AP CHEMISTRY LECTURE NOTES: Unit 1
CHAPTER 1: SOME BASIC CONCEPTS
Please scan through this chapter. It contains information that should be review for
you, including basic terminology, and concepts of measurement and mathematics.
If you have any questions regarding the material in this chapter, please ask them
in class, otherwise I shall assume that you know and understand it and will not return to it
for review.
CHAPTER 2: ATOMS, MOLECULES AND IONS
2.1 Atomic theory
The idea of fundamental particles of matter is very old, dating back to the ancient Greek
philosophers.
1803: John Dalton proposes the Atomic Theory:
1. Each element is composed of very small particles called atoms.
2. All atoms of a given element are identical.
3. Atoms of different elements have different properties.
4. Atoms of elements are not fundamentally changed during chemical reactions,
nor are atoms created or destroyed (Law of Conservation of Mass).
5. Compounds are formed when atoms of different elements combine.
6. In a given compound, the relative number and kind of each atom is constant
(for example water is always H2O, no matter where it comes from.)
Dalton's theory helps explain the Law of Multiple Proportions: If different elements
conbine in more than one way to form different compounds, the ratio of their masses in
the different compounds will be a small whole numbers. Ex: the mass ratio of O to H in
water is 8g:1g, whereas in hydrogen peroxide, the ratio is 16g:1g. When the oxygen is
compared in the 2 compounds, the ratio of oxygen in hydrogen peroxide to that in water
is exactly 2:1. This idea was of great help in the determination of many chemical
formulas.
2.2 Discovery of Atomic Structure
Even tho' we can't see individual subatomic particles, we know quite a bit about them.
Cathode rays: a stream of electrons produced in a high voltage vacuum tube.
Watching their deflection in a magnetic field showed they are (-) charged.
The rays are the same, no matter what the cathode is made of: conclusion, the
electron
is a fundamental particle of all matter.
1897 J. J. Thomson: determined the charge/mass ratio of the electron: 1.76 x 108 C/g (C
= coulomb, a unit of electrical charge)
1909 Millikan performs the famous "oil droplet" experiment and discovers that the
electric charge tiny oil droplets subjected to a magnetic field is a multiple of 1.60 x 10-19
C. He reasoned that this must be the charge of a single electron.
Electron mass = 1.60 x 10-19 C / 1.76 x 108 C/g = 9.11 x 10-28g
Radioactivity:
1896: Henri Becquerel discover natural radioactivity.
Roentgen discover X-rays
Late 1890's Marie and Pierre Curie isolate radium and polonium
Ernest Rutherford discover alpha, beta and gamma radiation.
By subjecting radioactive emissions to a magnetic field, he discovered that
• alpha ( particles are relatively large, with a +2 charge and low penetrating ability.
An alpha particle is a helium nucleus with no electrons. It consists of 2 protons and 2
neutrons.
• beta (particles are much like high speed electrons with a -1 charge and 100x
greater penetrating ability than alpha particles.
• gamma (rays are non-particulate waves of very high energy (much like an X-ray,
with 1000x greater penetrating ability than alpha.
1911: Rutherford shot alpha particles at a thin layer of gold foil.
Most of the particles passed straight through, a few were refracted as they passed
through, and a very few were reflected back. Rutherford reasoned that atoms are mostly
empty space.
Basic Forces
Gravitational forces: attractive force proportional to mass.
Electromagnetic forces: attract or repel based on electric or magnetic charge of objects
involved.
Strong nuclear forces: nuclear binding energy- holds nuclear particles together despite
electromagnetic repulsions.
Weak nuclear forces: poorly understood, but apparent based on observations of
radioactive materials.
2.3 Modern Atomic Theory
The true nature and structure of matter at the atomic level is still being vigorously
studied, and is still open to much debate. However, atoms are believed composed of
primarily 3 types of particles:
•
protons (p+) charge +1
•
neutrons (no) no charge
•
electrons (e-) charge -1
Atomic diameters range from about 1-5 Å (angstroms 1 Å - 10-10 m or 0.1 nm)
See sample exercise 2.1, p. 43.
