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Boundless Biology: 2.1: Atoms, Isotopes, Ions, and Molecules

Boundless Biology
2.1: Atoms, Isotopes, Ions, and Molecules
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table of contents
  1. 1: The Study of Life
    1. 1.1: The Science of Biology
      1. 1.1.0: Introduction to the Study of Biology
      2. 1.1.1: Scientific Reasoning
      3. 1.1.2: The Scientific Method
      4. 1.1.3: Basic and Applied Science
      5. 1.1.4: Publishing Scientific Work
      6. 1.1.5: Branches and Subdisciplines of Biology
    2. 1.2: Themes and Concepts of Biology
      1. 1.2.0: Properties of Life
      2. 1.2.1: Levels of Organization of Living Things
      3. 1.2.2: The Diversity of Life
  2. 2: The Chemical Foundation of Life
    1. 2.1: Atoms, Isotopes, Ions, and Molecules
      1. 2.1.0: Overview of Atomic Structure
      2. 2.1.1: Atomic Number and Mass Number
      3. 2.1.2: Isotopes
      4. 2.1.3: The Periodic Table
      5. 2.1.4: Electron Shells and the Bohr Model
      6. 2.1.5: Electron Orbitals
      7. 2.1.6: Chemical Reactions and Molecules
      8. 2.1.7: Ions and Ionic Bonds
      9. 2.1.8: Covalent Bonds and Other Bonds and Interactions
      10. 2.1.9: Hydrogen Bonding and Van der Waals Forces
    2. 2.2: Water
      1. 2.2.0: Water’s Polarity
      2. 2.2.1: Water’s States: Gas, Liquid, and Solid
      3. 2.2.2: Water’s High Heat Capacity
      4. 2.2.3: Water’s Heat of Vaporization
      5. 2.2.4: Water’s Solvent Properties
      6. 2.2.5: Water’s Cohesive and Adhesive Properties
      7. 2.2.6: pH, Buffers, Acids, and Bases
    3. 2.3: Carbon
      1. 2.3.0: The Chemical Basis for Life
      2. 2.3.1: Hydrocarbons
      3. 2.3.2: Organic Isomers
      4. 2.3.3: Organic Enantiomers
      5. 2.3.4: Organic Molecules and Functional Groups
  3. 3: Biological Macromolecules
    1. 3.1: Synthesis of Biological Macromolecules
      1. 3.1.0: Types of Biological Macromolecules
      2. 3.1.1: Dehydration Synthesis
      3. 3.1.2: Hydrolysis
    2. 3.2: Carbohydrates
      1. 3.2.0: Carbohydrate Molecules
      2. 3.2.1: Importance of Carbohydrates
    3. 3.3: Lipids
      1. 3.3.0: Lipid Molecules
      2. 3.3.1: Waxes
      3. 3.3.2: Phospholipids
      4. 3.3.3: Steroids
    4. 3.4: Proteins
      1. 3.4.0: Types and Functions of Proteins
      2. 3.4.1: Amino Acids
      3. 3.4.2: Protein Structure
      4. 3.4.3: Denaturation and Protein Folding
    5. 3.5: Nucleic Acids
      1. 3.5.0: DNA and RNA
      2. 3.5.1: The DNA Double Helix
      3. 3.5.2: DNA Packaging
      4. 3.5.3: Types of RNA
  4. 4: Cell Structure
    1. 4.1: Studying Cells
      1. 4.1.0: Cells as the Basic Unit of Life
      2. 4.1.1: Microscopy
      3. 4.1.2: Cell Theory
      4. 4.1.3: Cell Size
    2. 4.2: Prokaryotic Cells
      1. 4.2.0: Characteristics of Prokaryotic Cells
    3. 4.3: Eukaryotic Cells
      1. 4.3.0: Characteristics of Eukaryotic Cells
      2. 4.3.1: The Plasma Membrane and the Cytoplasm
      3. 4.3.2: The Nucleus and Ribosomes
      4. 4.3.3: Mitochondria
      5. 4.3.4: Comparing Plant and Animal Cells
    4. 4.4: The Endomembrane System and Proteins
      1. 4.4.0: Vesicles and Vacuoles
      2. 4.4.1: The Endoplasmic Reticulum
      3. 4.4.2: The Golgi Apparatus
      4. 4.4.3: Lysosomes
      5. 4.4.4: Peroxisomes
    5. 4.5: The Cytoskeleton
      1. 4.5.0: Microfilaments
      2. 4.5.1: Intermediate Filaments and Microtubules
    6. 4.6: Connections between Cells and Cellular Activities
      1. 4.6.0: Extracellular Matrix of Animal Cells
      2. 4.6.1: Intercellular Junctions
  5. 5: Structure and Function of Plasma Membranes
    1. 5.1: Components and Structure
      1. 5.1.0: Components of Plasma Membranes
      2. 5.1.1: Fluid Mosaic Model
      3. 5.1.2: Membrane Fluidity
    2. 5.2: Passive Transport
      1. 5.2.0: The Role of Passive Transport
      2. 5.2.1: Selective Permeability
      3. 5.2.2: Diffusion
      4. 5.2.3: Facilitated transport
      5. 5.2.4: Osmosis
      6. 5.2.5: Tonicity
      7. 5.2.6: Osmoregulation
    3. 5.3: Active Transport
      1. 5.3.0: Electrochemical Gradient
      2. 5.3.1: Primary Active Transport
      3. 5.3.2: Secondary Active Transport
    4. 5.4: Bulk Transport
      1. 5.4.0: Endocytosis
      2. 5.4.1: Exocytosis
  6. 6: Metabolism
    1. 6.1: Energy and Metabolism
      1. 6.1.0: The Role of Energy and Metabolism
      2. 6.1.1: Types of Energy
      3. 6.1.2: Metabolic Pathways
      4. 6.1.3: Metabolism of Carbohydrates
    2. 6.2: Potential, Kinetic, Free, and Activation Energy
      1. 6.2.0: Free Energy
      2. 6.2.1: The First Law of Thermodynamics
      3. 6.2.2: The Second Law of Thermodynamics
      4. 6.2.3: Activation Energy
    3. 6.3: ATP: Adenosine Triphosphate
      1. 6.3.0: ATP: Adenosine Triphosphate
    4. 6.4: Enzymes
      1. 6.4.0: Enzyme Active Site and Substrate Specificity
      2. 6.4.1: Control of Metabolism Through Enzyme Regulation
  7. 7: Cellular Respiration
    1. 7.1: Energy in Living Systems
      1. 7.1.0: Transforming Chemical Energy
      2. 7.1.1: Electrons and Energy
      3. 7.1.2: ATP in Metabolism
    2. 7.2: Glycolysis
      1. 7.2.0: Importance of Glycolysis
      2. 7.2.1: The Energy-Requiring Steps of Glycolysis
      3. 7.2.2: The Energy-Releasing Steps of Glycolysis
      4. 7.2.3: Outcomes of Glycolysis
    3. 7.3: Oxidation of Pyruvate and the Citric Acid Cycle
      1. 7.3.0: Breakdown of Pyruvate
      2. 7.3.1: Acetyl CoA to CO2
      3. 7.3.2: Citric Acid Cycle
    4. 7.4: Oxidative Phosphorylation
      1. 7.4.0: Electron Transport Chain
      2. 7.4.1: Chemiosmosis and Oxidative Phosphorylation
      3. 7.4.2: ATP Yield
    5. 7.5: Metabolism without Oxygen
      1. 7.5.0: Anaerobic Cellular Respiration
    6. 7.6: Connections of Carbohydrate, Protein, and Lipid Metabolic Pathways
      1. 7.6.0: Connecting Other Sugars to Glucose Metabolism
      2. 7.6.1: Connecting Proteins to Glucose Metabolism
      3. 7.6.2: Connecting Lipids to Glucose Metabolism
    7. 7.7: Regulation of Cellular Respiration
      1. 7.7.0: Regulatory Mechanisms for Cellular Respiration
      2. 7.7.1: Control of Catabolic Pathways
  8. 8: Photosynthesis
    1. 8.1: Overview of Photosynthesis
      1. 8.1.0: The Purpose and Process of Photosynthesis
      2. 8.1.1: Main Structures and Summary of Photosynthesis
      3. 8.1.2: The Two Parts of Photosynthesis
    2. 8.2: The Light-Dependent Reactions of Photosynthesis
      1. 8.2.0: Introduction to Light Energy
      2. 8.2.1: Absorption of Light
      3. 8.2.2: Processes of the Light-Dependent Reactions
    3. 8.3: The Light-Independent Reactions of Photosynthesis
      1. 8.3.0: CAM and C4 Photosynthesis
      2. 8.3.1: The Calvin Cycle
      3. 8.3.2: The Carbon Cycle
  9. 9: Cell Communication
    1. 9.1: Signaling Molecules and Cellular Receptors
      1. 9.1.0: Signaling Molecules and Cellular Receptors
      2. 9.1.1: Forms of Signaling
      3. 9.1.2: Types of Receptors
      4. 9.1.3: Signaling Molecules
    2. 9.2: Propagation of the Cellular Signal
      1. 9.2.0: Binding Initiates a Signaling Pathway
      2. 9.2.1: Methods of Intracellular Signaling
    3. 9.3: Response to the Cellular Signal
      1. 9.3.0: Termination of the Signal Cascade
      2. 9.3.1: Cell Signaling and Gene Expression
      3. 9.3.2: Cell Signaling and Cellular Metabolism
      4. 9.3.3: Cell Signaling and Cell Growth
      5. 9.3.4: Cell Signaling and Cell Death
    4. 9.4: Signaling in Single-Celled Organisms
      1. 9.4.0: Signaling in Yeast
      2. 9.4.1: Signaling in Bacteria
  10. 10: Cell Reproduction
    1. 10.1: Cell Division
      1. 10.1.0: The Role of the Cell Cycle
      2. 10.1.1: Genomic DNA and Chromosomes
      3. 10.1.2: Eukaryotic Chromosomal Structure and Compaction
    2. 10.2: The Cell Cycle
      1. 10.2.0: Interphase
      2. 10.2.1: The Mitotic Phase and the G0 Phase
    3. 10.3: Control of the Cell Cycle
      1. 10.3.0: Regulation of the Cell Cycle by External Events
      2. 10.3.1: Regulation of the Cell Cycle at Internal Checkpoints
      3. 10.3.2: Regulator Molecules of the Cell Cycle
    4. 10.4: Cancer and the Cell Cycle
      1. 10.4.0: Proto-oncogenes
      2. 10.4.1: Tumor Suppressor Genes
    5. 10.5: Prokaryotic Cell Division
      1. 10.5.0: Binary Fission
  11. 11: Meiosis and Sexual Reproduction
    1. 11.1: The Process of Meiosis
      1. 11.1.0: Introduction to Meiosis
      2. 11.1.1: Meiosis I
      3. 11.1.2: Meiosis II
      4. 11.1.3: Comparing Meiosis and Mitosis
    2. 11.2: Sexual Reproduction
      1. 11.2.0: Advantages and Disadvantages of Sexual Reproduction
      2. 11.2.1: Life Cycles of Sexually Reproducing Organisms
  12. 12: Mendel's Experiments and Heredity
    1. 12.1: Mendel’s Experiments and the Laws of Probability
      1. 12.1.0: Introduction to Mendelian Inheritance
      2. 12.1.1: Mendel’s Model System
      3. 12.1.2: Mendelian Crosses
      4. 12.1.3: Garden Pea Characteristics Revealed the Basics of Heredity
      5. 12.1.4: Rules of Probability for Mendelian Inheritance
    2. 12.2: Patterns of Inheritance
      1. 12.2.0: Genes as the Unit of Heredity
      2. 12.2.1: Phenotypes and Genotypes
      3. 12.2.2: The Punnett Square Approach for a Monohybrid Cross
      4. 12.2.3: Alternatives to Dominance and Recessiveness
      5. 12.2.4: Sex-Linked Traits
      6. 12.2.5: Lethal Inheritance Patterns
    3. 12.3: Laws of Inheritance
      1. 12.3.0: Mendel's Laws of Heredity
      2. 12.3.1: Mendel's Law of Dominance
      3. 12.3.2: Mendel's Law of Segregation
      4. 12.3.3: Mendel's Law of Independent Assortment
      5. 12.3.4: Genetic Linkage and Violation of the Law of Independent Assortment
      6. 12.3.5: Epistasis
  13. 13: Modern Understandings of Inheritance
    1. 13.1: Chromosomal Theory and Genetic Linkage
      1. 13.1.0: Chromosomal Theory of Inheritance
      2. 13.1.1: Genetic Linkage and Distances
      3. 13.1.2: Identification of Chromosomes and Karyotypes
    2. 13.2: Chromosomal Basis of Inherited Disorders
      1. 13.2.0: Disorders in Chromosome Number
      2. 13.2.1: Chromosomal Structural Rearrangements
      3. 13.2.2: X-Inactivation
  14. 14: DNA Structure and Function
    1. 14.1: Historical Basis of Modern Understanding
      1. 14.1.0: Discovery of DNA
      2. 14.1.1: Modern Applications of DNA
    2. 14.2: DNA Structure and Sequencing
      1. 14.2.0: The Structure and Sequence of DNA
      2. 14.2.1: DNA Sequencing Techniques
    3. 14.3: DNA Replication
      1. 14.3.0: Basics of DNA Replication
      2. 14.3.1: DNA Replication in Prokaryotes
      3. 14.3.2: DNA Replication in Eukaryotes
      4. 14.3.3: Telomere Replication
