Download The Atom: History and Structure

THE ATOM
Unit 5
Chemistry
Langley
ATOMIC TIMELINE
THE FIRST ATOMIC
SCIENTIST
 DEMOCRITUS (400 BC)
 First scientist to hypothesize that all matter was
made of small particles
 Named the small particles atoms
 Atoms in Greek means indivisible
 Democritus thought these particles could not be
broken down into smaller units, indivisible and
indestructible
 Only theory, did not have “hard” evidence
 His ideas were challenged by Aristotle and Plato
 Aristotle (360 BC) argued that atoms did not exist
FOUNDATIONS
 LAW OF CONVERSATION OF MASS
 Antoinne Lavoisier clearly formulated it in
1789
 Mass is neither created nor destroyed
 LAW OF DEFINITE PROPORTIONS
 Joseph Proust
 Specific substances always combine in the same
ratio; i.e. a chemical compound always contains
the same proportion of elements by mass. (Water
is always 2 hydrogen and 1 oxygen)
 Experiments conducted between 1797 and 1804
 What John Dalton used beginning in 1803 to formulate
the basis of his atomic theory
FIRST ATOMIC THEORY
 JOHN DALTON (1808)
 English scientist who investigated Democritus claim
of atoms
 Dalton created the first “accepted” Atomic Theory by
using experimental methods
 All matter is made up of small particles called atoms
 Atoms of a given element are identical in size and mass
 Atoms of different elements can physically mix together or
chemically combine to form compounds
 In reactions, atoms combine, separate, or rearrange
 Atoms cannot be divided, created, or destroyed
MODERN ATOMIC THEORY
 2 KEY CHANGES TO DALTON’S
THEORY
 Atoms are divisible into smaller particles
 Given elements can have atoms of different
masses
PARTICLES IN THE ATOM
PARTICLES IN THE ATOM
 DISCOVERY OF THE ELECTRON
 J.J. Thomson (1897)
 Created an experiment to test sending electric currents
through gases at low pressure
 Cathode Ray Experiment
 One end, anode, became positively charged.
 Other end, cathode, became negatively charged.
 Once charged, a glowing beam appeared that traveled
from the cathode to the anode (cathode ray).
 Since the atoms are suppose to be neutral, then why
would they atoms be attracted to the positive plate?
 Hypothesized that there was some negative particle in
the atom that was attracted to the positive plate
 Named the negative particle a corpuscle
PARTICLES IN THE ATOM
 DISCOVERY OF THE ELECTRON
 J.J. Thomson
 With the discovery of the electron, developed the
Plum Pudding Model
 Pictured plum like negatively charged electrons
embedded in a sphere of positively charged pudding
like “goo”
PARTICLES IN THE ATOM
 CHARGE AND MASS OF THE
ELECTRON
 Robert Millikan conducting the infamous Oil
Drop Experiment in 1909 (and again in 1913
due to infighting with another physicist) was
able to calculate the specific mass of a
single electron and prove that it does in fact
have a negative electric charge
 Mass of an electron is 9.109 x 10-28 kg
PARTICLES IN THE ATOM
 DISCOVERY OF THE PROTON
 Eugen Goldstein (1886)
 Observed a cathode-ray tube and found that rays
could also travel in a tube toward the cathode
(negative end)
 Since opposite attracts, there had to be a positive
particle in the atom, called these things canal
rays
 Positive particles are protons and have a mass
about 1840 times larger than an electron
 (Sometimes because his colleagues did not
agree with him, he is often not given credit with
discovering the proton)
PARTICLES IN THE ATOM
 DISCOVERY OF THE NEUTRON
 James Chadwick (1932)
 Found high energy particles with no charge and
roughly the same mass as a proton (neutron is
slightly larger)
 Neutron has no charge
THE CENTER OF THE
ATOM
 THE NUCLEUS
 Ernest Rutherford (1911)
 Gold Foil Experiment
 A piece of gold foil was bombarded with alpha particles
(positively charged)
 Assumed the positively charged and electrons were
evenly distributed throughout the gold (Plum Pudding
Model)
 Of 8,000 alpha particles fired, 7,999 “stuck” to the foil
(because opposite attract and the + alpha particles
were attracted to the – electrons in the Au atoms
 1 in 8,000 particles bounced back. These alpha
particles must have found the protons in the Au atoms
THE CENTER OF THE
ATOM
 THE NUCLEUS
 Ernest Rutherford
 Gold Foil Experiment showed only a small amount of
space occupied by positive particles
 Concluded that almost all the mass and positive charge
of an atom is located in the center (the nucleus) and the
electrons are just scattered around and orbit the nucleus
like the planets orbit the sun
 Later found that nucleus contains protons (positive
particles) and neutrons (particles with no charge)
FORCES IN THE NUCLEUS
 Nucleus held together by short range
nuclear forces
 Short range forces include:
 Proton-Proton
 Proton-Neutron
 Neutron-Neutron
ACCEPTED THEORY
TODAY
 Based on Wave Mechanics
 Center of the atom is called the nucleus
 Contains the protons and the neutrons
 Electrons “float” in the electron cloud
 Electron cloud broken up into energy levels
 Electrons’ distance from the nucleus depends on their energy
 All matter is made up of atoms
 Atoms from different elements combine to form
compounds in chemical reactions; can also physically
combine
 Atoms rearrange, separate, combine in chemical
reactions
IDENTIFYING ATOMS
 Atomic NumberNumber of protons in
the nucleus of an atom of that element
 Periodic table is arranged according to
increasing atomic number
 The atomic number identifies the element
 Because atoms are electrically neutrally, the
number of protons equals the numbers of
electrons
IDENTIFYING ATOMS
 Mass NumberTotal number of protons and
neutrons
 IS NOT ON THE PERIODIC TABLE!!!!!!!!
 Atoms of the same element do NOT have to have
the same number of neutrons (isotopes)
 Isotope: atoms of the same element that have
different masses due to different numbers of
neutrons
 Number of neutrons = mass # - atomic #
 Example: Carbon with a mass number of 15
IDENTIFYING ATOMS
 Average Atomic Mass
 Since not all atoms of the same element have
the same mass (because of the existence of
isotopes), the mass on the periodic table (the
decimal number) is the average atomic mass
 This mass is given in atomic mass units (amu)
 To find the average atomic mass: identify all
known isotopes of an element and record their
masses and determine the average
IDENTIFYING ATOMS
 Average Atomic Mass
 Examples:
 Hydrogen
 Formula: (%abudance/100)(mass)+(%abundance/100)(mass)
 Neon-20 has a mass of 19.992 amu and Neo-22 has a
mass of 21.991 amu. In an average sample of 100 Neon
atoms, 90 will be Neon-20 and 10 will be Neon-22.
Calculate the average atomic mass.
NOMENCLATURE
Mass #
17
o
Chemical Symbol
8
O - 17
Atomic #
Mass #
MOLE CONVERSIONS
WITH COMPOUNDS
 MOLE REVIEW
 One mole of ANY element is equal to
6.02x1023 atoms of that element
 The mass (in grams) of one mole of any
element is found on the periodic table
underneath the symbol
 Example: 45 grams of Na = ? Moles
 Example: 7.8x1045 atoms of K = ? g of K
MOLE CONVERSIONS
WITH COMPOUNDS
 56 grams of Al2O3 = ? Moles of Al2O3
 Molar Mass of a Compound
 Determine how many of each element is present
in the compound
 Multiply the mass of the element times the
number of that element present
 Add the masses together
 Convert using the fence method
 45 moles of H2O = ? of H2O
 9.4x1035 atoms of AmO = ? g of AmO
NUCLEAR RADIATION
 DISCOVERY
 Henri Becquerel
 Uranium salts and photographic plates
 Conducted experiment twice, once on a sunny day and
once on a cloudy day
 Marie Curie
 Named the process by which particles give off rays
radioactivity
 Spent entire scientific career working on advancements
within nuclear radiation field
 Won two Nobel Prizes for work; one was shared with
husband and Becquerel, one was independently wond
RADIATION DECAY
 RADIOACTIVE DECAY
 When a nucleus spontaneously disintegrates
into a lighter, more stable element
 When this happens, x rays and radiation are
given off
 NUCLEAR STABILITY
 In nuclear reactions, the nuclei of unstable
isotopes (radioisotopes) gain stability by
undergoing changes. These changes will
continue until stability is reached.
RADIATION DECAY
 6 TYPES OF RADIOACTIVE DECAY






