Valence Electrons and Covalent Bonding

Valence electrons range from one to eight across the periodic table, excluding transition elements.

Lewis Dot Symbols:

• Aluminum (Group 13) has three valence electrons.
• Chlorine has seven valence electrons, needing one more to achieve a noble gas configuration (octet rule).

Ionic Compounds:

• Electrons are transferred to achieve a noble gas configuration (octet).
• Metals lose electrons; nonmetals gain electrons to achieve the nearest noble gas configuration.
• Example: Formation of sodium ($Na$) and chloride ($Cl$) ions. Zero dots represent an octet, matching the noble gas below the element; eight dots represent the noble gas above.

Bonding:

• Metals want to lose electrons to match a noble gas configuration; nonmetals want to gain electrons.
• Magnesium ($Mg$) and oxygen ($O$) combine, generating energy.

Covalent Bonding:

• Atoms share valence electrons to achieve a noble gas configuration (octet rule).
• "Co-" means together; "valent" refers to valence electrons.
• Hydrogen ($H_2$) shares a pair of electrons. Each hydrogen atom contributes one electron, so by sharing, each effectively has two electrons, resembling helium.

Representing Covalent Bonds:

• A dash between two atoms (e.g., H-H) represents a shared pair of electrons (a single bond).

Diatomic Molecules:

• Chlorine ($Cl_2$) has seven valence electrons and needs one more. Two chlorine atoms share a pair of electrons to achieve an octet.
• Fluorine, iodine, and bromine also form diatomic molecules.

Bonding Pairs and Lone Pairs:

• Bonding pair or shared pair: electrons shared between atoms.
• Lone pairs: pairs of electrons not shared, residing only on one atom.

Example: Hydrochloric Acid ($HCl$)

• Hydrogen has one valence electron; chlorine has seven. They share a pair to achieve configurations of helium and argon, respectively.
• Nonmetals don't give away electrons; sharing is the only way to bond.

Double Bonds: Oxygen ($O_2$)

• Oxygen has six valence electrons and needs two more.
• Sharing one pair of electrons results in seven valence electrons per oxygen atom, similar to fluorine but still unstable.
• Oxygen atoms share two pairs of electrons (double bond) using $s$ and $p$ orbitals.

Sharing $p$ orbitals, oxygen atoms get close enough to overlap and share another pair of electrons.

Double bond representation: $O=O$, with two sticks representing two shared pairs of electrons.

Triple Bonds: Nitrogen ($N_2$)

• Nitrogen has five valence electrons and needs three more.
• Nitrogen atoms share three pairs of electrons (triple bond) using $s$ and $p$ orbitals.

Nitrogen overlaps three $p$ orbitals to share three pairs of electrons.

Triple bond representation: $N≡N$, with three sticks representing three shared pairs of electrons.

Carbon can form up to three shared series of electrons.

When atoms share electrons, each nitrogen atom achieves an electron configuration similar to neon.

Lone atoms in space lack opportunities to form bonds and are considered "ugly."

Single bonds: one shared pair of electrons.

Double bonds: two shared pairs of electrons.

Triple bonds: three shared pairs of electrons.

Molecules with more than two atoms achieve octets through specific combinations (e.g., water, ammonia).

Water ($H_2O$)

• Oxygen has six valence electrons, needing two more.
• Each hydrogen shares one electron with oxygen.
• Hydrogen can't form new bounds
• Two hydrogen atoms bond with one oxygen atom, resulting in each hydrogen having two electrons and oxygen having eight.

Ammonia ($NH_3$): Nitrogen needing three more electrons, bonding with three hydrogen atoms.

Atoms bond to become more stable by filling orbitals symmetrically and balancing charge.

Noble Gases: Generally non-reactive because they already have complete valence shells (octets).

Xenon ($Xe$) and radon ($Rn$) can be tricked into forming compounds, but they prefer to exist independently.

Riddle: Where does today come before yesterday? (Answer: In the dictionary)