Atoms: The Building Blocks of Stars and Everything Else
Why Stars Shine & The Nature of Light
The lecture aims to explore why stars shine, investigating the nature of light and its production within stars.
Key questions include:
What is the "shine" of stars?
What is light?
What produces light in stars?
How do these things produce light?
How have we learned so much about this?
The Composition of Matter
The fundamental question is: What are things made of?
Scanning electron microscope image (15kV, X100, 100μm) visually introduces the concept.
The Quest for Fundamental Building Blocks
Democritus (~400 BCE): Proposed that all matter consists of small, indivisible, and invisible particles called "atoms" (Greek for indivisible).
Observational Evidence (19th century):
J. Dalton: Substances combine in specific weight ratios to form other substances.
R. Brown: Brownian motion – pollen grains in liquid move erratically under a microscope.
Brownian Motion
Robert Brown (1827): Observed the random motion of pollen grains.
This motion is not caused by living organisms but by random collisions from smaller particles.
A demonstration video illustrates Brownian motion (https://youtu.be/FAdxd_2Iv-UA?t=135).
Molecules: Functional Building Units
Small molecules: Examples include air, water, oil, and most drugs.
Large molecules: Examples include DNAs, proteins, and polyesters.
The concept of “pure” implies that identical molecules are the same.
Atoms: Elements of the World
Late 18th Century Chemistry: Chemists discovered that chemical reactions occur with fixed ratios, and total mass remains constant.
Molecules: Defined as groups of connected atoms.
Chemistry: Focuses on adding, deleting, and rearranging atoms within molecules.
Elements: Different elements possess distinct chemical properties.
Approximately 120 different element types exist in the world, organized in the periodic table.
The Periodic Table
Elements are arranged with different properties and increasing size.
Key elements include Hydrogen (H), Helium (He), Lithium (Li), Beryllium (Be), Sodium (Na), Magnesium (Mg), etc.
Dmitri Mendeleev: He was the first to see that the elements were related in properties when listed in sequence by atomic weight. His arrangement formed the basis of the modern periodic table. Prediction of undiscovered elements and their chemical properties.
The entire Earth is composed of these ~120 elements.
Abundance of Atoms
Atoms are incredibly numerous.
One breath of air (~1 liter) contains approximately 10^{22} molecules.
The Earth's atmosphere holds about 10^{22} liters of air.
Atoms are constantly in flux, moving from one substance to another.
Richard Dawkins quote: “Every time you drink a glass of water, you probably imbibe at least one atom that passed through the bladder of Aristotle…there are many more molecules in a glass of water than are glasses of water on the earth.”
Size of an Atom
A video illustrates the scale of an atom (https://youtu.be/yQP4UJhNn0I).
Estimated number of stars in the universe: ~ 10^{24}.
Estimated number of atoms in a human body: ~ 10^{27}.
Concept: A universe of atoms.
Composition of Atoms
J.J. Thomson (1856-1940): Discovered that cathode rays are lighter than atoms (electrons).
Ernest Rutherford (1871-1937): Student of Thomson; discovered the atomic nucleus through the gold foil experiment where most alpha particles passed through, but some were strongly scattered. Rutherford discovered protons in 1917.
J. Chadwick: (student of Rutherford) discovered neutrons in 1932.
Subatomic Particles: Quarks, Electrons
Electrons are considered fundamental particles with no known internal structure.
Protons and neutrons are thought to be made of “quarks”.
The vast majority of matter consists of up quarks (u), down quarks (d), and electrons (e).
Proton composition: 2 up quarks, 1 down quark.
Neutron composition: 1 up quark, 2 down quarks.
Electrons, Protons, Neutrons: Properties
Electrons have no internal structure.
Masses:
Electron: 9.11 mes 10^{-31} kg
Proton: 1.673 mes 10^{-27} kg
Neutron: 1.675 mes 10^{-27} kg
Protons and neutrons are approximately 1800 times heavier than electrons.
Over 99.9% of an atom's mass is concentrated in its nucleus.
Over 99.999999999999% of an atom's volume is empty space, with electrons moving around in a quantum mechanical manner.
Fundamental Forces Binding Matter
Strong Force: Binds the nucleus (protons and neutrons).
Electromagnetic Force: Binds atoms (electrons to the nucleus).
Weak Force: Involved in radioactive decay.
Gravitational Force: Binds the solar system.
The Four Fundamental Forces
(1) Gravity: Obeys the inverse square law: F_{gravity} = G mes frac{Mm}{r^2}
(2) Electromagnetism: Obeys the inverse square law: F{electric} = K mes frac{q1q_2}{r^2}, also described by Maxwell's equations.
(3) Strong Nuclear Force: Holds protons and neutrons together within the nucleus.
(4) Weak Force: Responsible for radioactive decay.
Electromagnetic Forces
Electrons have negative electric charges.
Protons have positive electric charges.
Like charges repel each other.
Opposite charges attract each other.
Radioactive Decay
Discovery: Radioactivity was discovered by Henri Becquerel in uranium mineral ores.
Marie Curie: Predicted that radiation originates from the atoms themselves, not the molecules, and is unaffected by chemical reactions.
New Elements (1898): Marie Curie discovered radium and polonium in pitchblende and torbernite.
Nobel Prizes: Marie Curie won Nobel Prizes in 1903 and 1911.
Radioactivity results from weak nuclear forces.
Radioisotope Dating
Radioactive atoms have a fixed probability of decay per unit time.
Every half-life, there is a 50% accumulated probability that a radioactive atom has decayed.
Different Clocks for Measuring Different Ages
Not all atoms are suitable for radioisotope dating.
Not all atoms are radioactive.
Different radioactive atoms possess different half-lives.
The half-life of the chosen atom must be appropriate for the age being measured (neither too short nor too long).
Different elements exhibit varying abundances in different objects.
Nuclear Energy
Mass is lost in most nuclear reactions.
The lost mass is converted into energy according to E = \Delta mc^2.
Fusion and fission are nuclear reactions.
Iron (Fe) is the most stable nuclei.
Nuclear Fusion
Small nuclei can fuse into larger nuclei upon collision.
Energy is released if the resultant nucleus is more stable.
Example: 4 protons fuse into 1 Helium nucleus + 2 positrons (occurs in most stars).
Example: Deuterium + Tritium fuse into Helium + neutron (used in hydrogen bombs).
Fusion power has not been achieved due to the difficulty of overcoming the electric repulsion between nuclei.
Nuclear Fission
A heavy nucleus splits into smaller nuclei, releasing energy.
Example: Uranium-235 (^{235}U), when struck by a neutron, splits and creates more neutrons, releasing energy and causing a chain reaction.
Fission is easier to achieve than fusion.
Applications: Atomic bombs, nuclear power plants, etc.
The Dangers of Radiation
Nuclear radiations are high-energy rays of particles and/or photons.
Damage occurs by breaking molecules.
DNA damage: Leads to mutations.
Protein damage: Leads to denaturation.
Radiation diseases can result.
Radiation exposure is often undetectable until it's too late.
Earth's atmosphere provides protection from radiation.
Further Reading
Openstax Astronomy: chapter 5.4
Openstax Physics: chapter 30 (ignore the equations!)