objectives: understand the experimental design … · the atom objectives: understand the...
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THE ATOM Objectives: Understand the
experimental design and conclusions
used in the development of modern
atomic theory, including Dalton’s
Postulates, Thomson’s discovery of
electron properties, Rutherford’s
nuclear atom, and Bohr’s nuclear
atom.
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Democritus
• Made his discovery
around the year 250
B. C.
• This was the first
discovery about the
atom, the next would
come in another 2000
years.
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The First Atom
• Democritus took a sea shell and broke it in
half.
• Than he broke it in half again.
• When the pieces got to small he use a
mortar and pestle to crush the shell.
• He finally believed he got to the smallest
piece possible and called it the ATOM;
which in Greek means INDIVISIBLE.
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John Dalton (1766-1844)
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A New System of Chemical
Philosophy (1808)
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Dalton’s Atom Model
1. All matter is made on atoms; and atoms are indivisible.
2. Atoms of the same element are all identical.
3. Compounds are formed by a combination of two or more different atoms and they always have the same proportion of elements. THE LAW OF DEFINITE COMPOSITION
4. A chemical reaction is a rearrangement of atoms and the atoms are neither created nor destroyed. THE LAW OF CONSERVATION OF MATTER
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J. J. Thomson (1856-1940)
Joseph John Thomson
• English physicist who in 1897 discovered a particle smaller than the atom ; the electron.
• Particle has a negative charge and is much smaller than the atom so must come from the inside of the atom.
• Electrons are scattered around the atom like raisins in pudding. (THE PLUM PUDDING MODEL)
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Thomson and Rutherford
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Rutherford’s Gold
Foil Experiment
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Rutherford’s Gold
Foil Experiment
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Ernest Rutherford (1871-1937)
• New Zealand born physicist; worked in England
• 1911 conducted the “Gold Foil Experiment” the proved the existence of a small positively charged center of the atom.
• Disproved the “Plum Pudding Model”
• THE NUCLEAR MODEL
• Discovered the proton.
• Thought that the electrons orbited the nucleus like planets orbited the sun.
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Millikan’s Oil Drop Experiment
• A fine mist of oil droplets is introduced into the chamber.
• The oil is ionized by x-rays.
• The electrons adhere to the oil drops.
• The value for the charge of the electron can be calculated.
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Niels Bohr (1885-1962)
• Danish physicist,
produced his model in
1911.
• Saw problems with
Rutherford’s model.
• If electrons “orbit” than
they are changing
direction so they are
accelerating.
• That would require
energy.
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The Orbital Model
• Electrons do not
“orbit” but are in
allowable ENERGY
LEVELS.
• When the electrons
stay in these levels,
which are at specific
distances from the
nucleus, they do not
give off energy.
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Bright Line Spectrum
• But, if the electron moves from one level to another it gives off or absorbs energy.
• These Bright Line Spectrums are produced when the electrons “fall back” to a lower energy level and give off energy.
• Every element has a unique Bright Line Spectrum.
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The Subatomic Particles
THE PROTON
• p+
• positively charged
• located in the nucleus
• relative mass = 1 atomic mass unit
• mass = 1.673 x 10-24 grams
• equal to atomic number
• number of protons “defines” the atom
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The Subatomic Particles
THE NEUTRON
• n0
• neutral (no electrical) charge
• located in the nucleus
• relative mass = 1 atomic mass unit
• mass = 1.675 x 10-24 grams
• equal to mass number minus atomic number
• mass number is protons + neutrons
• James Chadwick proposed the existence of the
neutron.
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The Subatomic Particles
THE NEUTRON • Isotopes – different atoms of the same element that
have the same number of protons but different numbers of neutrons
• some isotopes are radioactive – they emit energy when the nucleus of the atom breaks down spontaneously
• most radioactive isotopes are not dangerous
• to determine if an isotope is radioactive calculate the proton to neutron ratio
• if ratio is greater than or less than 1:1 for “small” atoms the isotope is unstable (smaller than Ca)
• if ratio is greater then 1:1.5 for “large” atoms the isotope is unstable
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The Subatomic Particles
THE ELECTRON • e- (negative electrical charge)
• located in the electron cloud which is divided into energy levels, sublevels, orbitals, and spins
• relative mass = 0 atomic mass units
• mass = 9.11 x 10-28 grams
• equal to the number of protons if atom is neutral
• atom becomes a charged ion if electrons are gained or lost
• positive ion = CATION
• formed by the loss of electron, happens to metals
• negative ion = ANION
• formed by the gain of electron, happens to nonmetals
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Location of Electrons
• Energy Levels
• Discovered by Niels Bohr
• # electrons = 2n2
• “n” is the energy level
• 1st level can hold 2 e-
• 2nd level can hold 8 e-
• 3rd level can hold 18 e-
• (eight if the outside energy level)
• 4th level can hold 32 e-
• (eight if the outside)
• The outside level is called the valance level and can never hold more than 8 electrons.
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NUCLEAR SYMBOLS
mass number ion charge
23 +1
Na p+ = 11
11 n0 = 12
atomic number e- = 10
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name symbol atomic
number
mass
number
ion
charge
number
of
protons
number
of
neutrons
number
of
electrons
atomic
mass
calcium 20 42 +2 40.08
19 -1
F
9
10 10 10 20.18
238 0
U
92
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name symbol atomic
number
mass
number
ion
charge
number
of
protons
number
of
neutrons
number
of
electrons
atomic
mass
potassium 19 40 +1 39.098
amu
15 -2
O
8
18 22 18 39.948
amu
56 0
Fe
26
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Objective
• Use isotopic composition to
calculate the average atomic
mass of an element.
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Mass Number vs. Atomic Mass
• mass number is given for an individual atom
• mass number is given in nuclear symbols
• atomic mass is an average mass for all isotopes for the element
• atomic mass is the number on the periodic table
• if you round the average atomic mass you will have the mass number of the most common isotope
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Average Atomic Mass
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