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Human Biology Notes: Chapters 1-6

Chapter 1: Human's Closest Relative: Chimpanzee

Genetic Similarities and Differences

  • Genetic Difference: Humans and chimpanzees share a 1.6\% genetic difference.
  • Study of Evolution: We study evolution to help pick animals for research, inferring shared traits and differences.

Shared Adaptations

  • Opposable Thumbs: Thumbs can touch other fingers, a key adaptation for manipulation.
  • Bipedalism: Moving or walking on two rear limbs or legs.

Human Adaptations

  • Endurance: Enhanced capacity for sustained physical activity.
  • Evaporative Cooling: Such as sweating, for temperature regulation.
  • Breathing While Running: Efficient respiration during exertion (unlike panting).
  • Brain Size and Folding: Increased surface area due to folding leads to more neurons and, consequently, greater intelligence.

Dramatic Consequences of DNA Sequence Changes

  • Alterations in DNA sequences can significantly impact an organism's development.
  • Human-Specific DNA Sequences: Two notable sequences in humans that differ from chimpanzees, gorillas, and orangutans (where these traits are present) are:
    1. One that appears to suppress the proliferation of neurons, potentially influencing brain development.
    2. One that directs the formation of penile spines, which are absent in humans but present in other great apes.
  • Gene Expression: The primary difference between humans and chimpanzees lies in gene expression, not just the genes themselves.

Hierarchy of Complexity

  • Organism: The complete living being.
  • Organ Systems: Groups of organs working together (e.g., digestive system).
  • Organs: Structures composed of multiple tissue types (e.g., heart, liver).
  • Tissues: Groups of similar cells and extracellular matrix performing a specific function.
  • Cells: The fundamental unit of life; nothing below cells is considered living.
  • Organelles: Subcellular structures with specific functions.
  • Molecules: Two or more atoms covalently bonded.
  • Atoms: The basic unit of matter.

Biological Variation

  • Anatomical Variation: No two humans are exactly alike in anatomy. This can manifest in:
    • Differences in muscles.
    • Variations in vertebrae.
    • Differences in organs.
    • Left/right reversal of organ organization (e.g., situs inversus).
  • Physiological Variation: Individuals exhibit different physiological measures, such as heart rates, blood pressure, etc.
    • Medical Implications: This variation affects medical practices, such as drug dosages, and influences histological observations.
  • Causes of Variation: A combination of both genetics and environment explains the expected variation among individuals.

Homeostasis

  • Definition: Maintaining stability or balance in a dynamic internal environment.
  • Mechanisms: These activate based on environmental conditions. Examples include:
    • Sweating: When too hot, to cool the body.
    • Shivering: When too cold, to generate heat.
  • Set Point: Internal conditions are maintained within a range of values around a set point (e.g., blood pH).
  • Breakdown of Homeostasis: Conditions like hyperglycemia (high glucose levels) in Type 1 and 2 diabetes can disrupt normal homeostatic mechanisms.
  • Body Heat Regulation Example (Nerve Cells in Brain Monitoring Blood Temp):
    • If too hot:
      1. Vasodilation: Widening of blood vessels to increase heat loss.
      2. Sweating: If still too hot.
    • If too cold:
      1. Vasoconstriction: Shortening of blood vessels to conserve heat.
      2. Shivering: If still too cold.

Chapter 1: Feedback Mechanisms

Negative Feedback

  • Function: Helps keep physiological variables close to their set point, reversing the effects of change.
  • Mechanism:
    • Body senses change.
    • Body reverses the effects of changes via an effector.
    • The effect of the effector shuts down the effector itself.
    • This prevents extremes in the opposite direction.
  • Example: When testosterone levels get too high, negative feedback mechanisms shut down its production.

Feedback Mechanism Loop (e.g., feeling dizzy from standing up too quickly)

  • Receptor: Senses the change (e.g., baroreceptors sensing drop in blood pressure).
  • Integration: Brain processes the information.
  • Effector: Initiates a response (e.g., heart rate increases, blood vessels constrict).
  • Negative Feedback: The response diminishes the original stimulus (e.g., blood pressure returns to normal, shutting down the effector).

Positive Feedback

  • Definition: A mechanism that amplifies change in the same direction.
  • Harmful Potential: Can be harmful as it drives the variable further from the set point.
  • NOT HOMEOSTASIS: It does not maintain stability.
  • Examples:
    • Fevers: Body temperature moves too far from the normal set point (which itself raises with fever).
    • Childbirth: Contractions intensify, pushing the process to completion.
    • Bleeding: The clotting cascade amplifies to quickly stop blood loss.

