Water Everywhere: A Most Precious Resource

Electronegativity

Electronegativity is a fundamental concept in chemistry that describes the ability of an atom to attract electrons in a chemical bond. It helps explain bond polarity and molecular properties.

• Definition: Electronegativity (EN) is a measure of the attraction of an atom for electrons in a chemical bond.

• Trends: Electronegativity increases across a period (left to right) and decreases down a group (top to bottom) in the periodic table.

• Most electronegative element: Fluorine (F) is the most electronegative element.

• Example: In an H–F bond, F is more electronegative than H, so the shared electrons are pulled closer to F.

• Order Example: For Ba, Ca, Ga, Se: Ba < Ca < Ga < Se (increasing EN).

Bond Polarity and Polar Covalent Bonds

Bond polarity arises from differences in electronegativity between two bonded atoms, resulting in partial charges.

• Polar Covalent Bond: A bond where electrons are shared unequally due to a difference in electronegativity.

• Partial Charges: The more electronegative atom gains a partial negative charge (δ–), and the less electronegative atom gains a partial positive charge (δ+).

• Example: In O–H, O is more electronegative (EN = 3.5) than H (EN = 2.1), so O is δ– and H is δ+.

• EN Difference:

Predicting Molecular Polarity

The polarity of a molecule depends on both the polarity of its bonds and its molecular geometry.

• Nonpolar Molecules: Molecules with only nonpolar bonds (e.g., H2, Cl2) or symmetrical arrangements of polar bonds (e.g., CO2).

• Polar Molecules: Molecules with polar bonds and an asymmetrical shape (e.g., H2O).

• Rule of Thumb: If a molecule has lone pairs on the central atom and/or different terminal atoms, it is likely polar.

• Examples: H2O (polar), CO2 (nonpolar), SO3 (nonpolar), CCl4 (nonpolar).

Hydrogen Bonding

Hydrogen bonding is a special type of intermolecular force that significantly affects the physical properties of substances, especially water.

• Definition: A hydrogen bond is an electrostatic attraction between a hydrogen atom bonded to a highly electronegative atom (O, N, or F) and a lone pair on another O, N, or F atom in a neighboring molecule.

• Strength: Hydrogen bonds are stronger than other intermolecular forces (like dispersion or dipole-dipole), but weaker than covalent bonds.

• Examples: H2O, NH3, HF all exhibit hydrogen bonding.

Table: Hydrogen Bonding and Physical Properties

Additional info: Table values for CH4, NH3, and CO2 inferred for comparison.

States of Water and Anomalous Properties

Water exhibits unique properties due to hydrogen bonding, such as high boiling and melting points, high specific heat, and lower density of ice compared to liquid water.

• States: Water exists as solid (ice), liquid, and gas (vapor) under standard conditions.

• Melting Point: 0°C; Boiling Point: 100°C at 1 atm.

• Density: Ice is less dense than liquid water due to the open hydrogen-bonded lattice structure.

• Specific Heat: Water has a high specific heat capacity (4.18 J/g·°C), allowing it to absorb/release large amounts of heat with little temperature change.

Types of Intermolecular Forces (IMFs)

IMFs are forces of attraction between molecules, influencing physical properties like boiling and melting points.

• Dispersion Forces: Present in all molecules; arise from temporary dipoles.

• Dipole-Dipole Forces: Occur between polar molecules with permanent dipoles.

• Hydrogen Bonding: Strongest type of dipole-dipole force; occurs when H is bonded to O, N, or F.

Table: Types of Intermolecular Forces

Concentration of Solutions

Concentration expresses the amount of solute dissolved in a given quantity of solvent or solution. It can be represented in various units.

• Percent by Mass:

• Parts per Million (ppm):

• Parts per Billion (ppb):

• Molarity (M):

Example: To prepare a 2.00 M solution of CuSO4 in 1.00 L, dissolve 319 g of CuSO4 (molar mass = 159.62 g/mol) in water to make 1.00 L of solution.

Solubility of Ionic Compounds

Solubility rules help predict whether an ionic compound will dissolve in water. Solubility depends on the ions present and their interactions with water molecules.

• Soluble Compounds: All nitrates (NO3–), alkali metal salts (Li+, Na+, K+), and ammonium (NH4+) salts are soluble.

• Insoluble Compounds: Most carbonates (CO32–), phosphates (PO43–), sulfides (S2–), and hydroxides (OH–) are insoluble, except those of alkali metals and NH4+.

• Exceptions: AgCl, PbSO4, BaSO4 are insoluble.

Table: Solubility Rules (Summary)

Electrolytes and Nonelectrolytes

Electrolytes are substances that conduct electricity when dissolved in water, while nonelectrolytes do not.

• Strong Electrolytes: Completely dissociate into ions (e.g., NaCl, HCl).

• Weak Electrolytes: Partially dissociate (e.g., acetic acid, NH3).

• Nonelectrolytes: Do not dissociate (e.g., sugar, ethanol).

Solubility of Molecular Compounds: Like Dissolves Like

Solubility of molecular compounds depends on the polarity of both solute and solvent. "Like dissolves like" is a guiding principle.

• Polar solutes dissolve in polar solvents (e.g., sugar in water).

• Nonpolar solutes dissolve in nonpolar solvents (e.g., oil in hexane).

• Surfactants: Molecules with both polar and nonpolar regions can help mix polar and nonpolar substances (e.g., soaps).

Acids and Bases

Acids and bases are classified by their ability to donate or accept protons (H+) in aqueous solution.

• Acid: Substance that releases H+ ions in water (e.g., HCl, HNO3).

• Base: Substance that produces OH– ions in water (e.g., NaOH, Ba(OH)2).

• Strong acids/bases: Completely dissociate in water.

• Weak acids/bases: Partially dissociate in water (e.g., acetic acid, NH3).

Acid-Base Neutralization

Neutralization is a reaction between an acid and a base to form water and a salt.

• General Equation:

• Example:

• Net Ionic Equation:

• Spectator Ions: Ions that do not participate in the reaction (e.g., Na+, Cl–).

The pH Scale

The pH scale measures the acidity or basicity of a solution based on the concentration of hydrogen ions.

• Definition:

• Neutral Solution: [H+] = [OH–] = M, pH = 7

• Acidic Solution: [H+] > M, pH < 7

• Basic Solution: [H+] < M, pH > 7

• Relationship: at 25°C

Example: If [H+] = M, then pH =

Table: pH Values of Common Substances

Buffer Systems

Buffers are solutions that resist changes in pH upon addition of small amounts of acid or base. They are essential in biological and environmental systems.

• Example: The carbonate buffer system in blood and oceans involves CO2, H2CO3, HCO3–, and CO32–.

• Key Equilibrium: