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 ( )
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 ( ) 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 ( ) determined the composition of many compounds
“A given compound always contains the same proportion of elements”

3. CO2 CO 3. Multiple Proportions
Dalton ( ) 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 gram N per gram of O
Compound B has gram N per gram of O
Compound C has 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

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 ( )
“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 = 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 C = the charge of a single electron
4. He used Thompsons charge to mass ratio to find the mass of the electron

8. 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

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
E. 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)