• They tend to have high vapor pressure and ignite easily and they do not dissolve well in water.

Covalent Compounds vs. Ionic Compounds

• Covalent compounds share electrons, while ionic compounds exchange electrons.

Examples of Covalent Compounds

Naming Covalent Compounds

• Start with the element with the lower group number.
• Use prefixes (mono-, di-, tri-, etc.) to indicate the number of atoms.

• The second element's ending is dropped, and "ide" is added.
• Vowels are often dropped for pronunciation.

Ionic Compound Formation

• Metals (e.g., sodium, Na) lose electrons to form cations.
• Nonmetals (e.g., chlorine, Cl) gain electrons to form anions.
• The ionic bond results from the attraction between positive and negative charges.
• Example: Sodium chloride (NaCl) is formed by the transfer of an electron from sodium to chlorine.

Properties of Ionic Compounds

• High melting and boiling points
• Crystal lattice structure
• Solids break easily into sheets
• Solids do not conduct electricity (good insulators)

Ionic Compounds Structure: Crystal Lattice

• Ionic compounds form a rigid, ordered crystal lattice.
• Positive and negative ions align in a strict fashion.
• Shifting the arrangement can break the crystal.

Lattice Energy of Ionic Compounds

• Lattice energy is the energy released when one mole of ions forms a crystal lattice.
• Adding this energy can break the bonds in the solid ionic compound.

Melting and Boiling Point of Ionic Compounds

• Ionic compounds have high melting and boiling points due to strong intermolecular forces.

Liquid and Solution Ionic Compounds: Good Conductors; Solid Ionic Compounds: Good Insulators

• Water dissolves ionic compounds by separating their constituent ions.
• The resulting solution is a good conductor of electricity.
• Melted ionic compounds are also good conductors.
• Dry solid ionic compounds are good insulators.