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GENCHM 113
Session 1.2 A
The Atomic Structure
The Discovery
5 B.C.
1799
1808
Greek philosopher Democritus expressed the belief that matter consists of very small,
indivisible particles, which he named Atomos (uncuttable, indivisible).
Joseph Proust published the laws on definite proportion and multiple proportions.
John Dalton formulated a precise definition of the indivisible building blocks of matter
called atoms.
“Atoms are the basic building blocks of matter.”
John Dalton’s Postulates:
1. Each element is composed of extremely small particles called atoms.
2. All atoms of a given element are identical; the atoms of different elements are different
and have different properties.
Element X
Element Y
3. Atoms of an element are not changed into different types of atoms by chemical reactions;
atoms are neither created nor destroyed in chemical reactions.
Compound X2Y
4. Compounds are formed when atoms of more than one element combined; a given
compound always has the same relative number and kind of atoms.
Laws of Chemical Combination
1. Law of constant composition
In a given compound, the relative numbers and kinds of atoms are constant.
2. Law of conservation of mass
The total mass of materials present after a chemical reaction is the same as the total
mass before the reaction.
3. Law of multiple proportions
If two elements A and B combine to form more than one compound, the masses of B
than can combine with a given mass of a are in the ratio of small whole numbers
The Subatomic Particles
The Electron
Electrons – the negatively-charged particles in cathode rays causing them to be repelled by the
plate bearing negative charges.
1906
Joseph John Thomson used the cathode ray tube and discovered the electron. The
determined the ratio of electric charge to the mass of an individual electron.
Ratio = -1.76 x 108 C/g
The Cathode Ray tube
A
N
Anode
Cathode
B
S
C
+
1917
Voltage
source
Robert Millikan found the charge of the electron.
Charge = -1.60 x 10-19 C
Therefore,
Mass = 9.09 x 10-28 g
Radioactivity
1895
Wilhelm Roentgen discovered the X rays (origin unknown)
1896
Henri Becquerel discovered the spontaneous emission of radiation of a uranium
mineral.
Radioactivity – the spontaneous emission of radiation
Types of Radioactive Rays
Alpha () rays – consist of positively-charged particles.
Beta () rays – consist of negatively-charged particles.
Gamma () rays – have no charge and are not affected by an external field.
The Nucleus and The Proton
1900
J.J.Thomson proposed that the atom consisted of a uniform positive sphere of matter in
which the electrons were embedded.
The plum-pudding model
The electrons are embedded in the
atom much like raisins in a pudding or
like seeds in a watermelon.
1910
Ernest Rutherford disproved Thomson’s model
The Gold Foil Experiment
Gold Foil
Detecting
Screen
-Particle
Emitter
Most of the  particles passed through the gold
foil with little or no deflection.
A few were deflected at wide angles.
Occasionally, an  particle was turned back
The atom’s positive charges are all concentrated in the
Nucleus. Whenever an  particle came close to a
nucleus, it experienced a large repulsive force and
therefore a large deflection. An  particle traveling
directly toward a nucleus would be completely repelled
and its direction would be reversed.
1911
Ernest Rutherford explained his observations
Rutherford’s Postulates
1. Most of the mass of the atom and all its positive charge reside in a very small, extremely
dense region, the nucleus.
2. Most of the total volume of the atom is empty space.
3. Electrons move around the nucleus.
1919
Ernest Rutherford discovered the protons.
1932
James Chadwick discovered the neutrons.
Table 1: Comparison of the Proton, Electron, and Neutron
Particle
proton
neutron
electron
Charge
+1
none
-1
Mass (amu)
1.0073
1.0087
5.486 x 10-4
The Modern View of the Atomic Structure
The protons and neutrons of an atom are packed in
an extremely small nucleus. Electrons are shown
as “clouds” around the nucleus.
Mass Relationship of Atoms
Atomic Number, Z – number of protons (also equal to the number of electrons) in the
nucleus of each atom in the element.
Mass Number, A – total number of neutrons and protons present in the nucleus of an atom
of an element.
Nuclide – way of denoting an element; contains the symbol of the element, the atomic number
and the mass number.
chemical symbol
atomic number
ZX
A
mass number
Isotopes – atoms that have the same atomic number but different mass numbers. Isotopes of
the same element have similar chemistries, forming the same types of compounds and
displaying similar reactivities.
H-1
Protium
1P
0N
H-2
Deuterium
1P
1N
H-3
Tritium
1P
2N
Atomic Mass – mass of the atom in atomic mass units.
Atomic Mass Unit (amu) – mass that is exactly equal to one-twelfth the mass of one carbon12 atom.
Sample Problems:
2.15
How many protons, neutrons, and electrons are in the following atoms:
a. 28Si
b. 60Ni
c. 85Rb
d. 128Xe
2.17
Fill in the gaps in the following table, assuming each column represents a neutral atom:
Symbol
Protons
Neutrons
Electrons
Mass No.
2.23
52Cr
33
42
77
20
20
86
222
193
The element lead (Pb) consists of four naturally occurring isotopes with masses
203.97302, 205.97444, 206.97587, and 207.97663 amu. The relative abundances of
these four isotopes are 1.4, 24.1, 22.1, and 52.4 %, respectively. From these data,
calculate the average atomic mass of lead.