Introduction to Biology and Chemistry

Page 3: Learning Goals for Biology Course

Goals:

• Describe atoms, their structure, and how they bond.

• Understand water's features that support life.

• Explain the structure and function of key biomolecules:

  • Carbohydrates

  • Lipids

  • Proteins

  • Nucleic acids

Page 4: Connection Between Chemistry and Biology

• Chemistry is linked to the evolution of life.

• Exploring fundamental concepts: - How are atoms bonded together in molecules?

  • Unique properties of water.

  • Role of carbon in life's evolution.

Page 5: Environmental Chemistry Themes

• Topics of concern: - Climate change

  • Acid rain

  • Health effects of pesticides and GMOs

  • Ozone layer, cancers, and UV light

Page 6: Cuyahoga River - Environmental Case Study

• Reference to the Cuyahoga River in Cleveland.

Page 7: Overview of Macromolecules

• Types of macromolecules: - Carbohydrates

  • Lipids

  • Proteins

  • Nucleic acids

Page 8: Structure of Atoms

• Atoms consist of: - Proton (positive charge)

  • Neutron (neutral charge)

  • Electron (negative charge)

• Example: Hydrogen Atom - 1 Proton, 1 Electron, 0 Neutrons

• Example: Carbon Atom - 6 Protons, 6 Neutrons, 6 Electrons

• Forces of attraction between charged particles maintain electron proximity to nucleus.

Page 9: Classical Elements in Different Cultures

• Classical elements included: - Western: Air, Water, Fire, Earth (Greek origin by Empedocles)

  • Chinese: Water, Metal, Earth, Wood, Fire

  • Japanese: Earth, Water, Fire, Air/Wind, Void/Sky/Heaven

Page 10: The Big Four Elements

• Significant elements for biological processes: - Hydrogen (H) - smallest and most abundant element.

  • Carbon (C) - forms various allotropes including charcoal, graphite, diamond.

Page 11: Continuation of Big Four Elements

• Nitrogen (N) - has both toxic forms and those that cause light-headedness.

• Oxygen (O) - primarily exists as O2 in nature.

Page 12: Atoms and Elements Defined

• Structure of elements:

• Carbon (C) - Atomic number = 6; Protons = 6; Neutrons = 6; Electrons = 6

• Nitrogen (N) - Atomic number = 7; Protons = 7; Neutrons = 7; Electrons = 7

• Phosphorus (P) - Atomic number = 15; Protons = 15; Neutrons = 16; Electrons = 15

Page 13: Trace Elements with Health Implications

• Health impacts of lack of essential trace elements: - Iodine deficiency causes goiter

  • Potassium deficiency

Page 14: Top Elements Found in Human Body

• Composition of human body by percentages: - Oxygen (65%)

  • Carbon (18.5%)

  • Hydrogen (9.5%)

  • Nitrogen (3%)

• Remaining 4% consists of: - Calcium

  • Phosphorus

  • Potassium

  • Sulfur

  • Sodium

  • Chlorine

• Trace elements found in <0.1% of body composition.

Page 15: Mnemonic for Remembering the Big 4 Elements

• Mnemonic: "S.P.O.N.C.H"

• Each letter represents a significant element: - Sulfur

  • Phosphorus

  • Oxygen

  • Nitrogen

  • Carbon

  • Hydrogen

  • Potassium

Page 16: Radioactive Atoms

• Some atomic nuclei are unstable and decay spontaneously.

• Characteristics of radioactive atoms include measurable energy release at a constant rate.

Page 17: Definition of Isotopes

• Isotopes are alternative versions of elements that differ in mass due to varying neutron numbers.

Page 18: Isotopes and Their Properties

• Isotopes with the same atomic number but different atomic masses: - Carbon-12 (C-12): 6 Protons, 6 Neutrons

  • Carbon-13 (C-13): 6 Protons, 7 Neutrons

  • Carbon-14 (C-14): 6 Protons, 8 Neutrons

Page 19: Stability of Isotopes

• Most isotopes are stable and do not lose particles; examples include: - C-12 and C-13 (stable)

  • C-14 (unstable)

Page 20: Carbon-14 Stability

• Carbon-14 (C-14) undergoes radioactive beta decay: - Half-life of C-14 is 5730 years

  • Neutrons can hit nitrogen atoms, converting them to carbon-14

Page 21: Carbon-14 Usage

• Carbon-14 is absorbed by living organisms with no health issues; ratio to C-12 is used for dating.

Page 22: Measuring Radiation Levels

• Radiation levels measured in: - Becquerel (BQ): SI unit

  • Gray (Gy): SI unit for absorbed radiation

  • Rad: conventional unit

  • Curie (Ci): conventional unit for radioactivity measurement

Page 23: Half-Life Examples

• Examples of half-lives for various isotopes: - Carbon-14 (C-14): 5730 $\pm$ 40 years

  • Potassium-40 (K-40): $\pm$ 1.3 billion years

  • Uranium-235 (U-235): $\pm$ 700 million years

  • Uranium-238 (U-238): $\pm$ 4.5 billion years

• In general, shorter half-lives indicate less accuracy over long periods.

Page 24: Chernobyl Disaster

• Date of incident: 26 April 1986.

Page 25: Radiation Reading Data

• Radiation level recorded at 461.5 mSv/m/hr $\sim$ 40 rads/hr

• Fatal dose approximated at $\sim$ 500 rads in 24 hours.

Page 26: Mururoa Atoll, Tahiti

• Geographic coordinates: - $21^{\text{o}}33'$

  • $21^{\text{o}}39'$

  • $21^{\text{o}}45'$

  • $W139^{\text{o}}27'$

  • $W139^{\text{o}}21'$

  • $W139^{\text{o}}09'$

  • $W138^{\text{o}}57'$

  • $W138^{\text{o}}45'$

  • $W138$

Page 27: Radioactive Decay and Half-Lives

• Explanation of radioactive decay and half-lives of elements.

Page 28: Half-Life Calculations

• Example context for half-life calculation with initial values: - Initial: 40 Kg

  • Time progression noted: 20 Kg, then 10 Kg until decay reaches half.

Page 29: U-238 Decay Series

• U-238 decay and its half-life down to lead-206 (Pb-206) spans 4.5 billion years.

Page 30: Take-Home Message 2.1

• Everything, living or non-living, is made from atoms, which are the smallest unit of matter.

• Every atom has a similar structure, comprising protons and neutrons in the nucleus and electrons orbiting.

Page 31: Periodic Table Explained

• Structure of the periodic table: - Vertical columns indicate electrons in outermost shell.

  • Horizontal rows indicate total number of electron shells.

Page 32: Dalton's Atomic Theory

• Key principles of Dalton's Theory: - Elements consist of tiny, identical particles called atoms.

  • Atoms of an element differ from those of others.

  • Atoms can combine to form compounds.

  • Atoms cannot be created or destroyed; they are rearranged in reactions.

Page 33: Plum Pudding Model

• Proposed by J.J. Thomson, suggesting electrons are embedded within a 'pudding' of positive charge.

Page 34: Thomson Model of Electrons

• Conducted experiments involving electric current through gases to identify negative particles.

Page 35: Rutherford's Correction to Thomson's Model

• Ernest Rutherford determined electron arrangement led to further understanding of atomic structure.

Page 36: Bohr Model

• Niels Bohr introduced that negatively charged electrons occupy defined orbits around a positively charged nucleus.

Page 37: Heisenberg Uncertainty Principle

• Werner Heisenberg explained limitations in predicting electron locations, introducing the concept of atomic orbitals where electrons are likely to be found.

Page 38: Heisenberg's Prediction of Electron Location

• Uncertainty revealed in predicting exact electron position results in defining probability fields.

Page 39: Practical Implications of Electron Uncertainty

• Challenges in point-to-point electron transport due to unpredictable positions.

