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Chemistry Unit 3

Chemistry Unit 3. Atomic Structure (Ch.3). 3-1 Early Models of Atoms. Democritus (450 BC) Proposed that all matter was made of tiny indivisible particles. He called these particles atomos (meaning indivisible). We call them atoms . Good looking guy!. Atom.

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Chemistry Unit 3

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  1. Chemistry Unit 3 Atomic Structure (Ch.3)

  2. 3-1 Early Models of Atoms • Democritus (450 BC) • Proposed that all matter was made of tiny indivisible particles. • He called these particles atomos (meaning indivisible). • We call them atoms. Good looking guy!

  3. Atom • An atom is the smallest particle of an element that retains the identity of that element. • If we repeatedly cut a piece of Al, the smallest possible piece is an atom of Al. Classic model of an atom

  4. Aristotle • Didn’t agree with Democritus. • Believed matter was continuous and made up of only one substance called “hyle” • It wasn’t until the 1700’s when his ideas were reexamined.

  5. Newton and Boyle (1600s) • Published articles stating their belief in the atomic nature of elements • They had no proof

  6. Antoine Lavoisier (1770’s) • French, The “Father of Modern Chemistry” • Discovered the law of conservation of matter. • Matter is neither created nor destroyed.

  7. Joseph Proust (1799) • French Chemist • Developed The Law of Definite Proportions • Compounds always contain elements in the same proportion by mass.

  8. Law of Definite Proportions • H20 (by mass is always) • 88.9% Oxygen, 11.1% Hydrogen • If we had an 80g sample of H20 how much is O? • .889 x 80 = 71g • How much is H? • .111 x 80 = 9g

  9. John Dalton (1803) • Proposed the atomic theory of matter, which is the basis for present atomic theory John Dalton, English schoolteacher

  10. Atomic Theory of Matter • Each element is composed of extremely small particles called atoms. • All atoms of a given element are identical, but differ from those of any other element. Which element is this?

  11. Atomic theory of matter • When elements unite to form compounds, they do so in a ratio of small whole numbers. This is called the Law of Multiple Proportions. • Ex: C and O can combine to form CO or CO2, but not CO1.8.

  12. Dalton’s Model of an Atom All matter is composed of tiny particles

  13. J.L. Gay-Lussac (early 1800s) • Observed that working with gas reactions at constant volume, temperature and pressure are directly related. • He named the discovery of this relationship Charles Law, which is represented by P1/T1=P2/T2.

  14. Amadeo Avogadro (early 1800s) – Italian Physicist • Explained Gay-Lussac’s work using Dalton’s theory: Equal volumes of gases at the same temp/pres have the same number of gas molecules.

  15. Michael Faraday (1839) • Suggested that atomic structure was related to electricity. • Atoms contain particles that have electrical charges. • Positive (+) • Negative (-) • Opposite charges attract • Like charges repel

  16. William Crookes (1870’s) • English Physicist • Developed the cathode ray tube to find evidence for the existence of particles within the atom.

  17. J.J. Thomson (1896) • Used a cathode ray tube (CRT) to identify negatively charged particles, called electrons. • Determined the ratio of an electron’s charge to its mass. • Developed the “plum pudding” model of an atom. Cathode ray bending toward a positive charge

  18. Plum Pudding Model - + - + + - + - + + + - - Atoms are composed of randomly arranged charged particles

  19. Robert Millikan (early 1900s) • US Physicist • Used the oil drop experiment to prove the electron has a negative charge • Was able to determine the charge of the electron

  20. Millikan’s Oil Drop Experiment

  21. Bothe/Chadwick (early 1930s) • English • Found high energy particles with no charge with the same mass as the proton called neutrons.

  22. Ernest Rutherford (1909) • Used the gold foil experiment to prove the atom is mostly empty space. • Rutherford concluded that all of an atom’s positive charge, and most of its mass is located in the center, called the nucleus. Analogy: thumb nail and the 50 yard line.

  23. 98% of the particles passes straight through 2% of the particles deflected off at varying angles 0.01% of the particles bounced straight back

  24. Rutherford’s Planetary Model of an atom - - + + + - + Positive charge and majority of mass located in the nucleus. Negatively charged electrons orbit the nucleus like planets. + + + - - - Most of an atom is empty space!

  25. Problem • He thought a moving electrical charge (-) in a curved path should lose energy (give off light). • If it did, it would fall into the (+) nucleus. • Why don’t the (-) electrons fall into the (+) nucleus?

