1. Structure & Properties of Matter
Lesson # 1: Early Atomic Theories and Origins of Quantum Theory

2. Early Theories
Democritus – all matter is composed of tiny particles called atoms
Lavoisier – recorded quantitative information about chemical reactions
Dalton – all matter is composed of atoms which cannot be created or destroyed; atoms of the same element have the same mass, size and other properties.

3. The Electron
Discovered by Thompson. He applied high voltage to a partially evacuated tube with a metal electrode at each end.
A ray of energy (now we know it was electrons) was produced that started from the negative electrode, called the cathode. This is now called a cathode ray tube. The negative end repelled the ray, and the positive attracted it.

4. The Electron & Cathode Ray Tube
He tested this with various metals and concluded that all atoms must contain electrons, but that they also must possess a positive charge so that they are neutral. He thought that the electrons were randomly embedded within the atom.

5. Radioactivity & Atomic Models
In the late 19th century, it was found that certain elements emitted high levels of energy. Becquerel noticed that uranium produced an image on a photography plate, which means it was emitting radiation. Any atom that is found to do this was considered radioactive.
Today radioactivity is defined as the spontaneous decay of the nucleus of atoms, which was first proposed by Rutherford.
There are three types of radiation – alpha, beta, and gamma (gamma being the strongest).

6. Types of Radiation

7. Rutherford’s Model
Rutherford also wanted to look at the model of the atom – he decided to aim positive alpha particles at a thin sheet of gold foil, expecting most rays to pass through with little deflection. Instead he found some rays reflected, some deflected, although most did pass through, meaning that the electrons could not simply be randomly scattered throughout the atom.
He reasoned that atoms must be composed of mostly empty space for so many to pass through, and that any deflections would be caused by coming near or striking a positively charged Any that bounced straight back would have directly hit a positive charge, which he called protons.

8. Gold Foil Experiment

9. Rutherford’s Model
He reasoned the protons were gathered in the center of the atom, which we now call the nucleus, with electrons moving around the atom with much space between them.
James Chadwick added to Rutherford’s experiment by noticing that the masses of the protons did not equal the mass of the atoms, so he reasoned there must be uncharged particles in the nucleus as well – the neutrons.

10. Isotopes
Two atoms that contain the same number of protons but a different number of neutrons are called isotopes. Example: 12C, 13C, 14C
Unstable isotopes are called radioisotopes, as they emit radiation as the nucleus decays to a more stable form, releasing gamma rays and subatomic particles.
Radioisotopes are useful in carbon dating, nuclear energy, and even medicine.

11. Isotopes of Hydrogen

12. Classical Light Theory
Light is electromagnetic radiation.
Early scientists thought light was a stream of particles, and later they discovered light showed wave-like properties
Maxwell, in the 19th century, said that light interacts with particles because it existed as an electromagnetic wave inside of magnetic and electric fields. This was called the classical theory of light.

13. Electromagnetic Spectrum

14. Classical Light Theory
Up until then, light and matter were not studied together. Light had no mass or specific position in space, as matter did, but when Hertz studied waves using induction coils, he discovered that shining a light on a metal surface caused the emission of electrons from the metal. This was called the photoelectric effect.

15. Classical Light Theory
According to classical theory, by increasing the intensity (brightness) of the light, the kinetic energy of the electrons emitted should increase, however Hertz found that it was the frequency of the light that determined emission.

16. Developing Quantum Theory
Planck also studied light and emission spectra of solid objects, which he called blackbodies. When heated, they started to glow – first red, then white, then blue, all depending on the object.
He plotted the intensity on a graph, and like Hertz, thought that as he heat increasing, the emission should increase, but instead he found that it peaked and then decreased over and over through each change in colour even though the temperature increased constantly.

