Basic Chemistry, Mixtures, and Chemical Bonding

Classification and Properties of Mixtures

• Matter exists predominantly in the form of mixtures, defined as two or more components that are physically intermixed without chemical bonding.

• Mixtures are categorized into three basic physical types based on solute particle size, optical properties, and settling characteristics:

  • Solutions: Homogeneous mixtures containing very tiny solute particles that do not settle out at the bottom over time and do not scatter light.

  • Colloids: Heterogeneous mixtures (also known as emulsions) containing solute particles that are larger than those in a true solution. These particles scatter light (giving the mixture a cloudy or milky appearance) but do not settle out.

  • Suspensions: Heterogeneous mixtures containing the largest solute particles. These visible particles settle out at the bottom of a container when left undisturbed and may scatter light.

Solutions and Concentration Measurements

• Homogeneous Nature: Solutions are uniform in composition, meaning that solute particles are evenly distributed throughout the mixture.

• Components of a Solution:

  • Solvent: The substance present in the greatest quantity within a solution. It serves as the dissolving medium and is usually a liquid, such as water.

  • Solute(s): The substance or substances dissolved in the solvent, present in smaller amounts.

  • Example: In blood sugar, glucose acts as the solute, while blood plasma acts as the solvent.

• Properties of True Solutions:

  • True solutions are typically transparent.

  • Examples include air (a gaseous solution where nitrogen serves as the solvent), salt solutions, and sugar solutions.

  • Most solutions within the human body are true solutions composed of gases, liquids, or solids dissolved in water.

• Quantitative Expressions of Solution Concentration:

  • Percent (Parts per 100 parts): Quantifies solute relative to 100 total parts. For example, dissolved solute consisting of $15$ parts of salt in $85$ parts of water forms a $15\%$ salt solution.

  • Milligrams per Deciliter ($\text{mg/dl}$): Measures solute mass in milligrams relative to solution volume in deciliters. For instance, normal human blood glucose concentration is approximately $80\,\text{mg/dl}$.

  • Molarity ($M$): Defined as the number of moles of solute per liter ($\text{L}$) of solvent (water).

  • One mole of any chemical compound is equal to its molecular weight (the sum of its atomic weights) expressed in grams.

  • Glucose ($\text{C}_6\text{H}_{12}\text{O}_6$) has a molecular weight of $180.12\,\text{amu}$. Therefore, adding $180.12\,\text{g}$ of glucose to enough water to make a total volume of $1\,\text{L}$ produces a $1\,M$ solution of glucose.

  • One mole of any substance consistently contains $6.02 \times 10^{23}$ molecules. This constant value is known as Avogadro's number.

  • Solute concentrations in biological fluids are often extremely low (such as $0.0001\,M$). As a result, biological molarities are expressed in millimoles ($\text{mM}$), where $1000\,\text{mM} = 1\,M$.

Colloids and Sol-Gel Transformations

• Heterogeneous Nature: Colloids, also termed emulsions, are heterogeneous mixtures in which particles are not distributed completely uniformly.

• Optical and Physical Characteristics:

  • Solute particles in colloids are larger than those in true solutions.

  • The particles scatter light rays, imparting a translucent, cloudy, or milky visual aspect to the mixture.

  • Despite their larger size, colloidal particles do not precipitate or settle to the bottom of the vessel.

• Sol-Gel Transformations:

  • Colloids possess the unique ability to undergo reversible or irreversible sol-gel transformations, transitioning from a fluid state (sol) to a solid or semi-solid gel state.

  • Examples: Gelatin dessert products (such as Jell-O) transform from a liquid mixture into a gel. The cytosol inside living cells is also a sol-gel solution that changes consistency to drive cellular processes.

Suspensions

• Definition and Behavior: Suspensions are heterogeneous mixtures containing large, easily visible solutes that do not remain permanently suspended.

• Sedimentation: Solute particles in a suspension will settle out to the bottom of a container over time if left undisturbed. Suspended particles may also scatter light.

