BIO 120 Lecture Notes - 100 Practice Flashcards

What is Science? (Transcript: Page 3)

• Science is inherently a systematic and rigorous process for acquiring knowledge and understanding about the natural world.

• It involves the structured search for information, explanations, and patterns, typically driven by specific questions or phenomena.

• Scientific inquiry is a cornerstone of this process, characterized by hypothesis-based testing designed to explain natural occurrences.

• A scientific theory represents a broader, extensively verified, and well-supported explanation for a wide range of natural phenomena. Unlike a hypothesis, a theory is a comprehensive framework built upon a substantial body of evidence and is subject to continuous refinement through ongoing experimentation and observation.

The Process of Science / The Scientific Method (Transcript: Page 4-5)

• Scientific Inquiry: This is the fundamental search for information and explanations, initiated by a specific question or an intriguing observation.

• The core of this method involves hypothesis-based testing, aiming to provide plausible explanations for observed natural phenomena.

• The Generalized Flow of Reasoning in the Scientific Method is as follows:

  • Observations: Detailed and objective observations of natural phenomena or experimental results.

  • Hypothesis: A testable, falsifiable, and proposed explanation for an observation. It is an educated guess based on prior knowledge and observations.

  • Predictions: Logical deductions based on the hypothesis, stating what specific results are expected if the hypothesis is true. These are typically in an "if…then…" format.

  • Experiments or new observations: Controlled tests designed to evaluate the predictions. These experiments generate data that either supports or refutes the hypothesis.

  • Theory: If the results consistently support the hypothesis across many experiments and related studies, the hypothesis may contribute to or evolve into a scientific theory.

• Iterative Process: If experimental results are not consistent with the predictions, the hypothesis must be rejected or revised. This suggests the initial explanation was incorrect or incomplete.

• Building a Theory: When a hypothesis is consistently supported by numerous experiments and observations, and withstands attempts at falsification, it gains significant credibility. If multiple related, well-supported hypotheses converge, they can form the basis of a comprehensive scientific theory.

• Conceptual loop: The process is cyclical; new experiments or observations continuously refine or challenge existing theories (Experiments or new observations $\, \rightarrow \,$ Theory (if consistent) $\, \infty$).

Visual/Reference Pages (Transcript: Page 6)

• Magazine references and links listed, serving as visual or supplementary resources:

  • magazine.jhsph.edu (Johns Hopkins School of Public Health magazine)

  • personalbrandingblog.com

  • "Freeze-factured Mimivirus-Infected Acanthamoeba" (Likely an image or research topic related to virology/microscopy)

  • Hathan Zoubertan and Eyal Shimoni (likely authors or researchers associated with the image/topic)

  • The Welmann (Institution or publication, possibly 'Weizmann Institute of Science' given the context)

  • baitute of Science (likely a misrendered title for an institute of science)

  • uwo.ca (University of Western Ontario domain)

  • biosingularity.wordpress.com

Characteristics of Life (Transcript: Page 7)

Living organisms exhibit several key characteristics that distinguish them from non-living matter:

1. Order: Highly organized and complex structures at various levels (cells, tissues, organs).

2. Regulation: Maintenance of a stable internal environment (homeostasis) despite external changes.

3. Energy Processing: Intake and transformation of energy to perform work (e.g., metabolism).

4. Growth and Development: Consistent growth and development guided by inherited genetic information.

5. Reproduction: Ability to produce offspring, either sexually or asexually.

6. Response to Environment: Reactions to stimuli from the environment.

7. Evolutionary Adaptation: Capacity to adapt over generations to their environment, leading to evolutionary change.

Periodic Table and Essential Elements (Transcript: Page 8-9)

• The Periodic Table of Elements is a fundamental organizational tool in chemistry, showcasing the arrangement of all known chemical elements based on their atomic number, electron configuration, and recurring chemical properties. This arrangement allows for predictions about their reactivity and behavior.

• Common Essential Elements in Living Matter: Life as we know it is primarily composed of a relatively small set of elements, vital for constructing biological molecules and carrying out life processes. These include:

  • Macroelements (major components): Carbon (C), Oxygen (O), Hydrogen (H), Nitrogen (N) – these four make up about 96% of living matter and form the backbone of organic molecules (carbohydrates, lipids, proteins, nucleic acids).

