BackStructure and Properties of Ionic and Covalent Compounds
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Chemical Bonding
Definition and Nature of Chemical Bonds
Chemical bonding refers to the force of attraction between any two atoms in a compound. This attractive force overcomes the repulsion of the positively charged nuclei of the two atoms participating in the bond. Interactions involving valence electrons are responsible for the chemical bond.
Chemical bond: The force that holds atoms together in compounds.
Valence electrons: Electrons in the outermost shell of an atom, which participate in bonding.
Octet rule: Atoms tend to gain, lose, or share electrons to achieve a stable configuration of eight valence electrons.
Lewis Symbols
Representation of Atoms and Valence Electrons
Lewis symbols are a way to represent atoms using the element symbol and dots for valence electrons. Only valence electrons participate in bonding, making it easier to work with the octet rule. The number of dots corresponds directly to the number of valence electrons in the outermost shell.
Lewis symbol: Element symbol surrounded by dots representing valence electrons.
Each of the four sides around the atomic symbol can have up to two dots, for a maximum of eight (octet).
Unpaired dots indicate electrons available for bonding.
Writing Lewis Symbols
Place one dot on each side until there are four dots around the symbol.
Add a second dot to each side in turn.
The number of valence electrons limits the number of dots placed.
Each unpaired dot is available to form a chemical bond.
Lewis Symbols for Representative Elements
The periodic table can be used to determine the number of valence electrons for representative elements, which is reflected in their Lewis symbols.
Group | Lewis Symbol Example |
|---|---|
IA (1) | Na· |
IIA (2) | Mg·· |
VIIA (17) | Cl······· |
VIIIA (18) | Ne········ |
Principal Types of Chemical Bonds
Ionic and Covalent Bonds
There are two principal types of chemical bonds: ionic and covalent. Some bonds have characteristics of both types and are not easily identified as one or the other.
Ionic bond: Attractive force due to the transfer of one or more electrons from one atom to another. The attraction is due to the opposite charges of the resulting ions.
Covalent bond: Attractive force due to the sharing of electrons between atoms.
Ionic Bonding
Formation and Properties
Representative elements form ions that obey the octet rule. Electrons are lost by a metal and gained by a nonmetal, resulting in each atom achieving a noble gas configuration. The resulting cation and anion are attracted to each other, creating the ionic bond.
Metals tend to form cations (positive ions).
Nonmetals tend to form anions (negative ions).
Ions of opposite charge attract each other, forming ionic compounds.
Ionic Bonding Example: NaCl
Consider the formation of sodium chloride (NaCl):
Sodium (Na) has a low ionization energy and readily loses its electron:
Chlorine (Cl) has a high electron affinity and readily gains an electron:
When sodium loses its electron, it gains the neon configuration; when chlorine gains an electron, it gains the argon configuration.
The resulting ions aggregate into a crystal lattice.
Element | Process | Resulting Ion |
|---|---|---|
Na | Loses 1 electron | Na+ |
Cl | Gains 1 electron | Cl- |
Essential Features of Ionic Bonding
Metals have low ionization energies and low electron affinities; they form cations.
Nonmetals have high ionization energies and high electron affinities; they form anions.
Ionic bonds are formed by the transfer of electrons and the resulting electrostatic attraction between ions.
Covalent Bonding
Formation and Properties
Covalent bonds form between atoms with similar tendencies to gain or lose electrons. Compounds containing covalent bonds are called covalent compounds or molecules. Diatomic elements (e.g., H2, O2, N2, F2, Cl2, Br2, I2) have completely covalent bonds with equal sharing.
Atoms share electrons to achieve noble gas configurations.
The shared electron pair is called a covalent bond.
Covalent compounds exist as discrete molecules in solid, liquid, and gas states.
Examples of Covalent Bonding
Water:
Methane:
Polar Covalent Bonding and Electronegativity
Polar Covalent Bonds
Polar covalent bonds are formed when electron pairs are unequally shared between atoms. The difference in electronegativity between the atoms determines the extent of bond polarity.
Electronegativity: A measure of the ability of an atom to attract electrons in a chemical bond.
Greater difference in electronegativity leads to greater bond polarity.
Example: In HF, electrons spend more time with fluorine, making it partially negative and hydrogen partially positive.
Bond | Electronegativity Difference | Polarity |
|---|---|---|
H-F | 4.0 - 2.1 = 1.9 | More polar |
H-Cl | 3.0 - 2.1 = 0.9 | Less polar |
Naming Compounds and Writing Formulas
Nomenclature Systems
Ionic compounds: Name the cation followed by the anion; the anion suffix is changed to -ide.
Covalent compounds: Use prefixes to indicate the number of each kind of atom (mono-, di-, tri-, tetra-, penta-, etc.).
Prefix | Number of Atoms |
|---|---|
Mono- | 1 |
Di- | 2 |
Tri- | 3 |
Tetra- | 4 |
Penta- | 5 |
Hexa- | 6 |
Hepta- | 7 |
Octa- | 8 |
Nona- | 9 |
Deca- | 10 |
Formulas of Compounds
The formula represents the fundamental compound using chemical symbols and numerical subscripts.
The number and type of atoms are identified by the formula.
For ionic compounds, the formula is the smallest whole-number ratio of ions that results in a net charge of zero.
