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Atoms and the Atomic Theory

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Title: Atoms and the Atomic Theory


1
Atoms and the Atomic Theory 
  • All the matters can be broken down into elements.
    Is matter continuously divisible into ever
    smaller and smaller pieces, or is there an
    ultimate limit? What is an element made of?
  • Greeks
  • Aristotle- Continuous Theory of Matter
  • Democritus- Discontinous Theory of Matter
  • Atomos- indivisible
  • Early Chemical Discoveries and Atomic
    TheoryThree important fundamental laws in
    chemistry1) Law of conservation of mass2) Law
    of constant composition3) Law of multiple
    proportions

2
Law of Conservation of Mass
  • 1774 Antoine Lavoisier showed heating the red
    power HgO causes it to decompose into the silvery
    liquid mercury and the colorless gas oxygen. 2HgO
    ? 2Hg O2 then show that oxygen is the key
    substance involved in combustion.
  • Furthermore, results of combustion reactions
  • Total mass of products total mass of reactants
  • (tin air sealed glassed vessel) ?? (tin oxide
    remaining air glass vessel)
  • Law of Conservation of Mass
  • The total mass of substances present after a
    chemical reaction is the same as the total mass
    of substances before the reaction. Matter is
    neither created nor destroyed in a chemical
    reaction.

3
E.g. A 0.382g sample of magnesium reacts with
2.652g of nitrogen gas. The sole product is
magnesium nitride. After reaction, the mass of
unreacted nitrogen is 2.505g. What mass of
magnesium nitride is produced? Mass before
reaction 0.382g Mg 2.652g N2 gas
3.034g Mass after reaction ?g Mg3N2 gas
2.505 N2 gas 3.034g 2.505g 0.529g
4
Law of Constant Composition
1799 Joseph Proust Law of Constant Composition
(Definite Proportion) All samples of a compound
have the same composition- the same proportion by
mass of the constituent elements. This means that
the relative amount of each element in a
particular compound is always the same,
regardless of the source of the compound or how
the compound is prepared. E. g. Water is made
up of two elements H and O. The two sample of
water below have the same proportions of the two
elements, expressed as percentages by mass. Every
sample of water contains 1 part hydrogen and 8
parts oxygen by mass. ____________________________
_____________ Sample A Composition Sample
B 10.000g 27.000g 1.119g H H 11.19
3.031g H 8.881g O O 88.81 23.979g O
5
Daltons Atomic Theory
How can the Law of conservation of mass and Law
of constant composition be explain? Why do
element behave as they do? 1803-1808 John Dalton
proposed a new theory of matter. 1. Each
chemical element is composed of minute,
indestructible particles called atoms. 2. All
atoms of a given element are identical to one
another in mass and other properties, but the
atoms of one element are different from the atoms
of other elements.
6
Daltons Atomic Theory
  • 3. Atoms of an element are not changed into atoms
    of a different element by chemical reactions
    atoms are neither created nor destroyed in
    chemical reactions. Chemical compounds are formed
    when atoms combine with each other.
  • If atoms of an element are indestructible, then
    the same remains unchanged. This explains the
    law of conservation of mass.

7
Daltons Atomic Theory
  • 4. In each of their compounds, different elements
    combine in a simple numerical ratio e.g. one
    atom of A to one of B (AB) or one atom of A to
    two of B (AB2).
  • If all atoms of an element are alike in mass
    (assumption 2) and if atoms unite in fixed
    numerical ratio (assumption 3), the percent
    composition of a compound must have a unique
    value, regardless of the origin of the sample
    analyzed. This explains the law of constant
    composition.

8
Law of Multiple Proportions
Daltons theory leads to a prediction- the law of
multiple proportions. If two elements form more
than a single compound, the masses of one element
combined with a fixed mass of the second are in
the ratio of small whole numbers. Same elements
to combine in different ratios to give different
substances.
9
  • E.g. Oxygen and carbon can combine either in a
    1 1.333 mass ratio to make a substance or in a
    1 2.667 mass ratio to make a substance.
  • first 1 g carbon per 1.333 g oxygen CO mass
    ratio 1 1.333
  • second 1 g carbon per 2.667 g oxygen CO mass
    ratio 1 2.667
  • comparison CO mass ratio in second sample
    (1 g C)/(2.667g O) 2 of CO ratios CO mass
    ratio in first sample (1 g C)/(1.333g O)
  • Compare two substances clearly the second
    substance contains exactly twice as much oxygen
    as the first for a given number of carbon. If
    the first oxide has the molecular formula CO then
    the second oxide will be CO2.

10
E.g.There are two compounds, both contain
nitrogen and hydrogen. Compound A contains 1.50g
of N and 0.216g H. Compound B contains 2.00g of
N and 0.144g H. If the formula of compound B is
N2H2, what is the formula of compound A? NH
ratio in A 1.50 0.216g 1.00 0.144 NH
ratio in B 200 0.144 1.00 0.0720 H in A
is (0.144/0.0720 2 ) twice as much in B IF B
is N2H2 then A is N2H4
11
Atomic Mass Ratio
Daltons theory enables us to set up a scale of
relative atomic masses. He cannot measure the
exact mass of atoms but relative mass. E.g.
Consider calcium sulfide, which consists of 55.6
calcium by mass and 44.4 sulfur by mass. Suppose
there is one calcium atom for each sulfur atom in
calcium sulfide. Because we know that the mass
of a calcium atom relative to that of a sulfur
atom must be the same as the mass in calcium,
we know that the ratio of the mass of a calcium
atom to that of a sulfur atom is mass of Ca
atom 55.6 1.25 mass of a S atom 44.4
or mass of a Ca atom 1.25 x mass of a sulfur
atom By continuing in this manner with other
compounds, it is possible to build up a table of
relative atomic masses. We define a quantity
called atomic mass ratio, which is the ratio of
the mass of a given atom to the mass of some
particular reference atom.
12
The Structure of AtomsWhat is an atom made of ?
Discovery of subatomic particle The Discovery of
Electrons 1897 J.J. Thomson- cathode ray
experiment Thomsons experiment involved the use
of cathode-ray tube. When a sufficiently high
voltage is applied across the electrode, an
electric current flows through the tube from
negatively charged electrode ( the cathode) to
the positively charged electrode (the anode).
13
Thomsons Experiment

-
Vacuum tube
Metal Disks
14
Thomsons Experiment

-
  • By adding an electric field

15
Thomsons Experiment

-
  • By adding an electric field he found that the
    moving pieces were negative

16
Thomsons Model
Spherical cloud of positive charge
  • Found the electron
  • Couldnt find positive (for a while)
  • Said the atom was like plum pudding
  • A bunch of positive stuff, with the electrons
    able to be removed
  • established the ratio of mass to electric charge
    for cathode ray m/e -5.6857x10-9
    g/coulomb.

Electrons
17
Millikans Oil-Drop Experiment Mass of Electron
1909 Robert Millikan determined the electronic
charge through a series of oil-drop experiments.
The currently accepted value of the charge of the
e is 1.6022x10-19C. Substituting into
Thomsons mass to charge ratio then gives the
mass of electron as 1/1836( 9.1094x10-28g).
18
X-Ray and Radioactivity
Ernest Rutherford identified two type of
radiation from radioactive materials, alpha (?)
and beta (?). ?-particles (?He2)carry two
fundamental units of positive charge and have
essentially the same mass as He atoms.
?-particles are negatively charged particles
produced by changes occurring within the nuclei
of radioactive atoms and have the same properties
as electrons. A third form of radiation, that
is not affected by an electric field was
discovered in 1900 by Paul Villard. This
radiation, called ?-ray, is not made up of
particles it is electromagnetic radiation of
extremely high penetrating power.   Properties of
the three radioactive emissions discovered
Original name Modern name Mass
(amu) Charge ?-ray ?-particle 4.00 2 ?-ray
?-particle (electron) 5.49x10-4 -1 ?-ray ?-ray
0 0_______
19
1909 Ernest Rutherford Scattering Experiment
used ? particle to study the inner structure of
atoms. directed a beam of ?-particles at a thin
gold foil
Florescent Screen
Lead block
Uranium
Gold Foil
20
Rutherford Expected
  • The alpha particles to pass through without
    changing direction very much
  • WHY?
  • The positive charges were spread out evenly.
    Alone they were not enough to stop the alpha
    particles

21
What he expected
22
Because
23
Rutherford thought the mass was evenly
distributed in the atom
24
Rutherford thought the mass was evenly
distributed in the atom?a particles should pass
through the low density model.
25
What he got
The majority of ?-particles penetrated the foil
undeflected. ? Some ? particles experienced
slightly deflections. ? A few (about one in
every 20,000) suffered rather serious
deflections as they penetrated the foil. ? A
similar number did not pass through the foil at
all, but bounced back in the direction from which
they had come.
26
How he explained it
  • Atom is mostly empty
  • Small dense, positive piece at center
  • Alpha particles are deflected by it if
    they get close enough

27
(No Transcript)
28
The Nuclear Atom Protons and Neutrons
  • 1911 Rutherford explained his results by
    proposing a model of the atom known as the
    nuclear atom and having these features.
  • Most of the mass and all of the positive charge
    of an atom are centered in a very small region
    called the nucleus. The atom is mostly empty
    space.
  • The magnitude of the positive charge is
    different for different atoms and is
    approximately one-half the atomic weight of the
    element.
  • There are as many electrons outside the nucleus
    as there are units of positive charge on the
    nucleus. The atom as a whole is electrically
    neutral.
  •  
  • Rutherfords nuclear atom suggested the existence
    of positively charged fundamental particles of
    matter in the nuclei of atoms- called protons. He
    predicted the existence in the nucleus of
    electrically neutral particles.
  • 1932 James Chadwick
  • verified that there is another type of
    particles in atom called neutron.

29
The Structure of Atoms
  • Therefore
  • Modern picture of an atom, then, consist of three
    types of particles-electrons, protons and
    neutron.
  • Electric Charge Mass
  • Particle SI (C ) Atomic SI
    (g) amu Located
  • Electron -1.602x10-19 -1
    9.109x10-28 5.49x10-4 outside
    nucleus
  • Proton 1.602x10-19 1
    1.673x10-24 1.0073 in nucleus
  • Neutron 0 0
    1.675x10-24 1.0087 in nucleus

30
Size of an atom
  • Atoms are small 10-10 meters
  • Hydrogen atom, 32 pm radius
  • Nucleus tiny compared to atom
  • IF the atom was the size of a stadium, the
    nucleus would be the size of a marble.
  • Radius of the nucleus near 10-15m.
  • Density near 1014 g/cm

31
Conclusion
  • Matter is composed, on a tiny scale, of particles
    called atoms. Atoms are in turn made up of
    minuscule nuclei surrounded by a cloud of
    particles called electrons. Nuclei are composed
    of particles called protons and neutrons, which
    are themselves made up of even smaller particles
    called quarks. Quarks are believed to be
    fundamental, meaning that they cannot be broken
    up into smaller particles.

32
Chemical Elements
  • Atomic number
  • What is that makes one atom different from
    another?
  • Elements differ from one another according to the
    number of protons in their nucleus
  • atomic number (Z) Number of proton in atoms
    nucleus
  • mass number (A) of protons (Z) of
    neutrons (N)

33
Isotopes
  • Contrary to what Dalton thought, we know that
    atoms of an element do not necessarily all have
    the same mass.
  • Isotope- atoms of the same element containing
    different numbers of neutrons and therefore
    having different masses.

34
Isotopes of Hydrogen
35
Mass Spectrometer- The most accurate means of
determining atomic and molecular weights.


36
Mass Spectrum of Elemental Carbon
This small peak represents the relative abundance
of C13 in nature.
When 12C and 13C are analyzed in a mass
spectrometer, the ratio of their masses is found
to be Mass13C 1.0836129 Mass12C Since the
atomic mass unit is defined such that the mass of
12C is exactly 12 amu, then on this same
scale, Mass13C (1.0836129)(12amu) 13.003355
amu
37
Average Atomic Mass
  • When considering atomic masses from the P-Table,
    recall that reported values are actually weighted
    averages of all the naturally occurring isotopes.
  • Average atomic mass ?( of each isotope)(atomic
    mass of each isotope)
  • 100
  • Boron has two isotopes 10B and 11B. They have
    the abundance 18.7 and 81.3 respectively.
    Determine the average atomic mass for Boron.

