Chemical Bonding and Ionic Solids

Vocabulary practice flashcards covering electronegativity, bond types, the octet rule, ionic bonding, crystal lattice structures, and physical properties of ionic solids.

Electronegativity (E.N)

The tendency of an atom to attract electrons from a neighbouring atom.

The Octet Rule

The rule stating that atoms bond to achieve the electron configuration of their nearest noble gas, becoming most stable when they have $8$ electrons in their valence shell (hydrogen and helium are exceptions because they only need $2$ electrons).

Electronegativity Difference (E.N.D)

The calculated difference obtained by subtracting the smaller electronegativity value from the larger electronegativity value, used to determine the type of chemical bond between atoms.

Ionic Bond (E.N.D)

A chemical bond formed when the Electronegativity Difference (E.N.D) between bonded atoms is greater than $1.8$.

Polar Covalent Bond (E.N.D)

A chemical bond formed when the Electronegativity Difference (E.N.D) between bonded atoms is between $0.4$ and $1.8$.

Pure Covalent Bond (E.N.D)

A chemical bond formed when the Electronegativity Difference (E.N.D) between bonded atoms is less than $0.4$.

END Scale and Bond Type

A continuum scale showing that an Electronegativity Difference (E.N.D) below $0.4$ forms a nonpolar covalent bond, an E.N.D between $0.4$ (least polar) and $1.8$ (most polar) forms a polar covalent bond, and an E.N.D above $1.8$ forms an ionic bond.

Ionic Bonding

A type of bonding usually occurring between a metal and a nonmetal, where the metal loses one or more electrons to form a positive ion (cation) and the nonmetal gains one or more electrons to form a negative ion, allowing both to achieve a stable full octet.

Cation

A positive ion formed when a metal loses one or more electrons to achieve a stable full octet.

Negative Ion

An ion formed when a nonmetal gains one or more electrons to achieve a stable full octet.

Crystal Lattice

A $3\text{D}$ pattern of alternating positive and negative charges that are strongly attracted to each other in an ionic solid.

Formula Unit

The smallest whole number ratio of ions present in an ionic crystal, such as $1\text{Na}^+:1\text{Cl}^-$ in $\text{NaCl}$.

Melting Point of NaF

The melting temperature of sodium fluoride, which is $993\text{ }^\frac{1}{1}\text{ }^\text{o}\text{C}$ ($993^\text{o}\text{C}$), reflecting the large amount of energy required to break its rigid lattice.

Melting Point of KCl

The melting temperature of potassium chloride, which is $770^\text{o}\text{C}$.

Melting Point of LiCl

The melting temperature of lithium chloride, which is $605^\text{o}\text{C}$.

Brittleness of Ionic Solids

The property of ionic solids breaking along a smooth edge when struck, caused by rearranged ion layers aligning like charges adjacent to one another so that they strongly repel.

Electrical Conductivity of Ionic Solids

Ionic compounds are non-conductors of heat and electricity in solid form, but conduct electricity when molten (e.g., at $800^\text{o}\text{C}$) or dissolved in water (e.g., at $20^\text{o}\text{C}$) because the charged ions are free to move.

Dissociation

The process where ionic compounds split into individual ions when dissolved in water, which are then attracted to polar (partially charged) water molecules.