CHEM 191: Aqueous Solutions and Hydrogen Bonding

• CHEM 191 is Module 1, Lecture 2: Aqueous Solutions, based on Brown (15th) Chapter 11.2
• This module aims to explain what is the role of hydrogen bonding, the interactions between water and solutes, and the difference between strong and weak electrolytes

Aqueous Solutions

• These are important in biological systems
• The solubility of substances in aqueous biological solutions is crucial for:
  • Transport of gases
  • Transport of inorganic materials
  • Application and transport of drugs

Water Molecule

• Consists of two O-H bonds; oxygen is more electronegative than hydrogen
• The oxygen atom has a partial negative charge (δ-), while the hydrogen atom has a partial positive charge (δ+)
• Water molecules are attracted to each other through dipole-dipole interactions via partial charges
• Strong dipole-dipole interactions in water are called hydrogen bonding
• Water is a polar molecule with extensive hydrogen bonding in solid and solution states
• Polar molecules and ionic solids tend to be soluble in water

Unusual Properties of Water

• Exhibits high melting and boiling points compared to similar molecules
  • HF: Tm = -83°C, Tb = 19.5°C
  • H2O: Tm = 0°C, Tb = 100°C
  • NH3: Tm = -77.7°C, Tb = -33.4°C
  • CH4: Tm = -183°C, Tb = -161°C
  • H2S: Tm = -85.5°C, Tb = -60.7°C
• High values are due to extensive H-bonding
• Becomes more dense when it melts, reaching maximum density at 4°C, which makes ice float
• Freezing point decreases as pressure increases
• Exists in 14 different ice forms; one form, ice X, melts at 2100°C under 800,000 times atmospheric pressure
• Possesses high surface tension and heat capacity
• Has a high concentration: 55.5 mol L-1 at 25°C
• In liquid water, molecules move closely, constantly forming and breaking hydrogen bonds
• In solid water (ice), molecules are rigidly held in a structure that maximizes hydrogen bonds, making ice less dense than liquid water

Water as a Solvent

• Excellent solvent for electrolytes, which form ions upon dissolution
  • Example: NaCl(s) → Na+(aq) + Cl-(aq)
• Dissolution occurs spontaneously, despite strong electrostatic forces in NaCl (melts at 801 °C)
• Ions are high-energy species, but solvation stabilizes them via interactions with water molecule dipoles
• Solvation involves the interaction of charged ions with water dipoles in the hydration sphere, where ions are surrounded by water molecules
• When water acts as a solvent, this is called hydration

Aqueous Solubility

• Water is a good solvent for polar molecules, which have polar bonds due to electronegativity differences
• Polar molecules and ionic solids are more soluble in water compared to non-polar substances
• Examples include:
  • Alcohols
  • Amines
  • Carboxylic acids
  • Amino acids
• Non-polar molecules like O2(g) and N2(g) are generally not very soluble, but have some solubility
• A permanent dipole of water can induce a dipole in non-polar molecules
• The magnitude of the induced dipole depends on the size and shape of the gas molecule.
• Non-polar gases are not very soluble in water
  • He: 0.00150 g kg-1 (298 K)
  • O2: 0.0393 g kg-1 (298 K)
  • N2: 0.0175 g kg-1 (298 K)
  • Cl2: 6.41 g kg-1 (298 K)
  • CO2: 1.45 g kg-1 (298 K)
  • NH3: 489 g kg-1 (298 K)
• Ammonia (NH3) can hydrogen-bond and is very soluble; it also chemically reacts with water

Definitions

• A solution forms when one or more chemical species dissolve in a liquid
• A solute is a substance that dissolves in a liquid (solid, liquid, or gas)
• A solvent is a liquid in which the solute dissolves
• A dilute solution has n(solute ) less than n(solvent)
• An electrolyte dissolves to give ions in solution
  • e.g. NaCl, AgNO3, CuSO4.5H2O
• A nonelectrolyte dissolves without forming ions
  • e.g. glucose (C6H12O6), urea ((NH2)2CO)

Classification of Solutes

• Can be divided into different types such as electrolytes and non-electrolytes

Electrolytes and Non-Electrolytes

• Sodium chloride (salt) is an electrolyte
• It is an ionic solid with Na⁺ and Cl⁻ occupying lattice sites
• Dissolves spontaneously in water: NaCl(s) → Na+(aq) + Cl-(aq)
• Na⁺ and Cl⁻ ions are uniformly distributed in H2O
• Electrolytes include all ionic solids, acids, bases
• Glucose is a non-electrolyte and is a molecular solid held together by hydrogen bonds
• Dissolves spontaneously in water; C6H12O6(s) → C6H12O6(aq)
• Molecules are uniformly distributed in H2O and unchanged in their bonding

Strong electrolytes

• These undergo complete dissociation in water, where only ions are created
• All dissolved species are ions; dissociation reaction ‘goes to completion’
• Some ionic solids which is not notably soluble in water are still strong electrolytes
  • e.g. AgCl dissolves completely into Ag+ and Cl- ions: AgCl(s) → Ag+(aq) + Cl-(aq)

Weak Electrolytes

• These undergo incomplete dissociation in water
• Most dissolved species are molecules, with only some dissociating into ions
• Dissociation reaction does not ‘go to completion’
• Some dissociate into ions Weak acids such as acetic acid are classic examples of these electrolytes