Download 4.1 Early Theories of Matter The Philosophers Democritus – Greek

4.1 Early Theories of Matter
The Philosophers
 Democritus – Greek philosopher (460-370 B.C)
 Proposed atomos
 tiny solid, homogeneous, indestructible and indivisible particles
 different atomos have different sizes and shapes
 changes in matter result from changes in grouping
 Was way ahead of his time
 Aristotle – Greek philosopher (384-322 B.C)
 Couldn’t understand what held atomos together
 Couldn’t grasp empty space
 Was more popular
John Dalton –English Schoolteacher (1766-1844)
 Revived Democritus’ ideas along with research from other prominent
scientists of the time
 Borrowed concept of atomos from Democritus
 Borrowed Law of Conservation of Matter from Antoine Lavoisier
 Borrowed Law of Definite Proportions from Joseph Proust
 Developed Law of Multiple Proportions
 Dalton’s Atomic Theory
 All matter is composed of extremely small particles called atoms
 All atoms of a given element are identical with same size, mass &
properties. Atoms of different elements are different in size, mass &
properties.
 Atoms cannot be created, divided or destroyed.
 Atoms combine in simple whole number ratios to form compounds.
 In chemical reactions, atoms are
separated, combined or
rearranged.
Defining the Atom
 Atom – the smallest particle of an
element that retains the properties
of the element
4.2 Subatomic Particles and the Nuclear Atom
Discovering the Electron
 The Background Discoveries
 Benjamin Franklin – American scientist (1706-1790)
 Discovers positive/negative nature of electrical charges
 Like charges repel – opposite charges attract
 Michael Faraday – English chemist (1791-1867)
 Suggests that atomic structure is related to electricity.
 The Cathode Ray Tube (CRT)






A glass vacuum tube with electrodes at the ends through which a current
is passed
Electricity enters through the cathode and travels across the tube to that
anode
The anode was coated with phosphor which glowed when the electrical
current was flowing
William Crookes – English Physicist (1832-1919)
 Suggested that phosphorescence resulted from rays crossing tube
 Named these rays cathode rays (since they come from the cathode)
Ongoing research with the CRT showed:
 Cathode rays are a stream of charged particles
 These particles carried a negative charge
J.J. Thomson – English Physicist (1856-1940)
 Discovered that the path of electrons could be directed by applying a
magnetic or electric field to the stream
Determined the charge to mass ratio of electrons (1.76 x 108 C/g)
 Found that the mass of these particles was much less than hydrogen
 Discovered first subatomic particles (the electron)
Robert Millikan – American Physicist (1868-1953)



Conducted the “oil drop experiment” in which he determined the mass
of an electron by observing the reaction of negatively charged oil
droplets falling between electrically charged plates.
Mass of electron = 9.11x10-28g (1/1840 of a H atom)
 Raised questions:
o Why matter is neutral if it’s made of negatively charged particles?
o What accounts for the all the mass if electrons are so light?
The Plum Pudding Model
 JJ Thomson proposed a model in which the electrons were distributed
in a uniformly and positively charged atom


The Nuclear Atom
 Ernest Rutherford – New Zealand Scientist (1871-1937)
 Performed “gold-foil experiment”
 Shot alpha particles at a sheet of gold foil surrounded by a screen that
registered alpha particle impacts as flashes of light
 Most particles went straight through as expected
 Some particles bounced back or were significantly deflected