  • Minor elements: Phosphorus (P) – crucial for nucleic acids (DNA, RNA) and ATP; Sulfur (S) – found in certain amino acids and proteins; Calcium (Ca) – vital for bone structure, muscle contraction, and nerve function; Potassium (K) – essential for nerve impulses and fluid balance; Chlorine (Cl) – involved in fluid balance and digestion; Magnesium (Mg) – a cofactor in many enzymatic reactions and central to chlorophyll.

  • Trace elements: Elements required in very minute quantities (e.g., Iron (Fe), Copper (Cu), Zinc (Zn)), often serving as cofactors for enzymes or components of complex molecules.

• Roles of these elements for organisms:

  1. Growth: Providing the building blocks necessary for cell proliferation and overall increase in size and complexity.

  2. Maintenance of structure and function: Sustaining existing cellular and physiological processes, including membrane integrity, enzymatic activity, and signal transduction.

  3. Repair: Replacing damaged cells and biomolecules to ensure the organism's continued health and function.

• Note: The specific ratios and availability of these elements are critical for healthy biological systems, as imbalances can lead to various physiological dysfunctions.

Atoms and Elements (Transcript: Page 10)

• An atom is the smallest unit of matter that retains an element's chemical identity. It consists of a dense central nucleus surrounded by a cloud of negatively charged electrons.

• The nucleus contains two types of subatomic particles:

  • Protons: Positively charged particles (charge of $+\text{1.602} \times \text{10}^{-\text{19}}\text{ C}$), with a mass of approximately $1 \text{ atomic mass unit (amu)}$. The number of protons defines the atomic number (Z), which uniquely identifies an element.

  • Neutrons: Electrically neutral particles, with a mass very similar to protons (approximately $1 \text{ amu})$). The number of neutrons in an atom can vary, leading to isotopes of the same element, which have the same atomic number but different mass numbers.

• Electrons: Negatively charged particles (charge of $-\text{1.602} \times \text{10}^{-\text{19}}\text{ C}$), with a negligible mass (approx. $1/\text{1836}$ of a proton's mass) compared to protons and neutrons. Electrons move rapidly in specific energy levels called orbitals or shells around the nucleus.

• An element is a pure substance consisting only of atoms that all have the same number of protons (i.e., the same atomic number). Therefore, all atoms of a specific element are chemically identical in their fundamental properties.

Electron Shells and Valence (Transcript: Page 11)

• Electron shells represent different energy levels that electrons can occupy around the nucleus. Electrons in shells farther from the nucleus possess higher energy.

• Electron shells and their typical behavior (illustrative examples):

  • First shell (K-shell): This is the innermost shell, positioned closest to the nucleus. It has the smallest capacity, holding a maximum of 2 electrons. For instance, Hydrogen (H) has 1 electron in this shell, and Helium (He) has 2, filling its first shell.

  • Second shell (L-shell): Located further from the nucleus, this shell has a higher energy level and can accommodate up to 8 electrons.

  • Third shell (M-shell) and beyond: These shells are progressively further from the nucleus, corresponding to even higher energy levels and typically holding 8 electrons (for elements up to Calcium, often referred to as the octet rule for stability in bonding, though larger elements can have more).

• Key concepts:

  • Valence electrons: These are the electrons residing in the outermost electron shell of an atom. They are the primary determinants of an atom's chemical properties and its ability to form chemical bonds.

  • The number of valence electrons dictates an atom's bonding potential and how many bonds it is likely to form (or whether it will gain or lose electrons) to achieve a stable electron configuration, typically a full outer shell (often 8 electrons, known as the octet rule, or 2 for the first shell).

  • Each shell has a specific electron capacity, and the rules governing the filling of these shells (e.g., Aufbau principle, Hund's rule, Pauli exclusion principle) profoundly influence an element's chemical behavior and reactivity.

Ionic Bonds (Transcript: Page 12)

• Ionic bonds are strong chemical bonds that form when there is a significant attractive force between two atoms for valence electrons, leading to a complete transfer of one or more electrons from one atom to another.

• This typically occurs between atoms with a large difference in electronegativity (often between a metal and a nonmetal). The more electronegative atom effectively "strips" electrons away from the less electronegative atom.