Chapter 1: Gradients and Flow

Movement Down Gradients

  • Principle: Movement of substances occurs down gradients.
  • Equation for Gradients: (\text{difference} / \text{distance}).
  • Examples:
    • Blood Flow: Flows from high pressure to low pressure through vessels in the body.
    • Chemiosmosis: Differences in temperature, pressure, electric charge, etc., drive movement.
  • ATP Synthase Gradients: Require energy to be put into the system for flow to occur, creating potential energy that can then be harnessed.

Simple Diffusion

  • Principle: Movement from a region of high concentration to a region of low concentration.
  • Characteristics:
    1. Eventually, there is no net flow when equilibrium is reached (no directionality).
    2. Initially, a high concentration of soluble sugar at the top, low at the bottom.
    3. Sugar molecules bounce off water molecules and each other.
    4. When sugar molecules are evenly dispersed, there is no net flow.

Chapter 1: Anatomical Terminology

Anatomical Positions

  • Anterior: Front; ventral.
  • Posterior: Back; dorsal.
  • Superior: Above.
  • Inferior: Below.
  • Medial: Toward the midline (centered).
  • Lateral: Away from the midline.
  • Proximal: Closer to the point of origin or attachment (e.g., Closer to midline for limbs).
  • Distal: Farther from the point of origin or attachment (e.g., Farther from midline for limbs).
  • Superficial: Close to the surface.
  • Deep: Close to bone or internal structures.
  • Convoluted: Twisted or folded.
  • Combined Terms: Terms can be combined, e.g., Posterosuperior refers to the back of a structure/body, around the top.

Planes

  • Function: Allows for imaging the body in a 2D way instead of 3D.
  • Sagittal Plane:
    • Divides the body into right and left halves.
    • A mid-sagittal plane goes right down the midline.
  • Frontal Plane (Coronal Plane):
    • Divides the body into anterior (front) and posterior (back) portions.
  • Transverse Plane:
    • Perpendicular to the long axis of the body.
    • Divides the body into superior (upper) and inferior (lower) portions.

Chapter 2: Chemical Basis of Life

Basic Components and Elements

  • DNA: Contains 5 carbons in its sugar-phosphate backbone.
  • Living Cells: Composed of ions, atoms, and molecules.
  • Steroids: Are a type of lipid; sugar-glucose is a monosaccharide.
  • Carbon: The backbone for all organic molecules due to its ability to form four stable bonds.
  • Bones: Contain mostly calcium phosphate.
  • Chemical Elements:
    • There are 91 naturally-occurring chemical elements.
    • 24 have physiological roles in the human body.
    • 6 elements (\text{O}, \text{C}, \text{H}, \text{N}, \text{Ca}, \text{Ph}) constitute 98.5\% of body weight.
    • The remaining are lesser and trace elements.
  • Top 4 Elements (C, H, N, O):
    • These 4 make up 96\% of an organism's mass.
    • They make up 99\% of the atoms and 90\% of all atoms in your body.
    • Contrast with Earth's Crust: These percentages are radically different from Earth's crust (e.g., hydrogen is less than 5\%).
    • Hydrogen in Human Body: Hydrogen constitutes approximately 63\% of the atoms but only 10\% of the mass in the human body.

Minerals

  • Definition: Inorganic substances, meaning they lack carbon (C) and C-H bonds.
  • Source: Extracted from soil by plants; we obtain minerals by eating plants or animals that have consumed plants.
  • Electrolytes: Mineral salts essential for muscle and nerve function.
    • Examples: Calcium (\text{Ca}^{2+}), Magnesium (\text{Mg}^{2+}), Potassium (\text{K}^{+}), Sodium (\text{Na}^{+}), Chloride (\text{Cl}^{-}).
    • '+' indicates electrons have been given up; '-' indicates an electron has been accepted.

Organic Molecules

  • Definition: Will always have carbon bonded to hydrogen somewhere within the molecule.
  • Example: Carbon dioxide (\text{CO}_2) is not an organic molecule because it lacks hydrogen (C-H bonds).

Ions

  • Definition: An atom or molecule that has gained or lost one or more electrons.
  • Cations: Possess a net positive ('+') charge (lost electrons).
  • Anions: Possess a net negative ('-') charge (gained electrons).

