Chapter 4 – Atomic Structure and the Periodic Table
4.1 Defining the Atom and 4.2 Structure of the Nuclear Atom
Democritus (460 BC – 370 BC)
Greek philosopher who believed that atoms were indivisible and indestructible
Dalton’s Atomic Theory (John Dalton 1766-1844)
1. All elements are composed of tiny indivisible particles called atoms.
2. Atoms of the same element are identical. Atoms of different elements have different
properties.
3. Chemical compounds are formed when atoms combine with each other in simple,
whole-number ratios. Law of multiple proportions.
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. Law of conservation of mass.
Dalton’s work, published in 1808, was the basis for modern atomic theory. His model
represented the atom as a simple sphere with no internal structure.
Atoms
http://www.pbs.org/wgbh/aso/tryit/atom/
 An atom is the smallest piece of an element that still has the properties of that
element. Also called “the building blocks of matter.”
o If I take a piece of gold and cut it and cut it and cut it, I finally get to a piece
that’s too small to cut. That’s an atom of gold. It’s still gold because I didn’t
change it chemically, just chopped it down to the smallest possible piece.
o Word “atom” comes from the Greek word atomos, which means uncuttable.
o How small are atoms? To give you an idea, there are 6.02 x 1023 hydrogen
atoms in one gram. As a result of this size, we usually don’t count atoms
separately, but in larger units called “moles.”
o Radii of most atoms is 5x10-11m – 2x10-10m
 Can now view individual atoms using a scanning tunneling microscope.
 Working on the technology to move individual atoms to create new devices.
o http://www-03.ibm.com/ibm/history/exhibits/vintage/vintage_4506VV1003.html
o http://mrsec.wisc.edu/Edetc/nanoquest/atom_manipulation/
o IBM manipulated Xenon atoms on a nickel surface – right here in Almaden
 Atoms have no net electrical charge, they are neutral. Charged atoms are called
ions.
 Parts of an atom: electron, neutron, proton (subatomic particles)
o http://en.wikipedia.org/wiki/Timeline_of_particle_discoveries
o Additional subatomic particles – quarks, bosons, gluons, hadrons…
Particle
Electron
Proton
Neutron
Table 4.1 Properties of Subatomic Particles
Symbol Charge Relative Mass Actual Mass
e-1
9.11 x10-28 g
1/1840  0
p+
+1
1
1.67 x 10-24 g
0
n
0
1
1.67 x 10-24 g
Electrons (e-)
 Negatively charged subatomic particles
 Mass 1/1840 the mass of a proton (H atom)


o Mass does not affect mass of the atom
Discovered by J.J. Thomson in 1897 (Cathode ray tube experiment)
o Electrical current was applied across two wires (called electrodes) in a gas
filled glass tube. A glowing beam of particles formed from the cathode end
(negatively charged electrode) towards the anode (positive electrode).
Called cathode rays.
o Using magnets, discovered that the cathode rays contained negatively
charged particles. Opposite charges attract, like charges repel. Cathode
ray deflected away from the negative magnet towards the positive magnet.
o Magnets can only affect matter, so particle beam must contain mass.
o http://www.aip.org/history/electron/jjhome.htm
Robert A. Millikan in 1909 conducted an oil-drop experiment and determined the
charge of an electron to be 1.602x10-19 Coulomb
Protons (p+)
 Positively charged subatomic particles
o Eugen Goldstein in 1886 observed that there were rays traveling opposite
to the cathode ray that were positively charged –
o Ernest Rutherford in 1919 noticed that hydrogen nuclei appeared when
alpha particles shot into nitrogen gas, though nitrogen gas must contain
hydrogen nuclei and therefore hydrogen nuclei are fundamental particles
made up of a single proton
Neutrons (n0)
 Subatomic particles with no charge
 Mass equal to a proton
 Discovered by James Chadwick in 1932
Atomic Nucleus
 Initially thought that electrons were evenly distributed throughout an atom filled
uniformly with positively charged material.
