2.1 The Building Blocks

Lesson Overview

  • Focus: Building Blocks: Atoms, Isotopes, Ions, and Molecules

  • Chemistry is integral to biological processes.

Objectives

  • Define matter and elements.

  • Describe the relationship between protons, neutrons, and electrons.

  • Compare electron donation and sharing between atoms.

  • Explain the combination of elements into molecules and biological structures (cells, tissues, organs).

Basic Concepts

  • Matter: Anything that occupies space and has mass.

  • Elements: Unique forms of matter with specific properties; combinations form matter, including living organisms.

  • Abundant Elements in Organisms:

    • Carbon (C)

    • Hydrogen (H)

    • Nitrogen (N)

    • Oxygen (O)

    • Sulfur (S)

    • Phosphorus (P)

  • Biological processes adhere to the laws of physics and chemistry.

The Building Blocks of Matter

  • Atoms: Smallest units retaining an element's chemical properties.

    • Composed of a nucleus (containing protons and neutrons) and orbiting electrons.

  • Subatomic Particles:

    • Proton:

    • Charge: +1

    • Mass: 1 u

    • Location: nucleus

    • Neutron:

    • Charge: 0

    • Mass: 1 u

    • Location: nucleus

    • Electron:

    • Charge: -1

    • Mass: 0 u

    • Location: orbitals

Atomic Structure

  • Atomic Number: Number of protons in an atom.

  • Mass Number: Sum of protons and neutrons in the nucleus.

  • Isotopes: Atoms with the same number of protons but different numbers of neutrons, leading to different masses.

  • Atomic Mass: Average mass of an element's isotopes.

Isotopes and Radioactivity

  • Stability: Some isotopes are stable; others are unstable and radioactive.

  • Radioactive Decay: Process where unstable isotopes emit energy and particles to become stable, which can transform one element into another.

Carbon Dating

  • Carbon-14 (14C): Radioactive isotope formed in the atmosphere, absorbed by living organisms; decays at a known rate after death.

  • Method: By comparing 14C/12C ratios, scientists can date organic materials up to 50,000 years.

  • Other Methods: For older samples, use other isotopes like 40K and 235U.

The Periodic Table

  • Organization: Displays elements by physical properties and chemical reactivity, indicating their ability to bond.

  • Elements in the same group have similar electron configurations.

Electron Shells and the Bohr Model

  • Bohr Model: Electrons orbit the nucleus at fixed distances, categorized as shells or energy levels.

  • Shells are filled from the innermost to outermost. The stability of an atom is often determined by the outermost shell (valence shell).

  • Octet Rule: Atoms are most stable with 8 electrons in the valence shell.

Electron Configuration and Chemical Reactivity

  • Elements in the same group share the same number of valence electrons:

    • Noble Gases (Group 18): Full outer shells, nonreactive.

    • Alkali Metals (Group 1): Tend to lose one electron readily.

Chemical Bonds and Molecules

  • Chemical Bonds: Form when atoms combine through electron interactions.

    • Types of Bonds:

    • Ionic Bonds: Formed between charged ions.

    • Covalent Bonds: Involve sharing of electron pairs, can be single, double, or triple bonds.

    • Example: Water (H₂O) has two hydrogen atoms bonded to one oxygen atom.

Chemical Reactions

  • Reactions: Occur when bonds are formed or broken, transitioning reactants into products.

  • Balanced Equations: Must satisfy the law of conservation of matter, where each element's number of atoms is conserved.

    • Example Equation: 2H+O→H2O2H + O → H₂O

  • Reaction Types:

    • Irreversible Reactions: Proceed in one direction.

    • Reversible Reactions: Can go in both directions, indicated by a double-headed arrow.

Ionic and Covalent Bonds

  • Ionic Bonds: Between oppositely charged ions, resulting from electron loss or gain.

  • Electrolytes: Essential ions for biological functions.

Covalent Bonds

  • Involve sharing electrons between atoms to satisfy the octet rule.

  • Important for forming organic molecules such as DNA and proteins.

  • Bond Strength: Increases with the number of shared electron pairs.

Polar and Nonpolar Covalent Bonds

  • Polar Covalent Bonds: Electrons are shared unequally, resulting in partial charges. Example: Water molecules have a partial negative charge on oxygen.

  • Nonpolar Covalent Bonds: Equally shared electrons, leading to neutral molecules. Example: Methane (CH₄).

Intermolecular Forces

  • Hydrogen Bonds: Form between polar molecules, essential for water's properties and the stability of proteins and DNA.

  • Van der Waals Interactions: Weak attractions contributing to the structure of biomolecules like proteins.

Summary

  • Matter Composition: Composed of 98 naturally occurring elements.

  • Elements combine into complex structures (molecules, cells, tissues, organs).

  • Bonds prevail through electron interactions including ionic, covalent, hydrogen bonds, and van der Waals interactions.