    4. 14.4: DNA Repair
      1. 14.4.0: DNA Repair
  15. 15: Genes and Proteins
    1. 15.1: The Genetic Code
      1. 15.1.0: The Relationship Between Genes and Proteins
      2. 15.1.1: The Central Dogma: DNA Encodes RNA and RNA Encodes Protein
    2. 15.2: Prokaryotic Transcription
      1. 15.2.0: Transcription in Prokaryotes
      2. 15.2.1: Initiation of Transcription in Prokaryotes
      3. 15.2.2: Elongation and Termination in Prokaryotes
    3. 15.3: Eukaryotic Transcription
      1. 15.3.0: Initiation of Transcription in Eukaryotes
      2. 15.3.1: Elongation and Termination in Eukaryotes
    4. 15.4: RNA Processing in Eukaryotes
      1. 15.4.0: mRNA Processing
      2. 15.4.1: Processing of tRNAs and rRNAs
    5. 15.5: Ribosomes and Protein Synthesis
      1. 15.5.0: The Protein Synthesis Machinery
      2. 15.5.1: The Mechanism of Protein Synthesis
      3. 15.5.2: Protein Folding, Modification, and Targeting
  16. 16: Gene Expression
    1. 16.1: Regulation of Gene Expression
      1. 16.1.0: The Process and Purpose of Gene Expression Regulation
      2. 16.1.1: Prokaryotic versus Eukaryotic Gene Expression
    2. 16.2: Prokaryotic Gene Regulation
      1. 16.2.0: The trp Operon: A Repressor Operon
      2. 16.2.1: Catabolite Activator Protein (CAP): An Activator Regulator
      3. 16.2.2: The lac Operon: An Inducer Operon
    3. 16.3: Eukaryotic Gene Regulation
      1. 16.3.0: The Promoter and the Transcription Machinery
      2. 16.3.1: Transcriptional Enhancers and Repressors
      3. 16.3.2: Epigenetic Control: Regulating Access to Genes within the Chromosome
      4. 16.3.3: RNA Splicing
      5. 16.3.4: The Initiation Complex and Translation Rate
      6. 16.3.5: Regulating Protein Activity and Longevity
    4. 16.4: Regulating Gene Expression in Cell Development
      1. 16.4.0: Gene Expression in Stem Cells
      2. 16.4.1: Cellular Differentiation
      3. 16.4.2: Mechanics of Cellular Differentation
      4. 16.4.3: Establishing Body Axes during Development
      5. 16.4.4: Gene Expression for Spatial Positioning
      6. 16.4.5: Cell Migration in Multicellular Organisms
      7. 16.4.6: Programmed Cell Death
    5. 16.5: Cancer and Gene Regulation
      1. 16.5.0: Altered Gene Expression in Cancer
      2. 16.5.1: Epigenetic Alterations in Cancer
      3. 16.5.2: Cancer and Transcriptional Control
      4. 16.5.3: Cancer and Post-Transcriptional Control
      5. 16.5.4: Cancer and Translational Control
  17. 17: Biotechnology and Genomics
    1. 17.1: Biotechnology
      1. 17.1.0: Biotechnology
      2. 17.1.1: Basic Techniques to Manipulate Genetic Material (DNA and RNA)
      3. 17.1.2: Molecular and Cellular Cloning
      4. 17.1.3: Reproductive Cloning
      5. 17.1.4: Genetic Engineering
      6. 17.1.5: Genetically Modified Organisms (GMOs)
      7. 17.1.6: Biotechnology in Medicine
      8. 17.1.7: Production of Vaccines, Antibiotics, and Hormones
    2. 17.2: Mapping Genomes
      1. 17.2.0: Genetic Maps
      2. 17.2.1: Physical Maps and Integration with Genetic Maps
    3. 17.3: Whole-Genome Sequencing
      1. 17.3.0: Strategies Used in Sequencing Projects
      2. 17.3.1: Use of Whole-Genome Sequences of Model Organisms
      3. 17.3.2: Uses of Genome Sequences
    4. 17.4: Applying Genomics
      1. 17.4.0: Predicting Disease Risk at the Individual Level
      2. 17.4.1: Pharmacogenomics, Toxicogenomics, and Metagenomics
      3. 17.4.2: Genomics and Biofuels
    5. 17.5: Genomics and Proteomics
      1. 17.5.0: Genomics and Proteomics
      2. 17.5.1: Basic Techniques in Protein Analysis
      3. 17.5.2: Cancer Proteomics
  18. 18: Evolution and the Origin of Species
    1. 18.1: Understanding Evolution
      1. 18.1.0: What is Evolution?
      2. 18.1.1: Charles Darwin and Natural Selection
      3. 18.1.2: The Galapagos Finches and Natural Selection
      4. 18.1.3: Processes and Patterns of Evolution
      5. 18.1.4: Evidence of Evolution
      6. 18.1.5: Misconceptions of Evolution
    2. 18.2: Formation of New Species
      1. 18.2.0: The Biological Species Concept
      2. 18.2.1: Reproductive Isolation
      3. 18.2.2: Speciation
      4. 18.2.3: Allopatric Speciation
      5. 18.2.4: Sympatric Speciation
    3. 18.3: Hybrid Zones and Rates of Speciation
      1. 18.3.0: Hybrid Zones
      2. 18.3.1: Varying Rates of Speciation
    4. 18.4: Evolution of Genomes
      1. 18.4.0: Genomic Similiarities between Distant Species
      2. 18.4.1: Genome Evolution
      3. 18.4.2: Whole-Genome Duplication
      4. 18.4.3: Gene Duplications and Divergence
      5. 18.4.4: Noncoding DNA
      6. 18.4.5: Variations in Size and Number of Genes
    5. 18.5: Evidence of Evolution
      1. 18.5.0: The Fossil Record as Evidence for Evolution
      2. 18.5.1: Fossil Formation
      3. 18.5.2: Gaps in the Fossil Record
      4. 18.5.3: Carbon Dating and Estimating Fossil Age
      5. 18.5.4: The Fossil Record and the Evolution of the Modern Horse
      6. 18.5.5: Homologous Structures
      7. 18.5.6: Convergent Evolution
      8. 18.5.7: Vestigial Structures
      9. 18.5.8: Biogeography and the Distribution of Species
  19. 19: The Evolution of Populations
    1. 19.1: Population Evolution
      1. 19.1.0: Defining Population Evolution
      2. 19.1.1: Population Genetics
      3. 19.1.2: Hardy-Weinberg Principle of Equilibrium
    2. 19.2: Population Genetics
      1. 19.2.0: Genetic Variation
      2. 19.2.1: Genetic Drift
      3. 19.2.2: Gene Flow and Mutation
      4. 19.2.3: Nonrandom Mating and Environmental Variance
    3. 19.3: Adaptive Evolution
      1. 19.3.0: Natural Selection and Adaptive Evolution
      2. 19.3.1: Stabilizing, Directional, and Diversifying Selection
      3. 19.3.2: Frequency-Dependent Selection
      4. 19.3.3: Sexual Selection
      5. 19.3.4: No Perfect Organism
  20. 20: Phylogenies and the History of Life
    1. 20.1: Organizing Life on Earth
      1. 20.1.0: Phylogenetic Trees
      2. 20.1.1: Limitations of Phylogenetic Trees
      3. 20.1.2: The Levels of Classification
    2. 20.2: Determining Evolutionary Relationships
      1. 20.2.0: Distinguishing between Similar Traits
      2. 20.2.1: Building Phylogenetic Trees
    3. 20.3: Perspectives on the Phylogenetic Tree
      1. 20.3.0: Limitations to the Classic Model of Phylogenetic Trees
      2. 20.3.1: Horizontal Gene Transfer
      3. 20.3.2: Endosymbiotic Theory and the Evolution of Eukaryotes
      4. 20.3.3: Web, Network, and Ring of Life Models
  21. 21: Viruses
    1. 21.1: Viral Evolution, Morphology, and Classification
      1. 21.1.0: Discovery and Detection of Viruses
      2. 21.1.1: Evolution of Viruses
      3. 21.1.2: Viral Morphology
      4. 21.1.3: Virus Classification
    2. 21.2: Virus Infections and Hosts
      1. 21.2.0: Steps of Virus Infections
      2. 21.2.1: The Lytic and Lysogenic Cycles of Bacteriophages
      3. 21.2.2: Animal Viruses
      4. 21.2.3: Plant Viruses
    3. 21.3: Prevention and Treatment of Viral Infections
      1. 21.3.0: Vaccines and Immunity
      2. 21.3.1: Vaccines and Anti-Viral Drugs for Treatment
    4. 21.4: Prions and Viroids
      1. 21.4.0: Prions and Viroids
  22. 22: Prokaryotes: Bacteria and Archaea
    1. 22.1: Prokaryotic Diversity
      1. 22.1.0: Classification of Prokaryotes
      2. 22.1.1: The Origins of Archaea and Bacteria
      3. 22.1.2: Extremophiles and Biofilms
    2. 22.2: Structure of Prokaryotes
      1. 22.2.0: Basic Structures of Prokaryotic Cells
      2. 22.2.1: Prokaryotic Reproduction
    3. 22.3: Prokaryotic Metabolism
      1. 22.3.0: Energy and Nutrient Requirements for Prokaryotes
      2. 22.3.1: The Role of Prokaryotes in Ecosystems
    4. 22.4: Bacterial Diseases in Humans
      1. 22.4.0: History of Bacterial Diseases
      2. 22.4.1: Biofilms and Disease
      3. 22.4.2: Antibiotics: Are We Facing a Crisis?
      4. 22.4.3: Bacterial Foodborne Diseases
    5. 22.5: Beneficial Prokaryotes
      1. 22.5.0: Symbiosis between Bacteria and Eukaryotes
      2. 22.5.1: Early Biotechnology: Cheese, Bread, Wine, Beer, and Yogurt
      3. 22.5.2: Prokaryotes and Environmental Bioremediation
  23. 23: Protists
    1. 23.1: Eukaryotic Origins
      1. 23.1.0: Early Eukaryotes
      2. 23.1.1: Characteristics of Eukaryotic DNA
      3. 23.1.2: Endosymbiosis and the Evolution of Eukaryotes
      4. 23.1.3: The Evolution of Mitochondria
      5. 23.1.4: The Evolution of Plastids
    2. 23.2: Characteristics of Protists
      1. 23.2.0: Cell Structure, Metabolism, and Motility
      2. 23.2.1: Protist Life Cycles and Habitats
    3. 23.3: Groups of Protists
      1. 23.3.0: Excavata
      2. 23.3.1: Chromalveolata: Alveolates
      3. 23.3.2: Chromalveolata: Stramenopiles
      4. 23.3.3: Rhizaria
      5. 23.3.4: Archaeplastida
      6. 23.3.5: Amoebozoa and Opisthokonta
    4. 23.4: Ecology of Protists
      1. 23.4.0: Protists as Primary Producers, Food Sources, and Symbionts
      2. 23.4.1: Protists as Human Pathogens
      3. 23.4.2: Protists as Plant Pathogens
  24. 24: Fungi
    1. 24.1: Characteristics of Fungi
      1. 24.1.0: Characteristics of Fungi
      2. 24.1.1: Fungi Cell Structure and Function
      3. 24.1.2: Fungi Reproduction
    2. 24.2: Ecology of Fungi
      1. 24.2.0: Fungi Habitat, Decomposition, and Recycling
      2. 24.2.1: Mutualistic Relationships with Fungi and Fungivores
    3. 24.3: Classifications of Fungi
      1. 24.3.0: Chytridiomycota: The Chytrids
      2. 24.3.1: Zygomycota: The Conjugated Fungi
      3. 24.3.2: Ascomycota: The Sac Fungi
      4. 24.3.3: Basidiomycota: The Club Fungi
      5. 24.3.4: Deuteromycota: The Imperfect Fungi
      6. 24.3.5: Glomeromycota
    4. 24.4: Fungal Parasites and Pathogens
      1. 24.4.0: Fungi as Plant, Animal, and Human Pathogens
    5. 24.5: Importance of Fungi in Human Life
      1. 24.5.0: Importance of Fungi in Human Life
  25. 25: Seedless Plants
    1. 25.1: Early Plant Life
      1. 25.1.0: Early Plant Life
      2. 25.1.1: Evolution of Land Plants
      3. 25.1.2: Plant Adaptations to Life on Land
      4. 25.1.3: Sporophytes and Gametophytes in Seedless Plants
      5. 25.1.4: Structural Adaptations for Land in Seedless Plants
      6. 25.1.5: The Major Divisions of Land Plants
    2. 25.2: Green Algae: Precursors of Land Plants
      1. 25.2.0: Streptophytes and Reproduction of Green Algae
      2. 25.2.1: Charales
    3. 25.3: Bryophytes
      1. 25.3.0: Bryophytes
      2. 25.3.1: Liverworts and Hornworts
      3. 25.3.2: Mosses
    4. 25.4: Seedless Vascular Plants
      1. 25.4.0: Seedless Vascular Plants
      2. 25.4.1: Vascular Tissue: Xylem and Phloem
      3. 25.4.2: The Evolution of Roots in Seedless Plants
