1. Mr. Rumwald Leo G. Lecaros
School of Chemical Engineering and Chemistry
Mapua Institute of Technology
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2. 1) Dalton’s Atomic Theory
2) Atoms and Molecules
3) Structure of the atom
4) Subatomic particles
5) Atomic Number, Mass Number and Isotopes
6) Ions and Ionic Compounds
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3.  Matter is composed of atoms.
◦ Atoms
 Composed of electrons, protons and neutrons.
◦ Molecules
 Combination of atoms.
◦ Ions
 Charged particles
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4.  John Dalton -1808
 Postulates
1) An element is composed of extremely small, indivisible
particles called atoms.
2) All atoms of a given element have identical properties
that differ from those of other elements.
3) Atoms cannot be created, destroyed, or transformed
into atoms of another element.
4) Compounds are formed when atoms of different
elements combine with one another in small whole-
number ratios.
5) The relative numbers and kinds of atoms are constant
in a given compound.
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5.  Law of conservation of mass
◦ Matter is conserved in chemical reactions
 Law of constant composition
◦ The elements that a compound is composed of are
present in fixed and precise proportion by mass.
 Law of multiple proportions
◦ When the same elements can form two different
compounds, the ratio of masses of one of the
elements in the two compounds is a small whole
number relative to a given mass of the other
element.
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6.  Dalton’s atom had no features other than
mass.
 Many discoveries have demonstrated that the
atom is not featureless or indestructible, but
is composed of other parts. (subatomic parts)
 J.J. Thomson and Ernest Rutherford
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7.  First evidence for subatomic particles came
from the study of the conduction of electricity
by gases at low pressures.
◦ J.J. Thomson, 1897
◦ Rays emitted were called cathode rays
◦ Rays are composed of negatively charged particles
called electrons.
◦ Electrons carry a negative charge (-1) and have a
very small mass (1/2000 the lightest atomic mass)
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8. 8

9.  J.J. Thomson modified the cathode ray tube in
1897 by adding two adjustable voltage
electrodes.
◦ Studied the amount that the cathode ray beam was
deflected by additional electric field.
 Thomson used his modification to measure
the charge to mass ratio of electrons.
◦ e/m = -1.75881 x 108 coulomb/g of e
◦ He named the cathode rays electrons
◦ “discoverer of electrons”
◦ TV sets and computer screens
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10.  Robert A. Millikan won the 1stAmerican Nobel
Prize in 1923 for his famous oil-drop
experiment.
 In 1909 Millikan determined the charge and
mass of the electron.
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11.  Millikan determined that the charge on a
single electron = -1.60218 x 10-19coulomb.
 Using Thomson’s charge to mass ratio we get
that the mass of one electron is 9.11x 10-28g.
◦ e/m = -1.75881 x 108coulomb
◦ e = -1.60218 x 10-19coulomb
◦ Thus m = 9.10940 x 10-28g
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12.  Eugene Goldstein noted streams of positively
charged particles in cathode rays in 1886.
◦ Particles move in opposite direction of cathode
rays.
◦ Called “Canal Rays” because they passed through
holes (channels or canals) drilled through the
negative electrode.
 Canal rays must be positive.
◦ Goldstein postulated the existence of a positive
fundamental particle called the “proton”.
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13. 13

14.  Ernest Rutherford directed Hans Geiger and
Ernst Marsden’s experiment in 1910.
◦ α-particle scattering from thin Au foils
◦ Gave us the basic picture of the atom’s structure.
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15.  Ernest Rutherford, 1911
 Bombardment of gold foil with α-particles
(helium atoms minus their electrons
◦ Expected to see the particles pass through the foil
◦ Found that some of the alpha particles were
deflected by the foil
◦ Led to the discovery of a region of heavy mass at
the center of the atom
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16.  Rutherford’s major conclusions from the -
particle scattering experiment
1. The atom is mostly empty space.
2. It contains a very small, dense center called
the nucleus.
3. Nearly all of the atom’s mass is in the
nucleus.
4. The nuclear diameter is 1/10,000 to
1/100,000 times less than atom’s radius.
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17.  James Chadwick in 1932 analyzed the results
of α-particle scattering on thin Be films.
 Chadwick recognized existence of massive
neutral particles which he called neutrons.
◦ Chadwick discovered the neutron.
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18. 18
 Solar system depiction
of atomic structure.
◦ Emphasizes proton,
neutron and electron
distribution; does not
accurately depict
current accepted model
of atomic structure.

