AP CHAPTER TWO OUTINE
ATOMS AND ELEMENTS
I.
Atomic Structure and Subatomic particles

Electrical charges of the same type repel one another, and charges of the
opposite attract one another
A. Radioactivity
1898 Marie Curie suggested that certain atoms, such as radium, spontaneously emit
rays which she named radioactivity.
1. alpha (α) rays: have a +2 charge and mass
2. beta (β) rays: have a +1 charge and mass
3. gamma (γ) rays: have no detectable charge or mass
B. Electrons
1. mass: 9.1094 x 10-28 g
2. charge: -1
C. Protons
1. mass: 1.6726 x 10-24 g
2. charge: +2
D. Neutrons:
1. mass: 1.6749 x 10-24g
2. charge: 0
II.
The Nuclear Atom
Nucleus: dense region of the atoms where the protons and neutrons are, occupies
about 1/10,000 of the atom (most of an atoms volume is occupied by the electrons)
Overview
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Protons and neutrons make up the nucleus, providing most of the atom’s
mass; the protons provide all of its positive charge.
The nuclear radius is approximately 10,000 times smaller than the radius of
the entire atom.
Negatively charged electrons outside the nucleus occupy most of the volume
of the atom, but contribute very little mass.
A neutral atom has no net electrical charge because the number of electrons
outside the nucleus equals the number of protons inside the nucleus.
Ions are atoms that have gained or lost electrons and are not neutral. A
negative ion, anion, has gained electrons; whereas a positive ion, a cation, has
lost electrons.
III.
The Sizes of Atoms and the Units Used to Represent Them
A. Metric System: always used in science Units of Measurement
SI units are used in science worldwide
1. Mass: a measure of the quantity of matter
a) The unit for mass is the kilogram. One kg weighs 2.2 pounds
b) weight is a measure of the gravitational pull of matter
2. Length: measured in meters
3. Temperature: SI unit is the Kelvin. Frequently Celsius is also used.
B. Prefixes:
Kilo (k) 1000 or 103
Hecta (h) 100 pr 102
Deka (da) 10
Unit: meter, gram, etc
Deci 0.1 or 10-1
Centi 0.01 or 10-2
Milli 0.001 or 10-3
C. Derived Units
1. Volume: the amount of space occupied by an object
D. SI unit is m3. A cm3 is equal to a mL
2. Density: the ration of mass to volume or mass divided by volume
D = mass/volume
** Because density is a ratio of two measured quantities with different units,
density is always expressed with two units. For example, the density of water is
1.0 g/mL
3. Conversion Factors: equalities between units that are used to covert from one unit
to another. For example, since 2.54 cm = 1 inch
10 inches x (2.54cm/1 inch) = 25.4 cm
IV.
Using Scientific Measurement
1.
Accuracy: refers to the closeness of measurements to the
accepted value, for example an accurate “shot” would hit
the “bulls-eye”
2. Precision: refers to the closeness of a set of like
measurements. For example, if several shots all were
clumped together – even if they all are far from the bullseye – that would be precise
3. Percent Error: the accuracy of an experimental value
compared to actual value
Percent Error = Value accepted – Value experimental
Value accepted
x 100
A. Significant Figures
In science, measured values are reported in terms of significant figures. Significant
figures in a measurement consist of all the digits known with certainty, plus one
uncertain digit.
B. Rules for Determining Significant Figures
1) ALL non-zero numbers (1,2,3,4,5,6,7,8,9) are ALWAYS significant.
2) ALL zeroes between non-zero numbers are ALWAYS significant.
3) ALL zeroes which are SIMULTANEOUSLY to the right of the decimal point
AND at the end of the number are ALWAYS significant.
4) ALL zeroes which are to the left of a written decimal point and are in a number
>= 10 are ALWAYS significant.
C. Addition or Subtraction with Significant Figures
When adding or subtracting decimals, the answer must have
the same number of digits to the right of the decimal point as
there are in the measurement having the fewest digits to the
right of the decimal point.
D. Multiplication and Division with Significant Figures
For multiplication or division, the answer can have no more
significant figures than are in the measurement with the fewest
number of significant figures.
*****In working problems on a calculator, you should do
the calculation using all the digits and round only at the
end of the problem. *****
E. Scientific Notation
a method in which numbers are written in the form M x 10n, where the factor
M is a number greater than one but less than ten, and n is a whole number.
Example: 3,210 = 3.21 x 103
Example: 0.0064 = 6.4 x 10-3
 Convert from Scientific Notation to Real Number:
5.14 x 105 = 514000.0
Scientific notation consists of a coefficient (here 5.14) multiplied by 10 raised to
an exponent (here 5). To convert to a real number, start with the base and multiply
by 5 tens like this: 5.14 x 10 x 10 x 10 x 10 x 10 = 514000.0. Multiplying by tens
is easy: one simply moves the decimal point in the base (5.14) 5 places to the
right, adding extra zeroes as needed.
 Convert from Real Number to Scientific Notation:
0.000345 = 3.45 x 10-4
Here we wish to write the number 0.000345 as a coefficient times 10 raised to an
exponent. To convert to scientific notation, start by moving the decimal place in the
number until you have a number between 1 and 10; here it is 3.45. The number of places
to the left that you had to move the decimal point is the exponent. Here, we had to move
the decimal 4 places to the right, so the exponent is -4.
A. Adding and Subtracting with Scientific Notation can only be done if the values
have the same exponent.
B. Multiplication: the M factors are multiplied, and the exponents are added
C. Division: the m factors are divided, and the exponent of the denominator is
subtracted from the exponent of the numerator
V.
Atomic Numbers and Mass Numbers
A. Atomic Number: atoms of the same element have the same number of protons in
their nuclei. This number is called the atomic number and is given the symbol Z.
B. Mass Number: the sum of all the protons and neutrons inside an atoms nucleus.
C. Atomic Mass Unit (amu): protons and neutrons are measured in amu’s. 1 amu is
= 1/12 the mass of a carbon atom having 6 protons and 6 neutrons in the nucleus.
Z
A
X
A : mass number X: element symbol Z: atomic #
D. Isotopes: atoms with the same atomic number, but different mass number (due to
differing numbers of neutrons). Examples are 12C and 13C.
VI.
Isotopes and Atomic Weight
 In nature, many atoms occur with different isotopes. The proportion of
atoms of each isotope is called the percent abundance, the percentage of
atoms of a particular isotope.
Percent abundance =
(number of atoms of a given isotope)/(total number of atoms of all isotopes of that
element) x 100%

The atomic weight of an element is the average mass of a representative
sample of atoms of the element expressed in atomic mass units