• As a result of this electron transfer, both atoms become charged ions:

  • The atom that loses electrons becomes a cation (positively charged ion) because it now has more protons than electrons.

  • The atom that gains electrons becomes an anion (negatively charged ion) because it now has more electrons than protons.

• These oppositely charged ions are then strongly attracted to each other by electrostatic forces, forming the ionic bond. This attraction holds them together in a stable ionic compound, often forming crystal lattices.

• Terms:

  • Cation: A positively charged ion (e.g., $Na^+$ from Na losing an electron).

  • Anion: A negatively charged ion (e.g., $Cl^-$ from Cl gaining an electron).

• Example: Sodium chloride (NaCl). Sodium (a metal) readily loses its single valence electron to Chlorine (a nonmetal), which readily accepts an electron. This forms $Na^+$ and $Cl^-$, which are then held together by ionic bonding.

Covalent Bonds (Transcript: Page 13)

• Covalent bonds are strong chemical bonds characterized by the mutual sharing of one or more pairs of valence electrons between two atoms.

• This sharing occurs when both atoms involved have similar (or moderately different) electronegativities, and it allows each atom to achieve a more stable electron configuration, typically by completing its valence shell (e.g., fulfilling the octet rule) without losing or gaining electrons entirely.

• The shared electrons are simultaneously attracted to the nuclei of both atoms, holding them together.

• Nonpolar covalent bonds form when valence electrons are shared equally between two atoms. This typically happens when the atoms are identical (e.g., $O<em>2, H</em>2$) or have very similar electronegativities, resulting in no significant charge separation across the bond.

• Example reference: Methane ($CH_4$) is a classic example of a molecule formed entirely by nonpolar covalent bonds, where carbon shares electrons equally with four hydrogen atoms. Each C-H bond is considered nonpolar because the electronegativity difference is minimal.

• Covalent bonds can also be classified by the number of shared electron pairs:

  • Single bond: Shares 1 pair of electrons (e.g., H-H in $H_2$).

  • Double bond: Shares 2 pairs of electrons (e.g., O=O in $O_2$).

  • Triple bond: Shares 3 pairs of electrons (e.g., N$\equiv$N in $N_2$).

Polar Covalent Bonds (Transcript: Page 14)

• Polar covalent bonds develop when valence electrons are shared unequally between two atoms due to a difference in their electronegativity (the atom's attraction for shared electrons).

• The more electronegative atom pulls the shared electron pair closer to its nucleus, spending more time in its vicinity. This creates an uneven distribution of electron density across the bond.

• In Water ($H_2O$), a prime example of a molecule with polar covalent bonds, electrons are strongly drawn toward the highly electronegative oxygen atom.

• This unequal sharing results in the formation of partial charges within the molecule:

  • The oxygen atom, being more electronegative, acquires a partial negative charge ($\delta^-$) because electrons spend more time near it.

  • The hydrogen atoms, having less electron density around them, acquire partial positive charges ($\delta^+$).

• These partial charges make the molecule polar, meaning it has distinct positive and negative poles, which is crucial for water's unique properties and its role as the solvent of life.

Hydrogen Bond (Transcript: Page 15)

• A hydrogen bond is a type of weak intermolecular (between molecules) or intramolecular (within the same molecule) attractive force.

• It forms specifically between a partially positive hydrogen atom (which is already covalently bonded to a highly electronegative atom like oxygen, nitrogen, or fluorine) and a different, highly electronegative atom (also typically oxygen, nitrogen, or fluorine) that carries a partial negative charge.

• These bonds are significantly weaker than ionic or covalent bonds (about 5-10% the strength of a typical covalent bond), but their cumulative effect can be very strong and biologically significant.

• Example described: The attraction between the partial positive charge ($\delta^+$) on the hydrogen atom of one water molecule ($H<em>2O$) and the partial negative charge ($\delta^-$) on the nitrogen atom of an ammonia molecule ($NH</em>3$).

• Conceptual note: Hydrogen bonds are absolutely critical for biological systems. They are responsible for many of water's unique properties (cohesion, high specific heat), stabilize the double helix structure of DNA, dictate the folding patterns of proteins, and facilitate molecular recognition and binding in countless biological processes.