19. 19
 Electrons are depicted
as clouds of negative
charge surrounding
the nucleus.
◦ The density of the
small dots is related to
the probability of
finding an electron at a
particular location.

20. 20
Particle mass (amu) charge
Proton 1.007 +
Neutron 1.009 0
Electron 0.00055 –

21.  Information regarding atomic structure is
written in scientific shorthand called the atomic
symbol.
◦ E is the atomic symbol for element
◦ Superscript A is the mass number.
◦ Subscript Z is the atomic number.
Z
A
E

23.  Entry for carbon on
the periodic table.
◦ Z = 6
◦ Relative atomic mass
= 12.011 (~99% carbon-
12)
◦ Element Symbol: C
6
12
C

24.  Isotopes are atoms of an element that differ
in the number of neutrons in their nucleus.
◦ same Z but different A
 Isotopic abundance is the mass percentage of
an isotope in a naturally occurring element.

25.  Mass spectrometers can measure the masses of
atoms, isotopes, and molecules.

26.  Mass spectrum showing carbon isotopes.

27.  One example of an isotopic series is the
hydrogen isotopes.
◦ 1H or protium is the most common hydrogen
isotope.
 one proton and no neutrons
◦ 2H or deuterium is the second most abundant
hydrogen isotope.
 one proton and one neutron
◦ 3H or tritium is a radioactive hydrogen isotope.
 one proton and two neutrons
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28.  1H, 2H, 3H
◦ Hydrogen, deuterium, tritium
◦ Different masses
 Note that some of the ice is at the bottom of
the glass –this is 2H2O
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29.  Relative atomic mass for an element is an
average of the atomic masses for the
naturally occurring isotopes for an element.

30.  Naturally occuring Cu consists of 2 isotopes.
It is 69.1% 63Cu with a mass of 62.9 amu, and
30.9% 65Cu, which has a mass of 64.9 amu.
Calculate the atomic weight of Cu to one
decimal place.
Atomic weight = Σ (relative abundance) (isotopic mass)
Atomic weight = (0.691)(639 amu) + (0.309)(649 amu)
Atomic weight = 63.5 amu

31.  Naturally occurring argon consists of three
isotopes, the atoms of which occur in the
following abundances: 0.34% 36Ar
(35.9676amu), 0.07% 38Ar (37.9627), and
99.59% 40Ar (39.9624). Calculate the atomic
weight of argon from these data.
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Atomic weight = 39.95 amu

32.  Ions are formed when the number of protons
and electrons in an atom are not equal.
◦ Ions with more protons than electrons are called
cations.
 net positive charge
◦ Ions with more electrons that protons are called
anions.
 net negative charge

33.  A monatomic ion is derived from a single
atom.
 A polyatomic ion is derived from a group of
atoms with an overall charge.
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35.  Groups of atoms may carry a charge; these
are the polyatomic ions
◦ OH-
◦ NH4+
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36.  An element and its ion have the same
chemical symbol but different properties.
◦ Sodium metal atoms lose an electron to form
sodium cations.
 Sodium metal reacts violently with water.
◦ Chlorine gas molecules gain electrons to form
chlorine anions (chloride).
 Chlorine gas reacts violently with sodium metal.
◦ Ionic compounds containing sodium cation and
chlorine anion dissolve in water without reacting.

37.  Determine the number of protons, electrons,
and neutrons in:
 (a)
 (b)
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