Chemical Bonds: Structure, Types, and Properties (CHEM 110 Study Notes)

Chemical Bonds

Introduction to Chemical Bonds

Chemical bonds are the forces that hold atoms together in molecules and ions in crystals. The vast diversity of chemical compounds arises from the ability of atoms to form bonds with many other types of atoms. Chemical bonds determine the three-dimensional shape of molecules and influence the physical and chemical properties of substances.

• Key Point: Chemical bonds are essential for the existence and properties of matter.

• Key Point: The arrangement of electrons, especially valence electrons, is fundamental to chemical bonding.

• Example: Carbon forms graphite (in pencils) and diamond, demonstrating how bonding affects material properties.

Valence Electrons and the Octet Rule

Valence Electrons

Valence electrons are the electrons in the outermost shell of an atom. They are responsible for chemical bonding and determine the reactivity of elements.

• Key Point: The number of valence electrons for main group elements equals the group number in the periodic table.

• Example: Sodium (Na) has 1 valence electron; Chlorine (Cl) has 7 valence electrons.

The Octet Rule

The octet rule states that atoms tend to gain, lose, or share electrons to achieve a stable configuration of eight valence electrons, similar to noble gases.

• Key Point: Metals lose electrons to form cations; nonmetals gain electrons to form anions.

• Key Point: The resulting ions have noble gas electron configurations.

• Example: Sodium loses one electron to become Na+; Chlorine gains one electron to become Cl-.

• Equation: $\text{Na} \rightarrow \text{Na}^+ + e^-$

• Equation: $\text{Cl} + e^- \rightarrow \text{Cl}^-$

Lewis Symbols and Structures

Lewis Symbols

Lewis symbols visually represent valence electrons as dots around the chemical symbol of an atom. They are useful for predicting bonding behavior.

• Key Point: Dots represent valence electrons; the symbol represents the atomic core.

• Example: Na: • (one dot); Cl: ••••••• (seven dots)

Lewis Structures for Molecules and Ions

Lewis structures show how atoms share or transfer electrons to achieve stable configurations. They are used for both simple molecules and polyatomic ions.

• Key Point: Lone pairs and bonding pairs are indicated by dots and lines, respectively.

• Example: Water: H:O:H or H—O—H

Ionic Bonds

Formation of Ionic Bonds

Ionic bonds are formed by the complete transfer of electrons from a metal to a nonmetal, resulting in oppositely charged ions that attract each other.

• Key Point: Ionic compounds are held together by electrostatic forces and exist as crystalline solids.

• Equation: $\text{Na} + \text{Cl} \rightarrow \text{Na}^+ + \text{Cl}^- \rightarrow \text{NaCl}$

• Example: Sodium chloride (NaCl), calcium chloride (CaCl2), potassium oxide (K2O)

Properties of Ions

• Cations: Positive ions formed by loss of electrons (e.g., Na+, Mg2+)

• Anions: Negative ions formed by gain of electrons (e.g., Cl-, O2-)

• Isoelectronic: Ions with the same electron configuration (e.g., Cl- and Ar)

Naming and Writing Formulas for Ionic Compounds

Binary ionic compounds consist of two elements: a metal and a nonmetal. The formula is written by balancing charges using the crossover method.

• Key Point: The cation is named first, followed by the anion (with -ide suffix).

• Example: NaCl = sodium chloride; MgO = magnesium oxide

• Equation: $\text{Ca}^{2+} + 2\text{Cl}^- \rightarrow \text{CaCl}_2$

Covalent Bonds

Formation of Covalent Bonds

Covalent bonds are formed when two nonmetal atoms share one or more pairs of electrons. This sharing allows each atom to achieve a stable electron configuration.

• Key Point: Covalent bonds can be single, double, or triple, depending on the number of shared electron pairs.

• Example: H2 (single bond), O2 (double bond), N2 (triple bond)

Naming Binary Covalent Compounds

Binary covalent compounds are named using prefixes to indicate the number of atoms of each element.

Prefix	Number of Atoms
Mono-	1
Di-	2
Tri-	3
Tetra-	4
Penta-	5
Hexa-	6
Hepta-	7
Octa-	8
Nona-	9
Deca-	10

• Example: CO2 = carbon dioxide; P2O3 = diphosphorus trioxide

Bond Polarity and Electronegativity

Electronegativity

Electronegativity is a measure of an atom’s ability to attract electrons in a chemical bond. Differences in electronegativity determine bond polarity.

• Key Point: Atoms with higher electronegativity attract electrons more strongly.

• Example: Oxygen is more electronegative than hydrogen, making O-H bonds polar.

Polar and Nonpolar Covalent Bonds

The difference in electronegativity between two atoms determines whether a bond is nonpolar covalent, polar covalent, or ionic.

Electronegativity Difference	Type of Bond
< 0.5	Nonpolar covalent
0.5 – 2.0	Polar covalent
> 2.0	Ionic

• Example: H-Cl is polar covalent; C-H is nonpolar covalent.

HONC Rule for Covalent Bonding

Predicting Number of Bonds

The HONC rule helps predict the number of covalent bonds formed by common nonmetals:

• Hydrogen (H): 1 bond

• Oxygen (O): 2 bonds

• Nitrogen (N): 3 bonds

• Carbon (C): 4 bonds

Polyatomic Ions

Definition and Examples

Polyatomic ions are groups of covalently bonded atoms with an overall charge. They participate in both ionic and covalent bonding.

• Key Point: Parentheses are used in formulas to indicate multiple polyatomic ions.

• Example: Calcium nitrate: Ca(NO3)2; Ammonium sulfate: (NH4)2SO4

Lewis Structures: Rules and Applications

Steps for Drawing Lewis Structures

Lewis structures are drawn using a systematic approach:

1. Count total valence electrons.

2. Sketch a skeletal structure.

3. Place electrons as lone pairs around outer atoms to fulfill the octet rule.

4. Subtract used electrons; place remaining electrons around the central atom.

5. If the central atom lacks an octet, convert lone pairs to bonding pairs as needed.

• Example: Methanol (CH3OH) has 14 valence electrons.

Free Radicals

Definition and Examples

A free radical is an atom or molecule with an unpaired electron. Free radicals are highly reactive and play important roles in chemical and biological processes.

• Example: NO, NO2, ClO

Molecular Shapes: VSEPR Theory

Valence Shell Electron Pair Repulsion (VSEPR) Theory

VSEPR theory predicts the shape of molecules and polyatomic ions based on the repulsion between electron pairs around the central atom.

• Key Point: Electron pairs arrange themselves as far apart as possible.

• Example: Linear (180°), Trigonal planar (120°), Tetrahedral (109.5°)

Number of Electron Sets	Geometry	Bond Angle
2	Linear	180°
3	Trigonal planar	120°
4	Tetrahedral	109.5°

Molecular Polarity

Conditions for Molecular Polarity

A molecule is polar if it contains polar bonds and the arrangement of bonds leads to a separation of charge (dipole moment).

• Key Point: Symmetrical molecules may be nonpolar even if they contain polar bonds.

• Example: Methane (CH4) is nonpolar; Water (H2O) is polar.

Summary Table: Types of Chemical Bonds

Bond Type	Formation	Example	Properties
Ionic	Transfer of electrons	NaCl	Crystalline, high melting point
Covalent	Sharing of electrons	H2O	Molecular, low melting point
Polar Covalent	Unequal sharing	HCl	Dipole moment
Nonpolar Covalent	Equal sharing	CH4	No dipole moment

Additional info:

• Some context and examples were inferred to clarify fragmented points and ensure completeness.

• Tables were reconstructed based on standard chemistry knowledge.