1. Ch. 4: Atoms/Atomic Theory

2. Atoms Definition - the smallest particle that has the properties of an element, basic unit of matter 119 distinct atoms as of 1999, form elements

3. Atomic Symbols -each element has its own name, accompanied by a symbol -usually one/two letters (first one is always capitalized) ex. Iron: Fe -Fe represents 1 atom of iron, -2Fe represents 2 atoms of iron etc… -can also be written as Fe 2

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

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

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

7. 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

8. 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

9. Dalton’s Atomic Theory 1. All mass is made of atoms 2. Atoms of same element have the same size, mass, and properties 3. Atoms can’t be subdivided, created or destroyed 4. Atoms of different elements combine in whole number ratios to make compounds 5. In chemical reactions, atoms are combined, separated, and rearranged.

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

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

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

13. Subatomic Particles  number of protons identifies the atom as a certain element  protons and neutrons are about same size  electrons are much smaller

14. Comparing Subatomic Particles

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

16. Discovery of Electron Noted Qualities of Ray Produced: 1. existed- there was a shadow on the glass when an object was placed inside 2. 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 3. negatively charged- the rays behaved the same way around a magnetic field as a conducting wire 4. negatively charged- were repelled by a negatively charged object

18. Discovery of Electron  All of these led scientists to believe there were negatively charged particles inside the cathode ray

19. Discovery of Electron  J.J. Thomson (English 1897) did more experiments to actually make the discovery  he found ratio of charge of this particle to this mass of the particle  since the ratio stayed constant for any metal that contained it, it must be the same in all of the metals

20. 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

21. 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

22. Discovery of Nucleus

23.  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 ResultsActual Results

24. Discovery of Nucleus Characteristics of “Powerful Force”: 1. dense- since it was strong enough to deflect particle 2. small- only 1/8000 hit the force dead on and bounced back 3. positively charged- since there was a repulsion between force and alpha particles

25. Atomic Math and Isotopes

26. Atomic Number  number of protons is the atomic #  It is the identity of an element.  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

27. Isotopes  atoms of the same element with different numbers of neutrons  most elements exist as a mixture of isotopes

28. B. Isotopes © Addison-Wesley Publishing Company, Inc.

29.  What do the Carbon isotopes below have in common? What is different about them?

30. Mass Number For Isotopes  sum of particles in nucleus  Mass number for isotope = #p + #n Hydrogen isotopes have special names: protium protium deuterium deuterium tritium tritium

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

32. Examples 1. 7 protons, 8 neutrons Nitrogen-15 2. 17 electrons, 19 neutrons Chlorine- 36

33. Examples 3. Z=5, 6 neutrons Boron- 11 3. A=75, 42 neutrons Arsenic- 75

34. Ch. 4 Atoms Average Atomic Mass

35. Review Subatomic Particles Most of the atom’s mass. Atomic Number equals the # of... equal in a neutral atom

36. 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 nuclide is the standard and is assigned a mass of exactly 12 amu

37. Relative Atomic Mass  atomic mass unit (amu)- one is exactly 1/12 th 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

38. 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)

39. Calculating Average Atomic Mass  Naturally occurring copper consists of: 69.71% copper-63 (62.929598 amu) 69.71% copper-63 (62.929598 amu) 30.83% copper-65 (64.927793 amu) 30.83% copper-65 (64.927793 amu) (0.6971 x 62.929598)+(0.3083 x 64.927793) =63.88 amu

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

41. Calculating Average Atomic Mass Neon