Nuclear diameters are only about 10-4 Å. Nearly all of the atom's mass is in the nucleus.
Nuclear matter is extremely dense, in the range of 1013-1014 g/cm3. At this density, a
cube of nuclear matter 1 inch on a side (about the size of a match box) would weigh
1.8 billion tons! It has been speculated that this may be what collapsed stars are made of,
and may be the cause of black holes in space.
An atom with a 2 cm. diameter nucleus would have an electron cloud with a diameter of
about 200 m!
p+ and no are about equal in mass.
e- are only 1/1835 the mass of the other two.
An element's identity depends on its number of protons.
Atoms of the same element (having the same number of protons) with differing numbers
of neutrons (and therefore different masses) are isotopes.
Example: there are several isotopes of carbon. The most common is C-12 (6 p+ and 6
n×), however, C-10, C-11, C-13, C-14, C-15 and C-16 are also known. All have 6 p+ but
vary in number of n× from 4-10.
Atomic number = number p+ .
Mass number = p+ + no.
Nuclear binding energy is the tremendous force that holds atomic nuclei together. It is
much stronger than the electrostatic repulsions formed between the positively charged
protons. It is thought that these forces may be the strongest in the known universe. They
are capable of unleashing huge amounts of energy, as witnessed in nuclear reactions and
thermonuclear explosions.
2.4 The Periodic Table
Developed in 1869 by Mendeléev, originally by order of increasing atomic mass, later by
increasing atomic number.
Elements with similar properties are placed in the same vertical column, known as a
group or family.
Regular patterns of repeating chemical and physical properties occur as we move from
left to right across a horizontal row, known as a series or period.
Unfortunately, different groups of scientists have different ways of numbering the
columns of elements, which leads to some confusion.
To my knowledge, the most current scheme simply numbers the columns in the main
body of the table as 1-18. Your text does not use this method, but uses instead, the more
traditional American scheme of numbering the first 2 and last 6 columns as IA-VIIIA,
and assigning B's to the columns in the middle section of the table
Many of the families of elements have common names:
•
IA- alkali metals
•
IIA- alkaline earth metals
•
VIIA- halogens
•
VIIIA- noble gases (inert or rare gases)
There is a staircase line dividing the table into:
• metals, on the left of the line, and including the 2 rows of elements below the main
body of the table.
•
nonmetals, on the right of the line and including element #1, hydrogen.
• metalloids or semimetals, the elements bordering the line, excepting Al, which is
considered a metal.
Chemistry at work: the Chernobyl accident. Please read this article on p. 47.
2.5 Molecules and Ions
Molecules: a tightly bound group of atoms, composed of more than 1 part, but viewed as
a single object, like a car or television set.
Some elements are molecular: H O F Br I N Cl are all diatomic in their pure states.
O3- ozone has different properties than O2.
Other molecular elements include S8 and P5.
Molecular compounds have more than 1 element: H2O and H2O2- same elements,
different properties.
Molecular vs. Empirical Formulas: molecular formulas give the number and type of all
the atoms in a molecule of the compound. They may even tell the atomic arrangement.
Empirical formulas give the lowest whole number ratio of atoms in a compound. The
formulas for all ionic compounds are empirical.
Poem:
Structural formulas: Lewis or electron dot formulas: - = : (an electron pair) Shows the
arrangement of atoms in a molecule. Abbreviations are often used because large organic
molecules can get very complex.
Ions: an atom or group of atoms which have had 1 or more electrons added or removed.
The result is an electrically charged particle.
Example:
Some ions are polyatomic: NO3- or SO4-2 You should already be familiar with the most
common of these, but there are a large number of polyatomic ions. Many ions with their
names and charges are listed inside the back cover of your text.