  26. Atom:The smallest particle of an element that has the properties of that element. • Make up of nucleus consists of protons and neutrons • Surrounded by an electron cloud Electron cloud

  27. Sub-Atomic Particles • Protons • Positively (+) charged • The number of protons in an atom refers to the atomic number (Z) • Composed of 3 quarks (2 up, 1 down) • Mass= 1.6726 x 10-27kg • Atomic mass1 amu (µ)

  28. Sub-Atomic Particles • Neutrons • Found in nucleus • Neutral (no) charge • composed of 3 quarks (1 up, 2 down) • Atomic mass 1 amu (µ) • Isotopes- atoms of the same element that have a different number of neutrons.

  29. Sub-Atomic Particles • Electrons • Found in electron clouds surrounding the nucleus. • Negative (-) charge • Mass = .00091 x 10-27 kg • 1800 times smaller than protons & neutrons • Mass  0 amu (µ)

  30. Sub-atomic particles • Electrons • Orbit the nucleus at very high speed in energy levels (electron clouds). • Negatively (-) charged • Have no mass (when compared to protons and neutrons)

  31. Atomic Number = Protons • The atomic number of an element is the number of protons an element has. • Located above the symbol of the element • The number of protons determines the identity of the element. • Each element has a different atomic number

  32. Improved Rutherford’s work by saying electrons do not lose energy in the atoms so they will stay in orbit Stated there are definite levels in which the electrons follow set paths without gaining or losing energy (Planetary Model) Each level has a certain amount of energy associated with it and the electrons can only jump levels if they gain or lose energy Could not explain why (-) electrons don’t fall into the (+) nucleus. Neils Bohr (1913)

  33. Energy Levels • In the ground state for an atom, electrons are at their lowest, most stable energy levels. • In the excited state, atoms require energy and electrons move to a higher energy level.

  34. How many electrons are in an atom? • For an atom to have an overall neutral charge the number of electrons must equal the number of protons. • #Protons=#electrons • What element is this?

  35. Mass number • The Mass number of an atom is the sum of the mass of protons and neutrons • Located below the symbol of the element • Atomic mass is measured in amu’s, (atomic mass units) • Based on Carbon having a mass of 12 Mass = Protons + Neutrons

  36. How many neutrons are in an atom? • Mass=Protons+Neutrons • 195= 78 + Neutrons • 195-78= Neutrons • Platinum has 117 Neutrons • Find the number of neutrons in: • Hydrogen Carbon • Helium Potassium • Boron Gold

  37. Mass =Protons + Neutrons • Hydrogen (H) 1 =1 + Neutrons • Hydrogen has 0 neutrons • Helium (He) 4 = 2 + Neutrons • Helium has 2 neutrons • Boron (B) 11 = 5 + Neutrons • Boron has 6 neutrons • Carbon (C) 12 = 6 + Neutrons • Carbon has 6 neutrons • Potassium (K) 39 = 19 + N • Potassium has 20 neutrons • Gold (Au) 197 = 79 + N • Gold has 118 neutrons

  38. Atomic Mass • The average mass of all of the isotopes of an element. • Aka: average atomic mass number, or atomic weight. • Isotopes:Atoms of the same element with different masses.

  39. Average Atomic Mass • Ne-20 has a mass of 19.992 amu (u), and Ne-22 has a mass of 21.991 amu (u). In any sample of 100 Ne atoms, 90 will be Ne-20. Calculate the average atomic mass of Ne. • .90 x 19.992 = 17.9928 • .10 x 21.991 = 2.1991 • avg mass = 20.1919 amu

  40. Ions Na 11 P 11 e- • An atom that has gained or lost an electron. • It acquires a net electrical charge. • If an atom loses an electron (oxidation) it has more protons than electrons and has a net positive charge. (cation) 11 P 10 e- Na+

  41. Ions • If an atom gains an electron (reduction) it has more electrons than protons and has a net negative charge.(anion) 7 valence e- Full octet

  42. Ionic Charges • Charge of ion = # protons - # electrons • What is the charge of a magnesium atom that loses 2 electrons? • Number of protons 12 • -Number of electrons 10 • charge of ion +2 • Mg2+ or Mg+2 • Charge is written to the upper right of the symbol.

  43. Representations of atoms A • General form: (Elemental Notation) • X = Element Symbol • A = Atomic Mass (P + N) • Z = Atomic Number (P) • Ionic Charge Charge X Z

  44. What is the atomic structure? • Determine the number of: • P = • N = • e = 23 Na + 11

  45. What is the atomic structure? • Determine the number of: • P = 11 • N = 12 • e-= 10 23 Na + 11

  46. What is the atomic structure? • Determine the number of: • P = • N = • e- = I - 127 53

  47. What is the atomic structure? • Determine the number of: • P = 53 • N = 74 • e- = 54 I - 127 53

  48. Put into elemental notation • Atomic # = 29 • Atomic Mass = 64 • Ionic charge = +2 ?

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