Chapter 2 Atoms, Molecules, and Ions

1. Chemistry, The Central Science, 10th edition Theodore L. Brown; H. Eugene LeMay, Jr.; and Bruce E. Bursten Chapter 2Atoms, Molecules,and Ions John D. Bookstaver St. Charles Community College St. Peters, MO  2006, Prentice Hall, Inc.

2. 2. 1The Atomic Theory of Matter

3. Democritus (400 BC) • This is the Greek philosopher Democritus • His theory: Matter could not be divided into smaller and smaller pieces forever • eventually the smallest possible piece would be obtained. • This piece would be indivisible. • He named the smallest piece of matter “atomos,” meaning “not to be cut.”

4. Democritus’s theory was forgotten for over 2000 years • The philosophers of the time, Aristotle and Plato, had a more respected, (and ultimately wrong) theory. • Aristotle and Plato favored the earth,fire, air and waterapproach to the nature of matter.

5. Atomic Theory of Matter • The theory that atoms are the fundamental building blocks of matter reemerged in the early 19th century, championed by John Dalton. • Known as the “Father of the Atom” • First scientist to support the existence of the atom with scientific evidence

6. Dalton’s Postulates 1) Elements are made up of atoms 2) Atoms of each element are identical (in mass and properties). Atoms of different elements are different. 3) Chemical reactions are a rearrangement of atoms. Atoms are not created or destroyed. 4) Compounds are formed when atoms of multiple elements combine. Each compound has specific numbers and kinds of atoms.

7. Dalton’s Atomic Theory • He deduced that all elements are composed of atoms. Atoms are indivisible and indestructible particles. • Atoms of the same element are exactly alike. • Atoms of different elements are different. • Compounds are formed by the joining of atoms of two or more elements.

8. Law of Constant Composition • From Dalton’s 4th postulate • Also known as the law of definite proportions. • The elemental composition of a pure substance never varies. • In a compound or element, the numbers and kinds of atoms are constant.

9. Law of Multiple Proportions • When chemical elements combine, they do so in a ratio of small whole numbers. • For example, carbon and oxygen react to form carbon monoxide (CO) or carbon dioxide (CO2), but not CO1.3. • Further, if two elements, A & B, combine to form more than one compound, the masses of B that combine with a given mass A are in the ratio of small whole numbers • Example: • Water H2OvsHydrogen Peroxide H2O2 • Mass O = 16 gMass O = 32 g • Mass H = 2 gMass H = 2 g • Oxygen 32 g:16 g= 2 • Hydrogen 2 g :2 g= 1

10. Law of Conservation of Mass • From Dalton’s 3rd postulate • The total mass of substances present at the end of a chemical reaction is the same as the mass of substances present before the reaction took place.

11. 2. 2Discovery of Atomic Structure

12. J.J. Thompson’s Experiment • Using a cathode ray tube Thompson found that passing an electric current makes a beam of radiation appear to move from the negative to the positive end • By adding an electric field to the cathode ray tube, he found that the beam deflected towards the positive field so he deduces that the components of the beam were negative

13. Cathode Rays & Electrons • Radiation was passed through a tube from the cathode (negative end) to the anode (positive end) • The radiation would fluoresce (give off light) • When an electric/magnetic field was applied, the cathode rays deflected in a manner consistent with a stream of negative particles • Since the gas was known to be neutral (having no charge) he reasoned that there must also be positively charged particles in the atom (but he could never find them). • (Old TV sets were cathode ray tubes)

14. The Electron • Cathode Ray • Streams of negatively charged particles were found to emanate from cathode tubes. • J. J. Thompson is credited with their discovery (1897). • Thompson measured the charge/mass ratio of the electron to be 1.76  108 coulombs/g.

15. Millikan’s Oil Drop Experiment • Robert Millikan put a charge onto a tiny drop of oil and dropped it into an electric field • When the field was manipulated, the oil drop’s free fall was affected • By measuring how strong an applied electric field had to be in order to stop the oil drop in mid air, he was able to work out the mass of the drop • He could calculate the force of gravity on one drop and thus the electric charge that the drop must have. • By varying the charge on different drops, he noticed that the charge was always a multiple of -1.6 x 10-19 C, the charge on a single electron

16. Millikan’s Oil Drop Experiment Millikan’s Oil Drop Experiment • Once the charge/mass ratio of the electron was known, determination of either the charge or the mass of an electron would yield the other.

17. Radioactivity: • The spontaneous emission of radiation by an atom. • First observed by Henri Becquerel. • Also studied by Marie and Pierre Curie.

18. Radioactivity • Three types of radiation were discovered by Ernest Rutherford: •  [alpha] particles (positive charge) •  [beta] particles are high speed e– (neg. charge) •  [gamma] rays are unaffected by electric field/charge

19. “Plum Pudding” Model • In 1897, the English scientist J.J. Thompson provided the first hint that an atom is made of even smaller particles. • Thompson concluded that the negative charges came from within the atom. • A particle smaller than an atom had to exist… the atom was divisible! • Positive sphere of matter with negative electrons imbedded in it

20. Rutherford’s Gold Foil Experiment • 1909 - Ernest Rutherford shot  [alpha] particles at a thin sheet of gold foil and observed the pattern of scatter of the particles. Gold Foil Experiment

21. Results of the Gold Foil Experiment • Top: Expected results: alpha particles (+) should pass through the plum pudding model of the atom undisturbed. • Bottom: • Observed results: • Some of the particles were deflected at large angles, indicating a small, concentrated positive charge.

