1. The History of the Atom

2. Atomic Theory Because we can not see atoms, we use models to teach and learn about atoms. The atomic theory has changed over time as new technologies have become available. –R–Remember: Scientific knowledge builds on past research and experimentation.

3. Ancient Greece Aristotle: –There are four elements: Earth, air, water, fire –The four elements combine in various ways to make all matter. –Matter is continuous. (There is not a “smallest” particle.)

4. Ancient Greece Democritus: –Matter is made of tiny particles called atoms. –The atom is the smallest piece. It is indivisible. (It can’t be divided any further.) –Atoms of an element have specific properties (smooth, spiky, etc.) that give the element its properties.

5. John Dalton- Father of the Atomic Theory Matter is made of indivisible particles called atoms. Atoms of the same element are identical, but differ from atoms of other elements. Atoms cannot be created or destroyed. Atoms of different elements can combine in simple, whole number ratios to form compounds. (The Law of Definite Proportions) Atoms of same element can combine in more than one ratio to form two or more compounds. (The Law of Multiple Proportions) The atom is the smallest unit of matter that can take part in a chemical reaction.

6. John Dalton The “Billiard Ball” model: –An atom is a solid, indivisible sphere

7. J.J. Thompson The “cathode ray tube” experiment Energizing matter in the tube removes charged particles and makes a “cathode ray” (beam of light). Magnets can deflect the cathode ray in a way that shows it is made of negative particles. The amount of deflection tells us how massive the particles are.

8. J.J. Thompson Thomson discovered that: –The atom is NOT “indivisible” –A small negatively-charged particle (the electron) can be removed from the atom. –He concluded that: The atom must have negative particles (electrons) in it (it’s not like a billiard ball). The electrons are evenly distributed amongst positive background “stuff” that makes up most of the atom.

9. J.J. Thompson This led to the “Plum Pudding” model (The electrons are like the plums in a pudding– or the chips in a chocolate chip cookie.) Plum pudding– plum-flavored pudding with plums distributed throughout it– used to be a popular thing people ate. We still call this the Plum Pudding Model even though we don’t eat plum pudding.

10. Ernest Rutherford The Gold Foil Experiment Alpha particles (which are positive) were shot at a piece of gold foil. They were expected to go straight through. Some particles were (unexpectedly) deflected. What deflected them??

11. (Positive) alpha particles are deflected by other positives. (Like charges repel.) The nucleus must be positive! The nucleus must have most of the mass, or it would get pushed around by the alpha particles. The nucleus must be ridiculously small because 10,000 alpha particles pass straight through for each one that is deflected.

12. Ernest Rutherford The Gold Foil Experiment: Conclusions –Atoms have a NUCLEUS It contains all the positive charge It contains almost all the mass It is TINY (1/10000 of the atom volume) –Everything else is mostly empty space

13. Ernest Rutherford The “nuclear model” separates the atom into two parts: –Nucleus: positive & massive & tiny –Electrons: negative, tiny, practically weightless. The Nuclear Model: This picture is not to scale! The nucleus should be WAY smaller than the rest.

14. Neils Bohr Since opposites attract, electrons are attracted to the (positive) nucleus. Something must prevent electrons from “falling into” the nucleus. Electrons are located in orbits around the nucleus. Like planets orbiting the sun, they are attracted, but if they stay at the correct distance, they don’t fall in. The distance from the nucleus determines the energy.

15. Neils Bohr Bohr “quantized” the atom. Only certain energies are allowed, which means only certain orbits are allowed. All other places are “forbidden” to the electrons. Electrons are allowed to be here. Or here! But not here. (This is forbidden because no orbit for the electron to “land on.”)

16. Niels Bohr Moving from one energy level to another requires the electrons to absorb or emit energy Because only certain orbits are allowed… The energy comes in specific colors of light for each element.

17. Albert Einstein The photoelectric effect: –Electrons can be ejected from metal by light (energy) of a certain frequency.

18. Photoelectric Effect If the light isn’t high enough energy, nothing happens. If it does have enough energy, electrons are removed.

19. Albert Einstein The photoelectric effect proved: –Different colors of light are worth different amounts of energy. –Electrons can be moved to different orbits (or out of the atom altogether) by energy in the form of light. –Light has momentum– therefore it is a particle (even though it has no mass). –Light particles are called photons.

20. Thomas Young The double-slit experiment Electrons were sent through two small openings and collected on a screen on the other side. The electrons created interference patterns. Interference patterns come from waves overlapping constructively and destructively. Therefore electrons are a wave (?!)

21. Louis de Broglie Particle-Wave duality –Einstein proved light could act like a particle –Young proved electrons could act like a wave –De Broglie concluded that “wave” and “particle” aren’t mutually exclusive. The electron can act as either one. –Electrons have particle-wave duality.

22. Louis de Broglie –Electrons-as-waves explains the fact that only certain orbits are “allowed” for each element. The orbit’s circumference has to be a multiple of the wavelength in order for it to be “allowed.”

23. Heisenberg: The Uncertainty Principle –Electrons can be moved by light. –We see things because light bounced off of them. –So every time you “see” an electron, the light that made it visible to you probably also caused it to move. –So it isn’t actually there anymore! –It is impossible to know the location and velocity of an electron at the same time.

24. The Uncertainty Principle –Means that electrons cannot be traveling in predictable circular orbits around the nucleus. –They move randomly and unpredictably. –We can estimate the probability of finding an electron somewhere, but we cannot say where it is with certainty. –(Let’s face it: if de Broglie is correct and the electron is partly a wave, it doesn’t necessarily “travel” as a particle anyway.)

25. Erwin Schrödinger Developed an equation to solve for the locations the electron probability is greatest. –This basically describes a cloud-like region where the electron is most likely to be found. –It cannot say with any certainty where the electron actually is at any point in time, but it describes where the electron could be.

26. Erwin Schrödinger The probable locations of the electron predicted by Schrödinger's equation happen to coincide with the locations specified in Bohr's model. The difference is that now everything is fuzzy because of the lack of certainty.

27. Modern Atom: the Quantum model Combining Heisenberg & Schrodinger (with all the previous discoveries) gives us: Orbitals, not orbits Cloud-like regions (fuzzy borders) where the electron probability is highest. Orbitals are 3-D (orbits were 2-D). Quantized energies Electrons can only exist at specific energies allowed by the orbitals. Electrons can absorb or emit energy to move to higher or lower energy orbitals.

28. The current (modern) atomic model: