1 / 27

UNIT 1: Structure and properties Atomic Theory

UNIT 1: Structure and properties Atomic Theory. Development of the Bohr Model of the Atom. Early Atomic Theory Can you keep chopping an apple forever?.

snana
Download Presentation

UNIT 1: Structure and properties Atomic Theory

An Image/Link below is provided (as is) to download presentation Download Policy: Content on the Website is provided to you AS IS for your information and personal use and may not be sold / licensed / shared on other websites without getting consent from its author. Content is provided to you AS IS for your information and personal use only. Download presentation by click this link. While downloading, if for some reason you are not able to download a presentation, the publisher may have deleted the file from their server. During download, if you can't get a presentation, the file might be deleted by the publisher.

E N D

Presentation Transcript

  1. UNIT 1: Structure and propertiesAtomic Theory Development of the Bohr Model of the Atom

  2. Early Atomic TheoryCan you keep chopping an apple forever? • Ancient Greeks (e.g. Democritis) proposed that when matter is divided into smaller and smaller pieces, a finite limit known as the atom is ultimately reached. • Democritis put forward the ideas: i. this tiny indivisible grain, or atom, only comes in several different kinds. ii. by re-arranging these microscopic “Lego bricks” in different ways, we can make everything that exists in the universe. *note: Democritis performed no experiments!

  3. Boyle (1660s)Is matter blob-like or like “angry bees”? • 1st experimental evidence that matter is grainy • Made a container that could change its volume. -the smaller the container is made, the more the gas would push against the walls of the container with pressure. Why? What is pressure? • the gas is made of tiny grains, buzzing around like “angry bees”. Pressure on the container wall is caused by the “angry bees” bumping into it. If you decrease the volume of the container, the “angry bees” will have less space in which to travel and will hit the container wall more often…..in other words there will be more pressure. • -this observation can only be explained by matter being like tiny grains.

  4. Lavoisier (1789)Do atoms come in different “flavours”? • provided the first evidence that the tiny grains of matter come in different kinds • compiled a list of 23 “elements”. • after his death the number climbed to more than 50. • today we know of 92 natural elements. So if these tiny grains are the building blocks of matter, why are there so many different kinds of them??

  5. Prout (1815)Is there something smaller than the atom? • proposed the idea that the tiny grains (atoms) are not the ultimate building blocks of matter but instead made up of smaller things. • he weighed all the elements and found that most come in exact multiples of the weight of hydrogen, the lights element • -a pattern is the repetition of a simpler, more basic building block. For example, the number 87878787878787878787878787 – the simpler building block from which it is made is the number 87. • -similarly, Prout found in the weights of elements something was going on – they must possess a deeper structure and be made of even more simpler building blocks. • -Prout believed nature’s building block is the hydrogen atom, and all the other elements are simply different numbers of hydrogen atoms glued together.

  6. Thomson’s Model of the Atom (1897)What are atoms made of? • Michael Faraday and Svante Arrhenius studied electricity and batteries, suggesting that electric charges are part of matter. • Thomson detected and measured the mass of a beam of negative particles in a vacuum tube (or cathode ray tube, CRT) . • He called the tiny particles that emerged from a metal cathode electrons.

  7. Thomson’s Plum Pudding Model • Thomson explained electrical conduction in metals by the movement of subatomic electrons in the solid. • Electrical conduction in solutions (e,g, batteries) was explained by the existence of charged atoms called ions.

  8. Rutherford’s Atomic Model • The discovery of radioactivity (,  and  rays) allowed Ernest Rutherford to probe inside the atom. • Based on Thomson’s model, Rutherford proposed that alpha rays (high energy helium ions, He2+) should pass through the positive pudding in a very thin sheet of gold foil. • The results were slightly different…

  9. Gold Foil Experiment Expected result: Actual result:

  10. Rutherford’s Planetary Model (1911) • Rutherford concluded that the atom consists of a very tiny, dense nucleus composed of the positive charge. • >99.99% of the atom consists of empty space. • Tiny electrons are found orbiting the nucleus. • A teaspoon on the atomic nuclei would weigh six billion tonnes (6 x 1012 kg)!!!! 900 x

  11. In 1914, Rutherford proposed that there existed a positively charged particle called a proton. • In 1932, Chadwick & Rutherford found that the nucleus also contained a neutral particle called a neutron.

