1 st - Atomic History & Atomic Structure

1. 1st- Atomic History & Atomic Structure

2. History • Democritus • named the most basic particle • atom- means “indivisible • Aristotle • didn’t believe in atoms • thought matter was continuous

3. History • by 1700s, all chemists agreed: • on the existence of atoms • that atoms combined to make compounds • Still did not agree on whether elements combined in the same ratio when making a compound

4. Law of Conservation of Mass • mass is neither created or destroyed during regular chemical or physical changes

5. Law of Definite Proportions • any amount of a compound contains the same element in the same proportions by mass No matter where the copper carbonate is used, it still has the same composition

6. Law of Multiple Proportions • applies when 2 or more elements combine to make more than one type of compound • the mass ratios of the second element simplify to small whole numbers

7. Law of Multiple Proportions

8. Dalton’s Atomic Theory • All mass is made of atoms • Atoms of same element have the same size, mass, and properties • Atoms can’t be subdivided, created or destroyed • Atoms of different element combine in whole number ratios to make compounds • In chemical reactions, atoms are combined, separated, and rearranged.

9. Modern Atomic Theory • Some parts of Dalton’s theory were wrong: • atoms are divisible into smaller particles (subatomic particles) • atoms of the same element can have different masses (isotopes) • Most important parts of atomic theory: • all matter is made of atoms • atoms of different elements have different properties

10. Structure of Atom • Nucleus: • contains protons and neutrons • takes up very little space • Electron Cloud: • contains electrons • takes up most of space

11. Subatomic Particles • includes all particles inside atom • proton • electron • neutron • charge on protons and electrons are equal but opposite • to make an atom neutral, need equal numbers of protons and electrons

12. Subatomic Particles • number of protons identifies the atom as a certain element • protons and neutrons are about same size • electrons are much smaller • Strong nuclear force- when particles in the nucleus get very close, they have a strong attraction • proton + proton • proton + neutron • neutron + neutron

13. Atomic Radius • size of atom • measured from center of nucleus to outside of electron cloud • expressed in picometers (1012 pm = 1 m) • usually 40-270 pm

14. 2nd- Atomic Discovery

15. Discovery of Electron • resulted from scientists passing electric current through gases to test conductivity • used cathode-ray tubes • noticed that when current was passed through a glow (or “ray”) was produced (eventually led to TV!)

16. Discovery of Electron Noted qualities of ray produced: • It existed- there was a shadow on the glass when an object was placed inside • It had mass- the paddle wheel placed inside, moved from one end to the other so something must have been “pushing” it

17. Discovery of Electron Noted qualities of ray produced: • It was negatively charged- the rays behaved the same way around a magnetic field as a conducting wire (wires conduct the flow of negative particles) • It was definitely negatively charged- were repelled by a negatively charged object

18. Discovery of Electron • J.J. Thomson (English, 1897) led the experiments that proved the cathode rays were composed of negative particles • found ratio of charge to mass of this particle • since the ratio stayed constant for any metal that contained it, it must be the same in all of the metals • named this particle the ELECTRON • Robert Millikan (American, 1909) proved Thomson’s work (definitively determined the mass and charge of the electron

19. Are electrons the only particles? • since atoms are neutral, something must balance the negative charge • since an atom’s mass is so much larger than the mass of its electrons, there must be other matter inside an atom

20. Discovery of Nucleus • Rutherford discovered the nucleus by shooting alpha particles (have positive charge) at a very thin piece of gold foil • he predicted that the particles would go right through the foil at some small angle

21. Discovery of Nucleus Instead of passing straight through with little deflection, this is what happened:

22. Discovery of Nucleus • some particles (1/8000) bounced back from the foil • this meant there must be a “powerful force” in the foil to hit particle back Predicted Results Actual Results

23. Discovery of Nucleus Characteristics of “Powerful Force”: • dense- since it was strong enough to deflect particle • small- only 1/8000 hit the force dead on and bounced back • positively charged- since there was a repulsion between force and alpha particles

24. 3rd- Counting Atoms

25. Atomic Number • (Z) number of protons • All atoms of the same element have the same atomic number • located above the symbol in the periodic table • order of the elements in the periodic table

26. Mass Number • sum of particles in nucleus • M# = #p + #n

27. Isotopes • atoms of the same element with different numbers of neutrons • most elements exist as a mixture of isotopes • What do the Carbon isotopes below have in common? What is different about them?

28. Designating Isotopes • Hyphen notation: • Name - mass number • ex. Carbon – 13 • Nuclear Symbol notation:

29. Examples • 7 protons, 8 neutrons Nitrogen-15 • 17 electrons, 19 neutrons Chlorine- 36

30. Examples • Z=5, 6 neutrons Boron- 11 • A=75, 42 neutrons Arsenic- 75

31. Examples

32. Examples

33. 4th- Average Atomic Mass& The Mole

34. Relative Atomic Mass • since masses of atoms are so small, it is more convenient to use relative atomic masses instead of real masses • to set up a scale, we have to pick one atom to be the standard • since 1961, the carbon-12 atom is the standard and is assigned a mass of exactly 12 amu

35. Relative Atomic Mass • atomic mass unit (amu)- one is exactly 1/12th of the mass of a carbon-12 atom • mass of proton= 1.007276 amu • mass of neutron= 1.008665 amu • mass of electron= 0.0005486 amu

36. Relative Atomic Mass • the mass number (A) and the relative atomic mass are very close but not the same because • relative atomic mass includes electrons • the proton and neutron masses aren’t exactly 1 amu

37. Average Atomic Mass • weighted relative atomic masses of the isotopes of each element • each isotope has a known natural occurrence (percentage of that elements’ atoms)

38. Calculating Average Atomic Mass • Naturally occurring copper consists of: • 69.17% copper-63 (62.929598 amu) • 30.83% copper-65 (64.927793 amu)

39. Calculating Average Atomic Mass • An element has three main isotopes with the following percent occurances: • #1: 19.99244 amu, 90.51% • #2: 20.99395 amu, 0.27% • #3: 21.99138 amu, 9.22% • Find the average atomic mass and determine the element.

40. Calculating Average Atomic Mass

41. The mole • a unit for measuring a very large amount- like number of atoms or molecules in a sample • like one dozen (1 dozen = 12 things) • except bigger: 1 mole = 6.022x1023 things • Why 6.022x1023 ? • 6.022x1023 is the number of atoms in exactly 12 g of carbon-12

43. The mole • 6.022x1023 is called Avogadro’s Number in honor of all of his contributions to chemistry • can be used as a conversion factor between a number of things and mole

44. Molar Mass • the mass of one mole of pure substance (g/mol) • numerically equal to average atomic mass • under the symbol on the periodic table • can be used as a conversion factor between moles and grams

45. Conversion Factors # Atoms Grams Moles Use Molar Mass: grams per mole Use Avog.’s Number: atoms per mole

46. Gram  Moles • use molar mass • Ex. 32.3 g Na = ? mol Na • Ex. 0.56 mol Fe = ? g Fe

47. # Atoms  Moles • use Avogadro’s Number • Ex: 1.40 mol Na = ? Na atoms • Ex: 3.4x1023 atoms Fe = ? mol Fe

48. Grams  # Atoms • use both: Avogadro’s # and molar mass • Ex: 0.0326 g N = ? atoms of N • Ex: 2.01x1041 atoms of H = ? g H