Back ©Bires, 2002 Slide 1 2.2 – The Discovery of Atomic Structure AP Chemistry Summer Work Chapter 2 Anyone who says that they can contemplate quantum.

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Presentation transcript:

Back ©Bires, 2002 Slide – The Discovery of Atomic Structure AP Chemistry Summer Work Chapter 2 Anyone who says that they can contemplate quantum mechanics without becoming dizzy has not understood the concept in the least. -Niels Bohr

Back ©Bires, 2002 Slide 2 Atomic History The Greek philosopher Democritus (400BC) coined the term atomon which means “that which cannot be divided.” Idea of an indivisible thing that made up all matter was refined by colorblind chemist John Dalton in 1803.

Back ©Bires, 2002 Slide 3 Dalton’s Atomic Theory 1.All matter is made of indestructible and indivisible atoms. 2.Atoms of a given element have identical physical and chemical properties. (all atoms of X will behave the same anywhere) 3.Different atoms have different properties. (X behaves differently than Y) 4.Atoms combine in whole-number ratios to form compounds. (two H’s and one O = Water (H 2 O) 5.Atoms cannot be divided, created or destroyed in chemical reactions.

Back ©Bires, 2002 Slide 4 Dalton’s Atom

Back ©Bires, 2002 Slide 5 The Laws: Constant Composition: Ratios of atoms in a compound is constant for that compound. Conservation of Mass: Mass is not created or destroyed in a chemical reaction. Multiple Proportions: Since atoms bond in small, whole number ratios to form compounds, their masses are small whole number ratios.

Back ©Bires, 2002 Slide 6 J. J. Thomson Dalton had no experimental evidence for his theory, so scientists began working to provide evidence for the existence of atoms. J. J. Thomson set up a cathode ray tube and began using magnets to study the charge that was being emitted in the tube. The beam bent when placed next to a magnet, so the atom must have negative charges, which Thomson called electrons.

Back ©Bires, 2002 Slide 7 Beam is attracted to positive magnet and repelled by a negative magnet

Back ©Bires, 2002 Slide 8 Thomson’s Plum Pudding Model Thomson’s negative charges in a sea of positive charge

Back ©Bires, 2002 Slide 9 Millikan’s Oil Drop Experiment We now know the existence of the electron, but what is it’s charge? Millikan’s Oil Drop Experiment was engineered to precisely tell us the charge of a single electron. Thomson’s experiment had determined the charge to mass ratio (1.76 x 10 8 C/g), but the exact charge was still unknown.

Back ©Bires, 2002 Slide 10 Millikan’s Experiment By varying the voltage and the amount of electrons, he was able to determine the specific charge of the electrons

Back ©Bires, 2002 Slide 11

Back ©Bires, 2002 Slide 12 Millikan’s Discovery Millikan’s classic oil-drop experiment allowed the charge of a single electron to be determined: 1.60 x C. Knowing the charge and the charge to mass ratio from Thomson, the mass could be determined:

Back ©Bires, 2002 Slide 13 Ernest Rutherford While studying radioactive elements, New Zealander Physicist Ernest Rutherford found that radioactive alpha particles deflected when fired at a very thin gold foil. This was known as the gold foil experiment, and it suggested that the atom was not a hard sphere as thought, but was mostly space, with a small concentration of mass. This concentration of mass became known as the nucleus. Link to experiment…

Back ©Bires, 2002 Slide 14 The Bohr Model Niels Bohr, (a student of Rutherford) rebuilt the model of the atom placing the electrons in energy levels. Bohr was one of the founders of quantum physics – a discipline that states that energy can be given off in small packets or quanta of specific size. More of this theory and the Bohr model to come in future chapters!

Back ©Bires, 2002 Slide 15 Adding the Neutrons James Chadwick, proved the existence of massive, neutral particles. These particles came to be called neutrons. This completed our model of the atom that still remains the accepted model (with more modifications to come!).

Back ©Bires, 2002 Slide 16 The Modern Model (not to scale) Chadwick’s neutrons Rutherford’s space Democritus and Dalton’s atom Bohr’s energy levels Thompson’s electrons