BIOL 1406-46 Exam Review Notes

Chapter 1: Evolution, Themes of Biology, and Scientific Inquiry

  • Concept of Fitness in Natural Selection

    • Definition: Fitness refers to an organism's ability to survive and reproduce in its environment, not merely physical strength or size.

    • Key Note: It encompasses adaptability to environmental changes and reproductive success rather than just brute strength.

  • Key Scientific Concept by Darwin and Wallace

    • Concept: Theory of Natural Selection.

    • Description: Both scientists proposed that organisms evolve through natural selection, where those with advantageous traits are more likely to survive and reproduce, passing those traits to future generations.

  • Peppered Moth Example During the Industrial Revolution

    • Explanation: The color variation in moths (light and dark) demonstrated how industrial pollution affected camouflage, leading to a population shift towards darker moths as they blended better with soot-covered trees, illustrating natural selection in action.

  • White and Black Beans in Natural Selection Simulation Lab

    • Purpose: These beans were used to simulate predation.

    • Biological Principle: Demonstrated how certain traits (color) affect survival and reproductive success, illustrating natural selection in a controlled environment.

  • Three Domains of Life

    • Domains:

    • Bacteria: Example - Escherichia coli (E. coli).

    • Archaea: Example - Methanogens.

    • Eukarya: Example - Homo sapiens (humans).

  • Unity of Life Definition

    • Explanation: The concept that all living organisms share common characteristics and biological processes due to common ancestry.

    • Relation to Common Ancestry: Indicates that diverse forms of life evolved from shared ancestors, underpinning the interconnectedness of life forms.

Chapter 2: The Chemical Context of Life

  • Charge and Location of Protons

    • Charge: Positive (+1).

    • Location: Found in the nucleus of an atom.

  • Isotopes Description

    • Definition: Variants of a chemical element that have the same number of protons but a different number of neutrons.

    • Difference from Standard Element: Isotopes have different mass numbers due to the varying number of neutrons.

  • Structure of an Atom

    • Components:

    • Protons: Positive charge, located in the nucleus, relative mass = 1 amu.

    • Neutrons: Neutral charge, located in the nucleus, relative mass = 1 amu.

    • Electrons: Negative charge, outside the nucleus, relative mass ~ 1/1836 amu.

  • Definition of a Compound

    • Definition: A chemical substance formed from two or more different elements bonded together.

    • Example: Water (H₂O) differs from its elements, hydrogen and oxygen, in properties and composition.

  • Definition of an Element

    • Definition: A pure substance that cannot be reduced to simpler substances by chemical means.

    • Reason: Elements consist of only one type of atom, defined by the number of protons in their nuclei.

  • Covalent Bonds in Carbon

    • Covalent Bonds: Carbon can typically form four covalent bonds.

    • Importance: Enables the formation of complex organic molecules essential for life.

  • Definition of an Ion

    • Definition: An atom or molecule that has gained or lost one or more electrons, resulting in a net electric charge.

    • Formation: An atom becomes a cation (positively charged) by losing electrons and an anion (negatively charged) by gaining electrons.

  • Mass Number Calculation

    • Given Element: Has 7 protons and 7 neutrons.

    • Mass Number: The mass number is calculated as the sum of protons and neutrons, resulting in 14.

    • Representation: Mass number indicates the total number of protons and neutrons in an atom's nucleus.

  • Atomic Number Definition

    • Definition: The atomic number is the number of protons in the nucleus of an atom.

    • Significance: It determines the element's identity and its position in the periodic table.

  • Mass Number of an Atom

    • Definition: Mass number is the total count of protons and neutrons in an atom.

    • Calculation: ( ext{Mass Number} = ext{Number of Protons} + ext{Number of Neutrons})

  • Types of Elements and Ionic Bonds

    • Elements likely to form ionic bonds: Metals (especially alkali and alkaline earth metals) because they typically lose electrons.

  • Ionic Bond Formation

    • Description: Ionic bonds form through the transfer of electrons from one atom to another, resulting in the attraction between oppositely charged ions.

    • Example: Sodium chloride (NaCl) is a compound with ionic bonds.

  • Nonpolar Covalent Bond

    • Definition: A bond where electrons are shared equally between two atoms, resulting in no charge separation.

    • Example: A molecule of methane (CH₄) exhibits this type of bond.