Molecular Shape and Function (Transcript: Page 16)

• Molecular mimicry: The precise three-dimensional shape of a molecule is not merely a consequence of its atomic composition but is a crucial determinant of its biological function and how it interacts with other molecules.

  • The specific types of atoms present and the geometry of their bonding (bond lengths, bond angles) dictate the overall spatial arrangement and therefore the shape of the molecule.

  • The shape of a molecule is paramount for its function; specifically, molecular recognition often relies on a complementary fit between molecules (like a lock and key).

  • This principle is evident in drug action, enzyme-substrate binding, receptor-ligand interactions, and the assembly of complex cellular structures.

  • Different molecules can exhibit the same function if their shapes are sufficiently similar, even if their atomic composition differs. This phenomenon is known as convergent functional shapes or molecular mimicry when one molecule's shape allows it to bind to the same receptor or enzyme as another, thereby eliciting a similar biological response (e.g., certain drugs mimicking natural neurotransmitters).

Chemical Reactions and Equilibrium (Transcript: Page 17)

• Chemical reactions are processes that involve the breaking of existing chemical bonds and the formation of new ones, leading to a rearrangement of atoms and a change in the composition of matter. Reactants are transformed into products.

• Law of conservation of mass: A fundamental principle stating that in any closed system, matter is neither created nor destroyed during a chemical reaction; it is only rearranged. This means the total mass of the reactants must equal the total mass of the products.

• Reversibility: Many chemical reactions are reversible, meaning they can proceed in both the forward direction (reactants to products) and the reverse direction (products back to reactants).

  • If the rates of the forward reaction and the reverse reaction become equal, the system is said to have reached chemical equilibrium.

  • Equilibrium is a dynamic state where the concentrations of reactants and products remain constant over time, because the rate of their formation equals the rate of their consumption. It is a balance point, not a static halt.

Water as Solvent and Solutions (Transcript: Page 18-19)

• A solvent is the dissolving agent, typically a liquid, within which other substances (solutes) are dispersed.

• A solute consists of the dissolved particles or substances within the solvent. These particles are generally present in lesser amounts than the solvent.

• A solution is a homogeneous mixture (meaning it has a uniform composition throughout) formed when a solute is completely dissolved in a solvent.

• In aqueous solutions, the solvent is specifically water ($H_2O$). Water's remarkable ability to dissolve a wide array of substances makes these solutions fundamental to all biological processes.

• Hydrogen bonds in water: The polarity of individual water molecules, with their partial positive hydrogen atoms ($\delta^+$) and partial negative oxygen atom ($\delta^-$), enables the formation of extensive hydrogen bonds between adjacent water molecules.

  • These hydrogen bonds form between the $\delta^+$ of one water molecule's hydrogen and the $\delta^-$ of another water molecule's oxygen.

  • Water's inherent polarity and its extensive network of hydrogen bonds bestow upon it many unique and crucial properties, including its high specific heat, cohesive and adhesive forces, and its capacity to act as the "solvent of life."

Characteristics of Water (Transcript: Page 20)

Water is essential for life due due to its unique physical and chemical properties:

• Polarity: Water molecules are polar due to the uneven sharing of electrons in their covalent bonds; the oxygen atom is more electronegative than hydrogen, leading to a partial negative charge on oxygen ($\delta^-$) and partial positive charges on hydrogen ($\delta^+$). This molecular polarity is the basis for many other unique properties.

• Cohesion: The strong attraction between water molecules, primarily due to numerous hydrogen bonds, causes them to stick together. This quality is responsible for surface tension and plays a key role in the transport of water up the xylem in plants.

• Adhesion: Water's ability to form hydrogen bonds with other polar surfaces (e.g., the walls of plant vessels) is called adhesion. This property, along with cohesion, underlies capillary action.

• Surface tension: This is a measure of how difficult it is to stretch or break the surface of a liquid. Water has an unusually high surface tension because of the strong cohesive forces (hydrogen bonds) between its surface molecules, allowing some insects to walk on water.

• High specific heat: Water has an exceptionally high specific heat, meaning it takes a large amount of thermal energy to raise its temperature by a given amount. This is primarily due to the extensive hydrogen bonding, which absorbs and releases heat energy, helping to stabilize temperatures in organisms and on Earth's climate.