• Examples:

  • A physical mixture of sand suspended in water.

  • Blood is considered a suspension; when placed in a test tube and left standing or centrifuged, red blood cells settle to the bottom, separating from the upper liquid plasma layer.

Distinctions Between Mixtures and Compounds

• Chemical Bonding: Components of a mixture are physically combined without forming chemical bonds. In contrast, components of a compound are chemically bonded together.

• Physical Separation: Mixtures can be separated into their constituent parts using physical techniques such as filtration, straining, or centrifugation. Compounds can only be separated into their constituent elements by breaking chemical bonds through chemical reactions.

• Homogeneity: Mixtures can be either heterogeneous or homogeneous in their spatial arrangement, whereas compounds are strictly homogeneous.

Nature of Chemical Bonds and Atomic Structure

• Definition of Chemical Bonds: Chemical bonds represent energy relationships between electrons of interacting atoms rather than physical structural connectors.

• Central Role of Electrons: Subatomic electrons are the primary participants in all chemical interactions. They determine whether a given chemical reaction will occur and dictate the specific type of chemical bond formed.

• Electron Shells and Energy Levels:

  • Electrons occupy concentric regions surrounding the atomic nucleus called electron shells or energy levels, reflecting the discrete kinetic and potential energy of the electrons in that region.

  • Depending on atomic size, an atom can possess up to $7$ distinct electron shells.

  • Shells fill sequentially starting from the innermost shell closest to the nucleus:

  • Shell 1: Holds a maximum of $2$ electrons.

  • Shell 2: Holds a maximum of $8$ electrons.

  • Shell 3: Holds a maximum of $8$ electrons.

Valence Shell and the Octet Rule

• Valence Shell: The outermost electron shell of an atom. Electrons located in the valence shell contain the highest potential energy because they are situated farthest from the positively charged nucleus, and they are the sole subatomic particles involved in forming chemical bonds.

• Octet Rule (Rule of Eights):

  • Atoms interact to attain $8$ electrons in their valence shell, achieving maximum thermodynamic stability.

  • Exceptions: Small atoms such as Hydrogen ($\text{H}$) and Helium ($\text{He}$) require only $2$ electrons to completely fill Shell 1.

  • The drive to achieve an octet of valence electrons serves as the fundamental force behind all chemical reactions.

  • Atoms possessing a complete outer shell of $8$ valence electrons (or $2$ for Helium) are chemically unreactive or inert.

  • Most biological atoms lack a naturally full valence shell; they must gain, lose, or share electrons with other atoms to achieve stability.

Types of Chemical Bonds

• Ionic Bonds:

  • Ion Formation: Ions are charged particles formed when neutral atoms either gain or lose valence electrons, causing the total number of protons to differ from the number of electrons.

  • Anions: Atoms that gain one or more electrons, taking on a net negative charge.

  • Cations: Atoms that lose one or more electrons, taking on a net positive charge.

  • Bond Mechanism: The electrostatic attraction between opposing negative and positive charges binds cations and anions together in an ionic bond ("give and take" transfer of valence electrons).

  • Crystalline Salts: Most ionic compounds exist as salts that form structured crystalline arrays when dry.

• Covalent Bonds:

  • Formed when two atoms share two or more valence shell electrons, allowing each participating atom to complete its valence shell for at least part of the time.

  • Multiplicity of Covalent Bonds:

  • Single Bond: Formed by sharing $2$ electrons ($1$ pair).

  • Double Bond: Formed by sharing $4$ electrons ($2$ pairs).

  • Triple Bond: Formed by sharing $6$ electrons ($3$ pairs).

• Nonpolar vs. Polar Covalent Bonds:

  • Nonpolar Covalent Bonds: Formed when shared electrons are distributed equally between bonded atoms, resulting in an electrically balanced molecule.

  • Polar Covalent Bonds: Formed when electrons are shared unequally between two atoms due to differences in electron-attracting ability (electronegativity).