  12. Summary of the Subatomic Particles amu = atomic mass unit = 1.66 x 10-24 g

  13. Problems with the Planetary Model • According to classical physics, electrons orbiting the nucleus should lose energy and emit light. This loss of energy would cause the electrons to spiral into the nucleus, resulting in the collapse of the atom. • Excited electrons should emit a continuous spectrum of white light when they are excited. Instead, the emission spectrum of elements are all unique.

  14. Emission and Absorption Spectra • Neils Bohr observed the line spectra produced by the excitation of gas state elements. • Excited gases produce a unique emission spectrum. • Cold gases will absorb produce an absorption spectrum.

  15. The Wave Theory of Light Light is a form of electromagnetic radiation. All electromagnetic radiation is made up of electric and magnetic fields. These fields oscillate in a wave pattern as electromagnetic radiation moves through space. Electromagnetic radiation travels at a constant speed in a vacuum, but the wavelength and frequency varies.

  16. Basic wave terms: c= speed of light = 3.00 x 108 m/s = lambda = wavelength (m) f = frequency (cycles per second = s-1 = hertz) short wavelength (blue) long wavelength (red) direction of movement

  17. The Spectrum of Electromagnetic Radiation Note: 1 nanometer (nm) = 1 x 10-9 m The “visible” part of the spectrum is between 390 nm and 750 nm.

  18. Note that the wavelength and frequency or inversely proportional: If , then f

  19. Planck’s Quantum Hypothesis (1900) • Max Planck realized that the spectrum produced by white light produced was not continuous. • Planck proposed that light was composed of small packets or quanta of energy called photons. • Different colours of light are composed photons with different “quantums” of energy. • The energy of a particular photon is proportional to the frequency of the radiation.

  20. The Particle Theory of Light The energy of a photon is some multiple of an energy quantum called Planck’s constant. Where: E = energy of a single photon (kJ) f = frequency (cycles/s, Hz, s-1) h = Planck’s Constant = 6.63 x 10-37kJs Note: The energy of a mole of photons can be found using:

  21. Sample Wave and Particle Theory Problems: • What is the wavelength of FM 99.1? (99.1 FM has a frequency of 99.1 MHz). • What is the energy of 1 photon of electromagnetic radiation emitted by FM 99.1?? • What is the energy of 1 mole of these photons?

  22. Bohr’s Model of the Atom (1913) Bohr developed a model of the atom which explained the line spectrum of hydrogen and why the atom doesn't collapse. His theory is made up of two postulates: Postulate 1: Electrons can only move in certain fixed orbits. Each orbit corresponds to a specific energy level and an electron can move within an orbit without losing any energy. Postulate 2: An electron can only move from one orbit (or energy level) to another when it gains or loses energy.

  23. The Hydrogen Spectrum Explained In a discharge tube, the electrons become temporarily excited and thus move to orbits that are farther from the nucleus. These transitions are only temporary, though, so the electrons return to their lowest (or ground) states. When they do so, they give off energy.

  24. Why Only Four Visible Lines? Since electrons can only undergo transitions between certain specific energy levels, only certain quantities (quantums) of energy are given off. Since the colour of light corresponds to the energy of possessed by a quanta (or photon), only certain coloured lines are observed. Animation of the Bohr Model of the Atom (Emission) Absorption in the Bohr Model Another Animation

  25. Bohr concluded that the permitted energy levels (n) correspond to quantized transitions; only certain energy changes are permitted as determined by the Rydberg formula.

  26. Significance of the Bohr Model • The Bohr model explained why different elements have unique spectra as the value of Endepends on the charge of the nucleus. • Bohr’s model also explains Mendeleev’s Periodic Law as the chemical and behaviour of elements is related to the filling of Bohr’s energy levels (2, 8, 8,18 etc).

  27. Problems with Bohr’s Model • Calculations work for 1 electron systems (H, He+, Li2+). • Calculations for other atoms do not agree with the experimental results. • The visible spectral lines of hydrogen can be split by electric or magnetic fields fields. • In conclusion, the Bohr model was a great leap in the understanding of the atom. It was the first step in quantum mechanics, the current atomic theory.

More Related