 Ernest Rutherford experiment in 1911
o Used alpha particles (He2+ atoms, lost 2 e- and have 2+ charge) directed
at thin sheet of gold foil.
o Alpha particles should have passed through the gold with only slight
deflection due to the spread out positive charge.
o Surprise! Most alpha particles passed straight through without deflection.
o Surprise! Some alpha particles bounced off at large angles.
o Proposed that an atom is mostly empty space with the positive charge
concentrated in the center of the atom, called the nucleus
o http://micro.magnet.fsu.edu/electromag/java/rutherford/
 Nucleus is the central core of an atom and contains protons and neutrons.
 Volume of the nucleus is one trillion times less than the volume of the atom, but the
nucleus contains almost all the mass.
 Diameter of an atom is 10,000 to 100,000 times greater than the diameter of the
nucleus.
 Electrons occupy a large region of space centered around a tiny nucleus, which
accounts for the volume of the atom.
In summary, an atom consists of a nucleus made of protons (p+) and neutrons
(n0) that is orbited by electrons (e-).
4.3 Distinguishing Between Atoms
Atomic Number
 Elements are different because they contain different numbers of protons
o The number of protons determines the element
o The number of n0 and e- does not determine what element you have
 The atomic number of an element is the number of protons in the nucleus
o Atomic number identifies an element
o Since atoms are neutral, # p+ = # eo The number of protons never changes for an element
o The number of electrons might change for an element (forms an ion)
o Elements on the periodic table are organized according to increasing
atomic number, or increasing number of protons
o Show where students can find the atomic number of an element
Mass Number
 Total number of p+ and n0 in an atom is called the mass number
o Show where students can find the mass number of an element
o Number on chart is atomic mass, not a whole number
o Mass number is a whole number
 Number of neutrons can be determined by…
o Number of n0 = mass number – atomic number
 Shorthand notation for an element
A
X = Element symbol from the periodic table
Z X
A = Mass Number (# p+ + # n0)
Z = Atomic Number (# of p+)
 Usually abbreviate the names of the element by saying the name of the element
followed by its atomic mass, “carbon-12.”
Emphasize differences between atoms and ion in terms of protons, neutrons, and
electrons
Isotopes
 Isotopes are atoms of the same element that have different masses. The number of
protons and electrons is the same for all isotopes of an element, but the number of
neutrons is different, causing each isotope to have a different atomic mass. Atomic
number does not change.
o Hydrogen-1 has no neutrons, mass number of 1
o Hydrogen-2 (deuterium) has one neutron, mass number of 2
o Hydrogen-3 (tritium) has two neutrons, mass number of 3
 Isotopes are chemically alike because they have identical numbers of protons and
electrons which are the subatomic particles responsible for chemical behavior
 Why can there be more than one possible number of neutrons in an atom? For many
atoms, there can be several different numbers of neutrons that serve to stabilize the
positive charge in the nucleus.
Atomic Mass Or Atomic Weight
 Atomic mass is measured in atomic mass units (amu), defined as 1/12 the mass of
a carbon-12 atom. (Carbon-12 is the standard reference.)
 1 amu = 1.67 x 10-24 g
 Atomic mass is not always a whole number because of the relative abundance of the
naturally occurring isotopes of the element.
 Most elements occur as a mixture of two or more isotopes
 Each isotope of an element has a fixed mass and natural percent abundance
 The atomic mass of an element is the weighted average mass of the atoms in a
naturally occurring sample of the element.
 How to calculate atomic mass
o Need the number of stable isotopes of the element
o The mass of each isotope
o The natural percent abundance of each isotope
 Example:
Chlorine-35 occurs 75.77% (34.969 amu)
Chlorine-37 occurs 24.23% (36.966 amu)
Convert percentages into decimals (75.77% = 0.7577)
(0.7577 x 34.969) + (0.2423 x 36.966) = 35.453 amu
Closer to the mass of Chlorine-35 since that’s the more abundant isotope.