38
When natural copper is vaporized and injected
into a mass spectrometer, the results shown below
are obtained. Use these data to compute the
average mass of copper. The mass values for 63Cu
and 65Cu are 62.93 amu and 64.93 amu respectively.
Computing Average Mass from Mass Spectrometer
Data Atomic Weight
39
Isotopes and Average Atomic Mass Questions
  • 1. Do either of the following pairs represent
    isotopes of one another?
  • 40K19 and 40Ar18 b. 90Sr38 and 94Sr38
  • 2. The nobel gas Neon, has three isotopes of
    masses, 22, 21and 20. If the isotopes have the
    abundance of 8.01, 1.99 and 90.00
    respectively, what is the average atomic mass of
    neon atoms?
  • 3. A naturally occurring sample of an element
    consists of two isotopes, one of mass 85 and one
    of mass 87. The abundance of these isotopes is
    71 and 29. Calculate the average mass of an
    atom of this element.
  • 4. If 69Ga and 71Ga occur in the s 62.1 and
    37.9, calculate the average atomic mass of
    gallium atoms.

40
Mass Spectrum of Chlorine Molecule
(35Cl-35Cl), (35Cl-37Cl), or (37Cl-37Cl)
41
Ions
Ion an electrically charged particle obtained
from an atom or chemically bonded group of atoms
lose or gain electrons. The charge on an ion is
equal to the of protons minus the of
electrons. An atom that gains extra electrons
becomes a negatively charged ion, called an
anion. An atom that loses electrons becomes
positively charged ion, called a cation.
E.g. Determine numbers of electrons in Mg2
cation and the S2- anion? Mg2 number e
? S2- number e ?
42
Introduction to Periodic Table
  • With discovery of many elements
  • 1869 Mendeleev and Meyer
  • independently proposed periodic table organized
    the elements
  • In modern periodic table, The periodic table of
    the elements is organized into 18 groups and 7
    periods. Elements are represented by one or
    two-letter symbols and are arranged according to
    atomic number.
  • a horizontal row of elements- a period
  • a vertical row of elements- a group or family

43
Periodic Table of Elements
 
 
 
44
It is customary also to divide the elements into
broad categories known as Metals Except mercury
(liquid), metals are solid s at room temperature.
They are generally malleable, ductile, good
conductors of heat and electricity, and have a
lustrous or shiny appearance. Nonmetals
generally have opposite properties of metals
e.g. poor conductors of heat and
electricity. Metalloid (semimetal) is an element
having both metallic and nonmetallic
properties. Or into three groups Main group
elements are those in groups 1, 2 and 13-18. when
form ions, group 1, 2 lose the same e as their
group group 13 lose group -10 group 14-18
gain 18-group . Transition elements from group
3 to 12, and because all of them are metals, they
are also called the transition metals. The of
electrons lost in TM is not related to their
group . Inner transition metals which include
Lanthanides and Actinindes.
45
Nuclear Chemistry
  • Nuclear reactions involve changes that originate
    in the nucleus of the atom.
  • Chemical changes involve changes in the electron
    cloud.
  • Uses
  • 60Co- gamma ray emitter- ionizing radiation for
    treatment of cancerous tumors.
  • 201Thallium stress test of heart muscle
  • Radiocarbon dating? 14C ½ life 5730 years
  • Nuclear power 20 of US electricity production

46
Radioactivity
  • Recall that all atoms of the same element have
    the same number of protons. The number of
    neutrons in the atoms nucleus, however, may be
    different from one atom to the next Isotopes.
  • Uranium- 234 Uranium-235 Uranium-238
  • 92 protons 92 protons 92 protons
  • 142 neutrons 143 neutrons 146 neutrons
  • Trace 0.7 99.3
  • Different isotopes have different abundances
  • Different isotopes have different stabilities

47
Patterns of Nuclear Stability
As the atomic number increases, the neutron to
proton ratio of the stable nuclei increases. The
stable nuclei are located in the shaded area of
the graph known as the belt of stability. The
majority of radioactive nuclei occur outside this
belt.
48
Nuclear Equations
  • Radionuclides are unstable nuclei that emit
    particles and electromagnetic radiation to
    transform into a stable nucleus.

238
234
U
4
Th

He
92
90
2
49
Nuclear Equations
  • Mass numbers and atomic numbers must be balanced
    in all nuclear equations.

50
What product is formed when thorium-232 undergoes
alpha decay?
51
Types of Radioactive Decay
Alpha decay- nucleus emits 2 protons and 2
neutrons (He nucleus)
Beta decay- a neutron in the nucleus decays into
a proton and an electron, the electron is emitted
Gamma- high energy, short wavelength
electromagnetic radiation- accompanies other
radioactive emissions.
52
Types of Radioactive Decay
  • Electron Capture- capture by the nucleus of an
    electron from the electron cloud surrounding the
    nucleus.
  • Positron- particle with the same mass as an
    electron, but an opposite charge? collides with
    an electron and produces gamma radiation.