      4. 25.4.3: Ferns and Other Seedless Vascular Plants
      5. 25.4.4: The Importance of Seedless Vascular Plants
  26. 26: Seed Plants
    1. 26.1: Evolution of Seed Plants
      1. 26.1.0: The Evolution of Seed Plants and Adaptations for Land
      2. 26.1.1: Evolution of Gymnosperms
      3. 26.1.2: Evolution of Angiosperms
    2. 26.2: Gymnosperms
      1. 26.2.0: Characteristics of Gymnosperms
      2. 26.2.1: Life Cycle of a Conifer
      3. 26.2.2: Diversity of Gymnosperms
    3. 26.3: Angiosperms
      1. 26.3.0: Angiosperm Flowers
      2. 26.3.1: Angsiosperm Fruit
      3. 26.3.2: The Life Cycle of an Angiosperm
      4. 26.3.3: Diversity of Angiosperms
    4. 26.4: The Role of Seed Plants
      1. 26.4.0: Herbivory and Pollination
      2. 26.4.1: The Importance of Seed Plants in Human Life
      3. 26.4.2: Biodiversity of Plants
  27. 27: Introduction to Animal Diversity
    1. 27.1: Features of the Animal Kingdom
      1. 27.1.0: Characteristics of the Animal Kingdom
      2. 27.1.1: Complex Tissue Structure
      3. 27.1.2: Animal Reproduction and Development
    2. 27.2: Features Used to Classify Animals
      1. 27.2.0: Animal Characterization Based on Body Symmetry
      2. 27.2.1: Animal Characterization Based on Features of Embryological Development
    3. 27.3: Animal Phylogeny
      1. 27.3.0: Constructing an Animal Phylogenetic Tree
      2. 27.3.1: Molecular Analyses and Modern Phylogenetic Trees
    4. 27.4: The Evolutionary History of the Animal Kingdom
      1. 27.4.0: Pre-Cambrian Animal Life
      2. 27.4.1: The Cambrian Explosion of Animal Life
      3. 27.4.2: Post-Cambrian Evolution and Mass Extinctions
  28. 28: Invertebrates
    1. 28.1: Phylum Porifera
      1. 28.1.0: Phylum Porifera
      2. 28.1.1: Morphology of Sponges
      3. 28.1.2: Physiological Processes in Sponges
    2. 28.2: Phylum Cnidaria
      1. 28.2.0: Phylum Cnidaria
      2. 28.2.1: Class Anthozoa
      3. 28.2.2: Class Scyphozoa
      4. 28.2.3: Class Cubozoa and Class Hydrozoa
    3. 28.3: Superphylum Lophotrochozoa
      1. 28.3.0: Superphylum Lophotrochozoa
      2. 28.3.1: Phylum Platyhelminthes
      3. 28.3.2: Phylum Rotifera
      4. 28.3.3: Phylum Nemertea
      5. 28.3.4: Phylum Mollusca
      6. 28.3.5: Classification of Phylum Mollusca
      7. 28.3.6: Phylum Annelida
    4. 28.4: Superphylum Ecdysozoa
      1. 28.4.0: Superphylum Ecdysozoa
      2. 28.4.1: Phylum Nematoda
      3. 28.4.2: Phylum Arthropoda
      4. 28.4.3: Subphyla of Arthropoda
    5. 28.5: Superphylum Deuterostomia
      1. 28.5.0: Phylum Echinodermata
      2. 28.5.1: Classes of Echinoderms
      3. 28.5.2: Phylum Chordata
  29. 29: Vertebrates
    1. 29.1: Chordates
      1. 29.1.0: Characteristics of Chordata
      2. 29.1.1: Chordates and the Evolution of Vertebrates
      3. 29.1.2: The Evolution of Craniata and Vertebrata
      4. 29.1.3: Characteristics of Vertebrates
    2. 29.2: Fishes
      1. 29.2.0: Agnathans: Jawless Fishes
      2. 29.2.1: Gnathostomes: Jawed Fishes
    3. 29.3: Amphibians
      1. 29.3.0: Characteristics and Evolution of Amphibians
      2. 29.3.1: Modern Amphibians
    4. 29.4: Reptiles
      1. 29.4.0: Characteristics of Amniotes
      2. 29.4.1: Evolution of Amniotes
      3. 29.4.2: Characteristics of Reptiles
      4. 29.4.3: Evolution of Reptiles
      5. 29.4.4: Modern Reptiles
    5. 29.5: Birds
      1. 29.5.0: Characteristics of Birds
      2. 29.5.1: Evolution of Birds
    6. 29.6: Mammals
      1. 29.6.0: Characteristics of Mammals
      2. 29.6.1: Evolution of Mammals
      3. 29.6.2: Living Mammals
    7. 29.7: The Evolution of Primates
      1. 29.7.0: Characteristics and Evolution of Primates
      2. 29.7.1: Early Human Evolution
      3. 29.7.2: Early Hominins
      4. 29.7.3: Genus Homo
  30. 30: Plant Form and Physiology
    1. 30.1: The Plant Body
      1. 30.1.0: Plant Tissues and Organ Systems
    2. 30.2: Stems
      1. 30.2.0: Functions of Stems
      2. 30.2.1: Stem Anatomy
      3. 30.2.2: Primary and Secondary Growth in Stems
      4. 30.2.3: Stem Modifications
    3. 30.3: Roots
      1. 30.3.0: Types of Root Systems and Zones of Growth
      2. 30.3.1: Root Modifications
    4. 30.4: Leaves
      1. 30.4.0: Leaf Structure and Arrangment
      2. 30.4.1: Types of Leaf Forms
      3. 30.4.2: Leaf Structure, Function, and Adaptation
    5. 30.5: Plant Development
      1. 30.5.0: Meristems
      2. 30.5.1: Genetic Control of Flowers
    6. 30.6: Transport of Water and Solutes in Plants
      1. 30.6.0: Water and Solute Potential
      2. 30.6.1: Pressure, Gravity, and Matric Potential
      3. 30.6.2: Movement of Water and Minerals in the Xylem
      4. 30.6.3: Transportation of Photosynthates in the Phloem
    7. 30.7: Plant Sensory Systems and Responses
      1. 30.7.0: Plant Responses to Light
      2. 30.7.1: The Phytochrome System and Red Light Response
      3. 30.7.2: Blue Light Response
      4. 30.7.3: Plant Responses to Gravity
      5. 30.7.4: Auxins, Cytokinins, and Gibberellins
      6. 30.7.5: Abscisic Acid, Ethylene, and Nontraditional Hormones
      7. 30.7.6: Plant Responses to Wind and Touch
    8. 30.8: Plant Defense Mechanisms
      1. 30.8.0: Plant Defenses Against Herbivores
      2. 30.8.1: Plant Defenses Against Pathogens
  31. 31: Soil and Plant Nutrition
    1. 31.1: Nutritional Requirements of Plants
      1. 31.1.0: Plant Nutrition
      2. 31.1.1: The Chemical Composition of Plants
      3. 31.1.2: Essential Nutrients for Plants
    2. 31.2: The Soil
      1. 31.2.0: Soil Composition
      2. 31.2.1: Soil Formation
      3. 31.2.2: Physical Properties of Soil
    3. 31.3: Nutritional Adaptations of Plants
      1. 31.3.0: Nitrogen Fixation: Root and Bacteria Interactions
      2. 31.3.1: Mycorrhizae: The Symbiotic Relationship between Fungi and Roots
      3. 31.3.2: Nutrients from Other Sources
  32. 32: Plant Reproduction
    1. 32.1: Plant Reproductive Development and Structure
      1. 32.1.0: Plant Reproductive Development and Structure
      2. 32.1.1: Sexual Reproduction in Gymnosperms
      3. 32.1.2: Sexual Reproduction in Angiosperms
    2. 32.2: Pollination and Fertilization
      1. 32.2.0: Pollination and Fertilization
      2. 32.2.1: Pollination by Insects
      3. 32.2.2: Pollination by Bats, Birds, Wind, and Water
      4. 32.2.3: Double Fertilization in Plants
      5. 32.2.4: Development of the Seed
      6. 32.2.5: Development of Fruit and Fruit Types
      7. 32.2.6: Fruit and Seed Dispersal
    3. 32.3: Asexual Reproduction
      1. 32.3.0: Asexual Reproduction in Plants
      2. 32.3.1: Natural and Artificial Methods of Asexual Reproduction in Plants
      3. 32.3.2: Plant Life Spans
  33. 33: The Animal Body: Basic Form and Function
    1. 33.1: Animal Form and Function
      1. 33.1.0: Characteristics of the Animal Body
      2. 33.1.1: Body Plans
      3. 33.1.2: Limits on Animal Size and Shape
      4. 33.1.3: Limiting Effects of Diffusion on Size and Development
      5. 33.1.4: Animal Bioenergetics
      6. 33.1.5: Animal Body Planes and Cavities
    2. 33.2: Animal Primary Tissues
      1. 33.2.0: Epithelial Tissues
      2. 33.2.1: Connective Tissues: Loose, Fibrous, and Cartilage
      3. 33.2.2: Connective Tissues: Bone, Adipose, and Blood
      4. 33.2.3: Muscle Tissues and Nervous Tissues
    3. 33.3: Homeostasis
      1. 33.3.0: Homeostatic Process
      2. 33.3.1: Control of Homeostasis
      3. 33.3.2: Homeostasis: Thermoregulation
      4. 33.3.3: Heat Conservation and Dissipation
  34. 34: Animal Nutrition and the Digestive System
    1. 34.1: Digestive Systems
      1. 34.1.0: Digestive Systems
      2. 34.1.1: Herbivores, Omnivores, and Carnivores
      3. 34.1.2: Invertebrate Digestive Systems
      4. 34.1.3: Vertebrate Digestive Systems
      5. 34.1.4: Digestive System: Mouth and Stomach
      6. 34.1.5: Digestive System: Small and Large Intestines
    2. 34.2: Nutrition and Energy Production
      1. 34.2.0: Food Requirements and Essential Nutrients
      2. 34.2.1: Food Energy and ATP
    3. 34.3: Digestive System Processes
      1. 34.3.0: Ingestion
      2. 34.3.1: Digestion and Absorption
      3. 34.3.2: Elimination
    4. 34.4: Digestive System Regulation
      1. 34.4.0: Neural Responses to Food
      2. 34.4.1: Hormonal Responses to Food
  35. 35: The Nervous System
    1. 35.1: Neurons and Glial Cells
      1. 35.1.0: Neurons and Glial Cells
      2. 35.1.1: Neurons
      3. 35.1.2: Glia
    2. 35.2: How Neurons Communicate
      1. 35.2.0: Nerve Impulse Transmission within a Neuron: Resting Potential
      2. 35.2.1: Nerve Impulse Transmission within a Neuron: Action Potential
      3. 35.2.2: Synaptic Transmission
      4. 35.2.3: Signal Summation
      5. 35.2.4: Synaptic Plasticity
    3. 35.3: The Nervous System
      1. 35.3.0: The Nervous System
    4. 35.4: The Central Nervous System
      1. 35.4.0: Brain: Cerebral Cortex and Brain Lobes
      2. 35.4.1: Brain: Midbrain and Brain Stem
      3. 35.4.2: Spinal Cord
    5. 35.5: The Peripheral Nervous System
      1. 35.5.0: Autonomic Nervous System
      2. 35.5.1: Sensory-Somatic Nervous System
    6. 35.6: Nervous System Disorders
      1. 35.6.0: Neurodegenerative Disorders
      2. 35.6.1: Neurodevelopmental Disorders: Autism and ADHD
      3. 35.6.2: Neurodevelopmental Disorders: Mental Illnesses
      4. 35.6.3: Other Neurological Disorders
  36. 36: Sensory Systems
    1. 36.1: Sensory Processes
      1. 36.1.0: Reception
      2. 36.1.1: Transduction and Perception
    2. 36.2: Somatosensation
      1. 36.2.0: Somatosensory Receptors
      2. 36.2.1: Integration of Signals from Mechanoreceptors
      3. 36.2.2: Thermoreception
    3. 36.3: Taste and Smell
      1. 36.3.0: Tastes and Odors
      2. 36.3.1: Reception and Transduction
    4. 36.4: Hearing and Vestibular Sensation
      1. 36.4.0: Sound
      2. 36.4.1: Reception of Sound
      3. 36.4.2: Transduction of Sound
      4. 36.4.3: The Vestibular System
      5. 36.4.4: Balance and Determining Equilibrium
    5. 36.5: Vision
      1. 36.5.0: Light
      2. 36.5.1: Anatomy of the Eye
      3. 36.5.2: Transduction of Light
      4. 36.5.3: Visual Processing
  37. 37: The Endocrine System
    1. 37.1: Types of Hormones
      1. 37.1.0: Hormone Functions
      2. 37.1.1: Lipid-Derived, Amino Acid-Derived, and Peptide Hormones
    2. 37.2: How Hormones Work
      1. 37.2.0: How Hormones Work
      2. 37.2.1: Intracellular Hormone Receptors
      3. 37.2.2: Plasma Membrane Hormone Receptors
    3. 37.3: Regulation of Body Processes
      1. 37.3.0: Hormonal Regulation of the Excretory System
      2. 37.3.1: Hormonal Regulation of the Reproductive System
