CH 2: Atoms, Molecules, and Ions

1. CH 2: Atoms, Molecules, and Ions

2. Chapter Outline • History of chemistry (2.1 – 2.4) • Chemical laws – start with slide 8 for content, 3-7 FYI slides • Path to the atom • Modern atomic structure (2.5) • Molecules vs. Ions (2.6) • Naming molecular and ionic compounds (2.8) • Introduction to the periodic table (2.7)

3. History of Chemistry Greek Philosophers: 5th Century BCE(BCE = before the common era - replaces BC) • The Greek philosophers were the first to reflect on the nature of matter.  • They proposed that all matter is made out of first 4 elements -- earth, air , water, fire.  Aristotle added a fifth element, plasma (also called ether). 

4. Greek Philosophers • Democritus had an alternate view of matter.  He proposed that matter was  made up of tiny particles called atoms.  • His "theory" was not well accepted at the time.

5. Alchemy Alchemy: ~600-1600's CE • (CE = common era, replaces AD) • Alchemy developed at about the same time in China, India, and Greece. It spread into Europe in the 8th century.

6. Alchemy Alchemists had two pursuits • Search for a means to convert “base” metals into gold • Search for the elixir of life • Substance that would lead to immortality

7. Alchemy Advances from Alchemy • Many new substances where identified • Plaster of Paris, nitric acid…. • New lab techniques and equipment developed • New medicines identified

8. Modern Chemistry, ~1600 on • First chemists/physicists to use scientific method • Boyle - elements • Lavoisier – law of conservation of matter • Proust - law of definite proportions • Dalton – law of multiple proportions, atomic theory • Avogadro - hypothesis • Thomson – charge to mass ratio for an electron • Millikan – charge on the electron • Bequerel and the Curies - radioactivity • Rutherford – nuclear atom

9. Modern Chemistry Robert Boyle: ~1660 • Proposed a substance to be an element unless it can be broken down into simpler substances. • Proposed one of the gas laws – CH 5

10. Lavoisier: ~1760 Law of Conservation of Matter • Matter is neither created nor destroyed in a chemical reaction.

11. Proust: late 1700s Law of Definite Composition/Proportions • A given compound always contains the same proportion of elements by mass.

12. John Dalton: ~1800 Law of Multiple Proportions • When two elements form more than one compound, the ratios of the masses of the second element that combines with one gram of the first element can always be reduced to small whole numbers.

13. Law of Multiple Proportions Example • Consider two 100.0 g samples 2 different compounds containing only C and H. Compound A: 27.2 g of C • Compound B: 42.9 g of C. Show how this data illustrates the law of multiple proportions.

14. Dalton also proposed the first table of atomic masses • Most masses later need revision • Dalton is best known for proposing Atomic Theory

15. Dalton’s Atomic Theory • Elements are made up of tiny particles called atoms. • Atoms are indivisible and indestructible • Atoms of a given element are identical • Atoms of different elements differ in some fundamental way(s)

16. Dalton’s Atomic Theory • Compounds form when atoms of different elements combine with each other. • A given compound always has the same relative number and types of elements.

17. Dalton’s Atomic Theory • Chemical reactions occur when atoms change how they are bound to each other. • Individual atoms are not changed, just rearranged

18. Avogadro: 1811 Avogadro's Hypothesis • At the same temperature and pressure equal volumes of gases contain the same number of particles. • Based on Guy-Lussac’s data • See page 41/42

19. From Dalton to Atomic Structure • Dalton’s atomic theory lead to much research on the nature of the atom. • This research showed the atom to made up of smaller particles.

20. J.J. Thomson • Thomson measured the deflection of a cathode ray beam in electrical and magnetic fields of known strengths. Applied electrical field Cathode ray + (+) Metal electrode (-) Metal electrode - Cathode ray tube experiment, pg 43

21. Thomson found the cathode rays were attracted by the positive charge and repelled by negative • These findings clearly indicated that the rays consisted of negatively charged particles. • Today we know these particles as electrons.

22. Thomson measured the deflection of the beam in a magnetic field and more! From his data he determined the charge:mass ratio for an electron e = - 1.76 x 108 C/g m e = charge on the electron in coulombs M = mass of an electron in grams Cathode ray tube experiment

23. Thomson also found that the cathode ray particles were identical regardless of source. • Concluded all elements contain these negative particles (electrons)

24. J.J. Thomson Thomson: • identified cathode ray beams as a stream of negatively charged particles • calculated the charge to mass ratio for these negatively charged particles • proposed the existence of positively charged particles • To balance the negative charge of the electrons

25. Millikan ~1909 • Millikan’s oil drop experiment allowed him to determine the charge on an electron • This charge can be plugged into Thomson’s formula and the mass of the electron calculated • Mass electron = 9.11 x 10-31 kg • Page 44

26. Radioactivity Becquerel, Marie and Pierre Curie: ~1896 • Henri Becquerel - observed the natural emission of energy/rays by uranium. • Marie and Pierre Curie studied “Becquerel's rays”. • The Curies’ findings suggested that matter was composed of smaller particles than atoms. • The Curies coined the term radioactivity to describe the rays emitted.

