Chemical Bonds

Types of Chemical Bonds

Chemical bonds are the forces that hold atoms together in molecules and compounds. The type of bond formed depends on the difference in electronegativity between the atoms involved.

• Covalent Bonds: Atoms share pairs of electrons. If the sharing is equal, the bond is nonpolar; if unequal, the bond is polar.

• Ionic Bonds: Electrons are transferred from one atom to another, resulting in oppositely charged ions that attract each other.

• Hydrogen Bonds: A weak attraction between a hydrogen atom covalently bonded to a more electronegative atom (like oxygen or nitrogen) and another electronegative atom.

Electronegativity is a measure of an atom's ability to attract electrons in a bond. The difference in electronegativity determines bond type:

• Nonpolar Covalent: ΔEN < 0.4

• Polar Covalent: 0.4 ≤ ΔEN ≤ 2.0

• Ionic: ΔEN > 2.0

Example: In water (H2O), the O-H bond is polar covalent because oxygen is much more electronegative than hydrogen, resulting in partial charges (δ- on O, δ+ on H).

Visual Summary Table: Bond Types and Electronegativity

Properties of Water

Structure and Polarity

Water is a polar molecule with a bent shape, resulting in partial positive (δ+) charges on hydrogen atoms and a partial negative (δ-) charge on oxygen. This polarity allows water to form hydrogen bonds with other molecules.

• Hydrogen Bonding: Each water molecule can form up to four hydrogen bonds, leading to unique properties essential for life.

Unique Properties of Water

• High Specific Heat: Water can absorb or release large amounts of heat with little temperature change, moderating Earth's climate and body temperature.

• Cohesion and Adhesion: Water molecules stick to each other (cohesion) and to other surfaces (adhesion), aiding in processes like water transport in plants.

• Expansion Upon Freezing: Ice is less dense than liquid water, allowing it to float and insulate aquatic environments.

• Versatile Solvent: Water dissolves many polar and ionic substances, making it the "universal solvent" for biological reactions.

• Dissociation: Water can dissociate into hydronium (H3O+) and hydroxide (OH-) ions.

Example: Water's high specific heat helps stabilize ocean temperatures, creating a favorable environment for marine life.

Acids, Bases, and pH

pH Scale and Ion Concentrations

The pH scale measures the concentration of hydrogen ions (H+) in a solution. It is defined as:

• Acids: Substances that increase H+ concentration (pH < 7).

• Bases: Substances that decrease H+ concentration, often by increasing OH- (pH > 7).

• Neutral: Pure water has [H+] = M, so pH = 7.

Biological Importance of pH

• Enzyme activity and molecular structure are sensitive to pH changes.

• Organisms use buffers to maintain stable internal pH.

Example: Blood contains bicarbonate buffer to maintain pH around 7.4.

Table: Common Substances and Their pH

Carbon and the Diversity of Life

Properties of Carbon

Carbon is the backbone of organic molecules due to its ability to form four covalent bonds, allowing for a variety of stable structures including chains, rings, and branches.

• Tetravalence: Carbon has four valence electrons, enabling it to bond with up to four other atoms.

• Versatility: Can form single, double, or triple bonds; can bond with many elements (H, O, N, S, P).

Example: Ethane (C2H6) is a simple hydrocarbon with a carbon-carbon single bond.

Functional Groups

Functional groups are specific groups of atoms within molecules that determine the chemical properties and reactions of those molecules.

• Hydroxyl (-OH): Alcohols; polar, forms hydrogen bonds.

• Carbonyl (C=O): Aldehydes and ketones; polar.

• Carboxyl (-COOH): Carboxylic acids; acidic.

• Amino (-NH2): Amines; basic.

• Sulfhydryl (-SH): Thiols; forms disulfide bonds.

• Phosphate (-PO42-): Organic phosphates; involved in energy transfer.

• Methyl (-CH3): Nonpolar; affects gene expression.

Example: The carboxyl group in amino acids gives them acidic properties.

Isomers

Isomers are compounds with the same molecular formula but different structures and properties.

• Structural Isomers: Differ in covalent arrangement of atoms.

• Stereo-isomers: Same covalent arrangement but differ in spatial arrangement.

• Enantiomers: Mirror-image isomers, important in pharmaceuticals.

Example: Glucose and fructose are structural isomers (C6H12O6).

Additional info:

• Some content was inferred and expanded for clarity and completeness, such as the full list of functional groups and the explanation of isomers.

• Tables were recreated based on the context of the images and notes.