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Ch. 4 Atomic Structure

Ch. 4 Atomic Structure. Early Models of the Atom. Democritus’s Atomic Philosophy He believed atoms were indivisible and indestructible. * Philosophy because there is no experimental support. Dalton’s Atomic Theory. John Dalton (1766-1844) Theory because he had experimental support.

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Ch. 4 Atomic Structure

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  1. Ch. 4 Atomic Structure

  2. Early Models of the Atom • Democritus’s Atomic Philosophy • He believed atoms were indivisible and indestructible. • *Philosophy because there is no experimental support

  3. Dalton’s Atomic Theory • John Dalton (1766-1844) • Theory because he had experimental support

  4. Dalton’s vs. Modern Atomic Theory • Elements are composed of tiny, indivisible atoms. • Atoms of same element are identical • Different elements can mix or react to make compounds. • When chemical reactions occur, elements rearrange but no new elements. • Elements are made of atoms BUT atoms are divisible. • Atoms are not absolutely identical. Isotopes – different mass but same properties. • Same as Dalton’s • Same as Dalton’s

  5. Sizing up the atom* • The radii of most atoms are 5x10-11 to 2x10-10 m • A pure copper coin the size of a penny contains about 2.4x1022 atoms

  6. 4.2 Structure of the Atom

  7. Location of Particles • Nucleus (positive charge) • Protons and Neutrons • Electron Cloud (negative charge) • electrons

  8. Discovery!! Electrons! • Electrons were discovered by J.J. Thomson through the use of a cathode ray tube. • One end was positive, the other end negative.

  9. The particles that were sent through the tube we pulled towards the positively charged plate. • Because of this, Thomson concluded that the beam was negatively charged since opposite charges attract

  10. Plum pudding model • Thomson called these particles electrons. He knew they were a part of the atom but wasn’t sure where they went. • Decided electrons were mixed in with the other particles

  11. Video • https://www.youtube.com/watch?v=IdTxGJjA4Jw

  12. Gold Foil Experiment • In 1909 Ernest Rutherford designed an experiment to study the parts of the atom • Rutherford aimed a beam of small positively charged particles at a thin sheet of gold foil.

  13. He put a coating behind the foil that glowed when hit with particles.The coating allowed Rutherford to see where the particles went after going through the foil.

  14. What did this show? • If the atom had particles all mixed together, the particles should have gone straight through • In 1911 he revised the atomic theory.He proposed the center of the atom is a tiny, dense positively charged area called a nucleus (an atom's central region, which is made of protons and neutrons) and electrons were surrounding it

  15. Because like charges repel each other, Rutherford reasoned that positively charged particles that passed close to the nucleus were pushed away. • This is why some of the particles were deflected.

  16. He was able to calculate the diameter of a nucleus to be 100,000 times smaller than the atom itself. • Imagine a pinhead in AT&T park • This means that atoms are mostly empty space!

  17. ~Videos~ • https://youtube.com/watch?v=XBqHkraf8iE • https://youtube.com/watch?v=wzALbzTdnc8

  18. 4.3 Distinguishing among atoms • Atomic number: The number of protons in the nucleus of an atom. It is the same for all atoms of that element. • This determines what the element is. • Example: Carbon has an atomic number of 6, and always has 6 protons

  19. Mass number:This is the combined number of protons and neutrons in the nucleus • May be written as X-##, where X is the element abbreviation and ## is the mass • Carbon with 6 protons and 6 neutrons would be written as C-12

  20. •Atomic Number: # of p (in neutral atom also # of e- •Mass Number: # of p + # of n •To find # of n = Mass # - Atomic # • How many p, n, e- does Fe have?

  21. Isotope:different types of atoms of the same element. They have the same number of protons but a different number of neutrons • Same number of proton=same element

  22. Difference between isotopes • Mass number shows the difference between isotopes (number of neutrons differs) • Remember, the number of protons will always stay the same for each element.

  23. Carbon-12 C-14 • 6 protons 6 protons • 6 neutrons 8 neutrons • Uranium-235 U-237 • 92 protons 92 protons • 143 neutrons 145 neutrons

  24. Learning check • How many neutrons are found in Lead-206. Lead has 82 protons • 124 neutrons • What is the mass number for neon with 10 protons and 10 neutrons? • Neon-20

  25. Learning check • How many protons and electrons are in each atom? • fluorine (atomic number 9) • calcium (atomic number 20) • aluminum (atomic number 13) • How many neutrons are in each atom: • O-16 • S-32 • Ag-108 • Br-80 • Pb-207

  26. Many elements have several isotopes • To determine an average weight of all the isotopes, we use the atomic mass. • This is given in atomic mass units (amu) • 1 amu=1/12 the mass of a C-12 atom

  27. Atomic mass • Copper-63=69% • Copper-65=31% • (63x0.69)+(65x0.31) • Weighted average: 63.55 AMU=atomic mass

  28. Boron has 2 isotopes: boron-10 and boron-11. Which is more abundant and why, given that the atomic mass of boron is 10.81 amu. Calculate the atomic mass of bromine. The 2 isotopes have atomic masses and relative abundances of 98.92 amu (50.69%) and 80.92 amu (49.31%)

  29. Periodic Table Intro

  30. The periodic table is a way to organize the elements by similar or repeating properties • A quick way to compare different elements • The elements are arranged in order of increasing atomic number, as shown centered above the symbol

  31. Each horizontal row is called a period. • 7 periods • The properties vary as you move down the periods

  32. Each column is called a group • Elements within a group have similar chemical and physical properties • Each group is classified with a number and either A or B

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