UNIT 1: STRUCTURE AND PROPERTIES ATOMIC THEORY Development of the Bohr Model of the Atom Early Atomic Theory • 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. • However the theory that matter was composed of 4 classical elements was accepted until the late Middle Ages. Dalton’s Atomic Theory (1803) Dalton revived the atomic theory to explain several chemical laws: • Law of Conservation of Mass (Lavoisier, late 1700’s) • Law of Definite Proportions (Proust, 1799) • Law of Multiple Proportions (Dalton, 1803) This led to the “Billiard Ball” model of indivisible atoms of elements. These combine to form all known compounds. Thomson’s Model of the Atom (1897) • 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. 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. 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… Gold Foil Experiment Expected result: Actual result: 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 • 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. Summary of the Subatomic Particles Particle Location Charge Mass (amu) Proton Nucleus +1 1.0073 Neutron Nucleus 0 1.0087 Electron Outside Nucleus -1 0.00055 amu = atomic mass unit = 1.66 x 10-24 g Problems with the Planetary Model 1) 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. 2) Excited electrons should emit a continuous spectrum of white light when they are excited. Instead, the emission spectrum of elements are all unique. 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. 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. Basic wave terms: c= speed of light = 3.00 x 108 m/s = lambda = wavelength (m) f = frequency (cycles per second = s-1 = hertz) c f short wavelength (blue) long wavelength (red) direction of movement 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. Name of Radiation Radio Waves Microwaves Radar Waves Infrared Light Visible Light * UV Light X-Rays Gamma Rays Colour Red Orange Yellow Green Blue Indigo Violet Wavelength (m) Frequency (c/s or Hz) 102 106 10-7 Increasing Wavelength 10-14 1014 Increasing Frequency 1022 Wavelength (m) 6.5 x 10-7 Frequency (c/s or Hz) 4.6 x 1014 4.1 x 10-7 7.3 x 1014 Note that the wavelength and frequency or inversely proportional: If , then f c f c f 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. The Particle Theory of Light The energy of a photon is some multiple of an energy quantum called Planck’s constant. E hf Where: E = energy of a single photon (kJ) f = frequency (cycles/s, Hz, s-1) h = Planck’s Constant = 6.63 x 10-37 kJs Note: The energy of a mole of photons can be found using: E hfN A Sample Wave and Particle Theory Problems: 1) What is the wavelength of FM 92.5? (92.5 FM has a frequency of 99.1 MHz). 2) What is the energy of 1 photon of electromagnetic radiation emitted by FM 92.5?? 3) What is the energy of 1 mole of these photons? Answers: 3.24 m; 6.13 x 10-29 kJ/photon; 3.69 x 10-5 kJ/mol 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. 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. 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 • Bohr concluded that the permitted energy levels (n) correspond to quantized transitions; only certain energy changes are permitted as determined by the Rydberg formula. CALCULATING TRANSITION ENERGY: The relative energy of any orbit (n) of the hydrogen atom can be calculated by: 1312 En n2 where En is the energy of any orbit (kJ/mol). The transition energy (energy gained or lost as electrons change energy levels) can therefore be calculated. This is equal to the energy emitted by 1 mole of excited hydrogen atoms. Etransition En f Eni Etransition 1312 1312 2 n n2 f i Significance of the Bohr Model • The Bohr model explained why different elements have unique spectra as the value of En depends 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). 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. Sample Problem: Calculate the frequency and wavelength of electromagnetic radiation emitted from a hydrogen atom if an electron moves from the third energy level to the first energy level.