22. The Nuclear Atom • Rutherford predicted that most of the volume of the Au atoms was open, empty space. The “plum pudding” model could not be correct. • Rutherford concluded that an atom must have a small, dense, positively charged center (the nucleus) that repelled his positive particles • He also concluded that the e– were outside of the nucleus

23. Other Subatomic Particles • Protons • discovered by Rutherford in 1919. • Neutrons • discovered by James Chadwick in 1932.

24. 2. 3The Modern View of Atomic Structure

25. Subatomic Particles • Protons and electrons • the only particles that have a charge. • Protons and neutrons • have essentially the same mass. • The mass of an electron is so small that we ignore it (p+ and no are 1800x mass of e–).

26. Inside the Atom • Protons and neutrons reside inside the nucleus of the atom • Though the nucleus is extremely small, it contains nearly all the mass of the atom http://ed.ted.com/lessons/just-how-small-is-an-atom

27. Atomic Mass & Size • Because atoms have such tiny masses (heaviest is 4.0 x 10-22 g) we use the atomic mass unit(amu) • 1 amu = 1.66 x10-24 g • Atoms are very small so another common unit of length used is the Angstrom(Å) • diameters between 1x10-10 m and 5x10-10 m = 100-500 pm = 1 – 5 Å

28. Atomic Numbers, Mass Numbers, and Isotopes

29. Atomic Number (Z) • the # of protonsan atom contains • is unique to each element • Thus, # of p+ tells the identity of the atom • Because atoms are neutral, every atom has an equal number of protons and electrons

30. Mass Number (A) • Mass # = protons + neutrons • Mass of atoms in amu • Atoms of an element can differ in the # of neutrons they contain • As a result, atoms of the same element can have different masses

31. Practice Counting p+, no, and e– • p+ and e– = 15 • no = 31-15 = 16 • p+ and e– =26 • no = 56 -26 = 30 • p+ and e– = 10 • no = 20-10 = 10

32. Atoms of the sameelement (thus same # protons) with differentmasses and different #s of neutrons. 12 6 13 6 14 6 11 6 C C C C Isotopes: • Standard notation: • Hyphenated notation: • Carbon -12, Carbon -14, etc

33. More Practice • How many protons, neutrons, and electrons are in the following? • 108Ag • Nickel-60 • 209Pb • Silicon-30 p+ = 47; no = 108-47 = 61; e– = 47 p+ = 28; no = 60-28 = 32;e– = 28 p+ = 82; no = 209-82 = 127; e– = 82 p+ = 14; no = 30-14 = 16; e– = 14

34. 2. 4Atomic Weights

35. 24 Cr 51.996 Atomic Mass • Carbon-12 is thestandardfor atomic masses of elements: • It was assigned an exact mass of 12.00 amu • All other atomic masses were determined in comparison to the mass of Carbon-12 • The periodic table gives the masses of an “average” atom of each element • must be an avg. because there are multiple isotopes of each element present in nature • known as the average atomic mass

36. Average Atomic Mass • Avg. atomic mass is determined by • massof each isotope AND • each isotope’s relative abundance(how much of that isotope is found in nature) • That means that it is a weighted average of all the naturally occurring isotopes of that element. • Thus, avg. atomic mass is not a whole # • Time for some math! • *Notes Handout – Weighted Average*

37. Weighted Average Example • On the first day of school, Maggie’s Geometry teacher told the class that all grades would be calculated by the weighted average system seen below: • Participation: 15% Homework: 20% Quizzes: 25% Tests: 40% • At the end of the MP1, Maggie’s not sure if she is passing the class. She needs to know what her overall weigthed class average is. Below are Maggie’s average grades during MP1. Let’s help her! • Maggie’s MP1 Averages Participation: 70 Homework: 74 Quizzes: 72 Tests: 66

38. Avg. Atomic Mass Calculations - application of a weighted average - Avg. atomic mass (AAM) = [(rel, abundance isotope A) x (mass isotope A)] + [(rel. abundance isotope B) x (mass isotope B)] + [(rel. abundance isotope C) x (mass isotope C)] …etc *** Relative Abundance = % abundance *** 100

39. Avg. Atomic Mass Practice • Magnesium has three naturally occurring isotopes: 78.70% of Mg exists as Magnesium-24 (23.985 amu), 10.03% exist as Magnesium-25 (24.986 amu) and 11.17% exist as Magnesium-26 (25.983 amu). • What is the average atomic mass of Magnesium? • Avg. atomic mass Mg= (0.7870)(23.985 amu) 18.876 amu + 2.506 amu + 2.902 amu = = 24.284 amu + (0.1003)(24.986 amu) + (0.1117)(25.983 amu)

40. 2. 5The Periodic Table

41. The Periodic Table of Elements

42. Periodic Table • Period/Series = horizontal rows (#1-7, lanthanides, actinides) • Families/Groups= vertical columns(#1-18, special names for some) • Elements in the same family have similar chemical and physical properties.

43. Periodicity & Periodic Law Looking at the elements in the periodic table, there is a reoccurring pattern of chemical and physical properties at intervals.

44. Periodic Table Nonmetals are on the right side of the periodic table (with the exception of H).

45. Periodic Table Metalloids border the staircase line.

46. Periodic Table Metals are on the left side of the chart.

47. 1A – Alkali Metals (form +1 charge)

48. 2A – Alkali Earth Metals (form +2 charge)

49. 1B – Coinage Metals

50. 6A – Chalcogens (form –2 charge)