Ionic and Covalent Compounds Flashcards

Vocabulary practice flashcards covering key concepts from Chapters 3 and 4 on Ionic and Covalent Compounds, nomenclature, bonding rules, periodic trends, Lewis structures, molecular geometry, and bond polarity.

Chemical Bonding

The joining of two atoms in a stable arrangement.

Noble Gases Reactivity

Group 8A elements that do not readily react to form bonds because their electronic configurations are exceptionally stable.

Ionic Bond

A chemical bond resulting from the transfer of valence electrons from one element to another, typically between a metal and a nonmetal.

Covalent Bond

A chemical bond resulting from the sharing of electrons between two atoms.

Molecule

A discrete group of atoms held together by shared electrons or covalent bonds.

Octet Rule

The principle stating that main group elements adjust their valence electrons through gaining, losing, or sharing to attain eight outer shell electrons like the nearest noble gas.

Cation

A positively charged ion formed when a neutral atom loses one or more valence electrons.

Anion

A negatively charged ion formed when a neutral atom gains one or more valence electrons.

Group 1A Ionic Charge

Metals in Group 1A lose 1 electron to form cations with a $+1$ charge.

Group 2A Ionic Charge

Metals in Group 2A lose 2 electrons to form cations with a $+2$ charge.

Group 3A Aluminum Ionic Charge

Aluminum in Group 3A forms a cation with a $+3$ charge ($\text{Al}^{3+}$).

Group 5A Nonmetal Ionic Charge

Nonmetals in Group 5A gain 3 electrons to form anions with a $-3$ charge.

Group 6A Nonmetal Ionic Charge

Nonmetals in Group 6A gain 2 electrons to form anions with a $-2$ charge.

Group 7A Halogen Ionic Charge

Halogens in Group 7A gain 1 electron to form anions with a $-1$ charge.

Main Group Cation Nomenclature

Main group cations are named using the element name from which they are formed (e.g., sodium ion for $\text{Na}^+$).

Variable Charge Metals

Metals, particularly transition metals and high atomic number metals, that can form more than one type of positive cation.

Systematic Cation Naming Method

A naming method that specifies the charge of a variable cation using a Roman numeral in parentheses immediately following the element name.

Common Cation Naming Method

A naming method that uses the suffix '-ous' for a cation with a lower charge and '-ic' for a cation with a higher charge.

Ferrous Ion

The common name for the iron cation with a $+2$ charge ($\text{Fe}^{2+}$), systematically named Iron(II).

Ferric Ion

The common name for the iron cation with a $+3$ charge ($\text{Fe}^{3+}$), systematically named Iron(III).

Cuprous Ion

The common name for the copper cation with a $+1$ charge ($\text{Cu}^+$), systematically named Copper(I).

Cupric Ion

The common name for the copper cation with a $+2$ charge ($\text{Cu}^{2+}$), systematically named Copper(II).

Chromous Ion

The common name for the chromium cation with a $+2$ charge ($\text{Cr}^{2+}$), systematically named Chromium(II).

Chromic Ion

The common name for the chromium cation with a $+3$ charge ($\text{Cr}^{3+}$), systematically named Chromium(III).

Mercurous Ion

The common name for the mercury ion $\text{Hg}_2^{2+}$, systematically named Mercury(I), containing two mercury atoms each with a $+1$ charge.

Mercuric Ion

The common name for the mercury cation with a $+2$ charge ($\text{Hg}^{2+}$), systematically named Mercury(II).

Stannous Ion

The common name for the tin cation with a $+2$ charge ($\text{Sn}^{2+}$), systematically named Tin(II).

Stannic Ion

The common name for the tin cation with a $+4$ charge ($\text{Sn}^{4+}$), systematically named Tin(IV).

Plumbous Ion

The common name for the lead cation with a $+2$ charge ($\text{Pb}^{2+}$), systematically named Lead(II).

