Chapter 2: Chemical Compounds and Bonding

Section 2.1 - Chemical Bonds and Ionic Bonding

Introduction to Bonding

• Bonding: electrostatic attraction between two atoms or ions.

• Formation of compounds results in a more stable, lower energy state.

Ionic Bonding

• Ionic Bond: formed by the complete transfer of electrons from one atom to another, resulting in ions.

• the electrostatic force between of attraction between a positive and negative ion.

• Isoelectronic: metals are isoelectronic to the noble gas with the smaller atomic number. metals are isoelectronic to the noble gas with the larger electronic number. Ex. sodium ion is isoelectronic to neon.

• Electrostatic attractive forces: ions have opposites charges that are attracted to each other.

• forming ions=different properties compared to initial properties

  • typically solids at room temp - referred to as salts

  • can also exist as crystals - crystal lattice.

• Formula Unit: the smallest unit of an ionic compound. Ex. calcium and chloride combine in a 1 to 2 ratio, the formula unit is CaCl2.

• Properties:

  • high melting and boiling points

  • conduct electricity when dissolved in water or melted, as the ions are free to move. - electrolytes

  • hard and brittle in their solid state due to the strong ionic bonds that hold the ions in fixed positions.

Section 2.2 - Covalent Bonding

Covalent Bonding

• Covalent Bond: involves sharing of electrons between atoms rather than complete transfer.

  • Molecular compound: a substance formed when two or more nonmetals share electrons.

• Molecule: the smallest unit of a covalent compound.

• Atoms can form diatomic and triatomic molecules (e.g., O2 and N2).

Double bond= 2 pairs of electrons shared between 2 atoms

Triple bond= 3 pairs of electrons shared between two atoms.

• Bonding capacity: determines how many bonds it will form.

• Exceeding the Octet Rule: third period and below can have expanded octets.

  • second period cannot.

• Properties of Covalent compounds:

  • all three states

  • lower melting and boiling points

  • do not produce conducting solutions

Section 2.3 - Electronegativity

Electronegativity

• Describes an atom's ability to attract bonding electrons.

• Electronegativity increases across a period as number of protons increases and decreases down a group due to shielding effect.

Polar Covalent Bonds

• In some covalent bonds, electrons are not shared equally due to differences in electronegativity.

• The atom with higher electronegativity gains a partial negative charge because they attract the shared electrons more strongly, while the other gets a partial positive charge (dipole formation).

• EN difference determines how polar a bond is.

• 0.0-0.3=non polar covalent

• 0.4-1.9 = polar

• 2.0-3.3= ionic

Intermolecular Forces

• Intramolecular Forces: attractions within molecules (ionic and covalent).

• Intermolecular Forces: attractions between molecules; includes London dispersion forces, dipole-dipole interactions, and hydrogen bonding.

• Hydrogen bonding is a strong force between molecules with significant partial charges, especially in polar compounds like water.

Properties of Water

• High melting and boiling points due to hydrogen bonding and lots of energy required to break them.

• Density of ice is less than liquid water due to open lattice structure from hydrogen bonds.