      3. 37.3.2: Hormonal Regulation of Metabolism
      4. 37.3.3: Hormonal Control of Blood Calcium Levels
      5. 37.3.4: Hormonal Regulation of Growth
      6. 37.3.5: Hormonal Regulation of Stress
    4. 37.4: Regulation of Hormone Production
      1. 37.4.0: Humoral, Hormonal, and Neural Stimuli
    5. 37.5: Endocrine Glands
      1. 37.5.0: Hypothalamic-Pituitary Axis
      2. 37.5.1: Thyroid Gland
      3. 37.5.2: Parathyroid Glands
      4. 37.5.3: Adrenal Glands
      5. 37.5.4: Pancreas
      6. 37.5.5: Pineal Gland and Gonads
      7. 37.5.6: Organs with Secondary Endocrine Functions
  38. 38: The Musculoskeletal System
    1. 38.1: Types of Skeletal Systems
      1. 38.1.0: Functions of the Musculoskeletal System
      2. 38.1.1: Types of Skeletal Systems
      3. 38.1.2: Human Axial Skeleton
      4. 38.1.3: Human Appendicular Skeleton
    2. 38.2: Bone
      1. 38.2.0: Bone
      2. 38.2.1: Cell Types in Bones
      3. 38.2.2: Bone Development
      4. 38.2.3: Growth of Bone
      5. 38.2.4: Bone Remodeling and Repair
    3. 38.3: Joints and Skeletal Movement
      1. 38.3.0: Classification of Joints on the Basis of Structure and Function
      2. 38.3.1: Movement at Synovial Joints
      3. 38.3.2: Types of Synovial Joints
      4. 38.3.3: Bone and Joint Disorders
    4. 38.4: Muscle Contraction and Locomotion
      1. 38.4.0: Structure and Function of the Muscular System
      2. 38.4.1: Skeletal Muscle Fibers
      3. 38.4.2: Sliding Filament Model of Contraction
      4. 38.4.3: ATP and Muscle Contraction
      5. 38.4.4: Regulatory Proteins
      6. 38.4.5: Excitation–Contraction Coupling
      7. 38.4.6: Control of Muscle Tension
  39. 39: The Respiratory System
    1. 39.1: Systems of Gas Exchange
      1. 39.1.0: The Respiratory System and Direct Diffusion
      2. 39.1.1: Skin, Gills, and Tracheal Systems
      3. 39.1.2: Amphibian and Bird Respiratory Systems
      4. 39.1.3: Mammalian Systems and Protective Mechanisms
    2. 39.2: Gas Exchange across Respiratory Surfaces
      1. 39.2.0: Gas Pressure and Respiration
      2. 39.2.1: Basic Principles of Gas Exchange
      3. 39.2.2: Lung Volumes and Capacities
      4. 39.2.3: Gas Exchange across the Alveoli
    3. 39.3: Breathing
      1. 39.3.0: The Mechanics of Human Breathing
      2. 39.3.1: Types of Breathing
      3. 39.3.2: The Work of Breathing
      4. 39.3.3: Dead Space: V/Q Mismatch
    4. 39.4: Transport of Gases in Human Bodily Fluids
      1. 39.4.0: Transport of Oxygen in the Blood
      2. 39.4.1: Transport of Carbon Dioxide in the Blood
  40. 40: The Circulatory System
    1. 40.1: Overview of the Circulatory System
      1. 40.1.0: The Role of the Circulatory System
      2. 40.1.1: Open and Closed Circulatory Systems
      3. 40.1.2: Types of Circulatory Systems in Animals
    2. 40.2: Components of the Blood
      1. 40.2.0: The Role of Blood in the Body
      2. 40.2.1: Red Blood Cells
      3. 40.2.2: White Blood Cells
      4. 40.2.3: Platelets and Coagulation Factors
      5. 40.2.4: Plasma and Serum
    3. 40.3: Mammalian Heart and Blood Vessels
      1. 40.3.0: Structures of the Heart
      2. 40.3.1: Arteries, Veins, and Capillaries
      3. 40.3.2: The Cardiac Cycle
    4. 40.4: Blood Flow and Blood Pressure Regulation
      1. 40.4.0: Blood Flow Through the Body
      2. 40.4.1: Blood Pressure
  41. 41: Osmotic Regulation and the Excretory System
    1. 41.1: Osmoregulation and Osmotic Balance
      1. 41.1.0: Introduction to Osmoregulation
      2. 41.1.1: Transport of Electrolytes across Cell Membranes
      3. 41.1.2: Concept of Osmolality and Milliequivalent
      4. 41.1.3: Osmoregulators and Osmoconformers
    2. 41.2: Nitrogenous Wastes
      1. 41.2.0: Nitrogenous Waste in Terrestrial Animals: The Urea Cycle
      2. 41.2.1: Nitrogenous Waste in Birds and Reptiles: Uric Acid
    3. 41.3: Excretion Systems
      1. 41.3.0: Contractile Vacuoles in Microorganisms
      2. 41.3.1: Flame Cells of Planaria and Nephridia of Worms
      3. 41.3.2: Malpighian Tubules of Insects
    4. 41.4: Human Osmoregulatory and Excretory Systems
      1. 41.4.0: Kidney Structure
      2. 41.4.1: Nephron: The Functional Unit of the Kidney
      3. 41.4.2: Kidney Function and Physiology
    5. 41.5: Hormonal Control of Osmoregulatory Functions
      1. 41.5.0: Epinephrine and Norepinephrine
      2. 41.5.1: Other Hormonal Controls for Osmoregulation
  42. 42: The Immune System
    1. 42.1: Innate Immune Response
      1. 42.1.0: Innate Immune Response
      2. 42.1.1: Physical and Chemical Barriers
      3. 42.1.2: Pathogen Recognition
      4. 42.1.3: Natural Killer Cells
      5. 42.1.4: The Complement System
    2. 42.2: Adaptive Immune Response
      1. 42.2.0: Antigen-presenting Cells: B and T cells
      2. 42.2.1: Humoral Immune Response
      3. 42.2.2: Cell-Mediated Immunity
      4. 42.2.3: Cytotoxic T Lymphocytes and Mucosal Surfaces
      5. 42.2.4: Immunological Memory
      6. 42.2.5: Regulating Immune Tolerance
    3. 42.3: Antibodies
      1. 42.3.0: Antibody Structure
      2. 42.3.1: Antibody Functions
    4. 42.4: Disruptions in the Immune System
      1. 42.4.0: Immunodeficiency
      2. 42.4.1: Hypersensitivities
  43. 43: Animal Reproduction and Development
    1. 43.1: Reproduction Methods
      1. 43.1.0: Methods of Reproducing
      2. 43.1.1: Types of Sexual and Asexual Reproduction
      3. 43.1.2: Sex Determination
    2. 43.2: Fertilization
      1. 43.2.0: External and Internal Fertilization
      2. 43.2.1: The Evolution of Reproduction
    3. 43.3: Human Reproductive Anatomy and Gametogenesis
      1. 43.3.0: Male Reproductive Anatomy
      2. 43.3.1: Female Reproductive Anatomy
      3. 43.3.2: Gametogenesis (Spermatogenesis and Oogenesis)
    4. 43.4: Hormonal Control of Human Reproduction
      1. 43.4.0: Male Hormones
      2. 43.4.1: Female Hormones
    5. 43.5: Fertilization and Early Embryonic Development
      1. 43.5.0: Fertilization
      2. 43.5.1: Cleavage, the Blastula Stage, and Gastrulation
    6. 43.6: Organogenesis and Vertebrate Formation
      1. 43.6.0: Organogenesis
      2. 43.6.1: Vertebrate Axis Formation
    7. 43.7: Human Pregnancy and Birth
      1. 43.7.0: Human Gestation
      2. 43.7.1: Labor and Birth
      3. 43.7.2: Contraception and Birth Control
      4. 43.7.3: Infertility
  44. 44: Ecology and the Biosphere
    1. 44.1: The Scope of Ecology
      1. 44.1.0: Introduction to Ecology
      2. 44.1.1: Organismal Ecology and Population Ecology
      3. 44.1.2: Community Ecology and Ecosystem Ecology
    2. 44.2: Biogeography
      1. 44.2.0: Biogeography
      2. 44.2.1: Energy Sources
      3. 44.2.2: Temperature and Water
      4. 44.2.3: Inorganic Nutrients and Other Factors
      5. 44.2.4: Abiotic Factors Influencing Plant Growth
    3. 44.3: Terrestrial Biomes
      1. 44.3.0: What constitutes a biome?
      2. 44.3.1: Tropical Wet Forest and Savannas
      3. 44.3.2: Subtropical Deserts and Chaparral
      4. 44.3.3: Temperate Grasslands
      5. 44.3.4: Temperate Forests
      6. 44.3.5: Boreal Forests and Arctic Tundra
    4. 44.4: Aquatic Biomes
      1. 44.4.0: Abiotic Factors Influencing Aquatic Biomes
      2. 44.4.1: Marine Biomes
      3. 44.4.2: Estuaries: Where the Ocean Meets Fresh Water
      4. 44.4.3: Freshwater Biomes
    5. 44.5: Climate and the Effects of Global Climate Change
      1. 44.5.0: Climate and Weather
      2. 44.5.1: Causes of Global Climate Change
      3. 44.5.2: Evidence of Global Climate Change
      4. 44.5.3: Past and Present Effects of Climate Change
  45. 45: Population and Community Ecology
    1. 45.1: Population Demography
      1. 45.1.0: Population Demography
      2. 45.1.1: Population Size and Density
      3. 45.1.2: Species Distribution
      4. 45.1.3: The Study of Population Dynamics
    2. 45.2: Environmental Limits to Population Growth
      1. 45.2.0: Exponential Population Growth
      2. 45.2.1: Logistic Population Growth
      3. 45.2.2: Density-Dependent and Density-Independent Population Regulation
    3. 45.3: Life History Patterns
      1. 45.3.0: Life History Patterns and Energy Budgets
      2. 45.3.1: Theories of Life History
    4. 45.4: Human Population Growth
      1. 45.4.0: Human Population Growth
      2. 45.4.1: Overcoming Density-Dependent Regulation
      3. 45.4.2: Age Structure, Population Growth, and Economic Development
    5. 45.5: Community Ecology
      1. 45.5.0: The Role of Species within Communities
      2. 45.5.1: Predation, Herbivory, and the Competitive Exclusion Principle
      3. 45.5.2: Symbiosis
      4. 45.5.3: Ecological Succession
    6. 45.6: Innate Animal Behavior
      1. 45.6.0: Introduction to Animal Behavior
      2. 45.6.1: Movement and Migration
      3. 45.6.2: Animal Communication and Living in Groups
      4. 45.6.3: Altruism and Populations
      5. 45.6.4: Mating Systems and Sexual Selection
    7. 45.7: Learned Animal Behavior
      1. 45.7.0: Simple Learned Behaviors
      2. 45.7.1: Conditioned Behavior
      3. 45.7.2: Cognitive Learning and Sociobiology
  46. 46: Ecosystems
    1. 46.1: Ecology of Ecosystems
      1. 46.1.0: Ecosystem Dynamics
      2. 46.1.1: Food Chains and Food Webs
      3. 46.1.2: Studying Ecosystem Dynamics
      4. 46.1.3: Modeling Ecosystem Dynamics
    2. 46.2: Energy Flow through Ecosystems
      1. 46.2.0: Strategies for Acquiring Energy
      2. 46.2.1: Productivity within Trophic Levels
      3. 46.2.2: Transfer of Energy between Trophic Levels
      4. 46.2.3: Ecological Pyramids
      5. 46.2.4: Biological Magnification
    3. 46.3: Biogeochemical Cycles
      1. 46.3.0: Biogeochemical Cycles
      2. 46.3.1: The Water (Hydrologic) Cycle
      3. 46.3.2: The Carbon Cycle
      4. 46.3.3: The Nitrogen Cycle
      5. 46.3.4: The Phosphorus Cycle
      6. 46.3.5: The Sulfur Cycle
  47. 47: Conservation Biology and Biodiversity
    1. 47.1: The Biodiversity Crisis
      1. 47.1.0: Loss of Biodiversity
      2. 47.1.1: Types of Biodiversity
      3. 47.1.2: Biodiversity Change through Geological Time
      4. 47.1.3: The Pleistocene Extinction
      5. 47.1.4: Present-Time Extinctions
    2. 47.2: The Importance of Biodiversity to Human Life
      1. 47.2.0: Human Health and Biodiversity
      2. 47.2.1: Agricultural Diversity
      3. 47.2.2: Managing Fisheries
    3. 47.3: Threats to Biodiversity
      1. 47.3.0: Habitat Loss and Sustainability
      2. 47.3.1: Overharvesting
      3. 47.3.2: Exotic Species
      4. 47.3.3: Climate Change and Biodiversity
    4. 47.4: Preserving Biodiversity
      1. 47.4.0: Measuring Biodiversity
      2. 47.4.1: Changing Human Behavior in Response to Biodiversity Loss
      3. 47.4.2: Ecological Restoration

