Ch 2: Chemical property and chemistry

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  • Syllabus quiz due by 11:59pm on January 21

  • Quiz 1 (covering Chapter 1 and 2) opens January 17, at 10pm and is due January 21 by 11:59pm

  • Exam 1 available online from February 1 to February 2

Chapter 2 - The Chemical Context of Life

Importance of Chemistry in Biology

  • Chemistry fundamental to biology due to the hierarchy of life

  • All organisms composed of one or more cells, which arise from pre-existing cells

  • Understanding cell components essential for grasping their functions and structures

  • NOTE: A background in chemistry is not required

Biological Organization

  • Life can be studied at different levels, including:

    • Atom:

      • Biosphere

      • Hydrosphere

      • Atmosphere

      • Lithosphere

    • Molecule

    • Ecosystem

    • Cell

    • Tissue

    • Organism

    • Population

    • Community

    • Organ System

Emergent Properties

  • The concept that the whole is greater than the sum of its parts

Order and Organization of Life

  • Chemistry fits into life at various levels:

    • Atoms

    • Molecules

    • Cells

    • Tissues (e.g., epithelium)

Energetic Tendencies

  • Essential rules regarding chemical processes:

    • Positive and negative charges tend to balance

    • Electrons typically occur in pairs

    • Electrons fill shells based on energetic preferences

The Significance of Atomic Structure

  • Atoms make up everything in the universe

  • Matter comprises atoms and elements:

    • Element: a fundamental substance consisting of only one type of atom.

    • 92 naturally occurring elements; Hydrogen, Carbon, Oxygen, and Nitrogen are most abundant in life

Atomic Structure

  • An atom consists of three fundamental particles:

    • Protons: positively charged

    • Neutrons: uncharged

    • Electrons: negatively charged

Position of Particles

  • Protons and neutrons located in the nucleus

  • Rapidly moving electrons orbit the nucleus in electron shells

Atomic Number and Elements

  • The atomic number (proton count) defines an element

  • Removing a proton changes the element

  • The periodic table organizes elements by atomic number

Isotopes

  • Atoms can have varied neutrons; these are called isotopes

  • Example:

    • Carbon Isotopes:

      • Carbon-12: 6 protons, 6 neutrons

      • Carbon-14: 6 protons, 8 neutrons

  • Isotopes behave similarly chemically but differ in mass

Isotope Types and Stability

  • Heavy isotopes have more neutrons than the most prevalent forms

  • Example: Hydrogen isotopes: Protium (0 neutrons), Deuterium (1 neutron), Tritium (2 neutrons)

  • Some isotopes might be unstable and can decay, emitting energy

Isotope Decay

  • Decay alters isotopes, often transforming one element into another, like Carbon-14 to Nitrogen-14

Applications of Isotopes

  • Isotope decay can emit radiation utilized in nuclear medicine for diagnostics (e.g., PET scans)

  • Radiation can also pose risks, such as DNA damage from nuclear accidents

Electron Configurations

  • Electrons can lose, gain, or share to form bonds

  • Each shell has defined atomic orbitals:

    • 1st shell: 2 electrons

    • Other shells: up to 8

Valence Electrons

  • Elements in the same group (columns) share similar valence electron counts and hence have similar properties

  • Valence Electrons by Group:

    • Group 1: 1

    • Group 2: 2

    • Group 13: 3

    • Group 14: 4

    • Group 15: 5

    • Group 16: 6

    • Group 17: 7

    • Group 18: 8

Stability of Atoms

  • Full outer orbitals enhance atom stability; atoms typically strive for 8 electrons (octet rule)

    • Helium is an exception, stable with 2 electrons

Noble Gases

  • Noble gases, located in the last column, have full valence shells, resulting in low reactivity

Electronegativity

  • Electronegativity defines an atom’s capability of attracting electrons

  • Influenced by:

    • Electrons in the outer shell

    • Proximity of electrons to the nucleus

Key Electronegative Elements

  • Highly electronegative elements include: Fluorine (F), Nitrogen (N), Oxygen (O)

  • High electronegativity facilitates hydrogen bonding

Formation of Molecules

  • Chemical bonds form through electrical attractions between atoms creating molecules and compounds

  • A compound encompasses at least two different elements

Covalent Bonds

  • Covalent bonds arise when two atoms share pairs of electrons

  • Bond strength can determine whether reactions will proceed

Representations of Covalent Bonds

  • Visual representations include:

    • Electron-shell diagrams

    • Structural formulas

    • Space-filling models

    • Ball-and-stick models

Polar vs. Nonpolar Bonds

  • Unequal sharing leads to polar covalent bonds; electrons are attracted more to the electronegative atom

  • Polar molecules like water exhibit dipole characteristics

Water Molecule Properties

  • Water (H2O) has polar covalent bonds due to the electronegativity of oxygen

  • This distributes charges, resulting in regions of high and low electron density

Ionic Bonds

  • Ionic bonds form when an electronegative atom steers electrons from a less electronegative atom

  • Resulting ions (cations and anions) exhibit electrostatic attractions

Strength of Ionic Bonds

  • Ionic compounds (like salt) form strong crystal arrangements but may weaken in water due to hydration effects

Dissolution of Ionic Compounds in Water

  • Water stabilizes ions, enabling ionic bond breakdown,

  • Polar properties of water facilitate the dissolution of ionic compounds

Dipole-Dipole Interactions

  • Weak interactions occur between slightly charged atoms due to temporary charge separation

  • Includes significant biological interactions: hydrogen bonds and van der Waals forces

Hydrogen Bonds

  • Result from attraction between hydrogen atoms bonded to electronegative atoms (O, N, F) and lone electronegative atoms

Van der Waals Interactions

  • Very weak attractions arise from transient dipoles due to electron movement

Relative Bond Strengths

  • General strength ranking: Ionic > Covalent > Hydrogen > Van der Waals

Energy in Chemical Reactions

  • Energy is crucial for bond formation and breaking; it's conserved and transforms during reactions

  • Plants convert sunlight into chemical energy via photosynthesis

Overview of Chemical Reactions

  • Reactants are transformed into products through bond breaking and forming, requiring activation energy

  • Reactant bonds must be broken for the reaction to proceed

Types of Reactions

  • Hydrolysis Reaction: Consumes water to break complex molecules

  • Condensation Reaction: Combines molecules and releases water

Activation Energy

  • Activation energy (Ea) is necessary to initiate reactions; influences reaction rates

Reaction Rates

  • The speed of reactions can be affected by:

    • Activation energy (higher Ea results in lower rate)

    • Temperature (higher temperatures increase rates)

    • Concentration (higher concentrations yield more collisions)

Reversibility and Equilibrium

  • Reactions can be reversible with one direction more favorable, reaching a state of equilibrium

Properties of Water and its Importance

  • Living organisms require liquid water, which consists of two hydrogen atoms and one oxygen atom

  • Water is critical for facilitating biological chemical reactions in various environments

Characteristics of Hydrogen Bonds in Water

  • Hydrogen bonds within liquid water constantly form and break

Water’s Polarity and Its Effects

  • Water's unique properties critical to life stem from its polar nature

    • Lower density in solid state (ice)

    • Cohesion and adhesion capabilities

    • High specific heat and vaporization

    • Acts as a solvent for life

    • Undergoes self-ionization

Ice Density

  • Ice has lower density than water due to a stable crystal structure; this insulates aquatic ecosystems

Cohesion and Adhesion

  • Cohesion leads to surface tension; adhesion promotes attraction to other polar substances

Surface Tension in Water

  • High surface tension allows water to form droplets and supports small entities on its surface

Specific Heat and Heat of Vaporization

  • Water has a high specific heat and heat of vaporization, buffer against temperature changes for organisms

Evaporative Cooling

  • As water evaporates, it removes heat, assisting in temperature regulation (e.g., sweating)

Water as the Universal Solvent

  • Water's ability to dissolve substances makes it a key resource in biological systems

    • Hydrophilic: substances dissolve in water

    • Hydrophobic: substances do not dissolve in water

Self-Ionization of Water

  • Water can ionize spontaneously, forming H+ and OH- ions; equal concentrations indicate neutrality, giving water a pH of 7

Acids and Bases

  • Acids increase H+ concentrations, lowering pH; bases decrease H+ concentrations, raising pH

Buffer Systems

  • Buffers stabilize pH levels, essential for homeostasis in biological systems

    • E.g., human blood maintains pH between 7.35 and 7.45

Next Steps

  • The chemistry discussed forms the basis for understanding biological molecule formation, structure, and function.