    • Electronegative Atoms: Atoms with strong electron-pulling capabilities gain a partial negative charge ($\delta^-$).

    • Electropositive Atoms: Atoms with weaker electron-pulling capabilities gain a partial positive charge ($\delta^+$).

    • Polar covalent molecules are structurally weaker than nonpolar covalent molecules.

    • Water Example: In $\text{H}_2\text{O}$, the strongly electronegative oxygen atom exerts a greater pull on shared electrons than the electropositive hydrogen atoms, giving oxygen a partial negative charge ($\delta^-$) and each hydrogen a partial positive charge ($\delta^+$).

• Hydrogen Bonds:

  • Mechanism: An attractive force between a partial electropositive hydrogen atom ($\delta^+$) of one polar molecule and an electronegative atom ($\delta^-$) of another molecule.

  • Hydrogen bonds are not true chemical bonds; they represent weak magnetic dipole attractions.

  • Functions:

  • Responsible for intermolecular attractions between water molecules, maintaining water as a liquid at room temperature.

  • Function as intramolecular bonds within large bio-molecules (such as proteins and DNA), stabilizing their specific three-dimensional functional shapes.

Chemical Reactions and Chemical Equations

• Definition: Chemical reactions occur whenever chemical bonds are formed, rearranged, or broken.

• Chemical Equations: Represent chemical reactions using symbolic notation.

  • Reactants: The starting substances that enter into a reaction together.

  • Products: The resulting end chemical substances produced by the reaction.

  • Chemical equations must be balanced to accurately show proportional amounts of reactants and products:

  • $\text{Reactant A} + \text{Reactant B} \rightarrow \text{Product C} + \text{Product D}$

  • $\text{AB} + \text{CD} \rightarrow \text{AC} + \text{B} + \text{D}$

• Molecular Formulas: Represent chemical compounds symbolically (e.g., $\text{H}_2\text{O}$, $\text{C}_6\text{H}_{12}\text{O}_6$, $\text{H}_2$, $\text{CH}_4$).

  • Subscripts: Indicate the exact number of atoms joined together by chemical bonds within a molecule.

  • Prefixes: Indicate the number of separate, unjoined atoms or distinct molecules:

  • $\text{H} + \text{H} \rightarrow \text{H}_2$ (Hydrogen gas)

  • $4\text{H} + 1\text{C} \rightarrow \text{CH}_4$ (Methane)

Classes of Chemical Reactions

• Synthesis Reactions:

  • Also called combination reactions. Involve atoms or smaller molecules combining to form a larger, more complex molecule ($\text{A} + \text{B} \rightarrow \text{AB}$).

  • Fundamental to anabolic (building) processes within living tissue that consume energy.

• Decomposition Reactions:

  • Involve breaking down a complex molecule into smaller constituent molecules or individual atoms ($\text{AB} \rightarrow \text{A} + \text{B}$).

  • Represent the reverse of synthesis reactions and drive catabolic (bond-breaking) pathways that release energy.

• Exchange (Displacement) Reactions:

  • Involve both synthesis and decomposition steps simultaneously; bonds are broken and new bonds are formed ($\text{AB} + \text{CD} \rightarrow \text{AD} + \text{CB}$).

  • Example: Conversion of phosphoenolpyruvate and $ADP$ (plus $\text{H}^+$) into pyruvate and $ATP$, catalyzed by the enzyme pyruvate kinase.

• Reduction-Oxidation (Redox) Reactions:

  • Critical exchange reactions in living organisms driven by electron transfers.

  • Oxidation: Occurs when an atom or molecule loses electrons (or loses hydrogen / gains oxygen).

  • Reduction: Occurs when an atom or molecule gains electrons (or gains hydrogen / loses oxygen).

  • Mnemonic: OIL RIG (Oxidation Is Loss, Reduction Is Gain).

