Chapter_2_bio

Page 1: Introduction to Matter and Elements

• Matter: Any substance that occupies space and has mass.

• Elements: Unique forms of matter with specific chemical and physical properties that cannot break down further.

  • 118 total elements known.

  • 98 elements occur naturally; others are unstable and synthesized in laboratories.

  • Each element has a chemical symbol:

    • Examples: C (carbon), O (oxygen), Ca (calcium).

    • Some symbols from Latin names: Na (sodium from natrium).

• Common Elements in Living Organisms:

  • Four elements comprise 96% of living matter:

    • Oxygen (O): 65%

    • Carbon (C): 18%

    • Hydrogen (H): 10%

    • Nitrogen (N): 3%

• Abundance in Various Environments:

  • Atmosphere: Rich in nitrogen and oxygen; low in carbon and hydrogen.

  • Earth’s Crust: Contains oxygen, less hydrogen; very little nitrogen and carbon.

Page 2: Atomic Structure

• Atom: Smallest unit of matter that retains an element's chemical properties.

  • Composed of:

    • Nucleus: Center containing protons and neutrons.

    • Surrounding Electrons: Orbit around the nucleus.

• Protons and Neutrons: Approximately equal in mass (1.67 × 10^-24 grams).

  • Protons: Positively charged.

  • Neutrons: No charge.

  • Electrons: Negatively charged, significantly lighter (1/1800 mass of protons).

• Isotopes: Variants of elements with the same protons but different neutrons.

Page 3: Atomic Structure Continued

• Most of an atom's volume is empty space.

  • Electrons repel each other, preventing solid objects from passing through one another.

• Atomic Number: Number of protons determines the element; affects mass number (protons + neutrons).

• Atomic Mass: Mean of the mass number for isotopes, expressed as a fraction (e.g. Chlorine: 35.45).

Page 4: Isotopes and Carbon Dating

• Example of Isotopes:

  • Carbon-12: 6 protons, 6 neutrons.

  • Carbon-14: 6 protons, 8 neutrons; used in carbon dating.

• Carbon-14 Formation: From nitrogen in the atmosphere by cosmic rays, allowing plants and animals to absorb it.

• Upon death, organisms cease to absorb carbon-14, and it decays into nitrogen-14, allowing age estimation based on ratios.

Page 5: Carbon Dating Process

• Half-life of Carbon-14: 5,730 years; useful for dating material up to 50,000 years old.

• Comparison with Other Isotopes:

  • Potassium-40 (40K): 1.25 billion years

  • Uranium-235 (235U): 700 million years

Page 6: The Periodic Table

• Organizes elements by atomic number and groups elements with similar properties.

  • Developed by Dmitri Mendeleev in 1869.

• Each element listed includes:

  • Symbol: For example, C for carbon.

  • Atomic Number: Identifies the number of protons.

  • Atomic Mass: Listed below the symbol.

Page 7: Electron Shells and the Bohr Model

• Electron Configuration:

  • Bohr Model: Visualizes electrons in shells around the nucleus.

  • Electrons fill lowest energy shells first (1n closest to nucleus).

• Electrons transition between shells upon gaining energy.

Page 8: Electron Filling Order

• Electrons fill orbitals first in the lowest energy levels:

  • Octet Rule: Stability achieved with 8 electrons in outer shell.

• Energy Levels:

  • 1s: Holds 2 electrons

  • 2s: Holds 2

  • 2p: Holds 6

Page 9: Achieving Stable Configurations

• Atoms interact to achieve full valence shell through losing, gaining, or sharing electrons.

  • Reactivity influenced by outer electron arrangement.

  • Group 1 elements typically lose one electron; Group 17 elements gain one.

Page 10: Electron Orbitals and Shell Structures

• Subshells: Differ in shape and spatial distributions (s, p, d, f).

  • 1s holds 2 electrons; 2s holds 2; 2p has 3 orbitals that can hold 6; and so on.

