Chemical Level of Organization (Chapter 2)

Inorganic Chemistry

  • Definition: Study of structure and interaction of matter, which occupies space and has mass.

Levels of Organization

  1. Chemical Level

    • Atoms (Examples: C, H, O, N, P)
    • Molecule (Example: DNA)

  2. Cellular Level

    • Not specified in detail.

  3. Tissue Level

    • Not specified in detail.

  4. Organ Level

    • Stomach
    • Serous membrane
    • Epithelial tissue
    • Smooth muscle tissue layers

  5. System Level

    • Digestive system components:
      • Esophagus
      • Liver
      • Stomach
      • Pancreas
      • Gallbladder
      • Small intestine
      • Large intestine

Organization of Matter

  • Chemical Elements
    • Total of 112 elements identified.
    • An element cannot be split into a simpler substance by ordinary chemical means.
    • Elements present in the human body: 26 total.
    • CHON (Carbon, Hydrogen, Oxygen, Nitrogen) compose 96% of body mass.
    • Additional elements: Calcium (Ca), Phosphorus (P), Potassium (K), Sulfur (S), Sodium (Na), Chlorine (Cl), Magnesium (Mg), Iron (Fe) contribute an added 3.6%.
    • Trace elements (0.2% of body mass) are essential despite being present in very small amounts.

Structure of Atoms

  • Atoms: The smallest particles of matter that maintain the properties of their element.
  • Structure of the Atom: Includes
    • Nucleus
      • Protons (p+)
      • Neutrons (nº)
    • Electrons (e¯)
    • Electron shells (depicted in Figure 2.1)

Atomic Number and Mass Number

  • Atomic Number: Number of protons in an atom.
    • Different elements possess different atomic numbers.
  • Mass Number: Sum of protons and neutrons in the nucleus of an atom.
  • Isotopes: Atoms of an element that have different mass numbers due to varying numbers of neutrons.
  • Radioactive Isotopes: Unstable isotopes that decay over time into more stable forms and can transform into different elements.
  • Half-Life: The duration required for half of the atoms of an isotope to decay into another form.
    • Uses of radioactive isotopes: Important in medicine and research.

Atomic Interactions

  • Atoms consist of: Protons, Neutrons, Electrons
  • Atoms can gain or lose electrons, causing:
    • An atom gaining or losing electrons becomes an ion of the same element.
    • An atom gaining or losing protons becomes a different element.
    • An atom gaining or losing neutrons becomes an isotope of the same element.

Examples of Atomic Structure

  • Hydrogen (H):
    • Atomic number = 1
    • Mass number = 1 or 2
    • Atomic mass = 1.01
  • Carbon (C):
    • Atomic number = 6
    • Mass number = 12 or 13
    • Atomic mass = 12.01
  • Nitrogen (N):
    • Atomic number = 7
    • Mass number = 14 or 15
    • Atomic mass = 14.01
  • Oxygen (O):
    • Atomic number = 8
    • Mass number = 16, 17, or 18
    • Atomic mass = 16.00
  • Sodium (Na):
    • Atomic number = 11
    • Mass number = 23
    • Atomic mass = 22.99
  • Chlorine (Cl):
    • Atomic number = 17
    • Mass number = 35 or 37
    • Atomic mass = 35.45
  • Potassium (K):
    • Atomic number = 19
  • Iodine (I):
    • Atomic number = 53
    • Mass number = 39, 40, or 41
    • Atomic mass = 126.90

Atomic Mass

  • Measured in daltons (amu).
  • The mass of a single atom is approximately equal to the sum of the masses of its protons, neutrons, and electrons.

Chemical Bonds

  • Types of Bonds:
    • Ionic Bonds: Formed through the transfer of valence electrons leading to the attraction of positively and negatively charged ions.
    • Cations: Ions with a positive charge.
    • Anions: Ions with a negative charge.
    • Covalent Bonds: Involves sharing of electrons, can be strong or weak depending on the number of electron pairs shared.
    • Nonpolar covalent bonds share electrons equally.
    • Polar covalent bonds result in unequal sharing of electrons and create regions of electronegativity.
    • Hydrogen Bonds: Formed between molecules due to polar covalent bonds; key in the shape and structure of large molecules such as proteins.

Chemical Reactions

  • Definition: Involves the making and breaking of chemical bonds with reactants transforming into products.
  • Metabolism: Refers to all chemical reactions occurring in the body.
  • Activation Energy: The energy threshold required to break bonds in reactants before a reaction can occur and can be influenced by concentration and temperature.

Inorganic Compounds and Solutions

  • Definition: Compounds that lack carbon.
  • Acid: Dissociates into hydrogen ions (H+) and one or more anions. Described as proton donors.
  • Base: Dissociates into hydroxide ions (OH-) and one or more cations. Known as proton acceptors.
  • Salt: In water, dissociates into cations and anions that are not H+ or OH-.

Properties of Water

  • Polarity: Essential for life; it is an excellent solvent for polar compounds and interacts with ionized substances.
  • High Heat Capacity: Water can absorb large amounts of heat without significant temperature changes, crucial for thermoregulation.
  • Cohesion and Surface Tension: Water molecules exhibit cohesion due to hydrogen bonding, resulting in high surface tension.

pH and Acid-Base Balance

  • pH Scale: Ranges from 0-14; a logarithmic measure of acidity or alkalinity.
    • pH = (−log[H+])
    • Normal pH of the body: 7.35-7.45
    • Common substances with pH values such as saliva (6.35-6.85), pure water (7.0), blood (7.35-7.45), etc.
  • Buffer Systems: Help maintain pH by converting strong acids or bases into weaker counterparts.

Enzymes as Biological Catalysts

  • Definition: Enzymes are catalysts in living cells that speed up chemical reactions.
  • Characteristics:
    • Highly specific to substrates.
    • Subject to cellular controls.
    • Speed up reactions by increasing frequency of collisions, lowering activation energy, and properly orienting molecules.
  • Composition: Enzymes consist of an apoenzyme (protein part) and a cofactor (non-protein part). Named typically using the suffix ‘-ase’.
  • Efficiency: Can speed up reactions by a factor of up to 10 billion times.
  • Clinical Applications: Disorders such as galactosemia highlight the importance of enzymes; in this case, the lack of the enzyme galactase leads to the harmful accumulation of galactose in the body.