2.4 Study Notes on Inorganic Compounds Essential to Human Functioning

Inorganic Compounds Essential to Human Functioning

Learning Objectives

  • Compare and contrast inorganic and organic compounds.

  • Identify the properties of water that make it essential to life.

  • Explain the role of salts in body functioning.

  • Distinguish between acids and bases, and explain their role in pH.

  • Discuss the role of buffers in helping the body maintain pH homeostasis.

Overview of Inorganic Compounds

  • Inorganic compounds are defined as substances that do not contain both carbon and hydrogen.

    • Many inorganic compounds do contain hydrogen, such as

    • Water (H₂O)

    • Hydrochloric acid (HCl)

    • Carbondioxide (CO₂) is one of the few inorganic compounds that contains carbon.

  • Organic compounds contain both carbon and hydrogen, synthesized via covalent bonds within living organisms.

    • Carbon and hydrogen are the second and third most abundant elements in the human body.

Essential Inorganic Compounds for Life

Water

  • Make up approximately 70% of an adult's body weight.

  • Contained in cells and the interstitial spaces between them.

  • Roles of Water in the Body:

    • As a Lubricant and Cushion:

    • Major component of lubricating fluids like synovial fluid for joints and pleural fluid for lung function.

    • Keeps food flowing in the digestive tract, reducing friction between organs.

    • Cushions organs against physical trauma (e.g., brain, eyes, fetus during pregnancy).

    • As a Heat Sink:

    • Water absorbs heat generated by chemical reactions without significant temperature increase.

    • When environmental temperature rises, stored body water helps keep the body cool by facilitating sweat evaporation.

    • As a Component of Liquid Mixtures:

    • Mixture Definition: Combination of two or more substances where each maintains its chemical identity (e.g., flour and sugar in a bowl).

    • Types of liquid mixtures involving water:

      • Solutions: Homogeneous mixtures where the solute is evenly distributed (e.g., sugar dissolved in water).

      • Colloids: Mixtures containing large molecules that scatter light (e.g., milk, thyroid hormone as a colloid).

      • Suspensions: Temporary mixtures that separate over time (e.g., blood settle test).

Concentrations of Solutes

  • Concentration defined as the number of particles of solute in a specific volume of solution.

  • Measurement Standards:

    • Glucose in blood typically measured in mg/dL, with a healthy adult average of approximately 100 mg/dL.

    • Molarity (M) is another measurement, defined as moles of solute per liter of solution.

  • Example:

    • Glucose Example:

    • Chemical formula: C₆H₁₂O₆.

    • Molecular weight calculated as follows:

      • Carbon (C): 12.011 g, total for 6 carbons = 72.066 g

      • Hydrogen (H): 1.008 g, total for 12 hydrogens = 12.096 g

      • Oxygen (O): 16.00 g, total for 6 oxygens = 96.00 g

      • Overall molecular weight = 180.156 g.

    • Avogadro's Number: 6.02 × 10²³ particles/mole.

The Role of Water in Chemical Reactions

  • Two key chemical reactions involving water:

    • Dehydration Synthesis:

    • Involves the release of a water molecule as two reactants combine (loss of H from one reactant and OH from another).

    • This reaction is also known as a condensation reaction.

    • Hydrolysis:

    • A water molecule disrupts a compound, breaking its bonds (water splits into H and OH).

    • This reaction is reversible and important in organic chemistry.

Salts

  • Salts form when ions engage in ionic bonding, where one atom loses electrons (becoming positively charged) and another gains electrons (becoming negatively charged).

    • A salt is any substance that, when dissolved in water, dissociates into ions other than H⁺ or OH⁻.

    • Example: Sodium chloride (NaCl) dissociates in water, with water molecules surrounding and stabilizing the dissociated ions.

  • Salts behave as electrolytes; they conduct electrical currents in solution.

  • Importance: Ions play critical roles in nerve impulse transmission and muscle contraction.

Acids and Bases

  • Acids

    • Defined as substances that release hydrogen ions (H⁺) in solution.

    • Strong Acids: Ionize completely in solution (e.g., hydrochloric acid, HCl).

    • Weak Acids: Partially ionize, some hydrogen ions remain bonded in compound (e.g., acetic acid).

  • Bases

    • Defined as substances that release hydroxyl ions (OH⁻) in solution or accept H⁺ ions already present.

    • Strong Bases: Release most (or all) hydroxyl ions; Weak Bases: Release only some or accept few H⁺ ions.

The Concept of pH

  • pH Definition: The negative base-10 logarithm of the hydrogen ion concentration in a solution.

    • pH is a measure of a solution's acidity or alkalinity; a pH of 7 is neutral.

    • Lower pH indicates higher acidity (H⁺ concentration), while higher pH indicates increased basicity (lower H⁺ concentration).

    • For example, a solution with pH 4 has ten times greater H⁺ concentration than a solution with pH 5.

Buffers

  • Blood pH normally ranges from 7.35 to 7.45, with a commonly referenced pH of 7.4.

  • Buffers are solutions of a weak acid and its conjugate base, which help maintain stable pH levels in bodily fluids.

    • If pH decreases (becomes more acidic), the buffer can bind excess H⁺ ions; conversely, if pH rises (becomes more basic), the buffer can release H⁺ ions.

  • Conditions of acidity in body fluids: Acidosis, caused by ineffective breathing or metabolic issues leading to excess acids.

  • Conversely, alkalosis is a condition where body fluids are too alkaline, caused by various factors like respiratory issues or prolonged vomiting.