Alpha Decay
Beta Decay
Positron Decay
Gamma Decay
Proton
Neutron
 The neutron-to-proton ration determines the
type of decay that occurs
RADIATION DECAY
TYPE
SYMBOL
Alpha (helium nucleus)
a, 4
He
2
Beta (Electron)
b, 0
e
-1
Positron (particle with the mass of an
electron but with a positive charge)
0
Gamma (high energy, electromagnetic)
g
Proton
1
e
+1
p
0
Neutron
1
n
0
TYPES OF RADIATION
(NUCLEAR REACTIONS)
 3 MAIN TYPES
 Alpha Radiation
 Consists of a helium nuclei (alpha particle) emitted from a
radioactive source; alpha particle emitted which contains
two protons, two neutrons and has a double positive
charge
 Beta Radiation
 An electron resulting from the breaking apart of a neutron
in an atom; the neutron breaks apart into a proton. All that
remains is a nucleus; the electron is released (beta
particle)
 Gamma Radiation
 High-energy photon; electromagnetic. Nuclei often emit
gamma rays along with alpha or beta particles during
radioactive decay; (gamma ray-particle) no mass or
electrical charge
NUCLEAR REACTIONS
 EXAMPLES
 32
P 
15
__
 14
C 
6
___
14
___ +
N
___
7
 176
1
__
28
__ +
Ra + n
88
0
Al
13
___