Properties of Ionic and Covalent Compounds
Physical State
Ionic compounds are usually solids at room temperature.
Covalent compounds can be solids, liquids, or gases.
Melting and Boiling Points
Ionic compounds have much higher melting and boiling points than covalent compounds.
Large amounts of energy are required to break the electrostatic attractions between ions.
Structure in the Solid State
Ionic compounds are crystalline.
Covalent compounds may be crystalline or amorphous (no regular structure).
Electrolytes and Nonelectrolytes
Ionic compounds often dissolve in water and dissociate into ions, making the solution conductive (electrolytes).
Covalent compounds usually do not dissociate and do not conduct electricity (nonelectrolytes).
Property | Ionic | Covalent |
|---|---|---|
Composition | Metal + nonmetal | 2 nonmetals |
Electron Sharing | Transfer | Sharing |
Physical State | Solid, crystalline | Any; crystal or amorphous |
Dissociation | Yes (electrolytes) | No (nonelectrolytes) |
Melting/Boiling Point | High | Low |
Drawing Lewis Structures
Guidelines for Lewis Structures
Write the skeletal structure with the least electronegative atom in the center.
Hydrogen and halogens usually occupy terminal positions.
Determine the total number of valence electrons for all atoms in the compound.
Connect the central atom to surrounding atoms with single bonds.
Complete octets for terminal atoms, then for the central atom.
If not enough electrons for the central atom's octet, form multiple bonds as needed.
Lewis Structure Example: CO2
Carbon is the central atom (less electronegative than oxygen).
Total valence electrons:
Connect C to each O with single bonds, then complete octets.
Form double bonds to satisfy the octet rule for carbon.
Bond Energy, Bond Length, and Resonance
Bond Energy and Bond Length
Bond energy: The energy required to break a bond between two atoms.
Triple bond > double bond > single bond (in terms of bond energy).
Bond length: The distance between the nuclei of two bonded atoms.
Single bond > double bond > triple bond (in terms of bond length).
Resonance Structures
Some molecules can be represented by more than one valid Lewis structure.
Resonance structures are average or hybrid representations of these possibilities.
Example: Carbonate ion (CO32-) has three resonance structures.
Exceptions to the Octet Rule
Incomplete Octet
Some atoms (e.g., Be in BeH2) have less than eight electrons around them.
Odd Electron Compounds
Compounds with an odd number of valence electrons (e.g., NO) cannot satisfy the octet rule for all atoms.
Expanded Octet
Elements in the third period or below can have more than eight electrons (e.g., PF5).
VSEPR Theory and Molecular Geometry
Valence Shell Electron Pair Repulsion (VSEPR) Theory
VSEPR theory is used to predict the shape of molecules. Electron pairs around the central atom arrange themselves as far apart as possible to minimize repulsion.
Linear: 2 bonded atoms, 180° bond angle (e.g., BeH2).
Trigonal planar: 3 bonded atoms, 120° bond angle (e.g., BF3).
Tetrahedral: 4 bonded atoms, 109.5° bond angle (e.g., CH4).
Trigonal pyramidal: 3 bonded atoms + 1 lone pair, 107° bond angle (e.g., NH3).
Bent: 2 bonded atoms + 2 lone pairs, 104.5° bond angle (e.g., H2O).
Bonded Atoms | Lone Pairs | Bond Angle | Molecular Geometry | Example |
|---|---|---|---|---|
2 | 0 | 180° | Linear | CO2 |
3 | 0 | 120° | Trigonal planar | BF3 |
4 | 0 | 109.5° | Tetrahedral | CH4 |
3 | 1 | 107° | Trigonal pyramidal | NH3 |
2 | 2 | 104.5° | Bent | H2O |
Molecular Polarity
Determining Polarity
Draw the Lewis structure and determine the geometry.
If the molecule has no lone pairs on the central atom and all terminal atoms are the same, it is nonpolar.
Molecules with one or more lone pairs on the central atom are usually polar.
Polar molecules align in an electric field and behave as dipoles, with one end positive and the other negative.
Intermolecular Forces and Physical Properties
Intramolecular vs. Intermolecular Forces
Intramolecular forces: Attractive forces within molecules (chemical bonds).
Intermolecular forces: Attractive forces between molecules, determining many physical properties.
Solubility
"Like dissolves like": Polar molecules are most soluble in polar solvents; nonpolar molecules are most soluble in nonpolar solvents.
Example: Ammonia (NH3) dissolves in water because both are polar.
Oil and water do not mix because oil is nonpolar and water is polar.
Boiling and Melting Points
Energy is required to overcome intermolecular forces, leading to phase changes.
Stronger intermolecular forces result in higher melting and boiling points.
Larger molecules have higher melting and boiling points than smaller ones.
Polar molecules have higher melting and boiling points than nonpolar molecules of similar mass.
Compound | Bonding Type | Melting Point (°C) | Boiling Point (°C) |
|---|---|---|---|
Sodium chloride (NaCl) | Ionic | 801 | 1413 |
Potassium bromide (KBr) | Ionic | 734 | 1435 |
Water (H2O) | Covalent | 0 | 100 |
Methane (CH4) | Covalent | -182 | -161 |
Additional info: These notes are based on textbook slides and cover foundational concepts in chemical bonding, molecular structure, and physical properties relevant to General Chemistry.