• Solvent of life: Water's polarity allows it to readily dissolve many ionic compounds and polar molecular compounds. It forms hydration shells around ions and hydrogen bonds with polar molecules, facilitating chemical reactions and acting as a transport medium for nutrients and waste in living systems.

Hydrophilic vs Hydrophobic (Transcript: Page 21)

• Hydrophilic ("water-loving"): These are substances that have an affinity for water and readily mix with, dissolve in, or are wetted by water. This characteristic arises from their ability to form hydrogen bonds with water molecules.

  • Mechanism: Hydrophilic substances are typically ionic compounds (like NaCl, where water molecules surround and effectively separate the $Na^+$ and $Cl^-$ ions, forming hydration shells) or polar molecular compounds (like NH$_3$, sugars, and proteins, which contain polar bonds or charged groups that can form hydrogen bonds with water).

  • Forming these favorable interactions stabilizes the dissolved state, making the dissolving process energetically favorable.

• Hydrophobic ("water-fearing"): These are substances that do not have an affinity for water, do not dissolve in water, and are repelled by water. They cannot form hydrogen bonds with water.

  • Mechanism: Hydrophobic substances are typically non-polar compounds (e.g., oils, fats, waxes). When non-polar molecules are introduced to water, they disrupt the energetically favorable hydrogen bond network of water, forcing water molecules to orient themselves around the non-polar substance in an ordered, cage-like structure (clathrate cage), which is energetically unfavorable.

  • To minimize this unfavorable interaction and maximize water-water hydrogen bonding, hydrophobic substances tend to aggregate together, away from water.

Water Properties Review (Transcript: Page 22)

This section serves as a summary of water's critical properties and their underlying explanations:

• Polarity: Regions of partial positive charge ($\delta^+$ on H) and partial negative charge ($\delta^-$ on O) attributed to oxygen's higher electronegativity and the bent molecular geometry.

• Cohesion: The strong mutual attraction between water molecules themselves, mediated by the extensive hydrogen bond network, giving water adhesive properties to itself.

• Adhesion: The attraction between water molecules and other distinct polar surfaces or charged molecules, also occurring via hydrogen bonding.

• High surface tension: A direct consequence of the strong cohesive forces (hydrogen bonding) that create a tough, elastic-like membrane at the water's surface, resisting external force.

• High specific heat: The significant amount of energy required to change water's temperature is due to hydrogen bonds absorbing kinetic energy when heated and releasing it when cooled. This moderates temperature fluctuations.

• Solvent of life: Its capacity to dissolve many ionic and polar substances is enabled by its polarity and ability to form hydrogen bonds, allowing it to surround and interact with other molecules, which is essential for biological processes.

Water Chemistry and pH (Transcript: Page 23-24)

• Water dissociation (Ionization of Water):

  • Although water is generally stable, a very small fraction of water molecules at any given time can spontaneously dissociate (ionize) into two highly reactive ions: a hydrogen ion ($H^+$) and a hydroxide ion ($OH^-$).

  • This reaction is reversible: $H_2O \rightleftharpoons H^+ + OH^-$

  • More accurately, the $H^+$ ion does not exist independently but immediately associates with another water molecule to form a hydronium ion ($H<em>3O^+$): $2 H</em>2O \rightleftharpoons H_3O^+ + OH^-$

  • Both $H^+$ (or $H_3O^+$) and $OH^-$ ions are highly reactive due to their inherent electrical charges, and their concentrations profoundly affect biological systems.

  • In pure water at 25°C, the concentrations of hydrogen ions and hydroxide ions are equal: $[H^+] = [OH^-] = 10^{-\text{7}} M$ (M = moles per liter).

• Acids and bases:

  • Acids: Substances that increase the relative concentration of hydrogen ions ($H^+$) in a solution. They do this by donating protons (Arrhenius definition) or by releasing $H^+$ ions (e.g., HCl dissociates into $H^+$ and $Cl^-$).

  • Bases: Substances that decrease the relative concentration of hydrogen ions ($H^+$) in a solution. They achieve this by either directly accepting $H^+$ ions (e.g., $NH<em>3$ + $H^+$ $\, \rightarrow \,$ $NH</em>4^+$) or by increasing the concentration of hydroxide ions ($OH^-$), which then combine with $H^+$ to form water (e.g., NaOH dissociates into $Na^+$ and $OH^-$; $OH^- + H^+ \rightarrow H_2O$).