Ionic Bonds

  • Definition: Electrostatic attractions between oppositely charged cations and anions.
    • Example: \text{Na}^{+} + \text{Cl}^{-} = \text{NaCl} (sodium chloride), which is an ionic compound.
    • In \text{NaCl}, sodium (\text{Na}) is the main element, staying in its ionic form.
    • Table salt (\text{NaCl}) is not considered a molecule because it lacks covalent bonds.
    • Formation: Sodium loses an electron (its outer shell) to become \text{Na}^{+}; chlorine gains an electron to fill its outer shell and becomes \text{Cl}^{-}. Chemical stability occurs when the outer shell is full of electrons.
  • Weakness in Water: Ionic bonds are weak in water because water's polarity can form hydration spheres around the ions, causing them to ionize and separate.
    • Salts are strong as solids but separate into weak ions in water.
  • Importance of Ions in Water:
    • Make up electrolytes.
    • Affect chemical activity.
    • Contribute to osmotic effects (e.g., in IVs).
    • Pull water (osmosis) through membranes, causing movement and current of liquids in the body.

Covalent Bonds

  • Definition: Involve the sharing of electrons between atoms.
  • Strength: They are the strongest type of chemical bond.
  • Molecule Formation: Molecules are formed when two or more atoms are joined by covalent bonds.
  • Stability: Atoms form covalent bonds to fill their outer electron shells, achieving greater stability.

Chapter 2: Covalent Bonds & Hydrogen Bonds

Types of Covalent Bonds

  • Single, Double, or Triple Covalent Bonds: Determined by the number of shared electron pairs (one, two, or three, respectively), influencing how many electrons fill the outer shell.

  • Carbon Dioxide (\text{CO}_2): Is not an organic molecule because it lacks hydrogen. It contains two double covalent bonds.

  • Nonpolar Covalent Bond:

    • Atoms have similar electronegativity (affinity for electrons).
    • Electrons are shared equally, resulting in no partial charges.

  • Polar Covalent Bond:

    • Atoms have different electronegativity.
    • Electrons are shared unequally, leading to partial positive and partial negative charges (\delta^{+} and \delta^{-}) at opposite ends of the bond (a dipole).
    • Pattern to Know (Electrons Attract Partial Charges):
      • Bonds between \text{O + H} or \text{O + C} are polar.
      • Bonds between \text{N + C} or \text{N + H} are polar.
      • Bonds between \text{C + H} or \text{C + N} are often considered nonpolar for biological contexts where the difference is smaller.
      • Generally, if there's a large electronegativity difference, it's polar (e.g., high # on top, low # on bottom in a simple diagram).

  • Single Covalent Bond:

    • Allows for rotation around the bond axis.
    • Provides flexibility and allows for changes in molecular shape.
    • Water (\text{H}_2\text{O}) is held together by two single covalent bonds and is a polar molecule.

  • Double Covalent Bond:

    • Shorter and stronger than single bonds.
    • More rigid, less flexible.
    • Does not allow for change of shape (rotation is restricted).

Hydrogen Bonds

  • Definition: An attraction between a slightly positive hydrogen atom (\delta^{+}\text{H}) on one polar molecule and a slightly negative oxygen (\delta^{-}\text{O}) or nitrogen (\delta^{-}\text{N}) atom on another polar molecule.
  • Individual Strength: Individually, hydrogen bonds are weak and can be easily broken (e.g., by kinetic energy from heat).
  • Collective Strength: When many hydrogen bonds are present, they collectively become very strong (e.g., in water, DNA, proteins).
  • Temperature Effects:
    • Hot temperatures (high kinetic energy) can separate hydrogen bonds.
    • However, heat cannot break the stronger covalent bonds within molecules.
    • Cooling water reduces its kinetic energy, allowing hydrogen bonds to become more stable and form ice.

Hydrophobic Substances

  • Characteristics:
    • No partial charges.
    • No hydrogen bonds with water.
    • Will not interact with water.

Summary of Chemical Bonds

  • Covalent bonds form molecules, which can be single or triple, polar or nonpolar; they are the strongest chemical bonds.
  • Bond within water molecules: Polar covalent bonds.
  • Bond between water molecules: Hydrogen bonds.

Chapter 2: Mixtures and Solutions

Mixtures

  • Definition: Components are blended but not chemically combined.
  • Body Fluids: Mostly consist of water. Cells are also mostly water.
  • Water Content: Water makes up about 50\%-75\% of body weight.

Water: The Universal Solvent

  • Characteristics:
    • A universal solvent, excellent due to its polar covalent bonds.
    • Has the ability to support life.
  • Solvent Property: Dissolves hydrophilic substances/molecules.
  • Importance for Metabolism: Crucial for hydrolysis (water-splitting) and dehydration reactions.
  • Adhesion: Can cling to membranes and reduce friction around organs.
  • Sugar: A polar molecule with partial charges, so it dissolves in water.