53
Penetrating Power of Radioactive Decay
54
Radioactive Decay Particles
Particle Nuclear Equation Example
Alpha 2 protons and 2 neutrons Nucleus ? 4He 226Ra ?222Rn 4He
Beta neutron converts to proton and a high energy electron 1n ? 1p 0e 131I?131Xe 0e
Electron Capture electron captured by nucleus 1p 0e ?1n 81Rb 0e ? 81Kr
88
86
2
2
-1
53
54
-1
1
0
37
36
-1
1
-1
0
55
Radioactive Decay Particles
Particle Nuclear Equation Example
Positronproton converted to a neutron and an electron 1p ? 1n 0e 11C ?11B 0e
Gamma electromagnetic radiation Not shown in equations, but almost always accompanies other decay.
6
5
1
1
0
1
Remember a positron has the same mass as an
electron, but the opposite charge
56
Radioactive Decay Particles
Particle Effect
Alpha Decrease atomic mass by ___ and atomic number by _____.
Beta Atomic number _______________.
Electron Capture Atomic number _______________.
Positron Atomic number _______________.
Gamma
57
Half Life- the time required for half of any
given quantity of a substance to react / decay.
(independent of initial quantity of atoms)
Half Life Simulation
Number of Th-232 atoms in a sample initially
containing 1 million atoms as a function of time.
Th-232 has a half-life of 14 billion years.
58
Half Life Problems
  •  Example
  • An isotope of cesium (cesium-137) has a half-life
    of 30 years. If 1.0 mg of cesium-137
    disintegrates over a period of 90 years, how many
    mg of cesium-137 would remain?
  •  

59
Half Life Problems
1. A 2.5 gram sample of an isotope of
strontium-90 was formed in a 1960 explosion of an
atomic bomb at Johnson Island in the Pacific Test
Site. The half-life of strontium-90 is 28 years.
In what year will only 0.625 grams of this
strontium-90 remain? 2. Actinium-226 has a
half-life of 29 hours. If 100 mg of actinium-226
disintegrates over a period of 58 hours, how many
mg of actinium-226 will remain? 3.
Thallium-201 has a half-life of 73 hours. If 4.0
mg of thallium-201 disintegrates over a period of
6.0 days and 2 hours, how many mg of thallium-201
will remain? 4. Sodium-25 was to be used in an
experiment, but it took 3.0 minutes to get the
sodium from the reactor to the laboratory. If
5.0 mg of sodium-25 was removed from the reactor,
how many mg of sodium-25 were placed in the
reaction vessel 3.0 minutes later if the
half-life of sodium-25 is 60 seconds? 5.
Selenium-83 has a half-life of 25.0 minutes. How
many minutes would it take for a 10.0 mg sample
to decay and have only 1.25 mg of it remain?
60
Uranium-238 an example of an unstable nucleus
decaying to form other unstable nuclei
Uranium-238 is radioactive, undergoing alpha
decay. But, the daughter nuclide is also
radioactive, undergoing beta decay, to produce
yet another radioactive nuclide, which decays.
The atom goes through a rather involved sequence
of radioactive decays (both alpha and beta),
until a stable isotope (lead-206) is reached.
61
Fission Reaction
Collision of a neutron with a U-235 nucleus can
cause the nucleus to split, creating two smaller
nuclides and three free neutrons. The three
neutrons may travel outward from the fission,
colliding with nearby U-235 nuclei, causing them
to split as well. Each split (fission) is
accompanied by a large quantity of energy.
62
Fission Chain Reaction
Collision of a neutron with a U-235 nucleus can
cause the nucleus to split, creating two smaller
nuclides and three free neutrons. The three
neutrons may travel outward from the fission,
colliding with nearby U-235 nuclei, causing them
to split as well. Each split (fission) is
accompanied by a large quantity of energy. If
sufficient neutrons are present, we may achieve a
chain reaction. If only one neutron were produced
with each fission, no chain reaction would occur,
because some neutrons would be lost through the
surface of the uranium sample.
Mousetrap Chain Reaction
63
Fission Reaction
64
Fusion Reaction
Tremendous energy needed to overcome the
repulsion between nuclei. Heat required for this
reaction is on the order of 40,000,000 K. The
energy from an atomic bomb could generate this
heat (hydrogen or thermonuclear weapon).
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