Chemical Bonds

Chemical bonds hold atoms together in molecules and compounds. The type of bond influences a compound's physical properties, reactivity, and how it interacts with other molecules.

Valence Electrons Valence electrons are the electrons in the outermost shell of an atom. They determine how an atom bonds and reacts. Atoms tend to gain, lose, or share electrons to achieve a stable electron configuration, often a full outer shell.

Classifying Elements: Metals, Metalloids, and Nonmetals Elements are classified as metals, metalloids, or nonmetals based on their physical and chemical properties.

Metals: Good conductors, shiny, malleable, ductile, and readily lose electrons.

Metalloids: Have properties intermediate between metals and nonmetals; often function as semiconductors.

Nonmetals: Poor conductors, brittle (if solid), and tend to gain or share electrons during bonding.

Electronegativity Electronegativity is an atom's ability to attract shared electrons within a chemical bond. The greater the electronegativity, the more strongly an atom attracts bonding electrons.

Fluorine is the most electronegative element.

Electronegativity generally increases from left to right across a period because effective nuclear charge increases. It generally decreases from top to bottom within a group because valence electrons are farther from the nucleus and experience greater shielding.

Ionization Energy Ionization energy is the energy required to remove an electron from a gaseous atom. Elements with low ionization energy lose electrons easily, while elements with high ionization energy hold onto their electrons more tightly.

It increases across the period from left to right as there is increasing effective nuclear charge. It decreases within a group from top to bottom as atomic size increases, creating further distance and shielding between the nucleus and valence electrons.

Bond Polarity and Dipoles Bond polarity depends on how equally electrons are shared between two bonded atoms. When one atom attracts electrons more strongly due to higher electronegativity, the bond becomes polarized. This creates regions of partial positive (δ⁺) and partial negative (δ⁻) charge known as a dipole. The more electronegative atom develops a partial negative charge (δ⁻), while the less electronegative atom develops a partial positive charge (δ⁺). A molecule can contain polar bonds but still be nonpolar overall if its bond dipoles cancel because of its molecular geometry.

Covalent Bonds Covalent bonds form when atoms share electrons. They are the dominant bond type in organic molecules. Covalent bonds can be:

Single bonds: one shared electron pair

Double bonds: two shared electron pairs

Triple bonds: three shared electron pairs

Nonpolar Covalent Bonds Nonpolar covalent bonds occur when two atoms have little or no difference in electronegativity and share electrons equally. Because the electrons are evenly distributed, no partial charges develop across the bond.

Polar Covalent Bonds Polar covalent bonds form when there is a moderate difference in electronegativity, causing electrons to be shared unequally. The more electronegative atom develops a partial negative (δ⁻) charge, while the other atom develops a partial positive (δ⁺) charge.

Examples include: O–H, N–H, H–Cl, and C=O.

Ionic Bonds Ionic bonding is the electrostatic attraction between oppositely charged ions. Ionic compounds commonly form when electrons are transferred from a metal to a nonmetal, producing positively charged cations and negatively charged anions.

Polarity affects a molecule's solubility, boiling point, and how it interacts with other molecules.

Metallic Character and Metallic Bonding Metallic character is an element's tendency to lose or share electrons during chemical reactions. Metals typically have low ionization energies, making it easier for them to give up their valence electrons.

Metallic bonding is the attraction between positively charged metal ions and a shared "sea" of freely moving valence electrons. The sea of electrons describes the pool of delocalized valence electrons shared by all atoms within a metal. These mobile electrons hold metal ions together while allowing electrons to move freely throughout the material. This gives metals their strength, flexibility, and high electrical and thermal conductivity.