ATOMIC THEORY “MORPHS”

1. ATOMIC THEORY“MORPHS”

2. Development of an Atomic Theory During Dalton’s time the first empirical evidence (observational evidence) for the existence of atoms was being collected. This evidence was summarized in the Laws of Chemical change as follows: Laws of Chemical Change: 1. Law of Conservation of Matter (Matter is not lost or gained in a chemical Reaction) 2. Law of Definite Composition (Elements combine in definite ratios by mass) 3. Law of Multiple Proportions (Element-masses combine in small whole number ratios)

3. John Dalton, proposed theory for the underlying principles supporting the Laws of Chemical Change must include the following postulates: Dalton’s Atomic Theory 1 All matter is made up of indivisible particles called atoms 2 All atoms of one element are identical. The atoms of any one element are different from those of all other elements. 3 Chemical change is the union or separation of atoms 4 Atoms combine in small whole number ratios to form compounds. Daltton’s model of the atom was therefore a Q-ball model.

4. Thompson’ s Atomic Research Using cathode ray tubes Thompson discovered that rays given off by metals were made of tiny negatively charged particles, he called them electrons. Milliken determined the charge of an electron to be (1-). The atom is no longer considered indivisible however yet invisible. This model of the atom was called the raisin bun model. It had subatomic structure, a positive sphere with negative embedded electrons. This is the raisin-bun model ----->

5. Rutherford Performs the Gold Foil Experiment. This provided evidence that the atom is mainly empty space with all the positive charge and mass being located in the center known as the nucleus. This is a diagram of Rutherford’s experiment

6. This is what actually happened and why the model was changed from the “raisin bun” model to, an atom with a nucleus.

7. These spectra came from excited gaseous samples of hydrogen.

8. BOHR'S MODEL OF THE HYDROGEN ATOM: • According to Bohr's theory, the electron can orbit the proton (nucleus) only in certain "ENERGY LEVELS" or principal quantum levels, numbered n=1, n=2, n=3, etc., with n=1 being the normal or GROUND STATE level for the electron closest to the nucleus. • When the electron is excited up to the n=2 or higher level, a photon of light is absorbed, giving an absorption spectrum. Since the energy levels have a definite spacing, this results in photons of a specific colour or wavelength being absorbed. • When the electrons "fall" from a higher to lower levels, light is emitted, giving an emission spectrum.

9. The BALMER SERIES is the visible spectrum series of emission lines for the hydrogen atom. These lines are the result of electron transitions from: [n = 3 to shell 2] red,[n = 4 to n = 2] blue green, [n = 5 to n = 2]blue, [n = 6 to n = 2] violet, all ending in n = 2 energy level. These transitions can be modeled as follows. Bohr’s model of the atom was therefore one with a positive nucleus and “fixed energy” orbits for electrons.

10. Visible Spectrum • Brilliant colours that are seen in an emission spectrum are: • Atomic: When light is emitted by atoms of an element in the gaseous phase each element has unique colour. • The excited state environment that is generated when either heat or electric energy causes electrons to jump to higher energy levels in atoms. • When these “excited” electrons fall back to a lower energy level they emit light of particular colours or wavelengths of ‘fixed energy”. • Bohr’s model of the atom was therefore one with a positive nucleus and “fixed energy” orbits for electrons.

11. Presented ByGator-chem ATOMIC THEORY “MORPH” Terry Knock 2000 Central Kings Rural High School Thanks to Jan Gough, Horton High for her valuable contributions.