Comprehensive Introduction to Chemistry and Solution Chemistry

Academic Requirements and Assessment

• Midterm Examinations and Projects:

  • There are a total of 3 midterm tests and 1 final project.

  • Distribution: 1 midterm in the Winter semester and 2 midterms in the Summer semester.

  • Project: A topic is assigned by the professor at the end of the summer semester.

  • Exemption Policy: Students who achieve a score of $4.50$ or higher across the three midterm tests are exempted from the Blackboard ‘stop test’ during the final exam.

• Final Exam Structure:

  • The final exam is comprised of three distinct parts:

    1. Blackboard test (known as the ‘stop-test’).

    2. Written exam (essay format).

    3. Oral exam.

  • Successful completion of all three parts is mandatory for passing the course.

Fundamental Chemical Definitions

• Chemistry: The scientific study of matter and the various changes it undergoes.

• Matter: Defined as anything that occupies physical space and possesses mass.

• Substance: A specific form of matter characterized by a definite (constant) composition and distinct properties. Examples include water, ammonia, table sugar, silver, and oxygen.

• Mixture: A combination of two or more substances in which the substances retain their distinct properties. Mixtures do not have a constant composition and can be separated via physical methods.

  • Homogeneous Mixture: The composition is uniform throughout the entire mixture (e.g., a tablespoon of sugar dissolved in water).

  • Heterogeneous Mixture: The composition is non-uniform throughout the mixture (e.g., cereal in milk).

• Element: A substance that cannot be broken down into simpler substances through chemical processes.

  • There are currenty $118$ known elements, $94$ of which occur naturally.

  • Smallest atomic mass: Hydrogen ($H$).

  • Largest atomic mass: Oganesson ($Og$).

• Compound: A substance made of atoms of two or more elements chemically united in fixed proportions. Compounds can be separated into their constituent elements by chemical methods.

• Atom: The basic unit of an element that cannot be divided into smaller parts by simple chemical processes.

• Subatomic Particles: Includes electrons, protons, and neutrons.

The Periodic Table and Living Systems

• Periodic Table Organization:

  • Elements are categorized into groups: Alkali metals, Alkaline earth metals, Transition metals, Lanthanides, Actinides, Halogens, and Noble gases.

  • Specific atomic data provided in the table includes the Atomic Number, Symbol, Name, and Atomic Mass ($u$).

• Distribution of Elements in Living Systems:

  • Trace Elements: Specifically Iron ($Fe$), Copper ($Cu$), Zinc ($Zn$), Iodine ($I$), and Cobalt ($Co$).

  • Mass Percentage: Trace elements collectively make up approximately $0.1\,\%$ of the body’s mass.

  • Biological Significance: These elements are essential for growth, metabolism (transport of oxygen), and defense against diseases. Deficiencies or excesses can lead to serious illness, retardation, or death.

Molecules, Ions, and Allotropes

• Molecule: An aggregate of at least two atoms in a definite arrangement, held together by chemical bonds.

  • Diatomic Molecules: $H_2$, $N_2$, $O_2$, $Cl_2$, $Br_2$, $I_2$, $HCl$, $CO$.

  • Polyatomic Molecules: $H_2O$, $NH_3$, $CH_4$, and others.

• Ion: An atom or group of atoms with a net positive or negative charge.

  • Cations (Positive): $Na^+$, $H_3O^+$, $Ca^{2+}$.

  • Anions (Negative): $OH^-$, $CH_3COO^-$, $S^{2-}$.

• Allotrope: One of two or more distinct forms of an element.

  • Forms of Oxygen: $O_2$ and $O_3$ (Ozone).

  • Forms of Carbon: Diamond, graphite, fullerene, graphene.

States of Matter and Dispersed Systems

• States of Matter: Matter exists in solid, liquid, gaseous, and plasma states, including intermediate states and phase transitions occurring in nature and technology.

• Dispersed System: A system where one substance (dispersed phase) is distributed as particles throughout another substance (dispersion medium/continuous phase).

• Classification by Physical State (Phase Examples):

  • Gas in Gas: Air, gas mixtures.

  • Gas in Liquid: $O_2$-water, $CO_2$-water, foam, soda drink.

  • Gas in Solid: $H_2$ in palladium, marshmallow.

  • Liquid in Gas: Fog, mist.

  • Liquid in Liquid: Ethanol-water, oil-water, emulsions, milk.

  • Liquid in Solid: Hydrated salts, mercury-gold (amalgam), butter.

  • Solid in Gas: Dust, smoke, aerosol.

  • Solid in Liquid: Sugar-water, salts-water.

  • Solid in Solid: Alloys, bronze.

• Classification by Particle Size:

  • Coarse Dispersion: Particle size $> 1 \times 10^3\,nm$. These are heterogeneous (e.g., suspensions like antibiotics for children, emulsions).

  • Colloidal Dispersion: Particle size between $1 - 1 \times 10^3\,nm$. These are micro-heterogeneous (e.g., blood, lymph, breastmilk, saliva).

  • True Solutions: Particle size $< 1\,nm$. These are homogeneous mixtures. The particle size is comparable to the size of the dispersion medium particles. They can be molecular or ionic (e.g., normal saline solution, $0.9\,\%\,w/v\,NaCl$).

Water as the Universal Solvent

• Solubility Terms:

  • Hydrophilic: "Water loving"; substances with an affinity for water.