Plumbic Ion

The common name for the lead cation with a $+4$ charge ($\text{Pb}^{4+}$), systematically named Lead(IV).

Aurous Ion

The common name for the gold cation with a $+1$ charge ($\text{Au}^+$), systematically named Gold(I).

Auric Ion

The common name for the gold cation with a $+3$ charge ($\text{Au}^{3+}$), systematically named Gold(III).

Monatomic Anion Nomenclature

Monatomic anions are named by replacing the ending of the element name with the suffix '-ide'.

Bromide

The anion formed from bromine, represented as $\text{Br}^-$.

Chloride

The anion formed from chlorine, represented as $\text{Cl}^-$.

Fluoride

The anion formed from fluorine, represented as $\text{F}^-$.

Iodide

The anion formed from iodine, represented as $\text{I}^-$.

Nitride

The anion formed from nitrogen, represented as $\text{N}^{3-}$.

Oxide

The anion formed from oxygen, represented as $\text{O}^{2-}$.

Phosphide

The anion formed from phosphorus, represented as $\text{P}^{3-}$.

Sulfide

The anion formed from sulfur, represented as $\text{S}^{2-}$.

Zero Overall Charge Principle

The rule stating that the sum of positive cation charges and negative anion charges in an ionic compound formula must equal zero.

Crisscross Method

A technique for writing ionic formulas where the numerical charge value of each ion becomes the subscript of the opposite ion, reduced to the simplest whole-number ratio.

Ionic Compound Physical State

Ionic compounds exist as crystalline solids at room temperature due to strong electrostatic attractions between oppositely charged ions.

Sodium Chloride Melting Point

The high temperature at which solid $\text{NaCl}$ melts, equal to $801\,^\circ\text{C}$.

Sodium Chloride Boiling Point

The temperature at which liquid $\text{NaCl}$ boils, equal to $1413\,^\circ\text{C}$.

Aqueous Ionic Conductivity

When ionic compounds dissolve in water, they dissociate into cations and anions, forming an aqueous solution that conducts electricity.

Polyatomic Ion

A cation or anion that contains more than one atom held together by covalent bonds.

Carbonate

A polyatomic anion with the formula $\text{CO}_3^{2-}$.

Hydrogen Carbonate

A polyatomic anion with the formula $\text{HCO}_3^-$, commonly referred to as bicarbonate.

Acetate

A polyatomic anion with the formula $\text{CH}_3\text{CO}_2^-$.

Cyanide

A polyatomic anion with the formula $\text{CN}^-$.

Nitrate

A polyatomic anion with the formula $\text{NO}_3^-$.

Nitrite

A polyatomic anion with the formula $\text{NO}_2^-$.

Hydroxide

A polyatomic anion with the formula $\text{OH}^-$.

Phosphate

A polyatomic anion with the formula $\text{PO}_4^{3-}$.

Hydrogen Phosphate

A polyatomic anion with the formula $\text{HPO}_4^{2-}$.

Dihydrogen Phosphate

A polyatomic anion with the formula $\text{H}_2\text{PO}_4^-$.

Sulfate

A polyatomic anion with the formula $\text{SO}_4^{2-}$.

Hydrogen Sulfate

A polyatomic anion with the formula $\text{HSO}_4^-$, commonly referred to as bisulfate.

Sulfite

A polyatomic anion with the formula $\text{SO}_3^{2-}$.

Hydrogen Sulfite

A polyatomic anion with the formula $\text{HSO}_3^-$, commonly referred to as bisulfite.

Tums Active Ingredient

Calcium carbonate ($\text{CaCO}_3$), an active antacid ingredient in Tums.

Maalox Active Component

Magnesium hydroxide ($\text{Mg(OH)}_2$), an active antacid component in Maalox.

FeSO4 Dietary Use

Iron(II) sulfate ($\text{FeSO}_4$), an iron supplement used to treat anemia.

Metal Properties

Elements that are usually lustrous, malleable, and good conductors of heat and electricity.

