Early Atomic Models

1. Early Atomic Models Week 2 Tuesday Unit 1

2. Models of Matter • Scientists have been wondering what matter is made of for over 2500 years. • Empedocles theorised that all substances were made from the combination of four fundamental elements: earth, air, fire, water.

3. Democritus proposed that matter could be broken into smaller and smaller pieces unit a single invisible particle was reached. He called this particle an Atom which means uncutable or invisible in Greek. • Neither ancient scientist did any experimentation when they developed these models.

4. During the mid 1500’s became fashionable, during this time period alchemists constantly tried to convert lead into gold with no success. They did on the other hand develop many methods and tools that are still used today.

5. Dalton’s Atomic Theory In the 1800’s a man by the name of John Dalton came up with a theory based on experiments from various scientists. • Matter consists of definite particles called atoms • Each element is made up of its own type of atom • Atoms of different elements have different properties • Atoms of two or more elements can combine in constant ratios to form new substances. • Atoms cannot be created, destroyed, or subdivided in a chemical change.

6. What theory does this sound like to that we learned about last Thursday? Particle Theory Which past scientist‘s ideas did this theory support? Democritus ‘ 2000 year old model.

7. Subatomic Particles • J.J. Thomson conducted experiments using the cathode ray tube. • He found that the atom is not the smallest particle. There were particles within the atom. • He theorized that an atom was a positively charged sphere with negative charges embedded in it. • It is known as the raisin bun model. The dough would be the positively charged sphere and the raisins would be the negative charges.

8. Subatomic Particles • The negative charges were known as electrons • Electron – Subatomic particle with a negative charge

9. Earnest Rutherford tested this model in his now famous gold foil experiment.

10. He found that the atom contains a positively charged nucleus surrounded by mostly empty space. Some of this empty space contained electrons. • Later he found that the nucleus contained positively charged particles called protons. • Proton- Subatomic particle with a positive charge

11. Rutherford’s model

12. James Chadwick then expanded on this model and suggested that the nucleus also contained neutrons. • Neutron – An uncharged subatomic particle found in the nucleus of an atom.

13. Isotopes • Atoms can contain different numbers of neutrons. These are called Isotopes. • Isotope – An atom of an element that has the same number of protons as the element, but different numbers of neutrons. • Ex Cl 35 and Cl 37 • (more on this later in the unit)

14. Problems • Rutherford’s model had two problems. • Opposite charges attract. • law of moving charges states that as an electron orbits the nucleus, it should emit energy in the form of electromagnetic radiation. As the electron runs out of energy it should then collapse into the nucleus. Matter is stable so this must not be happening.

15. Niels Bohr • Niels Bohr suggested: • electrons were similar to light in the way that they are like a particle but also like a wave. • He used quantum physics to figure this out.

16. He explained it by looking at a phenomenon that had been noticed for years. This phenomenon is known as line spectra.

17. Bohr suggested that electrons orbit the nucleus in fixed orbits of energy. Thus, the electrons are limited to certain energy levels and the energy of the electrons is quantized. • Quantized – Possessing a specific value or amount (quantity)

18. If energy is added and the electron moves up to a higher orbit it is said to be in an excited state. If the electron moves down to a lower orbit it must release the same amount of energy that was required to raise it. Therefore the lines spectra of hydrogen are the result of energy released when the hydrogen atoms electron falls to a lower state and releases energy.

19. The lowest possible orbit for an electron is said to be its ground state. • This Theory was a huge breakthrough but it could only explain the line spectra for hydrogen. • Little white lies I was talking about

20. Bohr also predicted that each energy level could hold a certain number of electrons. He theorised that the first energy level could hold 2 electrons. The second energy level could hold 8 electrons and the third level could hold 18. (think back to your Bohr-Rutherford diagrams from previous years)

21. Bohr's model of the hydrogen atom was only an intermediate step on the way to a precise theory of the atomic structure, which was made possible by quantum mechanics and quantum electrodynamics. (Way beyond the scope of this class)

22. Questions • State the reason why Rutherford’s model of the atom failed to describe the observed behaviour of matter. • Describe Bohr’s model of the atom. How is it different from Rutherford’s model? How is it similar to Rutherford’s model? • Electrons can be found in the ground state or in the excited state. What is different about an electron in each state? • Even though Bohr’s model was an improvement on Rutherford’s model of the atom, there were problems with it. What was a major problem with Bohr’s atomic model?

23. Reminder - There will be no class on Thursday. The Electromagnetic Spectrum

24. The Electromagnetic Spectrum • Electromagnetic energy is commonly known as light energy and is thought to move in the form of a wave. Light waves can differ in their frequency. • Frequency – the number of cycles that pass a particular point in one second.

25. Electromagnetic energy • A wave has a maximum and minimum value called crests and troughs. • The distance between successive crests or successive troughs is known as a wavelength. We normally measure visible wavelengths in nanometers (nm). The symbol we use to represent a wavelength is the Greek letter lamda ( λ).

26. Typical Transverse wave

27. No matter what frequency we are talking about, all light waves travel at the same velocity which is 3.0 x 108 m/s or 1.09 x109 km/h. • Random facts about light • Light only travels at this speed in a vacuum. • Last year, a scientist was able to use a medium slow light down in to a walking pace

28. Frequencies and wavelength • If a wave has a shorter wavelength it will have a higher frequency because more wavelengths will pass a particular point in one second. These waves are higher in energy.

29. Frequencies and wavelength • If a wave has a longer wavelength it will have a lower frequency because fewer wavelengths will pass a particular point in one second. These waves are lower in energy.

30. Energy in waves • High energy frequencies are dangerous to humans. Depending on the intensity and duration of exposure can cause ailments such as cancer or kill in a matter of seconds.

31. electromagnetic spectrum • The electromagnetic spectrum is a continuous sequence of transverse waves that have the same velocity but differ in their frequency and wavelength. • Waves include; radio waves, microwaves, radar, infrared, visible light, ultraviolet, x-rays, and gamma rays. • The "electromagnetic spectrum" of an object is the characteristic distribution of electromagnetic radiation emitted or absorbed by that particular object.

33. Humans and the Electromagnetic Spectrum Humans can detect two areas of the spectrum. The first is infrared light waves. This is detected as heat. The second is detected by the eye and is known as the visible wavelength. • Visible wavelength – • the region of the electromagnetic spectrum that the human eye can detect.

34. Visible Light • Visible light waves are in the range of 400mn to 700nm, each wavelength is seen as a different colour. A rainbow contains the entire visible spectrum and is an example of a continuous spectrum. • Continuous Spectrum – • an uninterrupted pattern of colours that is observed when a narrow beam of white light passes though a prism.

35. Qualitative Analysis • When the electrons in elements are excited they move to higher energy levels and are said to be excited. When the electron moves down to a lower energy level it gives off a photon of light in a particular wavelength. Each element gives off unique line spectra which allow us to identify what element we are dealing with.

36. -Line spectra – a discontinuous spectrum that is produced when light emitted by an element is directed through a prism or a diffraction grating; unique to an element.

37. To observe line spectra a tool called a spectroscope • Spectroscope- An optical instrument that separates light energy into its component wavelengths; used in qualitative analysis.

39. NOW FOR THE LAB • I forgot to add this one to your list • Argon • Class website: • mrhoover.weebly.com

40. Home work • Questions that we did in class about atom models • Read pages 29-35