The pH Scale (Transcript: Page 24)

• The pH scale is a logarithmic scale used to quantify the acidity or basicity (alkalinity) of an aqueous solution by measuring the concentration of hydrogen ions ($H^+$).

• Definitions Revisited:

  • Acids: Solutions with a pH value less than 7, meaning they have a higher $H^+$ concentration (and lower $OH^-$ concentration) than pure water.

  • Bases (Alkaline solutions): Solutions with a pH value greater than 7, meaning they have a lower $H^+$ concentration (and higher $OH^-$ concentration) than pure water.

  • A pH of 7 is neutral. This occurs when the hydrogen ion concentration is precisely equal to the hydroxide ion concentration ($[H^+] = [OH^-] = 10^{-\text{7}} M$ at $25°C$).

• Scale characteristics:

  • The pH scale is logarithmic, meaning each whole number change in pH represents a ten-fold change in the concentration of hydrogen ions.

  • For example, a solution with pH 3 is ten times more acidic (has ten times higher $[H^+]$) than a solution with pH 4. Similarly, a solution with pH 9 is ten times more basic (has ten times higher $[OH^-]$ or ten times lower $[H^+]$) than a solution with pH 8.

  • Therefore, there is a ten-fold difference in the $H^+$ and $OH^-$ concentration with each one-unit change in pH value.

• Appendix note: The explicit formula used to calculate pH from the hydrogen ion concentration is:

  • $\mathrm{pH} = -\log_{10}[H^+]$

  • This formula highlights the inverse relationship between $H^+$ concentration and pH value: as $[H^+]$ increases, pH decreases, indicating greater acidity.

Buffers (Transcript: Page 25)

• Buffers are chemical substances or mixtures that play a crucial role in maintaining pH homeostasis by minimizing significant changes in the concentrations of $H^+$ and $OH^-$ when acids or bases are added to a solution.

• A typical buffer system is composed of a weak acid and its corresponding conjugate weak base. This pair can reversibly bind and release hydrogen ions, thereby balancing the pH.

• Carbonic acid/bicarbonate system is a vital example of a biological buffer, particularly important for regulating blood pH in living organisms:

  • The equilibrium reaction is: $\mathrm{H<em>2CO</em>3} \; \rightleftharpoons \; \mathrm{HCO_3^-} + \mathrm{H^+}$

  • Where $\mathrm{H<em>2CO</em>3}$ (carbonic acid) is the weak acid, and $\mathrm{HCO_3^-}$ (bicarbonate ion) is its conjugate weak base.

• Roles of buffer forms in pH regulation:

  • Responding to increased basicity (excess $OH^-$): If the solution becomes too basic (e.g., due to the addition of an alkali, increasing $OH^-$), the $OH^-$ ions will react with $H^+$ ions, removing them from the solution. To counteract this drop in $H^+$ concentration, the buffer's acidic form (carbonic acid, $\mathrm{H<em>2CO</em>3}$) will dissociate further, donating more $H^+$ ions into the solution. This shifts the reaction equilibrium to the right ($\mathrm{H<em>2CO</em>3} \rightarrow \mathrm{HCO_3^-} + \mathrm{H^+}$), thereby replenishing $H^+$ and resisting the increase in pH.

  • Responding to increased acidity (excess $H^+$): If the solution becomes too acidic (e.g., due to the addition of an acid, increasing $H^+$), the buffer's basic form (bicarbonate, $\mathrm{HCO<em>3^-}$) will act as a proton acceptor. It will bind to the excess $H^+$ ions, forming carbonic acid ($\mathrm{HCO</em>3^-} + \mathrm{H^+} \rightarrow \mathrm{H<em>2CO</em>3}$). This shifts the reaction equilibrium to the left, removing $H^+$ ions from the solution and thereby resisting the drop in pH.

• Overall function: Buffers are indispensable for maintaining the narrow and stable pH range required for optimal enzymatic activity and other biochemical processes in biological fluids, preventing drastic and potentially lethal pH changes when metabolic acids or bases are produced or consumed.