Types of Mixtures in the Body

  • Solute: Particles mixed within a solvent. Typically translucent. Example: sugar in water.
  • Colloids: Mixtures of protein and water. Larger particle size than solutes, appear cloudy. Example: albumin in blood, breast milk (both high in protein).
  • Suspension: Even larger particles, appear cloudy/opaque. Cannot pass through membranes, and particles separate (settle out). Example: red blood cells (RBCs) in blood.
  • Emulsion: One liquid suspended (not mixed) in another liquid. Occurs because components like carbons and hydrogens (lipids) do not mix with water. Example: fat in breast milk.
  • Blood: Contains a combination of components:
    • Sugars (solution component).
    • Proteins (colloid component).
    • Larger particles (RBCs) (suspension component).

Chapter 2: Acids, Bases, and Buffers

Acids and Bases

  • Acid: A proton (\text{H}^{+}) donor.
    • Example Reaction: \text{HCl} + \text{NaOH} \rightarrow \text{H}_2\text{O} + \text{NaCl}.
  • Base: A proton (\text{H}^{+}) acceptor (can often release \text{OH}^{-}).
    • In the example reaction, \text{HCl} is the acid, \text{NaOH} is the base, \text{H}_2\text{O} is water, and \text{NaCl} is salt.

Buffers

  • Function: Help to maintain a constant pH, tolerating only small changes in pH.
  • Acid-Base Buffer System: Contains a weak acid and a weak base to maintain pH within a narrow range by either adding \text{H}^{+} or releasing it as needed.

Chemical Reactions

  • Hydrolysis: Uses water to break bonds; water (\text{H}_2\text{O}) is added to a molecule, splitting it.
  • Dehydration Synthesis: Removes water to form bonds; water (\text{H}_2\text{O}) is removed from reactants to form a larger molecule.

Organic Compounds and Functional Groups

  • Carbon Bonding: Each carbon atom can form up to 4 covalent bonds because it has 4 electrons in its outer shell.
  • Organic Compounds: Must contain carbon bonded to hydrogen (C-H bonds somewhere within).
    • Methane (\text{CH}_4): The simplest organic compound (\text{H - C - H} on all sides). It cannot form hydrogen bonds.
  • Carbon as the Backbone: Carbon is the backbone of organic molecules due because it forms 4 strong covalent bonds, allowing for many complex shapes and structures.
  • Chemical Groups (Functional Groups):
    • Atoms or clusters of atoms covalently bonded to the carbon backbone.
    • They give organic compounds their different properties.
    • Each type of functional group exhibits the same properties in all molecules in which it occurs, which is what fundamentally changes molecules' roles.

Chapter 2: Chemical Groups and Macromolecules

Examples of Chemical Groups

  • Amino (\text{R - NH}_2): Acts as a weak base, accepting \text{H}^{+}.
  • Carboxyl (\text{R - COOH}): Acts as a weak acid, giving away \text{H}^{+}.
  • Aldehyde (\text{R - CHO}): Often polar.
  • Hydroxyl (\text{R - OH}): Polar.
  • Methyl (\text{R - CH}_3): Non-polar, hydrophobic. Can silence genes when associated with DNA.
  • Phosphate (\text{R - PO}_4^{2-}): Polar and often involved in energy transfer or structural roles.

Small Organic Molecules (Monomers)

  • These are the building blocks or subunits for larger organic molecules (macromolecules).
  • Some, like sugars and fatty acids, also serve as energy sources.

Macromolecules

  • Definition: Large organic molecules formed from small organic molecules (monomers) covalently linked together.
  • Types: Polysaccharides, proteins, and nucleic acids (DNA, RNA).
    • Sugar monomers form Polysaccharides.
    • Amino acid monomers form Proteins.
    • Nucleotide monomers form Nucleic Acids.

Sugars (Carbohydrates)

  • Composition: Made of Carbon (\text{C}), Hydrogen (\text{H}), and Oxygen (\text{O}).

  • Polarity: Polar, so they dissolve well in water.

  • Functions:

    • Primary energy sources (e.g., consuming starches breaks them down into glucose for energy).
    • Cells are sugar-coated on the outside (glycocalyx).

  • Monosaccharides:

    • The simplest carbohydrates (single sugar molecule).
    • Subunits of carbohydrate chains.
    • Isomers: Glucose, galactose, and mannose are isomers, meaning they have the same chemical formula (\text{C}6\text{H}{12}\text{O}_6) but differ in the arrangement of groups around one or two carbon atoms.

  • Disaccharides:

    • Short chain of two monosaccharides.
    • Formed by condensation (dehydration) reactions, creating glycosidic bonds (covalent bonds).
    • Examples: Sucrose, lactose.

  • **Polysaccharides (