Unit 3 and Ch. 4 Notes
Atomic Structure and the Periodic Table
History of Atomic Structure:
400 B.C. Democritus – a Greek teacher suggested the existence of atoms
Nearly 2000 years later, finally experimental evidence was available to support the existence of atoms
In the early 1800’s, John Dalton – an English school teacher performed experiments and came up with Dalton’s
Atomic Theory
1897 – Thomson – discovered the e- using a cathode ray tube
1911 – Rutherford – discovered the nucleus and its charge
1913 – Bohr – developed the atomic model that shows electrons in arranged in
orbits around the nucleus
1926 – Schrodinger – developed the quantum mechanical or wave model of the
atom which uses mathematical probability to determine the
location of an electron within an electron cloud
1986 – Swiss scientists Binnig and Rohrer invented the scanning tunneling microscope, which took the first
picture of individual atoms
Dalton’s Atomic Theory:
1. All elements are composed of tiny indivisible particles called atoms.
Indivisible – means cannot be divided
Incorrect - because atoms are not indivisible, they can be divided by nuclear reactions only, but lose their
characteristic properties in the process.
2. Atoms of the same element are identical.
Incorrect – because of isotopes that differ in the number of neutrons
3. Atoms of different elements combine in simple whole number ratios to form compounds. Like
H2O or H2O2
4. Chemical (not nuclear) reactions occur when atoms are separated (bonds broken), joined (bonds
made) and rearranged. Atoms of one element cannot be changed into atoms of a different element
by chemical reactions.
The second sentence is true for chemical reactions but not for nuclear. Radioactive decay of Uranium
turns it into Lead over time through a natural process.
The second sentence also originated from the field of Alchemy, which actually gave rise to the field of
Chemistry. In early times, alchemists were seers, prophets, sorcerers, etc. that tried to turn common
materials into gold. Even though unsuccessful, they learned much about chemistry along the way.
Dalton’s atomic theory is still very relevant to chemical reactions today. It does not apply, however, to
nuclear reactions.
Just how small is an atom?
There are something like 15,000 Carbon atoms in a pencil dot.
In one penny, there are 2.4 X 1022 Copper atoms. By comparison, there are only 6 X 109 people on Earth.
If 100 000 000 copper atoms were lined up side by side, they would only form a line 1 cm long.
Remember:
Like charges repel.
Unlike charges attract.
1897 Thomson – an English physicist discovered the e- using a cathode ray tube, it is basically a glass tube
filled with gas with a positive metal plate at one end and a negative metal plate at the other and then connected
to high voltage electricity
1911 – Ernest Rutherford – shot alpha particles (+2 charge) at a sheet of gold foil, most of the alpha particles
went straight through indicating the atom is mostly empty space. However, a few alpha particles (+2) were
repelled by hitting the nuclei of the gold atoms. His 2 main findings were: the discovery of the nucleus and
the positive charge of the nucleus.
1913 – Niels Bohr – a student of Rutherford, proposed the Bohr model of the atom in which electrons are
arranged in orbits around the nucleus and that electrons have a fixed energy
Atom – the smallest particle of an element that still retains the properties of that element
Subatomic Particles – any various units of matter below the size of an atom
Nucleus - the dense central portion of an atom, composed of protons and neutrons
Proton - a positively charged subatomic particle found in the nucleus of an atom
Neutron - a subatomic particle with no charge and a mass of 1 amu; found in the nucleus of the atom
Electron - a negatively charged subatomic particle, in clouds outside the nucleus
Energy Level - a region around the nucleus of an atom where an electron is likely to be moving
Valence electrons - the outermost (highest energy level) electrons involved in bonding
Subatomic Particle
Location
Symbol
Charge
Mass (amus)
Proton
nucleus
p+
+1
1
1.67 X 10 -24
Neutron
nucleus
n0
0
1
1.67 X 10 -24
Electron
outside nucleus
e-
-1
1/1840
Mass (g)
9.11 X 10 -28
amu – atomic mass unit – a unit of mass equal to 1/12 the mass of a carbon-12 atom
The protons and neutrons in the nucleus of an atom accounts for most of the mass of an atom.
The space the electrons are moving in accounts for most of the volume of the atoms.
This is why most of the alpha particles went through the foil in Rutherford’s experiment.
Analogy, if an atom were the size of a professional football stadium, a 2000- ton tennis ball sitting on the center
of the 50 yard line would be the nucleus and the rest of the stadium would be the space the electrons are moving
in.
Atomic number – the number of protons in the nucleus of an element, shown on the periodic table, matches the
number of electrons, identifies the element
Atoms are electrically neutral. p+ = eCompound (cpd.) – a chemical combination of 2 or more elements
Ionic Compound - compound composed of 1 metal and 1 or more nonmetals, a positive ion and a negative ion,
ex. NaCl, Fe2O3, NaHCO3 (salt, rust, baking soda)
Molecular Compound - compound composed of 2 or more nonmetals, ex. H2O, C6H12O6, NH3 (water, glucose,
ammonia)
Ions – are atoms with a + or – charge, due to the loss or gain of electrons
Positive ions have more protons than electrons (have lost electrons)
ex. Na+1 (lost 1 e-) Fe+2 (lost 2 e-), etc.
Negative ions have less protons than electrons (have gained electrons)
ex. Cl-1(gained 1 e-) O-2(gained 2 e-), etc.
Development of the Periodic Table
Mid 1800’s – Dmitri Mendeleev – a Russian chemist developed the first periodic table by arranging the
elements in order of increasing atomic mass. He then put the elements with similar properties in the same
columns. He left blank spaces for elements he believed to exist but had not yet been discovered. He was later
proven correct.