___
___
NUCLEAR REACTIONS
 TRANSMUTATION
 The conversion of an atom of one element to an
atom of another element (via changing the number
of protons in the element)
 Two ways it occurs
 Radioactive decay
 Particles bombard the nucleus
 Allows chemists to produce elements that do not
occur naturally
 These elements have atomic numbers greater than 92
 Earliest transmutation took place in 1919 by Rutherford;
took Nitrogen-14 and formed an unstable isotope of
Fluorine (Fluorine-18); it was this experiment that led to the
discovery of the proton
NUCLEAR STABILITY
 NUCLEAR STABILITY
 Helium has 2 protons, 2 neutrons, 2 electrons
 Atoms only made of protons, neutrons, electrons
 To find the mass of one atom Helium, add up
masses of the 3 particles





Proton = 1.007 amu
Neutron = 1.009 amu
Electron = 0.001 amu
2(1.007) + 2(1.009) + 2(0.001) = 4.033 amu
If you put that same atom on the scale, the mass is only
4.022 amuMass Defect
NUCLEAR STABILITY
 MASS DEFECT
 Difference between adding up the masses of all the
particles in an atom and the mass of the actual atom
 The EINSTEIN CONNECTION
 E = mc2
 Mass moving at high speed can be converted to
energy
 Mass defect comes because some of the mass is
changed into energy
 NUCLEAR BINING ENERGY
 Energy released when nucleus is formed
 The higher the energy, the more stable the atom
NUCLEAR FISSION




Splitting of a nucleus into smaller fragments
Uranium-235 and plutonium-239
Started with neutron bombardment
Continued via chain reaction
 Neutrons produced react with other fissionable atoms,
producing more neutrons which react with still more fissionable
atoms
 Continues until nucleus stability is reached
 Control fission in a nuclear reaction by neutron
moderation and nuclear absorption
 Nuclear Power Plants
 High nuclear waste
 Atomic bomb
NUCLEAR FUSION
 Nuclei combine to produce a nucleus of greater
mass
 Produces more energy than fission reactions
 Sun
 High temperatures required and maintained
 Low waste
 H bomb
NUCLEAR REACTIONS
 HALF LIFE
 The time required for one-half of the nuclei
of a radioisotope sample to decay to
products
 After each half-life, half of the existing
radioactive atoms have decayed into atoms
of a new element
NUCLEAR REACTIONS
 HALF LIFE
 Example 1: If Polonium-32 has a half life of 14.3
days and you start with 4.0 mg, how many mg with
you have after 57.2 days?
Step 1: How many half lives with the atom go
through?
(time you wait)/(half life length) = 57.2/14.3 = 4.00
Step 2:
Amount left = starting amount * (1/2)#half lives
Amount left = (4.0)*(1/2)4.00 = ?????
NUCLEAR REACTIONS
 HALF LIFE
 Example 2: The half life of radon-222 is
3.824 days. If you wait for 15 days and find
50 grams of radon-222, how much did you
start with?