Concluded that the plum pudding model was
wrong due to these observations
 Developed a new model of the atom based on
the concept that the atom is mostly empty space
 Theorized that:
 The atom is mostly empty space
 A tiny very dense region in the center of the
atom caused the deflections which he called
the nucleus
 The nucleus contributed the positive charge
of the atom and nearly all of its mass
 Electrons move rapidly around nucleus, held in proximity by the
attraction of the positive nucleus
 Big ideas here:
 The atom is mostly empty space
 The nucleus accounts for virtually all of the
mass of an atom
 The movement of the electrons accounts for
virtually all of the volume of the atom
(nucleus : electron cloud = 1 : 10 000)
Completing the Atom – The Discovery of Protons and Neutrons
 Protons
 Rutherford continued to work with the nucleus and using research from
Henry Moseley developed the concept of the proton (a positively charged
particle in the nucleus)
 Protons have a +1 charge
 Neutrons
 James Chadwick – English Physicist (1891-1974)
 Accounted for the remaining mass of the atom in 1932
 Showed that the nucleus contained another particle which was called
the neutron
 Determined that the neutron had a mass similar to a proton, but no
charge
 The Modern Atomic Model
 Atoms are:
 generally spherical in shape
 mostly empty space
 comprised of 2 parts
o the nucleus (where the mass is 99.97% of it)
o the electron cloud (where the volume is 99.9999% of it)
 Atoms are composed of 3 primary types of particles: protons, neutrons
and electrons
 Protons
o A nucleon (i.e., located in nucleus)
Positively charged
o Atomic mass of 1 amu
+
o Symbol - (p )
o Are made of smaller particles called quarks (2 up, 1 down)
 Neutrons
o A nucleon (i.e., located in nucleus)
o No charge, zilch, zip
o Atomic mass of 1 amu (slightly heavier than a proton)
0
o Symbol - (n )
o Are made of smaller particles called quarks (1 up, 2 down)
 Electrons
o Located in area around nucleus
o Negatively charged
th
o Atomic mass of 0 amu (1/1840 of a proton)
o Symbol - (e )
o Are fundamental particles – no substructure and belong to a group
called the leptons
Atoms are by definition electrically neutral
 The number of protons = the number of electrons
o

4.3 How Atoms Differ
Atomic Number
 Henry Moseley – English scientist (1887-1915)
 Refined the idea of the nucleus by working with hydrogen
 Discovered that each element contains a unique positive charge that
occurs in a whole number ratio to that of hydrogen
 Developed the idea of atomic number in terms of this positive charge
 The amount of positive charge would later be described in terms of
particles called protons (by Rutherford)
 The elements are now identified by atomic number and placed in the periodic
table in order by the atomic number
 Atomic number = number of protons = number of electrons
Isotopes and Mass Number
 Atoms of a particular element are not necessarily identical
 Atoms of a particular element will always have the same number of
protons and electrons, but the number of neutrons may vary
 Isotopes – Atoms with the same number of protons, but different numbers of
neutrons (every atom is an isotope of its element)
 Example: Isotopes of Hydrogen
+
o
o Hydrogen - 1p , 0n , 1e (Hydrogen-1)
+
o
o Deuterium - 1p , 1n , 1e (Hydrogen-2)
+
o
o Tritium - 1p , 2n , 1e (Hydrogen-3)

Most elements are found as a mixture of isotopes
o The mixture for each element will always be the same
 Isotopes of an element differ in mass because they differ in the number of
neutrons
 Although isotopes of an element have different numbers of neutrons, they
have the same chemical properties because chemical properties relate to
the number and arrangement of electrons
 Mass Number
 Mass number = number of protons + number of neutrons
 To identify different isotopes, the mass number of the isotope is written
after the element’s name (hydrogen-1, hydrogen-2, hydrogen-3)
 The isotopes may also be written in shorthand using the following
notation:
23
2+
Charge
Mass Number
11
Atomic Number
Number
Na
Neutrons = mass number – atomic number
Mass of Individual Atoms
 Atomic Mass
 Masses of atoms are too small to work with using grams
 Scientists instead use a system based on the mass of carbon-12
 Scientists use atomic mass units (amu) or Daltons (Da) to describe the
mass of atoms (or combinations of atoms)
 Carbon-12 is designated as having a mass of exactly 12 amus
th
 An amu is, therefore, 1/12 of the mass of Carbon-12
 Protons & neutrons have masses slightly greater than 1 amu
 Atomic mass ≠ mass number
 Mass number = number of protons + number of neutrons
 Atomic mass = a weighted average of all isotopes of an element
present in nature
 Calculating Atomic Masses
 Multiply the abundance (%) of each isotope by its mass
 Add the contributed masses of each isotope to find the atomic mass