2.1: Atoms, Isotopes, Ions, and Molecules

2.1.1: Overview of Atomic Structure

Atoms are made up of particles called protons, neutrons, and electrons, which are responsible for the mass and charge of atoms.

Learning Objective

Discuss the electronic and structural properties of an atom

Key Points

  • An atom is composed of two regions: the nucleus, which is in the center of the atom and contains protons and neutrons, and the outer region of the atom, which holds its electrons in orbit around the nucleus.
  • Protons and neutrons have approximately the same mass, about 1.67 × 10-24 grams, which scientists define as one atomic mass unit (amu) or one Dalton.
  • Each electron has a negative charge (-1) equal to the positive charge of a proton (+1).
  • Neutrons are uncharged particles found within the nucleus.

Key Terms

proton

Positively charged subatomic particle forming part of the nucleus of an atom and determining the atomic number of an element. It weighs 1 amu.

atom

The smallest possible amount of matter which still retains its identity as a chemical element, consisting of a nucleus surrounded by electrons.

neutron

A subatomic particle forming part of the nucleus of an atom. It has no charge. It is equal in mass to a proton or it weighs 1 amu.

An atom is the smallest unit of matter that retains all of the chemical properties of an element. Atoms combine to form molecules, which then interact to form solids, gases, or liquids. For example, water is composed of hydrogen and oxygen atoms that have combined to form water molecules. Many biological processes are devoted to breaking down molecules into their component atoms so they can be reassembled into a more useful molecule.

Atomic Particles

Atoms consist of three basic particles: protons, electrons, and neutrons. The nucleus (center) of the atom contains the protons (positively charged) and the neutrons (no charge). The outermost regions of the atom are called electron shells and contain the electrons (negatively charged). Atoms have different properties based on the arrangement and number of their basic particles.

The hydrogen atom (H) contains only one proton, one electron, and no neutrons. This can be determined using the atomic number and the mass number of the element (see the concept on atomic numbers and mass numbers).

Structure of an atom

Structure of an atom

Elements, such as helium, depicted here, are made up of atoms. Atoms are made up of protons and neutrons located within the nucleus, with electrons in orbitals surrounding the nucleus.

Atomic Mass

Protons and neutrons have approximately the same mass, about 1.67 × 10-24 grams. Scientists define this amount of mass as one atomic mass unit (amu) or one Dalton. Although similar in mass, protons are positively charged, while neutrons have no charge. Therefore, the number of neutrons in an atom contributes significantly to its mass, but not to its charge.

Electrons are much smaller in mass than protons, weighing only 9.11 × 10-28 grams, or about 1/1800 of an atomic mass unit. Therefore, they do not contribute much to an element's overall atomic mass. When considering atomic mass, it is customary to ignore the mass of any electrons and calculate the atom's mass based on the number of protons and neutrons alone.

Electrons contribute greatly to the atom's charge, as each electron has a negative charge equal to the positive charge of a proton. Scientists define these charges as "+1" and "-1. " In an uncharged, neutral atom, the number of electrons orbiting the nucleus is equal to the number of protons inside the nucleus. In these atoms, the positive and negative charges cancel each other out, leading to an atom with no net charge.

Protons, neutrons, and electrons

Protons, neutrons, and electrons

Both protons and neutrons have a mass of 1 amu and are found in the nucleus. However, protons have a charge of +1, and neutrons are uncharged. Electrons have a mass of approximately 0 amu, orbit the nucleus, and have a charge of -1.

Volume of Atoms

Accounting for the sizes of protons, neutrons, and electrons, most of the volume of an atom—greater than 99 percent—is, in fact, empty space. Despite all this empty space, solid objects do not just pass through one another. The electrons that surround all atoms are negatively charged and cause atoms to repel one another, preventing atoms from occupying the same space. These intermolecular forces prevent you from falling through an object like your chair.

2.1.2: Atomic Number and Mass Number

The atomic number is the number of protons in an element, while the mass number is the number of protons plus the number of neutrons.

Learning Objective

Determine the relationship between the mass number of an atom, its atomic number, its atomic mass, and its number of subatomic particles

Key Points

  • Neutral atoms of each element contain an equal number of protons and electrons.
  • The number of protons determines an element's atomic number and is used to distinguish one element from another.
  • The number of neutrons is variable, resulting in isotopes, which are different forms of the same atom that vary only in the number of neutrons they possess.
  • Together, the number of protons and the number of neutrons determine an element's mass number.
  • Since an element's isotopes have slightly different mass numbers, the atomic mass is calculated by obtaining the mean of the mass numbers for its isotopes.

Key Terms

atomic mass

The average mass of an atom, taking into account all its naturally occurring isotopes.

mass number

The sum of the number of protons and the number of neutrons in an atom.

atomic number

The number of protons in an atom.

Atomic Number

Neutral atoms of an element contain an equal number of protons and electrons. The number of protons determines an element's atomic number (Z) and distinguishes one element from another. For example, carbon's atomic number (Z) is 6 because it has 6 protons. The number of neutrons can vary to produce isotopes, which are atoms of the same element that have different numbers of neutrons. The number of electrons can also be different in atoms of the same element, thus producing ions (charged atoms). For instance, iron, Fe, can exist in its neutral state, or in the +2 and +3  ionic states.

Mass Number

An element's mass number (A) is the sum of the number of protons and the number of neutrons. The small contribution of mass from electrons is disregarded in calculating the mass number. This approximation of mass can be used to easily calculate how many neutrons an element has by simply subtracting the number of protons from the mass number. Protons and neutrons both weigh about one atomic mass unit or amu. Isotopes of the same element will have the same atomic number but different mass numbers.

