Chemistry Exam 4

Hydrogen bonding in water is the electrostatic attraction between the positively charged hydrogen atom of one water molecule (H₂O) and the negatively charged oxygen atom of an adjacent water molecule.

• Strength of Bonds: 
    - Hydrogen bonds are weak compared to covalent bonds within the molecule but are strong enough to dictate water's unique physical properties.
      Example: In ice, hydrogen bonds hold water molecules in a structured lattice, making it less dense than liquid water.

• Surface Tension: 
    - Caused by intermolecular forces that keep water molecules together.
      Example: Water droplets form on a waxed car hood due to high surface tension, creating a bead-like shape.

• Capillary Rise: 
    - Upward movement of liquids in narrow spaces (capillaries) against gravity, driven by adhesive forces (liquid-to-solid) and surface tension (cohesion within the liquid).
       Example: Water rises in a thin straw due to capillary action.

• Specific Heat: 
    - Water has a high specific heat, requiring more energy to heat up or cool down, helping regulate Earth's environment.
       Example: Oceans regulate air temperature, absorbing heat during the day and releasing it at night.

• Dissociation in Chemistry: 
    - The process where charged particles in an ionic solid separate from one another (e.g., NaCl to Na⁺ and Cl⁻).
       Example: When table salt (NaCl) dissolves in water, it dissociates into sodium (Na⁺) and chloride ions (Cl⁻).

• Solubility Terms: 
     - Unsaturated: Less than the maximum solute
     - Saturated: Maximum solute
     - Supersaturated: More than the maximum solute.
       Example: A sugar solution becomes saturated when no more sugar can dissolve at a given temperature.

• Freezing Point Depression and Boiling Point Elevation: 
    - When a solute is dissolved in a solvent, the boiling point goes up, and the freezing point goes down.
       Example: Adding salt to ice lowers the freezing point, helping to melt ice on roads in winter.

• Osmosis: 
     - Flow of solvent molecules through a selectively permeable membrane driven by concentration differences.
       Example: Plant roots absorb water from the soil through osmosis.

• Factors Affecting Solubility: 
    - Higher gas pressure and/or higher temperature make it easier for gases to dissolve in liquids.
       Example: Carbon dioxide dissolving in soda is more effective under higher pressure.

• Colloids: 
    - Mixtures that appear uniform but are not on a microscopic level.
       Example: Milk is a colloid where fat particles are suspended in water.

• Heat of Solution: 
    - Total amount of heat energy absorbed or released when a solute dissolves in a solvent.
       Example: The temperature decrease when ammonium nitrate dissolves in water is a result of heat of solution.

• Tyndall Effect: 
    - Scattering of light by particles in a colloid or fine suspension, making a light beam visible as it passes through the mixture.
       Example: A beam of sunlight is visible in a dusty room.

• Acids and Bases: 
    - Acidic hydrogen can be transferred to water; when litmus is exposed to acid, it turns red and blue for base.
       Example: HCl in water will turn blue litmus paper red, indicating its acidic nature.

• Ionization: 
    - Process where an ionic compound breaks apart to form ions. Both acids and bases are electrolytes; bases produce hydroxide (OH⁻).
       Example: Dissociation of sulfuric acid (H₂SO₄) in water results in hydrogen ions (H⁺) and sulfate ions (SO₄²⁻).

• Types of Bases: 
    - When an ionic base dissolves in water, hydroxide ions are produced.
       Example: Sodium hydroxide (NaOH) dissociates to produce Na⁺ and OH⁻.
    - When a covalent base dissolves in water, hydroxide ions are formed by breaking apart water.
       Example: Ammonia (NH₃) reacts with water to form NH₄⁺ and OH⁻.

• Anhydrides: 
    - Compounds containing oxygen bonded to one other element.
       Example: Sulfur trioxide (SO₃) is a non-metal oxide and can form sulfuric acid (H₂SO₄) when combined with water.