Chapter 2 Atoms, Molecules, and Ions

1. Chapter 2 Atoms, Molecules, and Ions • I. Chemistry History • Early chemistry • Greek chemistry (~ 400 BC) • Matter consists of four elements: fire, earth, air, water • Atoms as indivisible particles proposed by Demokritos • Didn’t test their ideas with experiments • Alchemy (300 BC to 1500 AD) • Pseudoscience: chemistry mixed with mysticism • Ultimate goal was to turn lesser metals into gold • Discovered some elements (S, Hg) and mineral acids (HCl, H2SO4) • Metallurgy and Medicine (~ 1500) • Georg Bauer advanced the isolation of metals from ores • Paracelsus used minerals as medicines

2. 4. Robert Boyle (1627-1691) • Quantitative experimentation on the pressure and volume of air • Allowed for multiple elements: can’t be broken down further • Still believed metals could be changed to other elements • Phlogiston • Stahl (~ 1700) proposed that a substance flowed out of burning objects • Priestly (~ 1774) discovered dephlogisticated air (oxygen) • Fundamental Chemical Laws • Conservation of Mass • Lavosier (1743-1794) did careful weighing of reactants and products • He quantitatively showed that “mass is neither created nor destroyed” • He determined oxygen was involved in combustion and needed for life • Definite Proportions • Proust (1754-1826) determined the composition of many compounds • “A given compound always contains the same proportion of elements”

3. CO2 • 3. Multiple Proportions • Dalton (1766-1844) reexamined elements as composed of atoms • “Ratio of elements in compounds can be reduced to small whole #’s” • Example: We have three compounds made of N and O • Compound A has 1.750 gram N per gram of O • Compound B has 0.875 gram N per gram of O • Compound C has 0.4375 gram N per gram of O • A has 2x as much N as B and 4x as much N as C per gram of O • A = N2O; B = NO; C = NO2 or A = NO; B = NO2; C =NO4 etc… • Dalton couldn’t calculate the exact formulas, but he could the ratios • Dalton’s Atomic Theory • Dalton’s 1808 book presented his theory of atoms using the above laws • He prepared a table of atomic masses based on the elemental ratios • He assumed the simplest possibility: • If water = H2O (O = 16x as heavy as H) • If water = HO (O = 8x as heavy as H) • Many of his assumed masses were wrong, but this was a good start CO

4. 3. Specifics of Dalton’s Atomic Theory • Each element is composed of tiny, indestructible atoms • All atoms of a given element are the same, but different from the atoms of other elements • Atoms combine in simple whole number ratios to form compounds • Atoms of one element can’t change into another element • In chemical reactions, the atoms don’t change, but the way they are bound together does

5. 4. Gay-Lussac (1778-1850) measured the volumes of gases that reacted • 5. Avogadro (1776-1856) • “Equal volumes of different gases contain the same # of particles” • Hydrogen, Oxygen, and Chlorine must all be diatomic (H2, O2, Cl2)

6. II. Characterizing the Atom • The Electron • Cathode is a negative electrode giving off a ray under high potential • The “ray” coming from the cathode was repelled by negative charge so the particles were called Electrons • Thomson measured the charge to mass ratio of the electron as e/m = -1.76 x 108 C/g • Multiple metals gave off electrons, so electrons must be in all elements • Since atoms are neutral, there must be some source of positive charge in atoms • Plum Pudding Model of the atom

7. In 1909, Millikan calculated the charge and mass of the electron • Tiny charged oil drops were allowed to drop through an electric field • Measuring the size of the field needed to stop the drop from falling allowed the calculation of the total charge of each drop • He found that the charge was always a whole number multiple of • -1.60 x 10-19 C = the charge of a single electron • 4. He used Thompsons charge to mass ratio to find the mass of the electron

8. Radioactivity • Becquerel (1896) discovered that uranium produced an image on photographic film in the absence of light • This spontaneous emission of radiation was called radioactivity • Three types of radioactive emission were eventually discovered • a. Gamma rays (g) = high energy light wave • Beta particles (b) = high energy electron • Alpha particles (a) = particle with + charge 2x as large as electron and a mass 7300 times that of an electron • The Nucleus • Rutherford (1911) tested Thomson’s Plum Pudding model

9. The results were different than expected • Heavy alpha particles should have gone right through foil • Some were deflected by a large angles and some reflected back • Rutherford surmised atomic + charge must be in a heavy nucleus • D. The Modern Atomic Structure • 1. Tiny Nucleus surrounded by a cloud of electrons

10. Sizes: Nucleus = ball bearing; Atom = football stadium • Mass: essentially all the mass is at the nucleus (7300:1 ratio proton to electron mass) • The chemistry of the atom is primarily due to the electrons • Different elements have different numbers of protons and electrons • This leads to their different properties and reactivities • Isotopes = atoms with the same # of protons, but different #’s of neutrons • Sodium (Na) always has 11 protons and 11 electrons • Sodium atoms can have 12 or 13 neutrons (23Na and 24Na) • Mass number = 23 or 24 and is written as a superscript to the left • Isotopes have different masses, but identical reactivity, since reactivity is due to the number of protons/electrons

11. Examples of Isotopes • Ions: Losing and Gaining electrons • Cation: positively charged particle left when electon(s) are lost • Li -----> Li+ + e- (Lithium cation) • Anion: negatively charged particle formed when electrons are gained • F + e- -----> F- (Fluoride anion)