Chapter 2: The Chemical Basis of Life, I: Atoms, Molecules, and Water

Chapter 2: The Chemical Basis of Life

Key Concepts

  • Atoms

  • Chemical Bonds and Molecules

  • Properties of Water

  • pH and Buffers

Atoms

  • Definition: Atoms are the smallest functional units of matter that form all chemical substances.

  • Element: An atom is the smallest unit of an element that possesses the chemical properties of that element.

  • Chemical Elements: Each specific type of atom is a chemical element (e.g., Nitrogen, Oxygen, Helium).

Composition of Atoms

  • Atoms are composed of three subatomic particles:

    1. Protons: Positive charge (+), found in the nucleus.

    2. Neutrons: No charge (neutral), found in the nucleus.

    3. Electrons: Negative charge (−), found in orbitals surrounding the nucleus.

Distinctions Among Atoms

  • Atomic Number: The number of protons in an atom that also determines its position in the periodic table and is equal to the number of electrons in a neutral atom (exceptions exist in ions).

  • Atomic Mass: Approximately equal to the sum of the number of protons and the number of neutrons in an atom.

Characteristics of Subatomic Particles

  • Protons and electrons balance each other in number to give the atom a net charge of zero (in the absence of ions).

  • Ions: Atoms with a net charge due to the loss or gain of one or more electrons (e.g., Ca2+).

  • Isotopes: Variations of atoms where the number of neutrons differs (e.g., Carbon-12 vs. Carbon-14).

Electron Configuration

  • Orbitals: Electrons occupy orbitals, which are regions surrounding the nucleus where the probability of finding an electron is high. These are visualized as electron clouds.

Structure of Orbitals

  • s Orbitals: Spherical shape.

  • p Orbitals: Dumbbell or propeller shaped.

  • Each orbital can accommodate a maximum of 2 electrons.

Electron Shells and Their Capacities

  • 1st Shell: Contains 1 spherical orbital (1s), holding up to 2 electrons.

  • 2nd Shell: Contains 4 orbitals—1s (2 electrons) and 3p (up to 6 electrons), thus accommodating up to 8 electrons in total.

  • Example: Nitrogen Atom:

    • Has 7 protons and 7 electrons.

    • 2 electrons fill the 1s orbital, and the remaining 5 fill the 2nd shell (2 in 2s and 1 in each of the three 2p orbitals).

Periodic Table

  • Organization: The periodic table is arranged by increasing atomic number.

  • Rows (Periods): Indicate the number of electron shells (e.g., first row has 1 shell, second row has 2, etc.).

  • Columns (Groups): Indicate the number of valence electrons, which determines similar properties among elements within the same column.

  • Columns are labeled:

    • Column 1: 1 valence electron

    • Column 2: 2 valence electrons

    • Column 3: 3 valence electrons

Atomic Mass and Isotopes

  • Atomic Mass: Roughly equal to the combined number of protons and neutrons, with protons and neutrons being over 1,800 times heavier than electrons.

  • Example: Carbon has an average atomic mass of 12.011 due to different isotopes contributing to the average.

Mass vs. Weight

  • Mass: The amount of matter in a substance, which remains constant regardless of location.

  • Weight: The gravitational pull on a mass, which can vary (e.g., a person weighing 154 pounds on Earth weighs 25 pounds on the Moon but would weigh 21 trillion pounds on a neutron star).

Units of Measurement

  • Dalton (Da): A unit for atomic mass, approximately equal to the mass of one proton.

  • Mole: Represents a quantity containing the same number of atoms as there are in 12 grams of carbon-12, known as Avogadro’s number (approximately 6.022 imes 10^{23} atoms).

Key Elements in Living Organisms

  • Major Elements: Hydrogen, oxygen, carbon, and nitrogen make up about 95% of the atoms in living organisms.

    • Hydrogen and oxygen are primarily in water.

    • Nitrogen is found in proteins.

    • Carbon is the foundation for all living matter.

  • Mineral Elements: Less than 1% of living organisms.

  • Trace Elements: Less than 0.01%, crucial for normal growth and function.

Chemical Bonds and Molecules

  • Molecule: Composed of two or more atoms bonded together.

    • Molecular Formula: Indicates the chemical symbols of the elements in a molecule, where subscripts show the quantity of each atom (e.g., C6H{12}O_6 for glucose).