  • Cellular Respiration Example:

  • $\text{C}_6\text{H}_{12}\text{O}_6 + 6\text{O}_2 \rightarrow 6\text{CO}_2 + 6\text{H}_2\text{O} + \text{Energy}$

  • Glucose ($\text{C}_6\text{H}_{12}\text{O}_6$) is oxidized to carbon dioxide ($\text{CO}_2$).

  • Oxygen ($\text{O}_2$) is reduced to water ($\text{H}_2\text{O}$).

Energy Flow, Reversibility, and Factors Affecting Reaction Rates

• Energetics of Reactions:

  • Exergonic Reactions: Reactions that yield a net release of energy. The chemical products possess less potential energy than the original reactants. Catabolic and oxidative reactions are exergonic.

  • Endergonic Reactions: Reactions that require a net absorption of energy. The chemical products possess more potential energy than the reactants. Anabolic and reductive reactions are endergonic.

• Reversibility and Chemical Equilibrium:

  • Theoretical reversibility applies to all chemical reactions ($\text{A} + \text{B} \rightleftharpoons \text{AB}$).

  • Chemical Equilibrium: Reached when the rate of forward product formation equals the rate of reverse reactant regeneration, leaving net component concentrations static.

  • Many biological reactions are practically irreversible because the energy input required to run the reaction in reverse is prohibitive, or because biological products are continuously removed from the cell system.

• Factors Modulating Reaction Velocity:

  • Temperature: Raising environmental temperature increases particle kinetic energy and collision frequency, accelerating reaction rates.

  • Concentration of Reactants: Increasing reactant particle counts increases collision rates, accelerating reaction rates.

  • Particle Size: Smaller reactant particles move faster and collide more frequently than larger particles at equal temperatures, accelerating reaction rates.

  • Catalysts: Chemical substances that increase reaction rates without undergoing permanent chemical changes themselves or becoming part of the end product.

  • Enzymes: Biological catalysts that increase reaction rates by lowering the activation energy barrier required for reactants to proceed to products.

Questions & Discussion

• Question 1: Is blood a solution, a suspension, and/or a colloid?

  • Detailed Analysis: Blood exhibits characteristics of all three mixture types depending on the fluid component evaluated:

  • Solution: Blood plasma contains dissolved electrolytes (salts), nutrient gases, and glucose in a transparent aqueous medium.

  • Colloid: Blood plasma functions as a colloid because large, suspended plasma proteins (such as albumin and immunoglobulins) scatter light but do not settle out.

  • Suspension: Whole blood acts primarily as a suspension because cellular formed elements (red blood cells, white blood cells, and platelets) are large visible components that settle out when blood sits undisturbed or is centrifuged.

• Question 2: Could two atoms of oxygen engage in ionic bonding? Why or why not?

  • Detailed Analysis: No. Two identical oxygen atoms possess identical electronegativity values and equal electron-attracting abilities. Neither oxygen atom can strip electrons from the other to form a cation-anion pair. Instead, two oxygen atoms share two pairs of valence electrons equally to form a nonpolar double covalent bond ($\text{O}_2$).

• Question 3: $\text{AB} + \text{CD} \rightarrow \text{AD} + \text{BE}$ — Is this a legitimate example of an exchange reaction? Why or why not?

  • Detailed Analysis: No, this equation violates the fundamental law of conservation of mass. The product side introduces an entirely new element ($\text{E}$) that was not present among the reactants, and element $\text{C}$ present in the reactants disappears completely from the products. In a legitimate exchange reaction (such as $\text{AB} + \text{CD} \rightarrow \text{AD} + \text{CB}$), bonds are broken and formed to rearrange existing atoms without creating or destroying atomic identities.

• Timing Guidelines for Discussion Activity:

  • Individual Writing: $5$ minutes to draft responses.

  • Group Discussion: $5\text{--}10$ minutes to evaluate with peers.

  • Class Sharing: $5\text{--}10$ minutes for open class presentation.