History of Atomic Theory

1. History of Atomic Theory Atomic models from Dalton to Bohr

2. A ‘Model’… • is not a real thing, but is used to explain, mimic or simulate reality, • is used as a tool, • is used to predict what happens in the real world, • is changed or modified until it best fits new information, • may have some limitations or be valid only under certain conditions. Examples: globes, computer simulations, product prototypes

3. Historical Models of the Atom John Dalton ‘Billiard ball’ model (1803) • All matter consists of atoms • Each element has its own atom type • Atoms of different elements have different properties • Atoms of two elements can combine to form compounds • Atoms are never created, destroyed or subdivided

4. Historical Models of the Atom J. J. Thomson ‘Raisin bun’ model (1897) • First to include sub-atomic particles (electrons) that had been seen in cathode ray tube experiments • Model is of a positively charged sphere with negatively charged electrons embedded in it • Positive ‘dough’ and negative ‘raisins’ make up an atom that is neutral over all

5. Historical Models of the Atom Ernest Rutherford Nuclear (‘beehive’) model (1911) • Tested Thomson’s theory with the famous “gold foil experiment” • His results suggested that the atom was mostly empty space with a very dense positively-charged ‘nucleus;’ he later discovered that protons were the positively-charged part • In this model, negatively-charged electrons existed within the empty space **Neutrons were discovered much later by James Chadwick (in 1932); why so late?

6. Historical Models of the Atom ? Niels Bohr Refined nuclear model (1913) • Bohr knew that a new model was needed, primarily because of a major problem with the Rutherford model: Bohr and his contemporaries knew that if a charged particle accelerates, it must give off energy, likely in the form of light. The electrons, which are definitely charged (-) and accelerating (changing direction constantly) should therefore give off energy and eventually spiral in towards the nucleus and cause the atom to collapse. Problem? Yes: Atoms don’t collapse! • Bohr was very interested in the newly-developed quantum theory of light proposed by Einstein and Planck, and thought it could be applied to the problem with Rutherford’s model…so… more about quantum theory next, then back to Bohr…

7. The Quantum Model Bohr’s Inspiration for a Better Model of the Atom…but still not the best…

8. Light defined as a Wave • Light travels through space as an electromagnetic wave. Waves are characterized by their wavelength, λ, and frequency, f, and amplitude, A. A • Light colour is related to the wavelength (and frequency) of a wave. Red light has a longer wavelength and lower frequency than blue light.

9. Relationship Between Wavelength and Colour of Light A spectrum containing all colours of visible light is called a continuous spectrum. This is what we see if we pass white light through a prism.

10. The Quantum Model Quantum – a specific allowable value. Quanta – a set of specific allowable values. Example – A staircase is like a set of quanta. Each stair is an allowable position. Other than when travelling between steps, an individual step is the only place you may exist. Compare to a ramp.

11. Origins of Quantum Theory • Max Planck first hypothesized that the energy of an oscillating atom was not continuous (or wavelike) when he studied blackbody radiation • Albert Einstein stated that if the energy of the vibrating atoms was quantized, the light they emit must also be quantized • Einstein earned a Nobel Prize when he used this new ‘Quantum Theory of Light’ to explain the photoelectric effect; one quantum of light (called a ‘photon’) could release one electron from a metal surface; higher-energy photons were more likely to liberate electrons than low-energy photons

12. Light Defined as a Particle(quantum model) • Light is also thought to propagate through space as individual particles called photons. • Each photon has a specific amount of energy that is related to the wave characteristics and to the colour of light, that is, blue photons have more energy than red photons

13. Atoms and Light • An element in a gaseous state produces light when it is heated to a certain temperature • By passing this light through a prism, we can see its ‘bright-line’ or ‘emission’ spectrum

14. Bohr wanted his atomic model to explain the bright-line spectra of the elements • Since only certain distinct colours of light could be absorbed or emitted by atoms, Bohr reasoned that this related to distinct ‘energies’ of the electrons inside the atom Conclusion: Electrons have distinct energies, and are therefore ‘quantized’

15. And now we can finish the story… Niels Bohr Refined nuclear model (1913) a.k.a.“Bohr-Rutherford Model” • Nucleus containing protons (+) ((and neutrons)), • Electrons (-) are organized into specific energy levels orbiting the nucleus, and are thereby ‘quantized.’ • Bohr’s model allows the electrons to exist in specific allowable energy levels that are identified by the principle quantum number, n. The allowable values of n are 1,2,3, … • Electrons follow “occupancy rules” Although technically ‘historical,’ this model is very useful. It is still used daily by students and scientists alike.

16. Bohr-Rutherford Model Electron occupancy rules: 2n2 n = energy level * Electrons will always occupy the lowest energy level available. *

17. What Bohr proposed… • Electrons are arranged in fixed energy states. • When an element is heated, electrons are promoted to higher energy states (excited states). When electrons return to a lower energy state, energy is given off in the form of light (a photon is emitted). • The movement of an electron between energy levels is referred to as a transition. • The type (colour) of light emitted is related to the size of the transition. Many transitions produce photons that are not in the visible region. • It is important to note that Bohr primarily studied hydrogen.

18. Electron transitions of Hydrogen

19. Usefulness of the Bohr-Rutherford Model • Periodic Trends • Valence electrons = group # (A groups) • Common ion charges (A groups) • Ionization energy • Stability of Noble Gases and trends with successive ionization energies. • Elements in group 1 have an unusually high 2nd I.E ; those in group 2 have an unusually high 3rd I.E. etc.. The B-R model explains this by suggesting that once an element has achieved an octet, it is in a stable arrangement that matches a noble gas. • Atomic radii • Reactivity

20. Usefulness of the Bohr-Rutherford Model • Stability of Noble Gases; trends with successive ionization energies

21. Usefulness of the Bohr-Rutherford Model 2. Predictions For Compounds • Bonding ratios (MgF2, CaF2, SrF2, etc.) • Bond polarity (electronegativity) • Bonding types – covalent vs. ionic 3. Physical Properties • Solubility • Melting and boiling points • Viscosity