Page 40: Electron Shells and Atomic Stability

• Electrons occupy predefined levels, which influence atomic stability. - First shell can hold up to 2 electrons

  • Second shell can hold up to 8 electrons

• Trends in electron filling indicate stability.

Page 41: Stability in Atoms

• Atoms are stable when outermost shells are filled.

• Unstable atoms tend to react and bond with others.

Page 42: Properties of Metals

• Characteristics: - Good conductors of electricity

  • Ductile and malleable

  • Shiny appearance

  • Tendency to corrode in water

Page 43: Properties of Non-Metals

• Characteristics: - Poor conductors of electricity

  • Generally dull appearance

  • Not ductile or malleable

  • Often brittle when solid

Page 44: Properties of Metalloids

• Conductivity falls between that of metals and non-metals; can exhibit mixed traits consistent with both categories.

Page 45: Mendeleev's Contributions to Periodic Table

• Dmitri Mendeleev created the periodic table in 1869, organizing elements based on mass and properties while predicting undiscovered elements.

Page 46: Mendeleev's Accurate Predictions

• Mendeleev's predictions were upheld by discoveries that followed his initial proposals.

Page 47: Additional Chemical Information

• Various peculiarities and historical references related to elements can be noted in this unstructured data.

Page 48: Take-Home Message 2.2

• Chemical characteristics based on outer electron shell makeup greatly affect atom behavior.

• Atoms display stability and less reactivity when shells are filled.

Page 49: Bond Types in Biological Macromolecules

• Importance of understanding bonding properties in biological systems.

Page 50: Bonding Properties Overview

• Atom behavior is significantly influenced by electron configurations, particularly in the outermost shell (valence).

Page 51: Understanding Atom Behavior

• The information discussed earlier signifies the importance of electron distribution.

Page 52: Definition of Ions

• Ions are defined as electrically charged atoms resulting from gain or loss of electrons.

• Example: - Calcium atom: 20 Protons, 20 Neutrons, 20 Electrons

  • Calcium ion (Ca2+): 20 Protons, 20 Neutrons, 18 Electrons (loss of 2 electrons)

Page 53: Electron Shell Distribution

• Electron shell capacity: - First shell: 2 electrons (1 orbital)

  • Second shell: 8 electrons (4 orbitals)

Page 54: Electron Distribution and Shape

• Illustrates electronic orbital distribution among electron shells, highlighting shape and behavior significance in chemical interactions.

Page 55: Chemical Reactions and Shell Filling

• Chemical reactions are often driven by atoms' tendency to complete partially filled valence shells, aiming for stability.

Page 56: Electron Behavior - Noble Gases

• Atoms with full valence shells (noble gases) demonstrate low reactivity.

Page 57: Summary of Atom Interactions

• Illustrations relate to real-world applications and interactions in biological systems.

Page 58: Molecules and Compounds

• Definitions: - Molecules: Combinations of two or more atoms.

  • Covalent Bonds: Interaction where uncharged atoms share electrons (e.g., CO2).

  • Ionic Bonds: Charged atoms attracting each other due to electrical charges (e.g., NaCl).

Page 59: Sodium Element Characteristics

• Sodium is typically found in the ocean but is highly reactive with water in its pure form.

Page 60: Chlorine Element Characteristics

• Inhalation of chlorine gas can lead to irritation and burning of eyes/sinuses.

Page 61: Example of Ionic Bond Formation

• Reaction: - 2 Na(s) + Cl2(g) $\rightarrow$ 2 NaCl(s)

• The properties of NaCl differ from its constituent elements.

Page 62: Energy and Bonds

• Chemical bonding involves energy changes at different shell levels, absorbing and releasing energy through interactions.

Page 63: The Importance of Chemical Bonds

• Examining the significance of bonds in biology - considerations of strength, permanence, and reusability.

Page 64: Covalent Bonds Explained

• Structures showing covalent bonds (e.g., H2 or dihydrogen) indicate shared electron interactions leading to molecular stability.

