Chemistry I
Mr. Patel
SWHS
 Learn Major Ions
 Defining the Atom (4.1)
 Subatomic Particles (4.2)
 Atomic Structure (4.2)
 Ions and Isotopes (4.3)
 Nuclear Chemistry (25.1)
 Atom – the smallest particle of an element
that retains its identity
 Can not see with naked eye
 Nanoscale (10-9 m)
 Seen with scanning
tunneling electron
microscope
 Democritus was a Greek to first come up
with idea of an atom.
 His belief: atoms were indivisible and
indestructible. = WRONG!
 Atom comes from “atmos” - indivisible
 2000 yrs later, John Dalton used scientific
method to transform Democritus’s idea into
a scientific theory
 Dalton put his conclusions together into his
Atomic Theory (4 parts)
1. All elements are composted of tiny,
indivisible particles called atoms.
2. Atoms of the same element are identical.
Atoms of different elements are different
3. Atoms of different elements can physically
mix or chemically combine in whole
number ratios.
4. Chemical reactions occur when atoms are
separated, joined, or rearranged. Atoms of
one element can never be changed into
atoms of another element due to a
chemical reaction.
 Particle with negative charge
 Discovered by J.J. Thomson
 Used cathode ray (electron) beam and a
magnet/charged plate.
 Millikan found the charge and mass
 An atom is electrically nuetral
 If there is a negative particle then there must
be positive particle
 Proton – particle with positive charge
 Chadwick discovered neutron – neutral charge
 Electrons distributed in a sea of positive charge
 Plum Pudding Model
 Performed Gold-Foil Experiment
 Beam of Alpha particles with positive charge
shot at thin piece of gold foil
 Alpha particles should have easily passed
through with slight deflection due to positive
charge spread throughout.
 Results: Most particles went straight through
with no deflection. Some were deflected at
large angles.
 The nucleus is the central part of the atom
containing protons and neutrons
 Positive charge
 Most of the mass
 Electrons are located outside the nucleus
 Negative charge
 Most of the volume
 An element is defined only by the number of
protons it contains
 Atomic Number – number of protons
 Number of protons = number of electron
 For a neutral element
1. Zinc (Zn)
1. 30
2. Iron (Fe)
2. 26
3. Carbon (C)
3. 6
4. Uranium (U)
4. 92
 Nucleus contains most of the mass
 Mass Number – total protons and neutrons
Number of neutron = Mass # – Atomic #
1. Lithium
1.
3 p+ , 3 e - , 4 n 0
(MN = 7)
2. Nitrogen
2. 7 p+ , 7 e-, 7 n0
(MN = 14)
3. Fluorine
(MN = 19)
**MN = Mass Number
3. 9 p+ , 9 e-, 10 n0
 Different element: different number of protons
 Ions – same number of proton, different
number of electrons
 Isotope – same number of proton, different
number of neutrons
 Different Mass Numbers
 Nuclear Notation
 Write the element symbol
 On left side, superscript = Mass Number
 On left side, subscript = Atomic Number
 Isotope –Hyphen Notation
 Write full name of element
Hydrogen - 3
 On right side, put a dash
 On right side put Mass Number after dash
 Atomic Mass Unit (amu) – one-twelfth of the
mass of the carbon-12 atom
 Different isotopes have different amu (mass)
and abundance (percentage of total)
 Atomic Mass – weighted average mass of the
naturally occurring atoms.
 Isotope Mass
 Isotope Abundance
 Because abundance is considered, the most
abundant isotope is typically the one with a
mass number closest to the atomic mass.
 Example, Boron occurs as Boron-10 and Boron-11.
Periodic Table tells us Born has atomic mass of
10.81 amu.
 Boron-11 must be more
abundant
 Convert the Percent Abundance to Relative
Abundance (divide by 100)
 Multiple atomic mass of each isotope by its
relative abundance
 Add the product (from step above) of each
isotope to get overall atomic mass.
 Radioactivity – nucleus emits particles and rays
(radiation)
 Radioisotope – a nucleus that undergoes
radioactive decay to become more stable
 An unstable nucleus releases energy through
radioactive decay.
 Nuclear force – the force that holds nuclear
particles together
 Very strong at close distances
 Of all nuclei known, only a fraction are stable
 Depends on proton to neutron ratio
 This region of stable nuclei called band of
stability
 Half Life – the time required for one-half the
sample to decay
 Can be very short
or very long
Symbol
Element
Radiation
Half-Life
Decay
Product
U-238
Uranium-238
alpha
4,460,000,00
0 years
Th-234
Th-234
Thorium-234
beta
24.1 days
Pa-234
Pa-234
Protactinium
-234
beta
1.17 minutes
U-234
U-234
Uranium-234
alpha
247,000 years
Th-230
Th-230
Thorium-230
alpha
80,000 years
Ra-226
Ra-226
Radium-226
alpha
1,602 years
Rn-222
Rn-222
Radon-222
alpha
3.82 days
Po-218
Po-218
Polonium218
alpha
3.05 minutes
Pb-214
Pb-214
Lead-214
beta
27 minutes
Bi-214
Bi-214
Bismuth-214
beta
19.7 minutes
Po-214
Po-214
Polonium-214
alpha
1
microsecond
Pb-210
Pb-210
Lead-210
beta
22.3 years
Bi-210
Bi-210
Bismuth-210
beta
5.01 days
Po-210
alpha
138.4 days
Pb-206
none
stable
(none)
Po-210
Pb-206
Polonium210
Lead-206
 Alpha Radiation (Helium Atom)
 Low penetrating power
 Paper shielding
 Beta Radiation (Electron)
 Moderate penetrating power
 Metal foil shielding
 Gamma Radiation (Pure energy)
 Very high penetrating power
 Lead/concrete shielding
 Transmutation – conversion from one element to
another through a nuclear reaction
 Only occur by radioactive decay
 Only when nucleus bombarded with a particle
 Emissions – given off
 Alpha Emission, Beta Emission, Positron Emission
 Positron = beta particle with a positive charge
 Captures – taken in
 Electron Capture