Transcript Chapter_2x
Chemistry Chapter 3 Atoms: The Building Blocks of Matter Atom who? • Atom – The smallest particle of an element that retains the chemical properties of that element Law of Conservation of Mass Mass is neither created nor destroyed during chemical or physical reactions. Total mass of reactants = Antoine Lavoisier Total mass of products Dalton’s Atomic Theory (1808) All matter is composed of John Dalton extremely small particles called atoms Atoms of a given element are identical in size, mass, and other properties; atoms of different elements differ in size, mass, and other properties Atoms cannot be subdivided, created, or destroyed Atoms of different elements combine in simple whole-number ratios to form chemical compounds In chemical reactions, atoms are combined, separated, or rearranged Modern Atomic Theory Several changes have been made to Dalton’s theory. Dalton said: Atoms of a given element are identical in size, mass, and other properties; atoms of different elements differ in size, mass, and other properties Modern theory states: Atoms of an element have a characteristic average mass which is unique to that element. Modern Atomic Theory Dalton said: Atoms cannot be subdivided, created, or destroyed Modern theory states: Atoms cannot be subdivided, created, or destroyed in ordinary chemical reactions. However, these changes CAN occur in nuclear reactions! Discovery of the Electron In 1897, J.J. Thomson used a cathode ray tube to deduce the presence of a negatively charged particle. Cathode ray tubes pass electricity through a gas that is contained at a very low pressure. Some Modern Cathode Ray Tubes Mass of the Electron 1909 – Robert Millikan determines the mass of the electron. The oil drop apparatus Mass of the electron is 9.109 x 10-31 kg and the charge is 1.67 x 10-19 C Millikan's Oil Drop Experiment Conclusions from the Study of the Electron Electrons are negative. Cathode rays have identical properties regardless of the element used to produce them. All elements must contain identically charged electrons. Atoms are neutral, so there must be positive particles in the atom to balance the negative charge of the electrons Electrons have so little mass that atoms must contain other particles that account for most of the mass Rutherford’s Gold Foil Experiment Alpha particles are positively charged Particles were fired at a thin sheet of gold foil Particle hits on the detecting screen (film) are recorded Gold Foil Experiment Rutherford’s Findings Most of the particles passed right through A few particles were deflected VERY FEW were greatly deflected Conclusions: The nucleus is small The nucleus is dense The nucleus is positively charged The Structure of the Atom • Atoms consist of two regions – Nucleus • Very small region in the center. • Contains protons & neutrons. – Electrons Cloud • Mainly empty space. • Very large compared to the nucleus. • Contains electrons. • Subatomic particles – Protons, neutrons, and electrons Atomic Particles Particle Charge Mass (kg) Electron -1 9.109 x 10-31 Proton Neutron +1 1.673 x 10-27 Location Electron cloud Nucleus 0 1.675 x 10-27 Nucleus Atomic Number Atomic number (Z) of an element is the number of protons in the nucleus of each atom of that element. Identifies the atom. Element Carbon Phosphorus Gold # of protons 6 15 Atomic # (Z) 6 15 79 79 Mass Number Mass number is the number of protons and neutrons in the nucleus of an isotope. Mass # = p+ + n0 Nuclide p+ n0 e- Mass # Oxygen - 18 8 10 8 18 Arsenic - 75 33 42 33 75 Phosphorus - 31 15 16 15 31 Isotopes Elements occur in nature as mixtures of isotopes. Isotopes are atoms of the same element that differ in the number of neutrons Isotopes…Again (must be on the test) Isotopes are atoms of the same element having different masses due to varying numbers of neutrons. Isotope Protons Electrons Neutrons Hydrogen–1 (protium) 1 1 0 Hydrogen-2 (deuterium) 1 1 1 Hydrogen-3 (tritium) 1 1 2 Nucleus Atomic Masses Atomic mass is the average of all the naturally isotopes of that element. On Periodic Table Carbon = 12.0125 amu Isotope Symbol nucleus % in nature Carbon12 12C 98.89% Carbon13 13C Carbon14 14C 6 protons 6 neutrons 6 protons 7 neutrons 6 protons 8 neutrons 1.11% <0.01% Writing Nuclear Symbols He Mass # 3 Atomic # 2 (proton + neutrons) (proton) Atomic Symbol How many protons, electrons, and neutrons? 2 protons, 2 electrons, 1 neutron Mass # - Atomic # = # Neutrons Writing Isotopes Using Hyphen Notation Uranium-235, Helium-3, or Carbon-14 Name Mass # of atom How many proton, electrons, neutrons? 92 protons, 143 neutrons, 92 electrons Isotope problems Convert these hyphen notation to nuclear symbols. Uranium-235, Helium-3, or Carbon-14 U 235 3 92 2 He 14 6 C Chapter 22 – Nuclear Chemistry The Nucleus • Contains nucleons – Protons & Neutrons • Nuclear forces – Short-range proton-neutron, proton-proton, and neutronneutron forces hold the nuclear particles together. Nuclear Stability • Kinetic Stability – Describes the probability that a nucleus will decompose (radioactive decay) Nuclear Stability Decay will occur in such a way as to return a nucleus to the band (line) of stability. Number of Stable Nuclides Related to Numbers of Protons and Neurons Types of Radioactive Decay alpha production (a): helium nucleus • 238 4 234 U He Th 92 2 90 • 0 beta production (b): 1 e 234 234 90Th 91Pa 0 1 e 4 2+ 2 He Alpha Radiation Limited to VERY large nucleii. Beta Radiation Converts a neutron into a proton. Types of Radioactive Decay gamma ray production (g): • 238 4 U 92 2 He 234 90Th 0 positron production 1 e : • 22 0 22 11 Na 1e 2 00 g 10 Ne electron capture: (inner-orbital electron is captured by the nucleus) 201 80 Hg 0 201 e 1 79 Au 00 g Types of Radiation Half-Life of Nuclear Decay The Decay of a 10.0-g Sample of Strontium-90 Over Time QUESTION In the following nuclear equation, identify the missing product: 43 1 Ca + a __________ + H 1 20 ANSWER 2) 46 Sc 21 Section 18.1 Nuclear Stability and Radioactive Decay (p. 841) Make sure to memorize the abbreviations for the subatomic particles. QUESTION Identify the missing particle in the following equation: 238 4 U He + ? 92 2 ANSWER 2) 234 90 Th Section 18.1 Nuclear Stability and Radioactive Decay (p. 841) Just as chemical equations need the same number of each type of atom on each side, nuclear equations need the same number of each type of nucleon on each side. QUESTION The nuclide 12 7 N is unstable. What type of radioactive decay would be expected? ANSWER 2) b + Section 18.1 Nuclear Stability and Radioactive Decay (p. 841) According to the band of stability graph (Figure 18.1) this nuclide is neutron-poor, so it must do something to decrease the number of protons or increase the number of neutrons. QUESTION Nuclides with too many neutrons to be in the band of stability are most likely to decay by what mode? ANSWER 5) b – Section 18.1 Nuclear Stability and Radioactive Decay (p. 841) This process is the opposite of positron emission and allows the change of a neutron into a proton. QUESTION A radioactive element has a half-life of 1.0 hour. How many hours will it take for the number of atoms present to decay to 1/16th of the initial value? ANSWER 3) 4 Section 18.2 The Kinetics of Radioactive Decay (p. 846) 1 ½ ¼ 1/8 1/16. Each arrow indicates a half-life of 1.0 hour.