Pre-class 3: chemistry

Matter and States of Matter

  • Matter: anything that takes up space and has mass.
  • States of matter: solid, liquid, gas, or plasma.
  • Everyday examples: anything we touch, the water we drink, the air we breathe.
  • Elements as the basic building blocks of matter.

Elements and Atoms

  • An element is a substance that cannot be broken down into simpler substances by ordinary chemical means.
  • There are 94 naturally occurring elements in the known universe (appendix A reference).
  • Some elements have been artificially constructed by physicists and are typically not biologically important.
  • All matter, including Earth's crust and living organisms, is composed of elements, but life relies on a subset.
  • Six elements are basic to life and make up about 95% of body weight: carbon (C), hydrogen (H), nitrogen (N), oxygen (O), phosphorus (P), and sulfur (S).
  • Other elements important to life include potassium (K), calcium (Ca), iron (Fe), magnesium (Mg), and zinc (Zn).
  • Dalton’s atomic theory (early 1800s): elements consist of tiny particles called atoms; an atom is the smallest unit that displays the properties of an element; atoms of an element share the element’s name.
  • Atomic symbols: one or two letters representing the element name (e.g., H for hydrogen, Na for sodium, Rn for radon).

Subatomic Particles and Nuclear Structure

  • The three best-known subatomic particles: protons (positive charge), neutrons (no charge), and electrons (negative charge).
  • Protons and neutrons reside in the nucleus; electrons move around the nucleus.
  • In simple models, electron locations are shown as shading or shells to indicate probable location; electrons are often described by electron shells or orbitals.
  • Modern physics reveals that most of an atom is empty space; atom-sized analogy: if the atom were the size of a football field, the nucleus would be a small ball at the center, with electrons as tiny specks in the stands around the field.
  • These models indicate where electrons are most likely to be, not exact fixed positions.
  • High-energy experiments (e.g., Large Hadron Collider) reveal complex internal structure beyond the simple models.

Atomic Number, Mass Number, and Isotopes

  • Atomic number (Z): number of protons in the nucleus; determines the element and its chemical properties; for a neutral atom, Z also equals the number of electrons.
  • Mass number (A): total number of protons and neutrons in the nucleus; A = Z + N, where N is the number of neutrons.
  • Protons and neutrons each have an atomic mass unit (AMU); electrons have a negligible AMU in most calculations.
  • Isotopes: atoms of the same element (same Z) that differ in the number of neutrons (N); therefore, they have different mass numbers (A).
  • Example: carbon has naturally occurring isotopes ${}^{12} ext{C}$, ${}^{13} ext{C}$, and ${}^{14} ext{C}$.
  • Mass vs. weight: mass is constant; weight depends on the gravitational field (Earth vs Moon).
  • Atomic mass vs mass number:
    • Mass number A = Z + N (nucleon count).
    • Atomic mass is the average mass of all isotopes of an element, typically expressed in AMU; carbon’s atomic mass is closer to 12 because most carbon is ${}^{12} ext{C}$.
  • Determining neutrons from mass data:
    • If you know A and Z, then the number of neutrons is N=A−Z.N = A - Z.
    • If you are given an atomic mass (approximate), you can estimate neutrons by N ext{ (approx)} \,\≈\, A - Z, and take the closest whole number.
  • Atomic symbol notation on their own (when not on the periodic table): the atomic number is written as a subscript to the lower-left of the symbol, and the mass number as a superscript to the upper-left:
    • This is represented as ZAextX_{Z}^{A} ext{X} where X is the element symbol.
  • Electrons in neutral atoms contribute negligibly to AMU, so electron mass is considered zero in most calculations.

The Periodic Table

  • Periods: horizontal rows.
  • Groups: vertical columns.
  • Across a period, the atomic number Z increases by one from left to right.
  • Atoms within the same group share similar chemical properties and bond-forming behavior.
  • Noble gases (group 18, often referred to as group eight in older schemes): He, Ne, Ar, Kr are inert and rarely react with other atoms.
  • The periodic table organizes elements to reflect recurring chemical and physical characteristics.
  • The periodic table’s structure helps explain trends in reactivity and bonding.

Radioactivity, Isotopes, and Detection

  • Some isotopes are unstable or radioactive; over time they decay into other elements.
  • Example: carbon-14 decays to nitrogen-14, a stable isotope.
  • Radiation emitted during radioactive decay can be detected in various ways.
  • Geiger counter: a common instrument used to detect radiation.
  • Historical context: Becquerel discovered radioactivity in uranium (1896), leading to further study by Marie Curie.
  • Marie Curie discovered polonium and radium and contributed to developing the theory of radiation, which challenged the belief that the atom was indivisible.
  • Nobel Prizes:
    • 1903: Nobel Prize in Physics awarded to Becquerel, Pierre Curie, and Marie Curie for work on radioactivity.
    • 1911: Marie Curie won Nobel Prize in Chemistry, becoming the first person to win two Nobel Prizes of any gender, and the first woman to win a Nobel Prize alone.
  • Curie’s work helped lay groundwork for medical and industrial applications of radioactivity, including radiotherapy and imaging.

Marie Curie, Legacy, and Real-World Applications

  • Curie’s contributions transformed our understanding of atomic structure and radiation.
  • Her work helped establish the practical use of radiation in medicine, industry, and science.
  • She faced significant gender barriers and discrimination, yet made foundational contributions to science and technology.
  • Early radiography in wartime: mobile X-ray units helped treat soldiers in World War I.
  • Modern applications of radioactivity include medical therapies, imaging, archaeology (radiocarbon dating), and energy production, as well as nuclear power and security considerations.

Ethical, Philosophical, and Practical Implications

  • The dual-use nature of radioactivity: benefits in medicine and industry vs potential hazards (radiation exposure, nuclear weapons).
  • The importance of safety, ethics, and regulation in handling radioactive materials.
  • The ongoing need for inclusive and equitable access to science and recognition of contributions by underrepresented groups, as highlighted by Curie’s story.
  • The interplay between fundamental science (atomic theory) and real-world applications (medicine, energy, defense).

Connections to Foundational Principles and Real-World Relevance

  • Atomic theory explains the properties of elements and molecules, guiding chemistry, biology, and materials science.
  • The periodic table reflects recurring chemical properties and periodic trends that underpin predictions of element behavior.
  • Understanding isotopes and radioactivity informs dating methods, medical imaging, cancer therapies, and radiation safety.
  • Ethical considerations around scientific discovery influence policy, education, and societal development.

Key Concepts and Equations (Summary)

  • Atomic number and mass number definitions:
    • Z=extnumberofprotonsinthenucleusZ = ext{number of protons in the nucleus}
    • A=Z+NA = Z + N where N=extnumberofneutronsN = ext{number of neutrons}
  • Isotopes: same Z, different N; different A.
  • Atomic symbol notation in place:
    • ZAextX_{Z}^{A} ext{X}
  • Electron mass is negligible in AMU:
    • me≈0extum_e \approx 0 ext{ u}
  • Atomic mass units for nucleons:
    • Protons and neutrons each have mass ~1extu1 ext{ u}
  • Carbon isotopes as examples: 12extC, 13extC, 14extC^{12} ext{C}, \,^{13} ext{C}, \,^{14} ext{C}
  • Mass vs weight distinction: mass is constant; weight varies with gravity.
  • Model of the atom emphasizes that most of the atom’s mass is in the nucleus and that the nucleus contains protons and neutrons, with electrons orbiting or existing as a probabilistic cloud.
  • Notable historical figures and milestones:
    • Becquerel: discovery of radioactivity (1896).
    • Marie Curie: polonium, radium; two Nobel Prizes (Physics 1903, Chemistry 1911).
    • Applications: medical radiotherapy, imaging, industrial uses, and the development of nuclear energy, with associated ethical implications.

Quick Reference Figures and Terms

  • Figure 2.2: arrangement of subatomic particles in a helium atom; visualizes electrons and nucleus.
  • Figure 2.3: portion of the periodic table (reference to appendix a).
  • Appendix A: full periodic table reference.
  • Terms to know: matter, elements, atoms, subatomic particles, nucleus, electron shells, isotopes, atomic number, mass number, atomic mass unit (AMU), noble gases, radioactivity, half-life (implied by discussion of isotopes), Geiger counter.