1913 – Moseley – a British physicist developed the modern periodic table by arranging the elements in order of
increasing atomic number
Periods – horizontal rows of periodic table
Groups – vertical columns, elements in a group have very similar physical and chemical properties
Periodic Law – when the elements are arranged according to increasing atomic number there is a periodic
repetition of physical & chemical properties
A – representative elements
B – transition and inner transition metals (misbehavers)
Metals
High electrical conductivity
High luster (metallic)
Ductile (drawn into wire)
Malleable (hammered into thin sheets)
80% of all elements are metals
Nonmetals
Poor conductors
Nonlustrous
Solids are brittle, some liquids, most gases
Metals
Alkali 1A
Alkaline Earth 2A
Transition B’s
Inner Transition B’s
Nonmetals
Halogens F, Cl, Br, I 7A
Noble gases (inert or unreactive) 8A
Boron, Carbon and Silicon 3A & 4A
Metalloids – border the stair step and have properties of both metal and nonmetals
Si – semiconductor used in computer chips
Ge – used in solar cells
Alkali metal - any metal in Group 1A of the periodic table
Transition metal - Group B element characterized by addition of electrons to d suborbitals
Alkaline Earth Metals - any metal in Group 2A of the periodic table
Halogens - any member of the nonmetallic elements in Group 7A of the periodic table
Noble Gases - any member of a group of gaseous elements in Group 18 of the periodic table; the s and p
sublevels of their outermost energy level are filled
Group and Periodic Trends or Periodicity
Group (Vertical)
Period (Horizontal)
Transition metals are often skipped for these trends
Trends Include: Atomic radius, Ionic Radius, Electron Affinity, and Electronegativity
Atomic radius – size of the neutral atom
Group Trend: increases as you go down a group
Periodic Trend: decreases from L to R across a period
Ionic radius – size of an ion
Group Trend: increases as you go down a group
Periodic Trend: decreases from L to R across a period
Electronegativity – the tendency of an atom to attract electrons to itself even when bonded to another element;
F is the most electronegative (reactive) element of all, noble gases are excluded from this trend
Group Trend – decreases as you go down a group
Periodic Trend – increases from L to R across a period
Ionization Energy - the energy required to remove an electron from a gaseous atom
Group Trend – decreases as you go down a group
Periodic Trend – increases from L to R across a period
Isotopes Discussion
# of Protons: determines the identity of the element; and the atomic number
# of Electrons: determines the charge of an atom (if an ion is formed or not)
# of Neutrons: determines the atomic mass and mass number or isotope of the element
Protons = Electrons – the atom is neutral
Protons > Electrons – atom is a positively charged ion (lost electrons)
Protons < Electrons – the atom is a negatively charged ion (gained electrons)
Ion – an element that has lost or gained electrons and has a charge
To gain stability all elements except the noble gases (group 18) either:
lose electrons (metals)
gain electrons (nonmetals)
or
share electrons (nonmetals)
Isotopes – atoms of the same element that have different numbers of neutrons
All atoms of an element are isotopes of that element
Most elements have 1, 2 or 3 naturally occurring isotopes each occurring a certain percent of the time in nature
While there are approximately 100 elements, there are about 1000 different isotopes. This is why mass
numbers and neutron information are not given on the periodic table. There are too many isotopes.
Mass number – the total number of protons and neutrons in the nucleus of an atom, not shown on the periodic
table
atomic mass – a weighted average of the masses of all the isotopes of an element
percent abundance – how often a particular isotope occurs in nature
Hydrogen-1
Hydrogen-2 Deuterium (deuce or 2)
Hydrogen-3 Tritium (tri or 3)
Deuterium and Tritium are future sources of fuel for fusion. Found in “heavy” ocean water. There is enough
deuterium and tritium in our oceans to supply all future energy needs.
Isotopes are designated 2 ways:
1.
Hyphen Notation; element name – mass #
2. Nuclear Symbol
C
For example: the element carbon has three isotopes:
Mass # = protons + neutrons
Carbon-12 (90% abundance)
C
Carbon-13 (9% abundance)
C
Carbon-14 (1% abundance)
C
3 Ways Isotopes of the Same Element Differ:
1.
atomic mass
carbon-12
2. mass #
3. # of neutrons
To calculate average atomic masses:
Multiply the mass of a particular isotope times its percent abundance. Then add together all figures for all the
isotopes of that element.
Boron has two isotopes, Boron-10 and Boron-11. The percent abundance of Boron-10 is 19.91% and for
Boron-11 is 80.09%. Calculate the average atomic mass of Boron.
10X .1991 = 1.991
11 X .8009 = 8.8099 +
10.8009
average atomic mass
Light and Energy Notes
Electromagnetic Radiation – light from the sun that travels through space, exhibits wave-like behavior and
travels at the speed of light (3.00 X 108 m/s)
Speed of Light – the speed at which light travels in a vacuum (space);
3.00 X 108 m/s
Spectroscope – an instrument that measures light given off by a radiant source
3 Characteristics of Waves:
1.
wavelength (λ, unit m or nm) – the distance between two consecutive peaks or troughs in a wave
frequency (f, unit Hz[hertz] cycles/s or s-1) – the number of waves per second that pass a given point in
space
3. speed (3.00 X 108 m/s) in a vacuum
2.
Wavelength and Frequency have an inverse relationship.
As wavelength increases, frequency decreases and vice versa
Know this Order:
1.
2.
3.
4.
5.
6.
7.
Shortwave, FM (Radio Waves)
Microwaves
Infrared
Visible (White) Light
Ultraviolet
X-Rays
Gamma Rays
Photon - a quantum of light; a discrete bundle of electromagnetic energy that behaves as a particle
Light and Energy Formulas