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

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”

CO2 CO 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 CO2 CO

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

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)

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

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

C. 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 D. The Nucleus Rutherford (1911) tested Thomson’s Plum Pudding model

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 E. The Modern Atomic Structure 1. Tiny Nucleus surrounded by a cloud of electrons

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

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)