  • Polar Covalent Bond

    • Definition: A bond formed when electrons are not shared equally between two atoms, resulting in a partial positive and negative charge.

    • Example: Water (H₂O) has polar covalent bonds due to the electronegativity difference between hydrogen and oxygen.

  • Nonpolar Molecule Definition

    • Definition: A molecule that has an even distribution of electrical charge, resulting in no significant dipole.

    • Example: O₂ (oxygen gas), which is nonpolar because of equal sharing of electrons.

  • Determining Atomic Reactivity

    • Factors: An atom's chemical reactivity is determined by the number of valence electrons present in its outer shell.

    • Role of Valence Electrons: Atoms tend to react to achieve a full outer electron shell, typically following the octet rule.

  • Most Electronegative Element

    • Element: Fluorine (F) is the most electronegative element.

    • Implication: It tends to attract electrons strongly and is likely to form polar covalent or ionic bonds with less electronegative elements.

  • Valence Electrons Definition

    • Definition: Electrons in the outermost shell of an atom that participate in chemical bonding.

    • Importance: They are crucial for determining an element’s chemical behavior and bonding capacity.

  • Cations and Anions

    • Differences:

    • Cation: A positively charged ion formed by losing one or more electrons.

    • Anion: A negatively charged ion formed by gaining one or more electrons.

  • Electronegativity Definition

    • Definition: A measure of an atom’s ability to attract electrons in a chemical bond.

    • Influence on Bond Type: Higher electronegativity differences between atoms lead to polar covalent or ionic bonds.

  • Example of a Compound with Covalent Bonds

    • Example: Carbon dioxide (CO₂), which involves sharing of electrons between carbon and oxygen.

  • Valence Electrons in Carbon

    • Number: Carbon has 4 valence electrons.

    • Effect on Bonding: This allows carbon to form four covalent bonds, essential for creating complex organic molecules.

  • Valence Electrons in Oxygen

    • Number: Oxygen has 6 valence electrons.

    • Effect on Bonding: Oxygen tends to gain or share two electrons to fill its outer shell, often forming two covalent bonds.

  • Type of Bond with Complete Transfer of Electrons

    • Type: Ionic bond involves a complete transfer of electrons from one atom to another.

Chapter 3: Water and Life

  • H⁺ Concentration and pH Change

    • Explanation: When pH decreases from 7 to 6, the H⁺ ion concentration increases tenfold, as the pH scale is logarithmic.

  • Indication of pH of 8

    • Description: A pH of 8 indicates the solution is basic (alkaline), as it is above 7 on the pH scale.

  • Approximate pH of Human Blood

    • Range: Normal human blood pH is approximately 7.4.

    • Importance: This narrow range is crucial for proper physiological functions and enzyme activities.

  • pH Scale Measurement

    • Description: The pH scale measures the acidity or basicity of a solution based on the concentration of H⁺ ions.

    • Relationship: A decrease in pH corresponds to an increase in H⁺ ion concentration and vice versa.

  • Bonding in Water Molecules

    • Type: Hydrogen bonds attract one water molecule to another due to the polarity of water.

  • Specific Heat Definition

    • Definition: Specific heat is the amount of heat required to raise the temperature of a substance by one degree Celsius.

    • Water's Role: Water’s high specific heat allows it to stabilize temperature variations in living organisms.

  • Ice Floating on Water

    • Explanation: Ice floats on liquid water because it is less dense than liquid water, a unique property that is crucial for the survival of aquatic ecosystems during freezing conditions.

  • Essentiality of Water to Life

    • Reasons:

    • Solvent Properties: Water is an excellent solvent for ionic and polar substances, facilitating chemical reactions in organisms.

    • Temperature Regulation: Water's high specific heat enables organisms to maintain stable internal temperatures.

  • Cohesion vs. Adhesion in Water

    • Definitions:

    • Cohesion: The tendency of water molecules to stick to each other due to hydrogen bonding.

      • Example: Water droplets forming on a leaf surface.

    • Adhesion: The attraction between water molecules and other substances.

      • Example: Water climbing up a plant stem against gravity.

  • Buffers in Biology

    • Definition: Buffers are substances that minimize changes in pH by accepting or donating H⁺ ions.

    • Importance: They are essential for maintaining pH stability in biological systems, ensuring proper enzyme function and metabolic processes.