4.4 Unstable Nuclei and Radioactive Decay
Radioactivity
 Nuclear reactions – reactions that involve changes to the nucleus of an atom
 Nuclear reactions may cause one element to change into another, or can
simply change the atomic mass of an element
 Henri Becquerel – French Physicist (1852-1908)
 Accidentally discovered radioactivity while working with uranium salts
 Radioactivity – the phenomenon of rays or particles being produced
spontaneously by unstable atomic nuclei (also referred to as radiation)
 Radioactivity can accompany changes in the identities of elements
 Transmutation – the changing of one element into another
 Unstable nuclei become more stable by releasing energy (radiation)
 Radioactive decay is spontaneous
Types of Radiation
 The nature of radioactive emissions was determined by observing its
deflection as it passed through charged plates

Alpha Radiation (α)
 Highly stable chunks (particles) that break away from an unstable nucleus
 Alpha particles are:
 Helium nuclei
 Made of 2 protons and 2 neutrons (helium without electrons)

Beta Radiation (β)
 High energy, high speed electrons given off when a neutron turns into a
proton in an atom’s nucleus
 Beta particles:
 Have a -1 charge


Gamma Radiation (γ)
 Electromagnetic waves (rays) that usually accompany alpha & beta
radiation
 Gamma rays:
 Have no mass
 Have no charge
 Account for most of the energy in a decay process
 Are extremely hazardous
Nuclear Stability
 Nuclear stability has to do with the number of protons and neutrons in a
nucleus. Not all combinations are stable.
 Protons repel other protons in the
nucleus
 Neutrons help hold the nucleus
together (they’re the glue)
o Neutrons (and protons) have a
force called the strong nuclear
force that pulls them together
o The more protons that are
present in the nucleus, the more
unstable it is and the more
neutrons are required to hold it
together
 Nuclei of low atomic numbers with a
1:1 neutron-proton ratio are very stable.
 The most stable nuclei tend to contain
an even number of both protons and
neutrons.
 Magic numbers (very stable): 2, 8, 20, 28, 50, 82, and 126
 Nuclear Reactions
 Nuclear changes (reactions) may be represented by nuclear equations
 Mass numbers must balance on either side of the equation
 Charge numbers must balance on either side of the equation
 Alpha Decay
226
222
4
Ra →
Rn + α
88
86
2

Beta Decay
131
131
0
I →
Xe + β
53
54
-1

Gamma Decay
238
234
4
0
U→
Th + α + γ
92
90
2
0
Supplement
Radioactive Decay Rates
 Radioactive decay is measured in terms of half-lives
 Half-life = time required for ½ of the nuclei to decay
 Half-lives may range from microseconds to billions of years
 The amount of radioisotope remaining can be calculated using:
Amount remaining = (initial amount)(1/2)n
 May be used to determine the age of objects
Amount of Iodine-131 (g)
20
Half-life of iodine-131 is 8 days
15
1 half-life
10
2 half-lives
5
3 half-lives
4 half-lives
etc…
0
0
8
16
24
32
Time (days)
40
48
56
Nuclear Energy
 The Law of Conservation of Mass
 Mass can be neither created nor destroyed in a chemical reaction.
 The Law of Conservation of Energy
 Energy can be neither created nor destroyed in a chemical reaction.
 Mass-Energy Equivalence
 Nuclear reactions are NOT chemical reactions.
 Mass and energy may be converted back and forth through nuclear
reactions
E = mc2
Where E = change in energy (J), m = change in mass(kg) and c = the
speed of light (3.00 x 108 m/s)
A small change in mass that occurs when the nucleus changes releases a
huge amount of energy (mass defect)


Fission

Heavy nuclei would be more stable if they broke down into smaller pieces
(with mass numbers closer to 60)
 The splitting of the nucleus is called fission
 Fission is used in nuclear power plants to generate heat and thereby
energy
 When an unstable nucleus is struck by a free neutron it may split, causing
other neutrons to be released
 These additional neutrons cause other nuclei to split
 This causes a cascade effect (chain reaction)
 If the sample is large enough to sustain the chain reaction it has
critical mass
Production of heat
Production of electricity
Fusion


Fusion is the combination of smaller nuclei into larger nuclei
Nuclei with mass numbers less than 60 may be combined to form a more
stable nucleus
 Fusion reactions are capable of releasing huge amounts of energy
 The sun is powered by fusion of hydrogen into helium and other larger
elements