Bonding: Ionic and Covalent Compounds Study Notes
UNIT 4 BONDING: IONIC AND COVALENT COMPOUNDS
Overview of Bonding
Bonding is defined as the joining of two atoms in a stable arrangement.
There are two main varieties of bonding:
Ionic Bonds: Result from the transfer of electrons from one element to another.
Covalent Bonds: Result from the sharing of electrons between two atoms.
Ionic versus Covalent Bonding
Ionic Bonds Form Between:
A metal (located on the left side of the periodic table).
A nonmetal (located on the right side of the periodic table).
Covalent Bonds: Formed when two nonmetals combine.
A molecule is defined as a discrete group of atoms that share electrons.
Example: NaCl represents the ionic bond between sodium (Na) and chlorine (Cl).
Valence Electrons and Bonding
Valence Electrons: Electrons available for bonding, located in the outermost shell (the highest principal quantum number, n) of an atom.
Configuration of various elements shows the number of valence electrons:
Helium (He): $1s^2$ – 2 valence electrons
Neon (Ne): $1s^2 2s^2 2p^6$ – 8 valence electrons
Argon (Ar): $1s^2 2s^2 2p^6 3s^2 3p^6$ – 8 valence electrons
Oxygen (O): $1s^2 2s^2 2p^4$ – 6 valence electrons
Sulfur (S): $1s^2 2s^2 2p^6 3s^2 3p^4$ – 6 valence electrons
Manganese (Mn): $1s^2 2s^2 2p^6 3s^2 3p^6 4s^2 3d^5$ – 2 valence electrons
Lutetium (Lu): $1s^2 2s^2 2p^6 3s^2 3p^6 4s^2 3d^{10} 4p^6 5s^2 4d^{10} 5p^6 6s^2 4f^{14} 5d^1$ – 2 valence electrons
Group Trends:
Elements within the same group share the same number of valence electrons.
General formula for determining the number of valence electrons: Group # 1A – 8A corresponds to the number of valence electrons (except for Helium, which has 2).
Noble Gas Configuration
Noble Gases: Highly chemically stable; they do not readily react with other substances. Their valence electron configurations are:
Helium (He): 2
Neon (Ne): 8
Argon (Ar): 8
Krypton (Kr): 8
Xenon (Xe): 8
Radon (Rn): 8
Ionic Compounds
Formation of Ionic Compounds:
Ionic compounds consist of oppositely charged ions exerting a strong electrostatic attraction.
Ions: Charged species where the numbers of protons and electrons differ.
Cations: Positively charged ions, formed when a metal loses electrons.
Anions: Negatively charged ions, formed when a nonmetal gains electrons.
By achieving a filled outer shell, each ion becomes more stable based on the Octet Rule:
Octet Rule: Atoms gain, lose, or share enough electrons to achieve the electron configuration of a noble gas (i.e., typically 8 electrons).
Examples of Cations and Anions
Sodium Ion Formation (Na+):
Neutral sodium atom: 11 protons (11p), 11 electrons (11e–).
After loss of 1 electron: Na $^{+}$ = 11p + 10e–, configuration → 1s²2s²2p⁶ (isolectronic with Ne).
Magnesium Ion Formation (Mg²+):
Neutral magnesium atom: 12 protons, 12 electrons.
After loss of 2 electrons: Mg $^{2+}$ = 12p + 10e–, configuration → 1s²2s²2p⁶.
Size of Ions
Cations are Smaller than Neutral Atoms due to lost outermost electrons:
Example:
Sodium ion (Na+) radius: 102 pm.
Lithium ion (Li+): radius decreases.
Anions are Larger than Neutral Atoms due to gained electrons:
Example: Fluoride Ion (F−) radius: 133 pm.
Electron Gain and Loss Rules
Metals typically lose electrons to form cations.
Nonmetals typically gain electrons to form anions.
The number of gained or lost electrons rarely exceeds 3.
Isoelectronic Species
Defined as two or more species containing the same number of electrons:
Cations become isoelectronic with the preceding noble gas.
Anions become isoelectronic with the following noble gas.
Formation of Ionic Compounds Explained
Ionic bonds arise due to the electron configurations of ions, as in the following example:
Lithium (Li) gives away an electron to become Li⁺.
Fluorine (F) gains an electron to become F−. Both achieve stable configurations.
Group Charge Patterns
Group 1A (1): Charge = +1.
Group 2A (2): Charge = +2.
Group 3A (3): Charge = +3.
Group 6A (16): Charge = −2 (obtained from 8).
Group 7A (17): Charge = −1.
Determining Ion Charges
Cation Charge: Equal to the group number.
Anion Charge: Calculated as group number minus 8 (or 18).
Polyatomic Ions
Defined as charged entities composed of multiple atoms, with formulas and charges needing memorization:
Example:
Ammonium ion (NH₄⁺)
Hydroxide ion (OH⁻)
Writing and Naming Formulas for Ionic Compounds
Generic Steps for Ionic Formulas:
Identify the positive ion (cation) and the negative ion (anion).
Ensure charge balance by adjusting ratios of ions.
Naming Cations:
For main group cations, use the element name followed by "ion".
Example: Li → Lithium ion.
Naming Anions:
Use the stem of the nonmetal name with the suffix –ide.
Example: Cl⁻ → Chloride ion.
Writing and Naming Compounds Containing Polyatomic Ions
Naming follows simple conventions based on the type of ions involved and may include Roman numerals for variable charge metals.
First Type (Metal + Nonmetal): Use name of the metal followed by name of the non-metal with –ide suffix.
Second Type (Polyatomic): Use the entire name of the polyatomic ion.
Molecular Orientation and Polarity
Bonding shapes affect molecule polarity stemming from charge distribution.
Molecules with asymmetrical charge distribution or containing polar bonds will be polar; symmetrical distributions will result in nonpolar molecules.
Conclusion: Key Differences Between Ionic and Covalent Compounds
**Ionic Compounds: ** Formed through electron transfer, typically between metals and nonmetals, represented by empirical formulas.
Covalent Compounds: Formed through electron sharing, typically between nonmetals, represented by molecular formulas (indicating the actual number of atoms).