Atomic number, chemical symbol, and mass number

Atomic number, chemical symbol, and mass number

Carbon has an atomic number of six, and two stable isotopes with mass numbers of twelve and thirteen, respectively. Its average atomic mass is 12.11.

Scientists determine the atomic mass by calculating the mean of the mass numbers for its naturally-occurring isotopes. Often, the resulting number contains a decimal. For example, the atomic mass of chlorine (Cl) is 35.45 amu because chlorine is composed of several isotopes, some (the majority) with an atomic mass of 35 amu (17 protons and 18 neutrons) and some with an atomic mass of 37 amu (17 protons and 20 neutrons).

Given an atomic number (Z) and mass number (A), you can find the number of protons, neutrons, and electrons in a neutral atom. For example, a lithium atom (Z=3, A=7 amu) contains three protons (found from Z), three electrons (as the number of protons is equal to the number of electrons in an atom), and four neutrons (7 – 3 = 4).

2.1.3: Isotopes

Isotopes are various forms of an element that have the same number of protons, but a different number of neutrons.

Learning Objective

Discuss the properties of isotopes and their use in radiometric dating

Key Points

  • Isotopes are atoms of the same element that contain an identical number of protons, but a different number of neutrons.
  • Despite having different numbers of neutrons, isotopes of the same element have very similar physical properties.
  • Some isotopes are unstable and will undergo radioactive decay to become other elements.
  • The predictable half-life of different decaying isotopes allows scientists to date material based on its isotopic composition, such as with Carbon-14 dating.

Key Terms

isotope

Any of two or more forms of an element where the atoms have the same number of protons, but a different number of neutrons within their nuclei.

half-life

The time it takes for half of the original concentration of an isotope to decay back to its more stable form.

radioactive isotopes

an atom with an unstable nucleus, characterized by excess energy available that undergoes radioactive decay and creates most commonly gamma rays, alpha or beta particles.

radiocarbon dating

Determining the age of an object by comparing the ratio of the 14C concentration found in it to the amount of 14C in the atmosphere.

What is an Isotope?

Isotopes are various forms of an element that have the same number of protons but a different number of neutrons. Some elements, such as carbon, potassium, and uranium, have multiple naturally-occurring isotopes. Isotopes are defined first by their element and then by the sum of the protons and neutrons present.

  • Carbon-12 (or 12C) contains six protons, six neutrons, and six electrons; therefore, it has a mass number of 12 amu (six protons and six neutrons).
  • Carbon-14 (or 14C) contains six protons, eight neutrons, and six electrons; its atomic mass is 14 amu (six protons and eight neutrons).

While the mass of individual isotopes is different, their physical and chemical properties remain mostly unchanged.

Isotopes do differ in their stability. Carbon-12 (12C) is the most abundant of the carbon isotopes, accounting for 98.89% of carbon on Earth. Carbon-14 (14C) is unstable and only occurs in trace amounts. Unstable isotopes most commonly emit alpha particles (He2+) and electrons. Neutrons, protons, and positrons can also be emitted and electrons can be captured to attain a more stable atomic configuration (lower level of potential energy) through a process called radioactive decay. The new atoms created may be in a high energy state and emit gamma rays which lowers the energy but alone does not change the atom into another isotope. These atoms are called radioactive isotopes or radioisotopes.

Radiocarbon Dating

Carbon is normally present in the atmosphere in the form of gaseous compounds like carbon dioxide and methane. Carbon-14 (14C) is a naturally-occurring radioisotope that is created from atmospheric 14N (nitrogen) by the addition of a neutron and the loss of a proton, which is caused by cosmic rays. This is a continuous process so more 14C is always being created in the atmosphere. Once produced, the 14C often combines with the oxygen in the atmosphere to form carbon dioxide. Carbon dioxide produced in this way diffuses in the atmosphere, is dissolved in the ocean, and is incorporated by plants via photosynthesis. Animals eat the plants and, ultimately, the radiocarbon is distributed throughout the biosphere.

In living organisms, the relative amount of 14C in their body is approximately equal to the concentration of 14C in the atmosphere. When an organism dies, it is no longer ingesting 14C, so the ratio between 14C and 12C will decline as 14C gradually decays back to 14N. This slow process, which is called beta decay, releases energy through the emission of electrons from the nucleus or positrons.

After approximately 5,730 years, half of the starting concentration of 14C will have been converted back to 14N. This is referred to as its half-life, or the time it takes for half of the original concentration of an isotope to decay back to its more stable form. Because the half-life of 14C is long, it is used to date formerly-living objects such as old bones or wood. Comparing the ratio of the 14C concentration found in an object to the amount of 14C in the atmosphere, the amount of the isotope that has not yet decayed can be determined. On the basis of this amount, the age of the material can be accurately calculated, as long as the material is believed to be less than 50,000 years old. This technique is called radiocarbon dating, or carbon dating for short.

Application of carbon dating

Application of carbon dating

The age of carbon-containing remains less than 50,000 years old, such as this pygmy mammoth, can be determined using carbon dating.

Other elements have isotopes with different half lives. For example, 40K (potassium-40) has a half-life of 1.25 billion years, and 235U (uranium-235) has a half-life of about 700 million years. Scientists often use these other radioactive elements to date objects that are older than 50,000 years (the limit of carbon dating). Through the use of radiometric dating, scientists can study the age of fossils or other remains of extinct organisms.

2.1.4: The Periodic Table

Everything in the universe is made of one or more elements. The periodic table is a means of organizing the various elements according to similar physical and chemical properties.

Learning Objective

Discuss the organization of the periodic table

Key Points

  • All matter is made from atoms of one or more elements. Living creatures consist mainly of carbon, hydrogen, oxygen, and nitrogen (CHON).
  • Combining elements creates compounds that may have emergent properties.
  • The periodic table is a listing of the elements according to increasing atomic number that is further organized into columns based on similar physical and chemical properties and electron configuration.
  • As one moves down a column or across a row, there are some general trends for the properties of the elements.
  • The periodic table continues to expand today as heavier and heavier elements are created in laboratories around the world.

Key Terms

periodic table

A tabular chart of the chemical elements according to their atomic numbers so that elements with similar properties are in the same column.

emergent properties

Properties found in compound structures that are different from those of the individual components and would not be predicted based on the properties of the individual components.

element

Pure chemical substances consisting of only one type of atom with a defined set of chemical and physical properties.

Example

  • Salt (NaCl) is a good example of a compound with emergent properties. Sodium is a metal that is highly reactive with water, and chlorine is a gas that is poisonous to humans. Together they form an ionic compound, NaCl, or table salt, that is safely consumed by humans everyday.

Matter and Elements

Matter comprises all of the physical objects in the universe, those that take up space and have mass. All matter is composed of atoms of one or more elements, pure substances with specific chemical and physical properties. There are 98 elements that naturally occur on earth, yet living systems use a relatively small number of these. Living creatures are composed mainly of just four elements: carbon, hydrogen, oxygen, and nitrogen (often remembered by the acronym CHON). As elements are bonded together they form compounds that often have new emergent properties that are different from the properties of the individual elements. Life is an example of an emergent property that arises from the specific collection of molecules found in cells.

Elements of the human body arranged by percent of total mass

Elements of the human body arranged by percent of total mass

There are 25 elements believed to play an active role in human health. Carbon, hydrogen, oxygen, and nitrogen make up approximately 96% of the mass in a human body.

The Periodic Table

The different elements are organized and displayed in the periodic table. Devised by Russian chemist Dmitri Mendeleev (1834–1907) in 1869, the table groups elements that, although unique, share certain chemical properties with other elements. In the periodic table the elements are organized and displayed according to their atomic number and are arranged in a series of rows (periods) and columns (groups) based on shared chemical and physical properties. If you look at a periodic table, you will see the groups numbered at the top of each column from left to right starting with 1 and ending with 18. In addition to providing the atomic number for each element, the periodic table also displays the element's atomic mass. Looking at carbon, for example, its symbol (C) and name appear, as well as its atomic number of six (in the upper left-hand corner) and its atomic mass of 12.11.

The periodic table

The periodic table

The periodic table shows the atomic mass and atomic number of each element. The atomic number appears above the symbol for the element and the approximate atomic mass appears below it.

The arrangement of the periodic table allows the elements to be grouped according to their chemical properties. Within the main group elements ( Groups 1-2, 13-18), there are some general trends that we can observe. The further down a given group, the elements have an increased metallic character: they are good conductors of both heat and electricity, solids at room temperature, and shiny in appearance. Moving from left to right across a period, the elements have greater non-metallic character. These elements are insulators, poor heat conductors, and can exist in different phases at room temperature (brittle solid, liquid, or gas). The elements at the boundary between the metallic elements (grey elements) and nonmetal elements (green elements) are metalloid in character (pink elements). They have low electrical conductivity that increases with temperature. They also share properties with both the metals and the nonmetals.

The main group elements

The main group elements

Within the p-block at the boundary between the metallic elements (grey elements) and nonmetal elements (green elements) there is positioned boron and silicon that are metalloid in character (pink elements), i.e., they have low electrical conductivity that increases with temperature.

Today, the periodic table continues to expand as heavier and heavier elements are synthesized in laboratories. These large elements are extremely unstable and, as such, are very difficult to detect; but their continued creation is an ongoing challenge undertaken by scientists around the world.

2.1.5: Electron Shells and the Bohr Model

Niels Bohr proposed an early model of the atom as a central nucleus containing protons and neutrons being orbited by electrons in shells.

Learning Objective

Construct an atom according to the Bohr model

Key Points

  • In the Bohr model of the atom, the nucleus contains the majority of the mass of the atom in its protons and neutrons.
  • Orbiting the positively-charged core are the negatively charged electrons, which contribute little in terms of mass, but are electrically equivalent to the protons in the nucleus.
  • In most cases, electrons fill the lower-energy orbitals first, followed by the next higher energy orbital until it is full, and so on until all electrons have been placed.
  • Atoms tend to be most stable with a full outer shell (one which, after the first, contains 8 electrons), leading to what is commonly called the "octet rule".
  • The properties of an element are determined by its outermost electrons, or those in the highest energy orbital.
  • Atoms that do not have full outer shells will tend to gain or lose electrons, resulting in a full outer shell and, therefore, stability.

Key Terms

octet rule

A rule stating that atoms lose, gain, or share electrons in order to have a full valence shell of 8 electrons. (Hydrogen is excluded because it can hold a maximum of 2 electrons in its valence shell. )

electron shell

The collective states of all electrons in an atom having the same principal quantum number (visualized as an orbit in which the electrons move).

Electron Shells and the Bohr Model

As previously discussed, there is a connection between the number of protons in an element, the atomic number that distinguishes one element from another, and the number of electrons it has. In all electrically-neutral atoms, the number of electrons is the same as the number of protons. Each element, when electrically neutral, has a number of electrons equal to its atomic number.

An early model of the atom was developed in 1913 by Danish scientist Niels Bohr (1885–1962). The Bohr model shows the atom as a central nucleus containing protons and neutrons with the electrons in circular orbitals at specific distances from the nucleus . These orbits form electron shells or energy levels, which are a way of visualizing the number of electrons in the various shells. These energy levels are designated by a number and the symbol "n." For example, 1n represents the first energy level located closest to the nucleus.

Orbitals in the Bohr model

Orbitals in the Bohr model

The Bohr model was developed by Niels Bohr in 1913. In this model, electrons exist within principal shells. An electron normally exists in the lowest energy shell available, which is the one closest to the nucleus. Energy from a photon of light can bump it up to a higher energy shell, but this situation is unstable and the electron quickly decays back to the ground state. In the process, a photon of light is released.

