02 ATOMIC THEORY AND ATOMIC STRUCTURE

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Chapter 02 02 Atomic Theory And Atomic Structure

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Atom – came from the word “atomos” which means small, indivisible particles Democritus 440 B.C. 2.3 Development of Atomic Theory Aristotle

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2.3 Development of Atomic Theory Antoine Lavoisier 1774 Law of Conservation of Mass – mass can neither be created nor destroyed in chemical reactions 4.55 g + 2.02 g = 6.57 g 3.25 g + 3.32 g = 6.57 g

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2.3 Development of Atomic Theory Joseph Proust 1799 Law of Constant Composition – Different samples of a pure chemical substance always contain the same proportion of elements by mass. By mass, water is: 88.8 % oxygen 11.2 % hydrogen

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2.3 Development of Atomic Theory John Dalton 1808 Dalton’s Atomic Theory All matter consists of tiny particles called atoms. An atom cannot be created, divided, destroyed, or converted to any other type of atom. Atoms of a particular element have identical properties Atoms of different elements have different properties Atoms of different elements combine in simple whole-number ratios to produce compounds. Chemical change involves joining, separating, or rearranging atoms. Law of Multiple Proportions - Elements can combine in different ways to form different substances, whose mass ratios are small whole-number multiples of each other.

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2.3 Development of Atomic Theory Benjamin Franklin 1750 Proposed the presence of positive and negative charges

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2.3 Development of Atomic Theory William Crookes 1875 Connected two metal electrodes (metal discs connected to a source of electricity) at opposite ends of a sealed glass vacuum tube. When electricity is turned on, rays of light were observed to travel between the two electrodes. The stream of light were called cathode rays, because they traveled from the cathode (negative electrode) to the anode (positive electrode)

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2.3 Development of Atomic Theory John Joseph Thomson 1882 - Demonstrated the electrical and magnetic properties of cathode rays

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2.3 Development of Atomic Theory

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2.3 Development of Atomic Theory John Joseph Thomson 1882 Cathode rays are streams of negative particles of energy. He called it electrons.

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2.3 Development of Atomic Theory Eugene Goldstein 1886 besides the cathode rays there is another stream that travels in the opposite direction as the electron flow. Attracted to negatively charged electric plates. He called them protons.

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2.3 Development of Atomic Theory John Joseph Thomson 1897 Proposed the “Raisin Bread Model of the Atom”

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2.3 Development of Atomic Theory Ernest Rutherford 1901 Demonstrated that atoms spontaneously “decay” to produce 3 types of radiation: alpha particles (a) beta particles (b) gamma particles (g). This phenomena is known as natural radioactivity.

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2.3 Development of Atomic Theory Ernest Rutherford 1910 Along with his students, Ernest Madsen and Hans Geiger, performed the “Alpha-particle Scattering Experiment”

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2.3 Development of Atomic Theory Proposed the Nuclear Model of the Atom Ernest Rutherford 1910

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Electromagnetic Radiation (EMR) - Is energy that travels as waves through space. visible light is a kind of EMR. travel at a speed of 3.0 x 108 m/s Is described in terms of wavelength (λ) and frequency (ν). Wavelength and Energy are inversely proportional 2.4 The Relationship between Light and Atomic Structure

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2.4 The Relationship between Light and Atomic Structure Emission Spectrum - Lines of different colors formed when light from a heated element passes through a prism.

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2.5 The Bohr Atom - He hypothesized that electrons can only occupy certain fixed energy levels. Each level is defined by a spherical orbit around the nucleus, located at a specific distance. Concept of certain fixed energy levels is called quantiz ation. Energy levels are also called Q uantum Levels Neils Bohr 1914

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2.5 The Bohr Atom

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2.5 The Bohr Atom Bohr Model of the Atom Q uantum Numbers - a way to identify the position or energy level of an electron.

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2.3 Development of Atomic Theory demonstrated the existence of the neutron. James Chadwick 1932

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2.6 Modern Atomic Theory Due to the observed photoelectric effect, light, in certain situations, possessed particle like properties. Albert Einstein 1905

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2.6 Modern Atomic Theory Louie de Broglie, proposed that light can also have wave like properties. Louie de Broglie 1924

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2.6 Modern Atomic Theory -1926, Erwin Schrödinger proposed the quantum mechanical model of the atom which focuses on the wavelike properties of -the electron. - stated that it is impossible to know precisely where an electron is and what path it follows—a statement called the Heisenberg uncertainty principle. Modern Atomic Model Werner HeisenBerg 1927

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Atom – is the smallest unit of an element that retains the chemical properties of that element. Electron (e-) Proton (p+) Neutron (n0) Neutrinos, gluons, quarks Distinct Regions of the Atom The Composition and Structure of the Atom 2.2 Composition of the Atom

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The mass of the atom is primarily in the nucleus The charge of the proton is opposite in sign but equal to that of the electron 2.2 Composition of the Atom

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Chemical Nuclear Symbol X = Chemical Symbol of the Element A = Mass Number Z = Atomic Number X = C A = 12 Z = 6 2.2 Composition of the Atom

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Chemical Nuclear Symbol X = Chemical Symbol of the Element A = Mass Number = Number of Protons (p+) + Number of Neutrons (n0) Z = Atomic Number = Number of Protons (p+) = Number of Electrons (e-) C = Charge = Number of Protons (p+) - Number of Electrons (e-) X = Chemical Symbol = C A = Mass Number = 12 Z = Atomic Number = 6 Protons = Atomic Number = 6 Neutrons = Mass number – Protons = 12 – 6 = 6 Electrons = Protons – (Charge) = 6 – (0) = 6 2.2 Composition of the Atom

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Isotopes Isotopes – are atoms of the same element (same number of protons) having different atomic masses because they contain different number of neutrons. 2.2 Composition of the Atom

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Ions Ion – an atom or molecule which has lost or gained one or more valence electrons, giving it a positive or negative electrical charge. Anion - a negatively charged ion, which has more electrons in its electron shells than it has protons in its nuclei. Protons = Atomic Number = 17 Neutrons = Mass number – Protons = 35 -17 = 18 Electrons = Protons – (Charge) = 17 – (-1) = 18 Cation - a positively-charged ion, which has fewer electrons than protons. Protons = Atomic number = 11 Neutrons = Mass number – Protons = 23 – 11 = 12 Electrons = Protons – (Charge) = 11 – (+1) = 10 2.2 Composition of the Atom