Page 65: Multiple Covalent Bonds

• Concept that pairs of atoms can share multiple pairs of electrons: - Double bonds: share 2 pairs

  • Triple bonds: share 3 pairs

Page 66: Nonpolar Covalent Bonds

• Example of a nonpolar covalent bond with equal sharing of electrons in hydrocarbons (e.g., methane CH4).

Page 67: Polar Bonding Dynamics

• Attraction dynamics that occur due to differences in electronegativity between atoms.

Page 68: Polar Covalent Bonds Detailed

• Example involving water (H2O) is a polar molecule due to unequal sharing of electrons, influenced by oxygen's greater electronegativity.

Page 69: Structure and Properties of Water

• Unique properties arise from water's structure, such as high boiling point, high melting point, and capability for hydrogen bonding.

Page 70: Hydrogen Bonding Mechanism

• Polar water molecules produce weak bonds through hydrogen attraction between molecules.

Page 71: Stability of Hydrogen Bonds

• Maximum stability occurs when donor, hydrogen, and acceptor are aligned in a straight line for effective bond formation.

Page 72: Liquid State of Water

• The fluid behavior of water results from transient clusters known as “flickering clusters.”

Page 73: Biological Importance of Hydrogen Bonds

• Examples include interactions between water and hydroxyl groups, carbonyl groups, polypeptide bonds, and complementary DNA bases.

Page 74: Ionic Bond Definition

• Ionic bonds occur through the full transfer of electrons, leading to the formation of charged ions (e.g., NaCl).

Page 75: Effects of Adding Water to Ionic Compounds

• Adding water disrupts the crystal lattice of ionic compounds as water surrounds sodium and chloride ions.

Page 76: Formation of Hydration Shells

• Hydration shells form when water molecules cluster around ions, preventing the reformation of crystal lattice structures.

Page 77: Geckos and Climbing Mechanisms

• Mechanisms used by geckos (and spiders) to climb walls, significantly tying biological design to molecular interactions.

Page 78: Unstable Bonds Explanation

• Unstable bonds arise through molecular adhesion formed via asymmetric electron distributions leading to van der Waals interactions, demonstrated through the structure of gecko toes.

Page 79: Visual Representation of Gecko Climbing

• Image demarcating the structural adaptations in geckos enabling them to climb surfaces through microscopic mechanisms.

Page 80: Signal Transmission at Cellular Level

• Overview of signal transmission molecules relevant in cellular communication and networks like synapses.

Page 81: Research on Parkinson's Disease

• Update or note on Parkinson's disease treatments and the ongoing research into its causes and management.

Page 82: Practical Details (Post-It Notes)

• Specifications on post-it notes dimensions highlighted: - 3 in x 3 in / 76.2 mm x 76.2 mm

Page 83: Summary of Three Types of Bonds

• Types and characteristics of bonding in molecules: 1. Covalent Bonds: Atoms share electrons; strongest bond type.

  2. Ionic Bonds: Attraction between oppositely charged ions.

  3. Hydrogen Bonds: Weak attraction involving hydrogen atoms.

Page 84: Summary Take-Home Message 2.3

• Summary of bonding types reflecting the strength and mechanisms exemplified in different biological and chemical contexts.

Page 85: Shape's Relationship to Function

• The notion that molecular shape is crucial to function, with structural integrity being vital in biological organisms.

Page 86: Introduction to Functional Groups

• Mention of key functional groups in biochemistry: alcohols, acetic acid, fatty acids, sugars, and amino acids that form complex biomolecules.

Page 87: Functional Group Anatomy

• Example: Configuration of an endorphin shows interplay of carbon, nitrogen, oxygen, hydrogen, and sulfur atoms in functional groups.

Page 88: Form Impacts Function - Thalidomide Example

• Discusses thalidomide as a sedative and teratogen, reflecting on how specific molecular structures yield different biological effects