Chapter 2 Atoms, Molecules, and Ions
I. Chemistry History
A. Early chemistry
1. Greek chemistry (~ 400 BC)
a. Matter consists of four elements: fire, earth, air, water
b. Atoms as indivisible particles proposed by Demokritos
c. Didn’t test their ideas with experiments
2. Alchemy (300 BC to 1500 AD)
a. Pseudoscience: chemistry mixed with mysticism
b. Ultimate goal was to turn lesser metals into gold
c. Discovered some elements (S, Hg) and mineral acids (HCl, H2SO4)
3. Metallurgy and Medicine (~ 1500)
a. Georg Bauer advanced the isolation of metals from ores
b. Paracelsus used minerals as medicines
4. Robert Boyle (1627-1691)
a. Quantitative experimentation on the pressure and volume of air
b. Allowed for multiple elements: can’t be broken down further
c. Still believed metals could be changed to other elements
5. Phlogiston
a. Stahl (~ 1700) proposed that a substance flowed out of burning objects
b. Priestly (~ 1774) discovered dephlogisticated air (oxygen)
B. Fundamental Chemical Laws
1. Conservation of Mass
a. Lavosier (1743-1794) did careful weighing of reactants and products
b. He quantitatively showed that “mass is neither created nor destroyed”
c. He determined oxygen was involved in combustion and needed for life
2. Definite Proportions
a. Proust (1754-1826) determined the composition of many compounds
b. “A given compound always contains the same proportion of elements”
3. Multiple Proportions
CO2
a. Dalton (1766-1844) reexamined elements as composed of atoms
b. “Ratio of elements in compounds can be reduced to small whole #’s”
c. Example: We have three compounds made of N and O
i. Compound A has 1.750 gram N per gram of O
CO
ii. Compound B has 0.875 gram N per gram of O
iii. Compound C has 0.4375 gram N per gram of O
iv. A has 2x as much N as B and 4x as much N as C per gram of O
v. A = N2O; B = NO; C = NO2 or A = NO; B = NO2; C =NO4 etc…
d. Dalton couldn’t calculate the exact formulas, but he could the ratios
C. Dalton’s Atomic Theory
1. Dalton’s 1808 book presented his theory of atoms using the above laws
2. He prepared a table of atomic masses based on the elemental ratios
a. He assumed the simplest possibility:
i. If water = H2O (O = 16x as heavy as H)
ii. If water = HO (O = 8x as heavy as H)
b. Many of his assumed masses were wrong, but this was a good start
3. Specifics of Dalton’s Atomic Theory
a. Each element is composed of tiny, indestructible atoms
b. All atoms of a given element are the same, but different from the
atoms of other elements
c. Atoms combine in simple whole number ratios to form compounds
d. Atoms of one element can’t change into another element
e. 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)
a. “Equal volumes of different gases contain the same # of particles”
b. Hydrogen, Oxygen, and Chlorine must all be diatomic (H2, O2, Cl2)
II. Characterizing the Atom
A. The Electron
1. Cathode is a negative electrode
giving off a ray under high
potential
2. The “ray” coming from the
cathode was repelled by
negative charge so the particles
were called Electrons
3. Thomson measured the charge
to mass ratio of the electron as
e/m = -1.76 x 108 C/g
4. Multiple metals gave off
electrons, so electrons must be
in all elements
5. Since atoms are neutral, there
must be some source of positive
charge in atoms
6. Plum Pudding Model of the
atom
B. In 1909, Millikan calculated the charge and mass of the electron
1. Tiny charged oil drops were allowed to drop through an electric field
2. Measuring the size of the field needed to stop the drop from falling
allowed the calculation of the total charge of each drop
3. 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
1g
 28
31
(1.60x10 19 C)  (
)

9.10x10
g

9.10x10
kg
 1.76x10 8 C
B. Radioactivity
1. Becquerel (1896) discovered that uranium produced an image on
photographic film in the absence of light
2. This spontaneous emission of radiation was called radioactivity
3. Three types of radioactive emission were eventually discovered
a. Gamma rays (g) = high energy light wave
b. Beta particles (b) = high energy electron
c. Alpha particles (a) = particle with + charge 2x as large as electron and
a mass 7300 times that of an electron
C. The Nucleus
1. Rutherford (1911) tested Thomson’s Plum Pudding model
2.
D.
The results were different than expected
a. Heavy alpha particles should have gone right through foil
b. Some were deflected by a large angles and some reflected back
c. Rutherford surmised atomic + charge must be in a heavy nucleus
The Modern Atomic Structure
1. Tiny Nucleus surrounded by a cloud of electrons
2. Sizes: Nucleus = ball bearing; Atom = football stadium
3. Mass: essentially all the mass is at the nucleus (7300:1 ratio proton to
electron mass)
4. The chemistry of the atom is primarily due to the electrons
a. Different elements have different numbers of protons and electrons
b. This leads to their different properties and reactivities
5. Isotopes = atoms with the same # of protons, but different #’s of neutrons
a. Sodium (Na) always has 11 protons and 11 electrons
b. Sodium atoms can have 12 or 13 neutrons (23Na and 24Na)
c. Mass number = 23 or 24 and is written as a superscript to the left
d. Isotopes have different masses, but identical reactivity, since reactivity
is due to the number of protons/electrons
6.
Examples of Isotopes
A
Z
7.
X  126 C, 136C,
14
6
C
Ions: Losing and Gaining electrons
a. Cation: positively charged particle left when electon(s) are lost
Li -----> Li+ + e- (Lithium cation)
b. Anion: negatively charged particle formed when electrons are gained
F + e- -----> F- (Fluoride anion)