Topic 6: The Universal Solvent - Water Study Guide

Fundamentals of Water as a Universal Solvent

• Water is widely regarded as the universal solvent primarily because it is an excellent solvent for many of the chemicals that compose human bodies.

• Solvent Definition: A substance that is capable of dissolving other substances.

• Solute Definition: The substances that are dissolved within a solvent.

• Aqueous Solutions: These are specific types of solutions where water serves as the solvent. An example provided is the dissolution of table salt ($NaCl$) in water.

Advantages and Disadvantages of Water as a Solvent

• Advantages for Selection:

  • Economic: It is inexpensive and widely available.

  • Stability: It is a stable chemical compound.

  • Safety: It is non-toxic and a non-irritant.

  • Physical Characteristics: It is a polar solvent, colorless, odorless, and tasteless.

  • Storage: It is easy to store and non-flammable.

  • Physical Properties: It possesses low viscosity.

  • Versatility: It acts as a good solvent for many compounds and is miscible with many other polar solvents.

• Disadvantages:

  • Microbial Growth: Water may aid the growth of various microorganisms.

  • Solubility Limits: It is not an effective solvent for non-polar compounds.

  • Reactivity: It may promote chemical reactions such as hydrolysis and oxidation.

Types of Water

• Potable water: Water that is safe to drink.

• Purified water: Water that has been purified through processes like distillation or other treatments (e.g., activated carbon/C etc.).

• Water for preparation: This refers to freshly boiled purified water.

• Water for injections: This is defined as sterilized distilled water.

The Phase Diagram of Water

• The phase diagram of water is a graphical representation of the physical states of water under different conditions of temperature and pressure. The diagram is divided into three distinct regions:

  • Region WABD: Represents Ice ($H_2O_{(s)}$).

  • Region DAE: Represents Liquid Water ($H_2O_{(l)}$).

  • Region WACE: Represents Water Vapour ($H_2O_{(g)}$).

Analysis of Curves and Equilibrium Lines

• The Curve ACE (Liquid/Vapour Line):

  • This line illustrates the variation of the saturated vapour pressure of liquid water as a function of temperature.

  • It indicates the boiling point of water at various pressures.

  • At a pressure of $1\,\text{atm}$, the boiling point of water is exactly $100^\circ\text{C}$. At these specific coordinates, liquid water and water vapour exist in dynamic equilibrium.

  • At a reduced pressure of $0.1\,\text{atm}$, the boiling point of water drops to $50^\circ\text{C}$.

• The Curve WA (Solid/Vapour Line):

  • This represents the saturated vapour pressure of ice (solid water) at different temperatures.

  • The vapour pressure of ice results from the sublimation of water molecules from the surface of the ice.

  • This curve shows the sublimation temperature against pressure. For example, a sample of ice at $0.006\,\text{atm}$ is solid at $0^\circ\text{C}$. If the temperature increases above $0.01^\circ\text{C}$ at this pressure, the solid will not melt but will sublime directly into gas.

• The Curve DA (Solid/Liquid Line):

  • This line exhibits a negative gradient, representing the melting point of water at different pressures.

  • The negative slope indicates that the melting point of ice decreases as pressure increases. (Note: For most other solids, the melting point increases as pressure increases).

  • This anomaly occurs because when ice melts, its volume decreases due to the unique nature of hydrogen bonding. Consequently, it is easier for ice to melt at high pressure compared to low pressure.

  • At a standard pressure of $1\,\text{atm}$, ice melts (or water freezes) at $0^\circ\text{C}$.

Significant Points in the Phase Diagram

• Triple Point (Point A):

  • This marks the unique conditions of temperature and pressure ($0.01^\circ\text{C}$ and $0.006\,\text{atm}$) where ice, liquid water, and water vapour all coexist in dynamic equilibrium.

  • The triple point is distinct from the melting point; the melting point is the temperature at which solid and liquid coexist specifically under atmospheric pressure (usually $1\,\text{atm}$).

• Critical Point (Point E):

  • This occurs at pressures above $200\,\text{atm}$ and temperatures approximately at $374^\circ\text{C}$.

  • It is defined as the maximum temperature at which a sample of water vapour can be condensed into a liquid by increasing pressure alone.

  • Any sample of water vapour at temperatures exceeding $374^\circ\text{C}$ cannot be liquified by pressure increase alone. To achieve condensation, the temperature must first be decreased to $374^\circ\text{C}$ or lower.

Molecular Structure and Hydrogen Bonding

• Polarity of the Water Molecule:

  • Water is a polar molecule. The oxygen (O) atom carries a partial negative charge ($\delta^-$) and the hydrogen (H) atoms carry partial positive charges ($\delta^+$).

  • This charge separation results in an uneven distribution of the electron cloud.

• Definition of Hydrogen Bonding:

  • Because opposite charges attract, a partially positive-charged H atom of one water molecule is attracted to the partially negative-charged region (the nonbonding electron pairs) on the O atom of an adjacent water molecule.

  • A hydrogen bond is an intermolecular force occurring between molecules. It specifically occurs when a hydrogen atom is covalently bonded to a small, highly electronegative atom (such as $F$, $O$, or $N$) and is attracted to a lone pair of electrons on another highly electronegative atom.

  • Each $H_2O$ molecule possesses two hydrogen atoms and two nonbonding pairs of electrons, allowing for multiple intermolecular attractions.

Importance of Hydrogen Bonding

• Boiling and Melting Points:

  • Hydrogen bonds are significantly stronger than other types of intermolecular forces.

  • The presence of these bonds increases the boiling and melting points of water because substantial energy is required to break these bonds.

  • If hydrogen bonds in water were weaker, water would have a much lower boiling point.

• Biological Significance:

  • Protein Structure: Hydrogen bonding stabilizes the shapes of large biological molecules. In proteins (the components of skin, hair, and muscle), hydrogen bonding occurs between hydrogen atoms and oxygen or nitrogen atoms.

  • DNA Stability: The coiled, double-helical structure of Deoxyribonucleic acid (DNA) is stabilized by thousands of hydrogen bonds formed between specific segments of the linked DNA strands. This process is essential to life processes.

• Density of Ice:

  • Hydrogen bonding explains why ice cubes and icebergs float in liquid water.

  • In the solid state (ice), each oxygen atom is rigidly bonded to four hydrogen atoms: two via covalent bonds and two via hydrogen bonds.

  • This arrangement forms a tetrahedron structure. This geometry pushes the molecules further apart than they are in the liquid state, which is why water expands when it freezes and becomes less dense as a solid.

Practice Question: Phase Diagram Analysis

Based on the provided diagram with pressures in $mm\,Hg$:

• Standard Values identified:

  • Normal freezing point: $0^\circ\text{C}$.

  • Triple point: $0.01^\circ\text{C}$ at $4.58\,mm\,Hg$.

  • Normal boiling point: $100^\circ\text{C}$ at $760\,mm\,Hg$.

• Task: Explain the changes from $a \rightarrow c$ and after point $c$.

  • At point $a$, the substance is in the solid phase.

  • Moving from $a \rightarrow b$, the temperature increases until it hits the sublimation curve (or melting curve depending on the specific pressure line), transitioning toward the vapor phase.

  • At point $c$, as temperature continues to rise at a specific pressure, the phase transitions completely into the vapor region.