Section 4.2
Objectives
 State the three subparticles of atoms
 State the charges of the subparticles
 Describe the composition of an atom
Atom
 Fundamental particles
which make up matter
 The smallest particle of
an element that retains
(keeps) its identity in a
chemical reaction
Democritus’s Atomic
Philosophy
 460 B.C. to 370 B.C.
 First suggested the existence of atoms
 400 B.C.
 Believed atoms were:
 Indivisible
 Indestructible
 Theory faults
 Did not explain chemical behavior
 Lacked experimental support

not based on the scientific method
Dalton’s Atomic
Theory
 2000 years after Democritus
 John Dalton (1766-1844)
 By using experimental methods, he transformed Dem.
ideas on atoms into a scientific theory.
 Studied the ratios in which elements combine in
chemical rxns
Dalton’s Atomic Theory
All elements are composed of tiny indivisibly
particles called atoms.
2. Atoms of the same element are identical. The atoms
of any one element are different from tose of any
other element.
3. Atoms of different elements can physically mix
together or can chemically combine in simple wholenumber ratios to form compounds.
4. Chemical reactions occur when atoms are separated,
joined, or rearranged. Atoms of one element,
however are never changed into atoms of another
element as a result of a chemical reaction
1.
Dalton’s Atomic
Theory
Example of a Timeline
Sizing Up the
Atom
 Atoms are VERY small
 100,000,000 copper atoms side by side = 1 cm
 Radii of an atom
 5 x 10-11 m to 2 x 10-10 m
 Individual atoms can be seen with instruments like a
scanning tunneling microscope
 Have the ability to move around and arrange them in
patterns
Scanning Tunneling Microscope
Quarks
Atoms
 Composed of:
 Contains all of the mass
 Electrons
of the atom
 Contains protons
 Particles with a charge
of +1
 Neutrons
 Particles with no
charge
 Negatively charged
particles
 Nucleus
 Located in the center of
the atom
 Positively charged
Properties of Subatomic Particles
Particle
Symbol
Location
Electron
e-
Proton
p+
In the space 1surrounding
the nucleus
In the
1+
nucleus
Neutron
nO
In the
nucleus
Charge
0
Electrons
 Fast moving
 Travel through the space around the nucleus
 Held within the atom because of the attraction to the
positive nucleus
Objectives
 Explain the role of the atomic number in determining
the identity of an atom
 Define an isotope and explain why atomic masses are
not whole numbers
 Calculate the number of electrons protons and
neutrons in an atom given its mass number and atomic
number.
Atomic Number
 Atoms of an element have a unique positive charge in
their nuclei
 Number of protons determines the properties of an
element
 Atomic number
 = # of protons = # of electons
 Determines the elements position on the periodic table
Isotopes
 Atoms with the same number of protons but different
number of neutrons
 Differ in mass
 More neutrons = more mass
 Same chemical behavior
 Chemical behavior is determined by the # of electrons
 Have a number after the element to distinguish one
from another
 Ex. Potassium-39 vs potassium-40
Mass Number
 Sum of the protons and neutrons
 Number of neutrons = mass # - atomic number
Mass Spectrometer
 Instrument used to find
the actual masses of
individual atoms
 Example:
 Fluorine atom
 3.155 x 10-23 g
 Small and impractical to
work with
 NASA-The Molecule
Dissector-Mass
Spectrometer
Atomic Mass
 Compare the masses of an atoms to an isotope as a
standard
 Carbon-12
 6 protons/6 neutrons
 Assigned a mass of 12 atomic mass units
 Atomic Mass Unit (amu)
 1 amu
 1/12 of the mass of a carbon-12 atom
 The mass of one proton or one neutron
Practice
Element
Number of
Protons
Number of
Neutrons
Helium
2
2
Nitrogen
7
7
Sulfur
16
16
Predicted
Atomic Mass
Actual Atomic
Mass
Why not use whole numbers?
 Most elements occur as a mixture of two or more
isotopes
 Each isotope has a different abundance
 Atomic Mass of an Element
 A weighted average mass of the atoms in a naturally
occurring sample of the element
 Takes into account:
 Mass of isotope
 Relative abundance in nature
Calculating Atomic Mass
 Multiply the mass of each isotope by it’s natural
abundance (expressed as a decimal)
 Add the products
Example
 Element X has two natural isotopes. The isotope with
a mass of 10.012 amu (10X) has a relative abundance of
19.91%. The isotope with a mass of 11.009 amu (11X) has
a relative abundance of 80.09%. Calculate the atomic
mass of this element.
Practice
 The element copper has naturally occurring isotopes
with mass numbers of 63 and 65. The relative
abundance and atomic masses are 69.2% for the mass
of 62.93 amu and 30.8% for the mass of 64.93 amu.
Calculate the average atomic mass of copper.
Practice
 Calculate the atomic mass of bromine. The two
isotopes of bromine have atomic masses and relative
abundance of 78.92 amu (50.69%) and 80.92 amu
(49.31%.
A Preview
The Periodic Table
 An arrangement of elements  separated into groups
based on similar properties
 Allows you to compare one element to the next
 Arranged in order of increasing atomic number
 Period:
 Horizontal row
 Group or Family
 Vertical column
Practice Makes Perfect
 Take a minute to look over the names and
abbreviations of the elements on the periodic table.
 Create 6 flashcards
 On one side put the elements name
 On the other side put the element’s symbol
Pick from the following elements
for you flashcards:
 Hydrogen
 Lithium
 Sodium
 Potassium
 Rubidium
 Barium
Sulfur
Nitrogen
Carbon
Gold
Silver
Copper
 Magnesium
 Calcium
 Tin
 Helium
 Neon
 Argon
 Krypton
 Florine
 Chlorine
 Bromine
Iron
Nickle
 Iodine
 Oxygen
Cobolt
Homework/Class work
 Create a concept
map on the back
of the element
symbols WS
using the
following terms:
 Atom
 Isotope
 Neutron
 Atomic mass
 Nucleus
 Mass number
 Atomic number
 Proton
 Electron