Page 11: Electron Configuration Examples

• Elements fill orbitals based on their atomic structure.

  • Hydrogen: 1s1 (1 electron)

  • Helium: 1s2 (2 electrons)

  • Neon: 1s22s22p6 (10 electrons).

Page 12: Chemical Reactions and Molecules

• Atoms form molecules through bonding—either covalent or ionic.

  • Molecules result from shared or transferred electrons.

  • Chemical Equation: Represents reactants (starting materials) → products (resulting substance from a reaction).

Page 13: Chemical Reactions Explained

• Conservation of Matter: Total number of atoms remains constant in a chemical reaction.

• Examples of Reactions:

  • Decomposition of hydrogen peroxide: 2H2O2 → 2H2O + O2

Page 14: Reversible and Irreversible Reactions

• Reversible Reactions: Can convert back to reactants under certain conditions.

  • Example of equilibrium in blood chemistry with carbonates.

Page 15: Ionic Bonds and Ions Overview

• Ions: Charged atoms from losing or gaining electrons.

  • Cations: Positively charged atoms (lose electrons).

  • Anions: Negatively charged atoms (gain electrons).

Page 16: Sodium and Chlorine Bonding

• Sodium (Na): Loses one electron → forms Na+ cation.

• Chlorine (Cl): Gains one electron → forms Cl- anion.

• These ions bond to form sodium chloride (NaCl).

Page 17: Covalent Bonds and Their Importance

• Covalent Bonds: Atoms share electrons to achieve stable configurations.

  • Can form single, double, or triple bonds, affecting strength and properties of molecules.

Page 18: Polar and Nonpolar Covalent Bonds

• Polar Covalent Bonds: Unequal sharing of electrons leads to partial charges.

  • Example: Water molecules are polar (H₂O).

• Nonpolar Covalent Bonds: Equal sharing, as seen in molecular oxygen (O₂).

Page 19: Career Connection: Pharmaceutical Chemist

• Role in drug discovery, testing, and approval process.

• Responsibilities include developing safer versions of drugs and testing for effectiveness.

Page 20: Life's Foundation: Water

• Water: Composes 60-70% of the human body; essential for life.

• Polarity enables unique properties and hydrogen bonding vital for life.

Page 21: Water's Unique Properties

• Polarity of Water: Causes attraction between molecules, making it an excellent solvent.

• High specific heat: Absorbs and releases heat with minimal temperature change.

Page 22: Temperature Effects on Water

• Water can exist in three states (solid, liquid, gas).

  • Unique behavior due to hydrogen bonding.

Page 23: Ice Formation and its Impact

• Ice's lower density allows it to float, insulating aquatic life beneath.

• Prevents freezing and damage by creating an insulating layer above water.

Page 24: Heat of Vaporization of Water

• High heat of vaporization serves as cooling mechanism (e.g., evaporation from skin).

Page 25: Water as a Solvent

• Water's Solvent Properties: Dissolves ionic and polar molecules, forming hydration shells.

Page 26: Cohesion and Surface Tension

• Water’s cohesive nature leads to surface tension, allowing small objects to float.

Page 27: Importance of Cohesion and Adhesion

• Vital for transporting water in plants via capillary action.

Page 28: Understanding pH: Acids and Bases

• pH Scale: Indicates acidity or basicity of a solution based on hydrogen ion concentration.

Page 29: Acids, Bases, and Buffer Systems

• Buffers: Help maintain stable pH levels in biological systems (e.g., blood).

Page 30: Macromolecules: Building Blocks of Life

• Macromolecules: Proteins, nucleic acids, carbohydrates, and lipids, with carbon as the backbone.

Page 31: Carbon’s Role in Macromolecules

• Hydrocarbons: Organic compounds of carbon and hydrogen.

• Essential in various life processes.

