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Unit 2

Unit 2. Atomic Structure. Atomic Theory. Greek Model. “To understand the very large, we must understand the very small.”. Greek philosopher Idea of ‘democracy’ Idea of ‘ atomos ’ Atomos = ‘ indivisible ’ ‘Atom’ is derived

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Unit 2

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  1. Unit 2 Atomic Structure

  2. Atomic Theory

  3. Greek Model “To understand the very large, we must understand the very small.” • Greek philosopher • Idea of ‘democracy’ • Idea of ‘atomos’ • Atomos = ‘indivisible’ • ‘Atom’ is derived • No experiments to support idea, his ideas were forgotten for thousands of years. Democritus Democritus’s model of atom No protons, electrons, or neutrons Solid and INDESTRUCTABLE

  4. Alchemy • After that, chemistry was ruled by alchemy. • They believed that that could take any cheap metals and turn them into gold. • Alchemists were almost like magicians. • Elixirs: physical immortality • Philosopher’s stone: change lead to gold

  5. GOLD SILVER COPPER IRON SAND . . . . . . . . . . . . . . . Alchemy Alchemical symbols for substances… transmutation: changing one substance into another D In ordinary chemistry, we cannot transmute elements.

  6. Contributions of alchemists: • Information about elements • - the elements mercury, sulfur, and antimony were discovered • - properties of some elements • Develop lab apparatus / procedures / experimental • techniques • - alchemists learned how to prepare acids. • - developed several alloys • - new glassware

  7. The Atomic Theory of Matter • In 1803, Dalton proposed that elements consist of individual particles called atoms. • His atomic theory ofmatter contains four hypotheses: 1. All matter is composed of tiny indivisible particles called atoms. 2. All atoms of an element are identical in mass and chemical properties. 3. A chemical compound is a substance that always contains the same atoms in the same ratio. 4. In chemical reactions, atoms from one or more compounds or elements rearrange in to form one or more new compounds. Atoms themselves do not change of identity in chemical reactions. Some parts of his theory were later proven wrong.

  8. Modern atomic theory Subatomic particles • Electrons: • Negatively charged subatomic particles • Discovered by J.J. Thomson in 1897 • Discovered by observing deflection of cathode rays. • Occupy most of the volume of an atom • YouTube - Thomson

  9. Electron Microscopy • The electron microscope is a type of microscope that uses a beam of electrons to create an image of the specimen. It is capable of much higher magnifications and has a greater resolving power than a light microscope, allowing it to see much smaller objects in finer detail. • Typical light microscope magnifies up to 1000 x and an electron microscope magnifies over 100,000 x • SEM - Image Gallery

  10. Protons : positively charge particles • Neutrons: no charge particles, mass similar to protons. • The atomic nucleus: • Protons and neutrons are located in the center of the atom • Through the gold-foil experiment, Rutherford determined: http://www.mhhe.com/physsci/chemistry/essentialchemistry/flash/ruther14.swf • The atom is mostly empty space • Positive charge and most of atom’s mass is concentrated in a small region, called the nucleus (composed of protons and neutrons)

  11. Distinguishing Among Atoms • Elements are different because they contain different number of protons. • Atomic number: indicates the number of protons in the nucleus of an element. • Since atoms are electronically neutral: # protons= # electrons Atomic number: Carbon has 6 protons Carbon has 6 electrons 6 C

  12. Learning Check: complete the following table

  13. Mass number • Most of the mass of an atom is concentrated in its nucleus and depends on the number of protons and neutrons. • Mass number: total number of protons and neutrons in an atom. • Mass number= protons + neutrons • # neutrons= mass number – protons • Representing atoms: Mass number Mass number 197 Au gold-197 Atomic number 79

  14. Learning check: determine the number of protons, neutrons and electrons for the following atoms. • Carbon-12 • Fluorine-19 • Beryllium-9

  15. Isotopes • Isotopes are atoms of the same element that have the same number of protons, but different number of neutrons. • Isotopes of an element have the same atomic number but different number of neutrons, thus have different mass numbers. • Hydrogen has 3 isotopes: • Hydrogen-1 or simply hydrogen • Hydrogen-2 or deuterium • Hydrogen-3 or tritium

  16. Classwork: isotope notation handout

  17. Atomic Mass • Mass of proton or neutron: 1.67x10-24g • Mass of electron : 9.11x10-28g • These values are impractical to work with, so scientists compare relative masses of atoms using a reference isotope : carbon-12 • An atomic mass unit (amu) is defined as 1/12 the mass of carbon-12

  18. Atomic mass (continued) • The atomic mass of an element is not a whole number because the isotopes of an element and its natural abundance is taken in consideration. • The atomic mass of an element is a weighted average mass of the atoms in a naturally occurring sample of the element.

  19. Atomic mass (continued) • Ex. 1 The atomic mass of copper is 63.546 amu. Which of copper’s two isotopes is more abundant: copper-63 or copper-65? • Since 63.546 is closer to 63 than 65, the most abundant isotope is copper-63.

  20. Atomic mass (continued) • Ex. 3 There are 3 isotopes of silicon; they have mass numbers of 28, 29, and 30. The atomic mass of silicon is 28.086 amu. Comment on the relative abundance of these 3 isotopes. • Silicon-28 must be the most abundant.

  21. Atomic mass (continued) 1. Divide the percentages by 100: 0.5069 and 0.4931 • Ex. 2 Calculate the atomic mass of bromine. The two isotopes of bromine have atomic masses and relative abundance of 78.92 amu (50.69%) and 80.92 amu (49.31%) atomic mass= (0.5069x78.92) + (0.4931x80.92) =79.9062 = 79.91 amu

  22. Atomic mass (continued) • Classwork p103 # 12,13; p 105 #33-34

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