In ionic compounds, large numbers of oppositely charged ions (+) and (-) are arranged in
a regular geometric pattern forming a crystal lattice. In many ways, such a structure
could be viewed as a giant molecule. Since each crystal would have a different number
of total ions and therefore a different molecular formula, ionic compounds are always
represented by their empirical formulas.
+ and - ions are shown in their lowest whole number ratio in the crystal lattice.
Naming Inorganic Compounds
Many chemical substances (there are more than 11 million known) have been known for
a long time, and have common names. Ex. water, ammonia, alum and lime.
But for most materials, a systematic naming approach is required.
Ionic cpds (compounds):
•
+ ion (cation) is always first, - ion (anion) always second.
Some ions have only 1 possible ionic charge, some have more than 1 possible and are
called variable charge ions.
There is a new and an older style naming system for such ions. You should be familiar
with both.
Examples:
new system
old system
Fe+2 = iron (II)
*ferrous ion
*Note use of the Latin roots for iron and
copper.
Fe+3 = iron (III)
ferric ion
Other ex: plumbus, nicklous, manganous,
etc.
Cu+1 = copper (I)
*cuprous ion
Cu+2 = copper (II)
cupric ion
"ous ending = lower possible charge.
"ic" ending = higher possible charge.
Remember: "O is low, I is high."
Another notable example are the mercury ions: Hg+2 = mercury (II) or mercuric ion and
Hg2+2 = mercury(I) or mercurous ion. **Note that the mercury (I) ion is diatomic.
All compounds with monatomic anions (single atom negative ions) have names ending
with "ide."
Three polyatomic ions end with "ide": CN-1 cyanide, OH-1 hydroxide and O2-2 peroxide.
All other polyatomic anions have names ending with "ate" or "ite."
You should know the names and charges of all the ions found on Table 2.4, p. 57 in your
text.
Oxyions: Parent ion name ends with "ate."
1 less oxygen atom changes the name to "ite."
Prefixes include: "per" = 1 more oxygen than parent ion.
"hypo" = 2 less oxygens than parent.
The prefix "bi" = hydrogen in the ion. Ex: HCO3-1 = hydrogen carbonate or
bicarbonate
HSO3-1 = bisulfite Bisulfites are used as a preservative in salad bars and other
places to
keep lettuce and fresh fruits from turning brown. Some people are allergic to
these
compounds.
The prefix "thio" = sulfur in the ion.
If more than 1 polyatomic ion is present in the empirical formula of a cpd, the ion
formula is enclosed in parentheses.
Naming Aqueous Acids:
Binary or non-oxyacids: cpd name ends with "ide"- acid name = "hydro____ic"
Ex. HCl = hydrogen chloride = hydrochloric acid.
Oxyacids: cpd name ends with "ate"- acid name changes to "ic."
Ex. H3PO4 = hydrogen phosphate = phosphoric acid
cpd name ends with "ite"- acid name changes to "ous."
Ex. HNO2 = hydrogen nitrite = nitrous acid.
Inorganic Molecular compounds: names end with "ide." (as does any binary cpd.)
Prefixes tell the number of each type of atom in the molecule: mono, di, tri, tetra,
penta, hexa, hepta, octa, nona, deca.
Drop the "a" from the prefix if the second element in the cpd begins with a vowel
(oxygen or iodine.)
If no prefix appears, mono is assumed. Mono usually dropped if first element is a
single atom.
Ex. CO2 technically monocarbon dioxide, but usually carbon dioxide.
*Generally treat metalloids as non-metals when naming their compounds. Ex. As2S3
= diarsenic trisulfide.
**And remember: There are many exceptions and inconsistencies in naming
compounds.
Here's a paradox for you to gnaw on: "The only rule to which there is no exception is
that there is an exception to every rule."
Your homework assignments are found on the web. Please do the selected questions
from Chapters 1 and 2.