Download Chapter 4 Structure of the Atom

Chemistry
Chapter 4 Atomic Struture
Chapter 4 Structure of the Atom
I. Early Theories of Matter (pgs. 87-91)
A. Greek Philosophers
1. Democritus & Leucippus (500 BC) all matter is made up of tiny particles (atomos - indivisible) - not based
on experimental evidence; based on thinking and reasoning
2. Aristotle (400 BC) Matter is continuous and can be divided infinitely. Matter is made up of four elements
(Earth, wind, fire & water).
B. Alchemists started experimentation.
C. Lavoisier & Proust (1770)
1. Law of Conservation of Matter – (Lavoisier)
• importance of measurement!
• matter can not be created or destroyed
• The mass of the products equals the mass of the reactants
2. Law of Definite Proportions (Proust)
• a compound contains the same elements in exactly the same proportions by mass regardless of size
of the sample or source of the compound.
3. Law of Multiple Proportions
• the mass ratio for one of the elements that combines with a fixed mass of the other element can be
expressed in small whole numbers
• pairs of elements can form more than one type of compound
II. Atomic Theory - Dalton 1807 Atomic Theory of Matter
made atomic theory quantitative, elements have own atoms, shape
and size and mass
1. All matter is made of indivisible and indestructible atoms.
2. Atoms of a given element are identical in the physical and chemical properties
3. Atoms of different elements have different physical and chemical properties
4. Atoms of different elements combine in simple, whole-number ratios to form chemical compounds.
5. Atoms cannot be subdivided, created, or destroyed when they are combined, separated, or rearranged in
chemical reactions.
III. Modern Atomic Theory
1. Atoms have detailed sub-structure
2. Atoms can be changed from one element to another, but not by chemical reactions (nuclear only)
3. Atoms of the same element can have different masses, but their average mass is the same.
Periodic table was based on average masses
IV. Finding Subatomic Particles (pgs. 92-107)
A. Radioactivity - the phenomenon of rays being produced spontaneously by unstable atomic nuclei
B. x-rays & radioactivity
1. Wilhelm Roentgen
a. shot cathode rays at metal and radiation was given off that would darken photographic film
b. the rays could pass through, flesh, but not bone to expose
photo plate
c. rays were unknown, so they were given the name "x-rays"
2. Becquerel - obtained a sample of uranium ore (pitchblende)
a. he noticed uranium also darken photographic paper
b. Radioactive Decay
c. Transmutation - change of one element into another
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Chemistry
Chapter 4 Atomic Struture
3. Marie Curie - noticed polonium and radium also gave off radiation
a. radiation of radium is two million times greater than uranium
b. it conducted electricity in air "ionized air"
1. ionizing radiation has sufficient energy to change atom and molecules into ions
2. non-ionizing radiation does not, such as radio waves
C. Electrons
1897 (found by studying electric current)
1. Glass Discharge Tube (crookes tube (CRT) 1870)
2. J.J. Thomson - Plum Pudding Model
• showed deflection of rays with magnet,
amount of deflection was dependent upon
the particles charge and mass
• determined discrete particles bearing a
negative charge (electrons)
• ratio of charge to mass determines
deflection
charge (e) = 1.76x108 Coulombs/gram
mass (m)
Coulomb = unit of electric charge
J.J. Thomson constructed a Cathode Ray Tube like the one above,
filled with gas at very low pressure. When high voltage is applied to
the electrodes, a glowing beam is projected toward the fluourescent
screen, creating a pinpoint glow. The position of the glow can be
deflected when a magnet is placed near the tube.
3. Millikan's Oil Drop Experiment (1909)
• tried to determine the charge of a single electron
• determined charge, then used Thompson's ratio
to determine mass
e = 1.6022x10-19 C
mass = 9.1096x10-28 g
4. Rutherford - 1911
• gold foil experiment
• bombarded thin film of foil
• 90% of particles pass with
no deflection
• atom was not solid, mostly
empty space
Rutherford’s Gold Foil
Experiment
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Chemistry
Rutherford’s Gold Foil Experiment
Chapter 4 Atomic Struture
Problems with Rutherfords Model?
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D. Subatomic Particles & the Nuclear Atom
1. The Nucleus (discovered by Rutherford)
• contains positively charged protons and neutral neutrons
• VERY dense! If a nucleus were the size of the dot in the
exclamation point at the end of this sentence, its mass would
be approximately as much as that of 70 automobiles!
2. Protons (discovered by Rutherford)
• Carries a charge of 1+, equal but opposite that of an electron
• Relative mass is 1 amu (actual mass is 1.673x10-24 g).
• Located in the nucleus
• Refered to as atomic number
some questions arose about why the
electrons don't fall into the nucleus (after
all, opposite charges should attract)
Rutherford said that the motion of the
electrons kept them from falling into the
nucleus, but...
according to the laws of physics,
charged particle moving in a curved path
should give off energy, and if it did, it
should fall toward the nucleus
Niels Bohr proposed a model in which
orbiting electrons don't lose energy
to understand this, we first have to
understand the nature of light and other
electromagnetic radiation
4. Electrons (discovered by Thomson)
• Negative charge (1-)
• Relative mass is 1/1840 amu (actual mass is
-28
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9.115x10 g).
Located in the space surrounding the nucleus
Number is electron is the same as the
number of protons in a neutral atom.
number of protons = number of electrons
Atomic number = number of protons = number of electrons
5. Atomic number
(How atoms differ 4.3 pgs. 98-104)
• Refers to the number of protons
• Determines the elements position on the
3. Neutrons (discovered by Chadwick)
• Neutral particle, carries no electrical charge
• Relative mass is 1 amu (actual mass is 1.675x10-24 g).
• Located in the nucleus
• Number is neutrons can be found by subtracting the atomic
periodic table
Atomic number =
number of protons = number of electrons
number from the mass number.
Number of neutrons = mass number – atomic number
6. Isotopes and Mass Number (How atoms differ 4.3 pgs. 98-104)
• All atoms of a particular element have the same number of of protons and electrons, but neutrons may differ these types of
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elements are called Isotopes.
Mass number = protons
The sum of the number of protons and neutron is the Mass Number.
The atomic mass of an element is the weighted average mass of the isotopes of that element.
39
Neutrons=
93.25% = ?g
40
Neutrons=
5.7302% = ?g
3
41
+ neutrons
Neutrons=
0.127% = ?g