17. Developing Quantum Theory
Planck said that matter must be able to gain or lose energy, only in whole-number multiples according to the equation: E = nhf
n = any whole number
f = frequency of radiation
ħ = Planck’s constant = 6.63 x J.s.
Planck also knew that radiation was emitted as atoms vibrated back and froth, and he thought that atoms must emit energy in bursts, or specific quantities, instead of in constant streams. One burst is now known as a quantum of energy.
This led to the discovery that light and matter are related.

18. Developing Quantum Theory
Einstein suggested that electromagnetic radiation could be viewed as a stream of particles called photons. A photon is a unit of light energy.
He said that electrons are emitted from the surface of a metal because a photon had collided with it, and some of the energy from the photon transferred to it, allowing it to break away from the atom.
It needs to have the right quantity of energy for this to happen, or the electron will remain in place, despite the amount of photons aimed directly at the metal. Again – this is linked to the frequency of light.
All of these discoveries have led to the quantum theory we now use to describe light and matter.

19. Atomic Spectra
Spectroscopy is the scientific study of spectra in order to determine properties of the source of its spectra.
Spectrometers measure the intensity of light at different wavelengths – light passes through a sample, it is dispersed by a prism or diffraction grating, and it forms a spectrum. A detector calculates the amount of light absorbed or transmitted at each wavelength.
Since all atoms interact differently with light, using standard line spectra can help identify unknown compounds.

20. Basic Spectrometer

21. Emission Spectra
When a high-energy spark is applied to hydrogen, H2, the molecules absorb energy and the H-H bond breaks. The resulting atoms are called “excited” as they contain excess energy. They release this energy by emitting light at various wavelengths. We can observe this light through a spectrometer, and we call it an emission spectrum.
There are two types of emission spectra – a continuous spectrum – where all colours in that electromagnetic spectrum are produced, and line spectrum – where only particular wavelengths are visible

22. Emission Spectrum
A continuous spectrum is like when light is appears from a prism, whereas a line spectrum is what is observed when electrons emit energy. Each element has its own unique line spectrum.
This applies to atomic theory because it demonstrated that the electron in hydrogen can exist only at discrete energy levels based on the fact that only single wavelengths of colour are produced – the energy is quantized.

23. Determining Wavelength
c = λ•f
c = speed of light in a vacuum (2.998x108m.s-1)
λ = wavelength (in m)
f = frequency (in s-1)
Example: Calculate the wavelength of infrared radiation that has a frequency of 9.73 x 1013 s-1.

24. Determining Energy E = ħf OR E = ħc/ λ E = energy (in J)
f = frequency of radiation (in s-1)
ħ = Planck’s constant = 6.63 x J.s.
c = speed of light in a vacuum (2.998x108m.s-1)
λ = wavelength (in m)

25. Example
Calculate the energy of one photon of red light with a wavelength of 669 nm.

26. Bohr’s Model of the Atom
A limitation to Rutherford’s model was that he said electrons were in constant motion in orbits. Physics tells us that changes in direction mean acceleration, which would cause it to emit radiation, lose energy, and spin in towards the nucleus. We know that this does not happen!
Bohr used the knowledge of emission spectra and the hydrogen atom to develop a new model of the atom. He said that electrons exist in particular orbitals of increasing energy levels around the nucleus.

27. Bohr’s Model of the Atom
When an atom gained more energy it could move into an orbit farther from the nucleus. This is called the atoms “excited state”. They release that energy as light as they move back to their original energy level, called the “ground state”

28. Limitations to Bohr’s Model
Bohr’s model also assigned a maximum number of electrons to each orbital based on his study of emission spectra. 2 in the first shell, 8 in the second, 18 in the third, etc (2n2, with n being the orbital number).
This only works with the first 20 elements when drawing Bohr-Rutherford diagrams, but then it gets a bit too complicated to draw them (hence you never had to do beyond calcium in grade 9 and 10).
Another limitation to Bohr’s theory is that emission spectra for elements higher than hydrogen do not exactly match line spectra as Bohr would have assumed for what he saw based on hydrogen, as these atoms have more than one electron. There must have been something missing in his theory!