1. WHAT’S A THEORY?

2. Atomic Theory

3. The Ancient Greeks Democritus and other Ancient Greeks were the first to describe the atom around 400 B.C. The atom was nature’s basic particle that makes up all matter. The word atom comes from the Greek word “atomos” meaning indivisible.

4. Aristotle The Greek philosopher Aristotle did not believe in atoms, and instead believed all matter was continuous. Aristotle’s belief was accepted for ~2000 years, as there was no evidence to support either theory.

5. Antoine Lavoisier First to state the law of conservation of mass, which states that mass is neither created nor destroyed during chemical or physical changes.

6. Joseph Louis Proust Law of Definite Proportions = chemical compounds contain the same elements in the same proportions by mass regardless of the source or the amount. i.e. Sodium Chloride (table salt) is always 39.34% Na and 60.66% Cl by mass.

7. “The BIG 3” 3 Major Theories on the ATOM –Dalton –Thompson –Rutherford

8. John Dalton’s Atomic Theory 1)All matter is composed of extremely small particles called atoms. 2)Atoms of an element are identical to each other and different from atoms of other elements.

9. 3) Atoms cannot be created or destroyed. 4) Atoms of different elements combine in simple whole-number ratios to form chemical compounds. 5) Atoms are combined, separated or rearranged when chemical reactions occur.

10. Impact of Dalton’s Atomic Theory By relating atoms to mass, Dalton turns Democritus’ ideas into a scientific theory which could be tested. Dalton’s model of the atom is a small, solid, and indivisible particle. (known as the “ball-bearing” model)

11. Flaws in Dalton’s Theory Not all parts of Dalton’s Theory have proven to be correct. Today we know that… 1.Atoms are divisible into even smaller particles 2.Atoms of the same element can have different masses.

12. Cathode-Ray Tube Experiments Electric current was passed through gases at low pressures. It was observed that the end of the tube near the cathode glowed and a ray traveled from there to the anode when current was passed through the tube.

13. In these experiments, it was also observed that the rays were deflected away from magnetic fields and negatively charged objects.

14. J.J. Thomson Carries out experiments that support the hypothesis that cathode rays are made of negatively charged particles. Measured the charge to mass ratio of cathode ray particles and determined that it was the same regardless of the metal or gas used. Thomson concludes that all cathode rays are composed of identical negatively charged sub-atomic particles, which he calls electrons.

15. Thomson’s Model of the Atom Thomson’s experiments show that (1) atoms are divisible and (2) they include negatively charged electrons. Thomson’s Plum-Pudding Model of the Atom has negative particles spread evenly throughout a solid, positively charged sphere.

16. Robert Millikan Conducted an Oil Drop Experiment that measured the charge of the electron.Oil Drop Experiment Using this charge and Thomson’s charge to mass ratio, the mass of the electron is also discovered. Mass of 1 electron = 9.109 x10 -31 kg

17. Ernest Rutherford Conducted the Gold- Foil Experiment Positively charged alpha particles were fired at a thin piece of gold foil. It was expected that the alpha particles would pass through with only a slight deflection.

18. However, it was discovered that about 1 out of every 8,000 particles were actually deflected back toward the source. Rutherford concludes that there must be a very small and densely packed region of positive charge in the atom that would repel these alpha particles. Rutherford is credited with discovering the nucleus.

19. The Nuclear Model of the Atom Nucleus = very small and dense center of the atom, where protons and neutrons are found. Protons have a positive charge. Neutrons are electrically neutral. Negatively charged electrons orbit the nucleus like the planets orbit the sun. Most of the atom is empty space.

20. Subatomic Particles ProtonNeutronElectron Charge+0— Mass*1 amu ~ 0 LocationNucleus Electron Cloud *Mass of protons and neutrons = 1.67 x10 -27 kg compared to the electron’s mass of 9.109 x10 -31 kg

21. The Electromagnetic Spectrum

22. λ (lambda) = wavelength –The distance between 2 corresponding points on adjacent waves. ν (nu) = frequency –The number of waves that pass a point during a period of time. –Units = Hertz (Hz) = 1 wave/second

23. All forms of EM radiation move at a constant speed of c = 3.00 x10 8 m/s. The relationship between frequency and wavelength can be expressed mathematically: c = λ ν

24. Before 1900, scientists believed that light behaved as a wave. However, the Photoelectric Effect could not be explained by this theory. In this experiment, light shines on a metal causing electrons to be emitted from the metal. However, if the frequency of light is below a certain minimum, it was observed that no electrons were emitted. Scientists could not explain why there was a minimum frequency of light that would cause the Photoelectric Effect to occur.

25. Max Planck Proposes that objects can only emit energy in small and specific amounts. A quantum = minimum amount of energy that can be gained or lost by an atom. Planck discovers the following relationship: E = energy (measured in Joules) h = Planck's constant = 6.626 x10 -34 J*s ν = frequency E = h ν

26. Albert Einstein In 1905, Einstein expands on Planck’s ideas with his theory on the dual wave-particle nature of light. Light behaves like both a wave and a stream of particles. These particles (called photons) each carry a quantum of energy and have zero mass.

27. Einstein explains the photoelectric effect… –EM radiation is only absorbed in whole #s of photons. –Electrons are only emitted from a metal when struck by a photon with a minimum amount of energy.

28. Electromagnetic Radiation c = λ ν E = hν c h λ E =

29. 1)Calculate the frequency of a light wave with a wavelength of 515 nm. 2)Calculate the frequency of a wave with an Energy of 3.75 x10 -6 J. 3) A near infrared wave has a wavelength of 1.00 x10 -6 meters. Calculate the energy of this wave.

30. Flame Test Why do different elements emit a different color of light when exposed to a flame?

31. Line Spectra When the light from the flame is passed through a prism a spectrum is formed:

32. A spectrum that contains only certain colors (wavelengths) is called a line spectrum. When different elements are vaporized in the flame, they each produce a unique spectrum: this is a kind of “atomic fingerprint” that can be useful in identifying elements. But why does this occur? And where does this energy come from?

33. The Bohr Model Niels Bohr combines the ideas of: –Rutherford (planetary model with a nucleus and circling electrons) –Planck (quantized amounts of energy) Bohr concludes that the energy of the electrons in an atom has to be quantized. This means that electrons can only have certain amounts of energy.

34. Each electron in an atom can only have a certain orbit / energy level (n). The lowest energy level (n=1) includes the orbit closest to the nucleus. This is referred to as the ground state. The excited states represent orbits farther from the nucleus that have more energy. When an electron absorbs the right amount of energy it jumps to a higher energy level and enters an excited state (n=2, n=3, n=4).

35. An electron in the ground state is closest to the nucleus (n=1). When it absorbs energy, it enters an excited state and moves to a higher energy level.

36. Bohr’s Model of the H atom When an electron in the excited state falls back to the ground state, energy is emitted in the form of radiation. This emission of radiation produces the line spectra observed for hydrogen. Bohr uses his model and Planck’s equation to calculate the frequencies for the line spectrum of hydrogen.

37. Limitations of the Bohr Model Bohr’s model is able to explain the line spectrum for hydrogen, but could not explain the spectra of more complex atoms that have more electrons.