Electrons fill orbit shells in a consistent order. Under standard conditions, atoms fill the inner shells (closer to the nucleus) first, often resulting in a variable number of electrons in the outermost shell. The innermost shell has a maximum of two electrons, but the next two electron shells can each have a maximum of eight electrons. This is known as the octet rule which states that, with the exception of the innermost shell, atoms are more stable energetically when they have eight electrons in their valence shell, the outermost electron shell. Examples of some neutral atoms and their electron configurations are shown in . As shown, helium has a complete outer electron shell, with two electrons filling its first and only shell. Similarly, neon has a complete outer 2n shell containing eight electrons. In contrast, chlorine and sodium have seven and one electrons in their outer shells, respectively. Theoretically, they would be more energetically stable if they followed the octet rule and had eight.

Bohr diagrams

Bohr diagrams

Bohr diagrams indicate how many electrons fill each principal shell. Group 18 elements (helium, neon, and argon are shown) have a full outer, or valence, shell. A full valence shell is the most stable electron configuration. Elements in other groups have partially-filled valence shells and gain or lose electrons to achieve a stable electron configuration.

An atom may gain or lose electrons to achieve a full valence shell, the most stable electron configuration. The periodic table is arranged in columns and rows based on the number of electrons and where these electrons are located, providing a tool to understand how electrons are distributed in the outer shell of an atom. As shown in , the group 18 atoms helium (He), neon (Ne), and argon (Ar) all have filled outer electron shells, making it unnecessary for them to gain or lose electrons to attain stability; they are highly stable as single atoms. Their non-reactivity has resulted in their being named the inert gases (or noble gases). In comparison, the group 1 elements, including hydrogen (H), lithium (Li), and sodium (Na), all have one electron in their outermost shells. This means that they can achieve a stable configuration and a filled outer shell by donating or losing an electron. As a result of losing a negatively-charged electron, they become positively-charged ions. When an atom loses an electron to become a positively-charged ion, this is indicated by a plus sign after the element symbol; for example, Na+. Group 17 elements, including fluorine and chlorine, have seven electrons in their outermost shells; they tend to fill this shell by gaining an electron from other atoms, making them negatively-charged ions. When an atom gains an electron to become a negatively-charged ion this is indicated by a minus sign after the element symbol; for example, F-. Thus, the columns of the periodic table represent the potential shared state of these elements' outer electron shells that is responsible for their similar chemical characteristics.

2.1.6: Electron Orbitals

Electron orbitals are three-dimensional representations of the space in which an electron is likely to be found.

Learning Objective

Distinguish between electron orbitals in the Bohr model versus the quantum mechanical orbitals

Key Points

  • The Bohr model of the atom does not accurately reflect how electrons are spatially distributed around the nucleus as they do not circle the nucleus like the earth orbits the sun.
  • The electron orbitals are the result of mathematical equations from quantum mechanics known as wave functions and can predict within a certain level of probability where an electron might be at any given time.
  • The number and type of orbitals increases with increasing atomic number, filling in various electron shells.
  • The area where an electron is most likely to be found is called its orbital.

Key Terms

electron shell

The collective states of all electrons in an atom having the same principal quantum number (visualized as an orbit in which the electrons move).

orbital

A specification of the energy and probability density of an electron at any point in an atom or molecule.

Although useful to explain the reactivity and chemical bonding of certain elements, the Bohr model of the atom does not accurately reflect how electrons are spatially distributed surrounding the nucleus. They do not circle the nucleus like the earth orbits the sun, but are rather found in electron orbitals. These relatively complex shapes result from the fact that electrons behave not just like particles, but also like waves. Mathematical equations from quantum mechanics known as wave functions can predict within a certain level of probability where an electron might be at any given time. The area where an electron is most likely to be found is called its orbital.

First Electron Shell

The closest orbital to the nucleus, called the 1s orbital, can hold up to two electrons. This orbital is equivalent to the innermost electron shell of the Bohr model of the atom. It is called the 1s orbital because it is spherical around the nucleus. The 1s orbital is always filled before any other orbital. Hydrogen has one electron; therefore, it has only one spot within the 1s orbital occupied. This is designated as 1s1, where the superscripted 1 refers to the one electron within the 1s orbital. Helium has two electrons; therefore, it can completely fill the 1s orbital with its two electrons. This is designated as 1s2, referring to the two electrons of helium in the 1s orbital. On the periodic table, hydrogen and helium are the only two elements in the first row (period); this is because they are the sole elements to have electrons only in their first shell, the 1s orbital.

Second Electron Shell

The second electron shell may contain eight electrons. This shell contains another spherical s orbital and three "dumbbell" shaped p orbitals, each of which can hold two electrons . After the 1s orbital is filled, the second electron shell is filled, first filling its 2s orbital and then its three p orbitals. When filling the p orbitals, each takes a single electron; once each p orbital has an electron, a second may be added. Lithium (Li) contains three electrons that occupy the first and second shells. Two electrons fill the 1s orbital, and the third electron then fills the 2s orbital. Its electron configuration is 1s22s1. Neon (Ne), on the other hand, has a total of ten electrons: two are in its innermost 1s orbital, and eight fill its second shell (two each in the 2s and three p orbitals). Thus, it is an inert gas and energetically stable: it rarely forms a chemical bond with other atoms.

Diagram of the S and P orbitals

Diagram of the S and P orbitals

The s subshells are shaped like spheres. Both the 1n and 2n principal shells have an s orbital, but the size of the sphere is larger in the 2n orbital. Each sphere is a single orbital. p subshells are made up of three dumbbell-shaped orbitals. Principal shell 2n has a p subshell, but shell 1 does not.

Third Electron Shell

Larger elements have additional orbitals, making up the third electron shell. Subshells d and f have more complex shapes and contain five and seven orbitals, respectively. Principal shell 3n has s, p, and d subshells and can hold 18 electrons. Principal shell 4n has s, p, d, and f orbitals and can hold 32 electrons. Moving away from the nucleus, the number of electrons and orbitals found in the energy levels increases. Progressing from one atom to the next in the periodic table, the electron structure can be worked out by fitting an extra electron into the next available orbital. While the concepts of electron shells and orbitals are closely related, orbitals provide a more accurate depiction of the electron configuration of an atom because the orbital model specifies the different shapes and special orientations of all the places that electrons may occupy.

2.1.7: Chemical Reactions and Molecules

Chemical reactions occur when two or more atoms bond together to form molecules or when bonded atoms are broken apart.

Learning Objective

Describe the properties of chemical reactions and compounds

Key Points

  • Atoms form chemical bonds with other atoms thereby obtaining the electrons they need to attain a stable electron configuration.
  • The substances used in the beginning of a chemical reaction are called the reactants and the substances found at the end of the reaction are known as the products.
  • Some reactions are reversible and will reach a relative balance between reactants and products: a state called equilibrium.
  • An arrow is typically drawn between the reactants and products to indicate the direction of the chemical reaction.

Key Terms

reaction

A transformation in which one or more substances is converted into another by combination or decomposition

molecule

The smallest particle of a specific compound that retains the chemical properties of that compound; two or more atoms held together by chemical bonds.

reactant

Any of the participants present at the start of a chemical reaction.

Chemical Reactions and Molecules

According to the octet rule, elements are most stable when their outermost shell is filled with electrons. This is because it is energetically favorable for atoms to be in that configuration. However, since not all elements have enough electrons to fill their outermost shells, atoms form chemical bonds with other atoms, which helps them obtain the electrons they need to attain a stable electron configuration. When two or more atoms chemically bond with each other, the resultant chemical structure is a molecule. The familiar water molecule, H2O, consists of two hydrogen atoms and one oxygen atom, which bond together to form water . Atoms can form molecules by donating, accepting, or sharing electrons to fill their outer shells.

Atoms bond to form molecules

Atoms bond to form molecules

Two or more atoms may bond with each other to form a molecule. When two hydrogens and an oxygen share electrons via covalent bonds, a water molecule is formed.

Chemical reactions occur when two or more atoms bond together to form molecules or when bonded atoms are broken apart. The substances used in the beginning of a chemical reaction are called the reactants (usually found on the left side of a chemical equation), and the substances found at the end of the reaction are known as the products (usually found on the right side of a chemical equation). An arrow is typically drawn between the reactants and products to indicate the direction of the chemical reaction. For the creation of the water molecule shown above, the chemical equation would be:

An example of a simple chemical reaction is the breaking down of hydrogen peroxide molecules, each of which consists of two hydrogen atoms bonded to two oxygen atoms (H2O2). The reactant hydrogen peroxide is broken down into water (H2O), and oxygen, which consists of two bonded oxygen atoms (O2). In the equation below, the reaction includes two hydrogen peroxide molecules and two water molecules. This is an example of a balanced chemical equation, wherein the number of atoms of each element is the same on each side of the equation. According to the law of conservation of matter, the number of atoms before and after a chemical reaction should be equal, such that no atoms are, under normal circumstances, created or destroyed.

Even though all of the reactants and products of this reaction are molecules (each atom remains bonded to at least one other atom), in this reaction only hydrogen peroxide and water are representative of a subclass of molecules known as compounds: they contain atoms of more than one type of element. Molecular oxygen, on the other hand, consists of two doubly bonded oxygen atoms and is not classified as a compound but as an element .

Irreversible and Reversible Reactions

Some chemical reactions, such as the one shown above, can proceed in one direction until the reactants are all used up. The equations that describe these reactions contain a unidirectional arrow and are irreversible. Reversible reactions are those that can go in either direction. In reversible reactions, reactants are turned into products, but when the concentration of product goes beyond a certain threshold, some of these products will be converted back into reactants; at this point, the designations of products and reactants are reversed. This back and forth continues until a certain relative balance between reactants and products occurs: a state called equilibrium. These situations of reversible reactions are often denoted by a chemical equation with a double headed arrow pointing towards both the reactants and products.

For example, in human blood, excess hydrogen ions (H+) bind to bicarbonate ions (HCO3-) forming an equilibrium state with carbonic acid (H2CO3). If carbonic acid were added to this system, some of it would be converted to bicarbonate and hydrogen ions.

In biological reactions, however, equilibrium is rarely obtained because the concentrations of the reactants or products or both are constantly changing, often with a product of one reaction being a reactant for another. To return to the example of excess hydrogen ions in the blood, the formation of carbonic acid will be the major direction of the reaction. However, the carbonic acid can also leave the body as carbon dioxide gas (via exhalation) instead of being converted back to bicarbonate ion, thus driving the reaction to the right by the chemical law known as law of mass action. These reactions are important for maintaining the homeostasis of our blood.

2.1.8: Ions and Ionic Bonds

Ionic bonds are attractions between oppositely charged atoms or groups of atoms where electrons are donated and accepted.

Learning Objective

Predict whether a given element will more likely form a cation or an anion

Key Points

  • Ions form from elements when they gain or lose an electron causing the number of protons to be unequal to the number of electrons, resulting in a net charge.
  • If there are more electrons than protons (from an element gaining one or more electrons), the ion is negatively charged and called an anion.
  • If there are more protons than electrons (via loss of electrons), the ion is positively charged and is called a cation.
  • Ionic bonds result from the interaction between a positively charged cation and a negatively charged anion.

Key Terms

ionic bond

A strong chemical bond caused by the electrostatic attraction between two oppositely charged ions.

ion

An atom, or group of atoms, bearing an electrical charge, such as the sodium and chlorine atoms in a salt solution.

Ions and Ionic Bonds

Some atoms are more stable when they gain or lose an electron (or possibly two) and form ions. This results in a full outermost electron shell and makes them energetically more stable. Now, because the number of electrons does not equal the number of protons, each ion has a net charge. Cations are positive ions that are formed by losing electrons (as the number of protons is now greater than the number of electrons). Negative ions are formed by gaining electrons and are called anions (wherein there are more electrons than protons in a molecule). Anions are designated by their elemental name being altered to end in "-ide". For example, the anion of chlorine is called chloride, and the anion of sulfur is called sulfide.

This movement of electrons from one element to another is referred to as electron transfer. As illustrated, sodium (Na) only has one electron in its outer electron shell. It takes less energy for sodium to donate that one electron than it does to accept seven more electrons to fill the outer shell . When sodium loses an electron, it will have 11 protons, 11 neutrons, and only 10 electrons. This leaves it with an overall charge of +1 since there are now more protons than electrons. It is now referred to as a sodium ion. Chlorine (Cl) in its lowest energy state (called the ground state) has seven electrons in its outer shell. Again, it is more energy efficient for chlorine to gain one electron than to lose seven. Therefore, it tends to gain an electron to create an ion with 17 protons, 17 neutrons, and 18 electrons. This gives it a net charge of -1 since there are now more electrons than protons. It is now referred to as a chloride ion. In this example, sodium will donate its one electron to empty its shell, and chlorine will accept that electron to fill its shell. Both ions now satisfy the octet rule and have complete outer shells. These transactions can normally only take place simultaneously; in order for a sodium atom to lose an electron, it must be in the presence of a suitable recipient like a chlorine atom.

Electron Transfer Between Na and Cl

Electron Transfer Between Na and Cl

In the formation of an ionic compound, metals lose electrons and nonmetals gain electrons to achieve an octet. In this example, sodium loses one electron to empty its shell and chlorine accepts that electron to fill its shell.

Ionic bonds are formed between ions with opposite charges. For instance, positively charged sodium ions and negatively charged chloride ions bond together to form sodium chloride, or table salt, a crystalline molecule with zero net charge. The attractive force holding the two atoms together is called the electromagnetic force and is responsible for the attraction between oppositely charged ions.

Certain salts are referred to in physiology as electrolytes (including sodium, potassium, and calcium). Electrolytes are ions necessary for nerve impulse conduction, muscle contractions, and water balance. Many sports drinks and dietary supplements provide these ions to replace those lost from the body via sweating during exercise.

2.1.9: Covalent Bonds and Other Bonds and Interactions

Covalent bonds result from a sharing of electrons between two atoms and hold most biomolecules together.

Learning Objective

Compare the relative strength of different types of bonding interactions

Key Points

  • A polar covalent bond arises when two atoms of different electronegativity share two electrons unequally.
  • A non-polar covalent bond is one in which the electrons are shared equally between two atoms.
  • Hydrogen bonds and Van Der Waals are responsible for the folding of proteins, the binding of ligands to proteins, and many other processes between molecules.

Key Terms

dipole

Any object (such as a magnet, polar molecule or antenna), that is oppositely charged at two points (or poles).

hydrogen bond

A weak bond in which a hydrogen atom in one molecule is attracted to an electronegative atom (usually nitrogen or oxygen) in the same or different molecule.

covalent bond

A type of chemical bond where two atoms are connected to each other by the sharing of two or more electrons.

Example

  • Adenosine triphosphate, or ATP, is the most-commonly used cofactor in all of biology. Its biosynthesis involves breaking the triple bond of molecular nitrogen, or N2, followed by the formation of several carbon-nitrogen single and double bonds.

Covalent Bonds and Other Bonds and Interactions

The octet rule can be satisfied by the sharing of electrons between atoms to form covalent bonds. These bonds are stronger and much more common than are ionic bonds in the molecules of living organisms. Covalent bonds are commonly found in carbon-based organic molecules, such as DNA and proteins. Covalent bonds are also found in inorganic molecules such as H2O, CO2, and O2. One, two, or three pairs of electrons may be shared between two atoms, making single, double, and triple bonds, respectively. The more covalent bonds between two atoms, the stronger their connection. Thus, triple bonds are the strongest.

The strength of different levels of covalent bonding is one of the main reasons living organisms have a difficult time in acquiring nitrogen for use in constructing nitrogenous molecules, even though molecular nitrogen, N2, is the most abundant gas in the atmosphere. Molecular nitrogen consists of two nitrogen atoms triple bonded to each other. The resulting strong triple bond makes it difficult for living systems to break apart this nitrogen in order to use it as constituents of biomolecules, such as proteins, DNA, and RNA.

The formation of water molecules is an example of covalent bonding. The hydrogen and oxygen atoms that combine to form water molecules are bound together by covalent bonds. The electron from the hydrogen splits its time between the incomplete outer shell of the hydrogen atom and the incomplete outer shell of the oxygen atom. In return, the oxygen atom shares one of its electrons with the hydrogen atom, creating a two-electron single covalent bond. To completely fill the outer shell of oxygen, which has six electrons in its outer shell, two electrons (one from each hydrogen atom) are needed. Each hydrogen atom needs only a single electron to fill its outer shell, hence the well-known formula H2O. The electrons that are shared between the two elements fill the outer shell of each, making both elements more stable.

Polar Covalent Bonds

There are two types of covalent bonds: polar and nonpolar. In a polar covalent bond, the electrons are unequally shared by the atoms because they are more attracted to one nucleus than the other . The relative attraction of an atom to an electron is known as its electronegativity: atoms that are more attracted to an electron are considered to be more electronegative. Because of the unequal distribution of electrons between the atoms of different elements, a slightly positive (δ+) or slightly negative (δ-) charge develops. This partial charge is known as a dipole; this is an important property of water and accounts for many of its characteristics. The dipole in water occurs because oxygen has a higher electronegativity than hydrogen, which means that the shared electrons spend more time in the vicinity of the oxygen nucleus than they do near the nucleus of the hydrogen atoms.

Polar and Nonpolar Covalent Bonds

Polar and Nonpolar Covalent Bonds

Whether a molecule is polar or nonpolar depends both on bond type and molecular shape. Both water and carbon dioxide have polar covalent bonds, but carbon dioxide is linear, so the partial charges on the molecule cancel each other out.

Nonpolar Covalent Bonds

Nonpolar covalent bonds form between two atoms of the same element or between different elements that share electrons equally. For example, molecular oxygen (O2) is nonpolar because the electrons will be equally distributed between the two oxygen atoms. The four bonds of methane are also considered to be nonpolar because the electronegativies of carbon and hydrogen are nearly identical.

Hydrogen Bonds and Van Der Waals Interactions

Not all bonds are ionic or covalent; weaker bonds can also form between molecules. Two types of weak bonds that frequently occur are hydrogen bonds and van der Waals interactions. Without these two types of bonds, life as we know it would not exist.

Hydrogen bonds provide many of the critical, life-sustaining properties of water and also stabilize the structures of proteins and DNA, the building block of cells. When polar covalent bonds containing hydrogen are formed, the hydrogen atom in that bond has a slightly positive charge (δ+) because the shared electrons are pulled more strongly toward the other element and away from the hydrogen atom. Because the hydrogen has a slightly positive charge, it's attracted to neighboring negative charges. The weak interaction between the δ+ charge of a hydrogen atom from one molecule and the δ- charge of a more electronegative atom is called a hydrogen bond. Individual hydrogen bonds are weak and easily broken; however, they occur in very large numbers in water and in organic polymers, and the additive force can be very strong. For example, hydrogen bonds are responsible for zipping together the DNA double helix.

Adenosine Triphosphate, ATP

Adenosine Triphosphate, ATP

Adenosine Triphosphate, or ATP, is the most commonly used cofactor in nature. Its biosynthesis involves the fixation of nitrogen to provide feedstocks that eventually produce the carbon-nitrogen bonds it contains.

Like hydrogen bonds, van der Waals interactions are weak interactions between molecules. Van der Waals attractions can occur between any two or more molecules and are dependent on slight fluctuations of the electron densities, which can lead to slight temporary dipoles around a molecule. For these attractions to happen, the molecules need to be very close to one another. These bonds, along with hydrogen bonds, help form the three-dimensional structures of the proteins in our cells that are required for their proper function.

2.1.10: Hydrogen Bonding and Van der Waals Forces

Hydrogen bonds and van der Waals interactions are two types of weak bonds that are necessary to the basic building blocks of life.

Learning Objective

Describe how hydrogen bonds and van der Waals interactions occur

Key Points

  • Hydrogen bonds provide many of the critical, life-sustaining properties of water and also stabilize the structures of proteins and DNA, the building block of cells.
  • Hydrogen bonds occur in inorganic molecules, such as water, and organic molecules, such as DNA and proteins.
  • Van der Waals attractions can occur between any two or more molecules and are dependent on slight fluctuations of the electron densities.
  • While hydrogen bonds and van der Waals interactions are weak individually, they are strong combined in vast numbers.

Key Terms

electronegativity

The tendency of an atom or molecule to draw electrons towards itself, form dipoles, and thus form bonds.

van der Waals interactions

A weak force of attraction between electrically neutral molecules that collide with or pass very close to each other. The van der Waals force is caused by temporary attractions between electron-rich regions of one molecule and electron-poor regions of another.

hydrogen bond

The attraction between a partially positively-charged hydrogen atom attached to a highly electronegative atom (such as nitrogen, oxygen, or fluorine) and another nearby electronegative atom.

Ionic and covalent bonds between elements require energy to break. Ionic bonds are not as strong as covalent, which determines their behavior in biological systems. However, not all bonds are ionic or covalent bonds. Weaker bonds can also form between molecules. Two weak bonds that occur frequently are hydrogen bonds and van der Waals interactions.  

Hydrogen bonds between water molecules

Hydrogen bonds between water molecules

The slightly negative oxygen side of the water molecule and the slightly positive hydrogen side of the water molecule are attracted to each other and form a hydrogen bond.

Hydrogen Bonding

Hydrogen bonds provide many of the critical, life-sustaining properties of water and also stabilize the structures of proteins and DNA, the building block of cells. When polar covalent bonds containing hydrogen form, the hydrogen in that bond has a slightly positive charge because hydrogen’s one electron is pulled more strongly toward the other element and away from the hydrogen. Because the hydrogen is slightly positive, it will be attracted to neighboring negative charges. When this happens, an interaction occurs between the δ+of the hydrogen from one molecule and the δ– charge on the more electronegative atoms of another molecule, usually oxygen or nitrogen, or within the same molecule. This interaction is called a hydrogen bond. This type of bond is common and occurs regularly between water molecules. Individual hydrogen bonds are weak and easily broken; however, they occur in very large numbers in water and in organic polymers, creating a major force in combination. Hydrogen bonds are also responsible for zipping together the DNA double helix. 

Applications for Hydrogen Bonds

Hydrogen bonds occur in inorganic molecules, such as water, and organic molecules, such as DNA and proteins. The two complementary strands of DNA are held together by hydrogen bonds between complementary nucleotides (A&T, C&G). Hydrogen bonding in water contributes to its unique properties, including its high boiling point (100 °C) and surface tension.

Water droplets on a leaf

Water droplets on a leaf

The hydrogen bonds formed between water molecules in water droplets are stronger than the other intermolecular forces between the water molecules and the leaf, contributing to high surface tension and distinct water droplets.

In biology, intramolecular hydrogen bonding is partly responsible for the secondary, tertiary, and quaternary structures of proteins and nucleic acids. The hydrogen bonds help the proteins and nucleic acids form and maintain specific shapes.

Van der Waals Interactions

Like hydrogen bonds, van der Waals interactions are weak attractions or interactions between molecules. Van der Waals attractions can occur between any two or more molecules and are dependent on slight fluctuations of the electron densities, which are not always symmetrical around an atom. For these attractions to happen, the molecules need to be very close to one another. These bonds—along with ionic, covalent, and hydrogen bonds—contribute to the three-dimensional structure of proteins that is necessary for their proper function.

Attributions

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  • Atomic Number and Mass Number
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  • Isotopes
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  • The Periodic Table
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  • Electron Shells and the Bohr Model
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  • Electron Orbitals
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  • Chemical Reactions and Molecules
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  • Covalent Bonds and Other Bonds and Interactions
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