Nonmetal Properties

Elements that are usually not lustrous, brittle, and poor conductors of heat and electricity.

Metalloid Properties

Elements that have properties intermediate between metals and nonmetals.

Hydrogen Periodic Exception

An element located in Group 1A that behaves chemically and physically like a nonmetal.

Atomic Radius

The distance from the nucleus to the outer boundary of an atom's electron cloud.

Atomic Radius Group Trend

Atomic radius increases from top to bottom down a group due to the addition of principal energy levels.

Atomic Radius Period Trend

Atomic radius decreases from left to right across a period due to an increased number of protons pulling electrons closer.

Cation Radius vs Neutral Atom Radius

A cation is smaller than its neutral atom because losing electrons removes an outer energy level and reduces electron repulsion.

Anion Radius vs Neutral Atom Radius

An anion is larger than its neutral atom because gaining electrons increases electron-electron repulsion, spreading the orbital out.

Ionization Energy

The amount of energy required to remove an electron from a neutral gaseous atom.

Ionization Energy Group Trend

Ionization energy decreases down a group because outer electrons are farther from the nucleus and less tightly held.

Ionization Energy Period Trend

Ionization energy increases across a period from left to right because atoms are smaller and valence electrons are held more tightly.

Electronegativity

A measure of the ability of an atom in a covalent bond to attract shared electrons toward itself.

Pauling Electronegativity Scale

A scale of relative electronegativities where fluorine is assigned the highest value of $4.0$.

Electronegativity Periodic Trends

Electronegativity increases from left to right across a period and decreases down a main group.

Lewis Structures

Diagrams that represent the valence electrons of atoms as dots surrounding element symbols.

Valence Electrons

Electrons located in the outermost energy level of an atom that participate in chemical bonding.

Main Group Valence Electron Rule

For main group elements (Groups 1A–7A), the group number gives the number of valence electrons.

Lone Pair

An unshared pair of nonbonding electrons belonging exclusively to one atom.

Single Bond

A covalent bond in which two atoms share 1 pair of electrons.

Double Bond

A covalent bond in which two atoms share 2 pairs of electrons.

Triple Bond

A covalent bond in which two atoms share 3 pairs of electrons.

Central Atom Selection Criteria

The central atom in a Lewis structure is never hydrogen, is usually the atom present in the smallest quantity, and is typically the least electronegative element.

Terminal Atoms

Outer atoms in a Lewis structure that are attached to the central atom.

VSEPR Theory

Valence Shell Electron Pair Repulsion theory, used to predict molecular geometry by minimizing electrostatic repulsion between electron pairs around a central atom.

Tetrahedral Geometry

The molecular shape adopted when a central atom has 4 bonding electron groups and 0 lone pairs.

Trigonal Pyramidal Geometry

The molecular shape adopted when a central atom has 3 bonding electron groups and 1 lone pair (e.g., $\text{NH}_3$).

Bent / V-Shaped Geometry

The molecular shape adopted when a central atom has 2 bonding electron groups and 2 lone pairs (e.g., $\text{H}_2\text{O}$).

Trigonal Planar Geometry

The molecular shape adopted when a central atom has 3 bonding electron groups and 0 lone pairs (e.g., $\text{BF}_3$).

Linear Geometry

The molecular shape adopted when a central atom has 2 bonding electron groups and 0 lone pairs (e.g., $\text{CO}_2$).

Prefix Mono-

A prefix used in nomenclature to indicate 1 atom of an element in a covalent molecule.

Prefix Di-

A prefix used in nomenclature to indicate 2 atoms of an element in a covalent molecule.

Prefix Tri-

A prefix used in nomenclature to indicate 3 atoms of an element in a covalent molecule.

Prefix Tetra-

A prefix used in nomenclature to indicate 4 atoms of an element in a covalent molecule.

Prefix Penta-

A prefix used in nomenclature to indicate 5 atoms of an element in a covalent molecule.