Page 32: Hydrocarbon Structures and Geometry

• Hydrocarbon Geometry: Affects molecular function and properties (e.g., chain vs. ring forms).

Page 33: Types of Hydrocarbons

• Aliphatic (linear) and aromatic (ringed) hydrocarbons have different chemical properties.

Page 34: Isomers in Organic Chemistry

• Isomers: Molecules with identical formulas but differing structures; affect chemical behavior.

Page 35: Functional Groups in Biological Molecules

• Functional Groups: Small groups of atoms that affect the properties and reactions of organic molecules.

Page 36: Fatty Acid Configurations: Saturated vs. Unsaturated

• Saturated Fats: No double bonds between carbon atoms.

• Unsaturated Fats: Contain double bonds, affecting fluidity and packing at room temperature.

Page 37: Enantiomers: A Special Case of Isomers

• Enantiomers: Molecules that are mirror images; essential in biochemical interactions.

Page 38: Functional Group Importance in Macromolecules

• Functional groups dictate the chemical nature and biological activity of macromolecules.

Page 39: Hydrogen Bonds in DNA Structure

• Hydrogen bonds stabilize the double helix structure of DNA, vital for genetic code preservation.

Page 40: Key Terms (Definitions)

• Acid: Donates hydrogen ions, increasing hydrogen ion concentration in a solution.

• Cohesion: Attraction between water molecules due to hydrogen bonding.

• Ionic Bond: Electrostatic attraction between charged ions.

Page 41: Chemical Reaction Glossary

• Chemical Bond: Interaction between two atoms forming a molecule.

• Electron Configuration: Arrangement of electrons in atomic orbitals.

Page 42: Summary of Key Terms

• Definitions and explanations of terms critical for understanding molecular biology and chemistry.

Page 43: Multiple Choice and Review Questions

• Self-assessment questions based on the material presented.

Page 44: Further Review Questions

• Additional questions for reinforcing understanding of key concepts in chemistry and biology.

Page 45: Continuing Questions for Understanding

• More questions to review factual accuracy and comprehension.

Page 46: Questions on Electron Configurations and Bonds

• Exploring elemental properties and bonding behaviors of selected atoms.

Page 47: Understanding Acids and Bases

• True/False questions evaluating understanding of acid-base chemistry as it pertains to living systems.

Introduction to Matter and Elements

• Matter: Any substance that occupies space and has mass.

• Elements: Unique forms of matter with specific chemical and physical properties that cannot break down further. There are 118 total elements known, with 98 elements occurring naturally, while others are unstable and synthesized in laboratories.

• Chemical Symbols: Each element has a unique chemical symbol (e.g., C for carbon, O for oxygen, Ca for calcium). Some symbols derive from Latin names, such as Na (sodium from natrium).

Common Elements in Living Organisms

• Four elements comprise 96% of living matter:

  • Oxygen (O): 65%

  • Carbon (C): 18%

  • Hydrogen (H): 10%

  • Nitrogen (N): 3%

Abundance in Various Environments

• Atmosphere: Rich in nitrogen (78%) and oxygen (21%); low in carbon dioxide and hydrogen.

• Earth’s Crust: Contains oxygen as the most abundant element, less hydrogen, and very little nitrogen and carbon.

Atomic Structure

• Atom: The smallest unit of matter that retains an element's chemical properties.

• Composed of:

  • Nucleus: Contains protons (positively charged) and neutrons (no charge).

  • Electrons: Negatively charged, orbit around the nucleus.

• Isotopes: Variants of elements with the same number of protons but different neutrons.

Key Concepts to Remember

1. Atomic Number: Defines an element and is equivalent to the number of protons.

2. Mass Number: Total number of protons and neutrons in the nucleus.

3. Electron Configuration: Determines the unique behavior of each element based on how electrons are arranged.

4. Chemical Reactions: involve the formation and breaking of bonds between atoms to create new substances, reflecting the conservation of matter.