27. Radioactivity • Three types of radioactivity were identified: • gamma rays - very high energy light • beta particles - high energy electrons • alpha particles - He+2 particles • 2 protons and 2 neutrons

28. Plum Pudding Model of the Atom • In the early 1900’s the accepted model of the atom was called the plum pudding model of the atom • Electrons (tiny and negatively charged) were pictured to be dispersed in a ‘cloud’ of positive charge. • Proposed by JJ Thomson and Lord Kelvin in 1904

29. Rutherford and the Nuclear Atom • In 1911 an experiment conducted in Ernest Rutherford’s lab showed the “plum pudding” model to be incorrect. • Experiment was conducted by Geiger and Marsden and the findings interpreted by Rutherford. • See page 45

30. Rutherford’s Atom • First to propose a nuclear atom. • An atom has a dense positive center containing all of positive charge and most of the mass of the atom – the nucleus • Electrons are scattered in the empty space around the nucleus • Electrons occupy a volume that is huge as compared to the size of the nucleus.

31. A New Model of the Atom Expected based on Plum pudding model Rutherford’s model Based on ”his” results

32. Modern Atomic Structure • Rutherford continued to study the atom and the positive matter of the atom. • 1919, + particle named the proton • ~1932 James Chadwick proposed the existence of a third subatomic particle, the neutron.

33. Subatomic Particles

34. Mass of Subatomic Particles • Protons and neutrons have ~ the same mass (in the range of 10 -27 kg). • Mass of each and of individual atoms is often expressed in amu rather than grams • Atomic mass unit (amu) – 1/12 the mass of a carbon-12 atom

35. Mass of Subatomic Particles • The mass of the electron (10-31 kg) is tiny as compared to that of the proton and neutron (10-27 kg) . • Therefore, the electron’s mass is considered to be ~0 amu when calculating the mass of an atom.

36. Subatomic Particles and the Elements • Each element has a unique number of protons. • Number of protons defines the element. Atomic # = # protons 6 C

37. Subatomic Particles • Since atoms are neutral, for every proton there is a/n _________. • When atoms interact to form compounds, it is their ___________ that interact.

38. Terms • Mass number = sum of the # of protons and the # neutrons in the nucleus of an atom • FOR MOST ELEMENTS THE MASS NUMBER IF NOT ON THE PERIODIC TABLE. • You will be given enough information to determine mass number or number of neutrons.

39. Terms • Isotopes = atoms of a given element that differ in mass number • Isotopes have the same number of _____________. • Isotopes differ in the number of _______.

40. Isotopes • Writing atomic symbols for isotopes • pg 46 11 Mass # B Symbol for element 5 Atomic #

41. FAQ - Isotopes • When is mass number found on the periodic table? • What’s the atomic mass? Is it the same as the mass number? C 84 (209) 6 12.0107 Po

42. Molecules and Ions (2.6) • Atoms of different elements combine to form compounds • Atoms in compounds are held together by chemical bonds. • Bonds involve interactions of the bonding atoms’ ________

43. Bonding There are two types of bonds: • Covalent bonds – bonding atoms share electrons • Atoms are always nonmetal atoms • Covalently bonded atoms form molecules • Ways to represent molecules • Chemical formula; H2O • Structural formula O H H

44. Bonding • Ionic bonds – attractive force among oppositely charged ions • Bond formed between metal cations and nonmetal anions • No molecules involved

45. Ions - Terms • Ion – charged atom or group of atoms • Formed when atoms gain or lose electrons • Cation – positively charged ion • Formed when an atom _______ electrons • Anion – negatively charged ion • Formed when an atom ______ electrons

46. Ions • Describing ion formation • Cation example: • Anion example:

47. Naming Binary Compounds • Binary compounds – compound composed of 2 elements • NaCl • CO • CO2

48. Types of Binary Compounds • Type I binary ionic compounds • Metal forms only one ion • Type II binary ionic compounds • Metal forms more than one ion • Use roman numerals to indicate the charge on the ion • Type III binary covalent compounds • Compound between 2 nonmetals

49. Types I Binary Compounds • Compound between a metal and a nonmetal • Metal forms only one ion • Name the cation and then the anion. • Name of the cation is the name of the element • Name of the anion is the name of the nonmetal with the ending changed to “ide”

50. Monoatomic cations to know