  • Hydrophobic: "Water fearing"; substances that repel water or lack affinity for it.

• Unique Properties of Water ($H_2O$):

  • Excellent Solvent: Facilitates transport of nutrients and waste; enables biological processes.

  • High Dielectric Constant: Results in high solubility and ionization of ionic substances.

  • Surface Tension: Higher than any other liquid; governs drop formation and physiological surface phenomena.

  • Transparency: Allows visible and long-wavelength UV light to reach depths for photosynthesis.

  • Density: Reaches maximum density as a liquid at $4^\circ C$. This allows ice to float and restricts vertical circulation in stratified water bodies.

  • Thermal Properties:

    • Highest heat of evaporation: Determines heat transfer.

    • High latent heat of fusion: Stabilizes temperature at the freezing point.

    • High heat capacity: Stabilizes temperatures of organisms and geographical regions.

Quantifying Solution Composition

• Qualitative Descriptions:

  • Dilute: Relatively little solute present.

  • Concentrated: Large amount of solute present.

• Concentration Units:

  • Mass Percent:     $\text{percent by mass}(\%) = \frac{\text{mass of solute}}{\text{mass of solute} + \text{mass of solvent}} \times 100$

  • Mole Fraction ($x_A$): A unitless value representing the ratio of moles of one component to the total moles.     $x_A = \frac{n_A}{n_A + n_B}$

  • Molarity ($M$): Number of moles of solute per liter of solution.     $M = \frac{n}{V}$     $V$ is in Liters ($1\,L = 1\,dm^3 = 1000\,cm^3$). Units: $mol/L$, $mol \cdot L^{-1}$, $mol/dm^3$, $M$.     Calculation formula: $M = \frac{m}{M \times V}$

  • Molality ($m$): Amount of substance dissolved in a specific mass of solvent.     $m = \frac{n}{\text{kilograms of solvent}}$     Mass is used because it does not change with temperature. Units: $mol/kg$, $mol \cdot kg^{-1}$.     Note: For dilute aqueous solutions, Molarity and Molality are nearly equal because the density is close to $1.0\,g/mL$.

The Mole Concept and Conversions

• Definition of the Mole: The amount of substance containing as many elementary entities (atoms, molecules, ions) as there are atoms in exactly $12\,g$ of Carbon-12.

• Avogadro’s Number: $6.0221415 \times 10^{23}$.

• Molar Mass ($M$): The mass in grams or kilograms of $1\,mole$ of a substance. Numerically equal to atomic or molecular mass.

• Key Formula: $n = \frac{m}{M}$

• Practical Examples for 1 Mole:

  • Water ($H_2O$): $M \approx 18\,g/mol$. $1\,mol = 18\,g \approx 18\,mL$.

  • Sulfuric Acid ($H_2SO_4$): $M \approx 98\,g/mol$. $1\,mol = 98\,g$. Given density $\rho = 1.83\,g/cm^3$, the volume $V = \frac{98}{1.83} \approx 53.5\,mL$.

  • Lead ($Pb$): $M = 207.2\,g/mol$. $1\,mol = 207.2\,g$. Given density $\rho = 11.342\,g/cm^3$, the volume is $\approx 18.3\,cm^3$.

Solvation and Solubility Processes

• Solvation: The interaction of a solvent with dissolved molecules. Solvent molecules surround solute ions in a concentric shell (solvation complexes).

  • Hydration: Solvation specifically involving water.

• Intermolecular Forces in Solvation:

  • Hydrogen Bonding: Interaction between Hydrogen and an electronegative atom ($N$, $O$, or $F$) of another molecule.

  • Van der Waals Forces: Electrostatic forces caused by fluctuating dipoles (dipole-dipole, dipole-induced dipole, London dispersion). These are nonspecific and short-lived but can be permanent in large numbers.

• Dissolution Mechanisms:

  • Ionic Compounds (e.g., $NaCl$): Negative oxygen ends of water attract $Na^+$; positive hydrogen ends attract $Cl^-$. Attactive forces pull ions apart.

  • Molecular Compounds (e.g., Sugar): Attraction between positive/negative ends of water and sugar molecules pull individual sugar molecules away from the crystal.

Factors Affecting Solubility

• Definition of Solubility: The maximum amount of substance that dissolves in a given solvent at a specific temperature under dynamic equilibrium (rate of dissolution equals rate of phase joining/precipitation).

• Solute-Solvent Interactions: "Like dissolves like." Polar solutes dissolve in polar solvents; non-polar in non-polar.

  • Example: Vitamin B6 contains multiple polar groups ($OH$), making it water-soluble.

• Common-Ion Effect: The decrease in solubility of an ionic compound when a salt containing a common ion is added to the equilibrium mixture.

• Temperature Effects:

  • Solids:

    1. Endothermic ($\Delta H_{solvation} > 0$): Increasing temperature shifts equilibrium toward the products, increasing solubility.

    2. Exothermic ($\Delta H_{solvation} < 0$): Increasing temperature shifts equilibrium toward the reactants, decreasing solubility.

  • Gases: Increased temperature leads to lesser solubility. Gas particles gain kinetic energy, increase motion, and are more likely to escape the liquid phase.

• Pressure Effects: Significant only for gases in liquids.

  • Henry's Law: At constant temperature, the solubility of a gas is directly proportional to its partial pressure ($p$).     $p = k_h \times c$     Where $k_h$ is the Henry's law constant and $c$ is the concentration of the gas.