  • Compound: Any molecule with two or more different elements (e.g., C6H{12}O6 is a compound, while N2 and O_2 are not compounds).

Types of Chemical Bonds

  1. Covalent Bond

    • Electrons are shared to fill valence shells.

    • Types:

    • Polar Covalent Bonds: Electrons shared unequally between atoms with different electronegativities.

    • Nonpolar Covalent Bonds: Equal sharing of electrons occurs between atoms with similar electronegativities.

  2. Hydrogen Bond

    • A hydrogen atom from one polar molecule is attracted to an electronegative atom of another molecule.

    • Generally weak individually, but collectively can be strong (e.g., holding DNA strands together).

  3. Ionic Bond

    • Occurs when electrons are transferred, forming oppositely charged ions that attract each other, such as in sodium chloride (NaCl).

Covalent Bonds

  • Atoms share pairs of electrons, forming strong bonds as the shared electrons behave as if they belong to both atoms.

  • Covalent bonds can be single (one pair), double (two pairs), or triple (three pairs).

    • Example:

    • 1 pair: H-F (single bond)

    • 2 pairs: O=O (double bond)

    • 3 pairs: N ext{≡} N (triple bond)

Octet Rule

  • Atoms are stable when their outer shell is full, typically with 8 electrons.

  • Exceptions: Hydrogen and helium, which fill their shell with 2 electrons.

Example: Oxygen Atom

  • Oxygen has 6 valence electrons. With 8 protons and electrons:

    • 2 electrons fill the 1s orbital.

    • 6 electrons in the 2nd shell (2 in 2s and 4 in 2p).

Nonpolar vs. Polar Covalent Bonds

  • Nonpolar Covalent Bonds:

    • Atoms have similar electronegativity and share electrons equally (e.g., O_2).

  • Polar Covalent Bonds:

    • Atoms have different electronegativities, causing an unequal sharing of electrons, resulting in a charge separation in the molecule (e.g., in water, H2O).

Water: The Polar Molecule

  • Characteristics of Water:

    • A classic example of polar covalent bonds due to the unequal sharing of electrons between oxygen and hydrogen atoms.

    • Leads to partial positive charges around hydrogen atoms and a partial negative charge around the oxygen atom.

  • Hydrogen Bonds: Formed between polar molecules, are weak individually but contribute to the stability of larger structures (e.g., DNA).

Ionic Bonds

  • Definition: An ion is an atom or molecule that has gained or lost electrons.

    • Cations: ions with a net positive charge (e.g., Na+).

    • Anions: ions with a net negative charge (e.g., Cl-).

  • Ionic bonds form when a cation binds to an anion through electrostatic attraction, resulting in ionic compounds or salts (e.g., NaCl).

Properties of Water

  • Water is not just a solvent but also participates in various biological functions:

    • Acts as a lubricant, provides support, and helps in evaporative cooling (e.g., sweat).

  • Water can dissolve ionic compounds and polar molecules, indicating its role as a solvent.

Chemical Reactions

  • Involves the transformation of reactants into products, generally requiring energy from chemical reactions and often needing enzymes as catalysts.

  • Equilibrium in reactions indicates a dynamic balance between reactants and products.

Acids, Bases, and pH

  • Acids: Release hydrogen ions (H+) when dissolved in water (e.g., HCl).

    • Strong acids release more H+ than weak acids (e.g., carbonic acid H2CO3).

  • Bases: Lower the H+ concentration by either releasing OH- ions or binding H+ ions (e.g., NaOH).

pH Scale

  • Ranges from 0 to 14, where:

    • pH < 7 is acidic (H+ concentration > OH-).

    • pH = 7 is neutral (H+ = OH-).

    • pH > 7 is alkaline (H+ concentration < OH-).

  • Acidic solutions, like human stomach fluid (pH 1.3), and alkaline substances like bleach (pH 12.5) reflect the acidity or basicity of certain substances.

Effects of pH

  • pH impacts molecular shape, reaction rates, binding abilities, and solubility.

Buffers

  • Buffers help maintain a constant pH in biological systems, typically found as acid-base pairs.

  • Human blood is a weakly alkaline buffer system, maintaining pH around 7.35 - 7.45.

  • Response of buffers to changes in pH includes removing or releasing H+ to stabilize the pH.

Note: The information presented in this guide is comprehensive and reflects all key concepts, definitions, examples, and important details relevant to understanding the chemical basis of life and its properties.