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Atomic Theory CP
History of the Atom; Modern Atomic Theory, Subatomic
Particles
Reasons to look at history
• Shows that present theory is the work of many scientists over
many years
• To become aware of the logic used in obtaining evidence to
support the theory
• To gain a better understanding of the properties of atoms
Greek Philosophers
• Democritus (460 – 371 B.C.)
• Stated that the world was made of tiny particles that are mostly
empty space and that they are the building blocks of all matter
• Tiny particles could not be divided
• Termed them atomos, or atoms
• Means “uncut” or “indivisible”
• Aristotle (384 – 322 B.C.)
• Proposed that matter was continuous (i.e. came from only four
elements – earth, air, fire, and water)
Dalton’s Atomic Theory
• John Dalton (1766 – 1844)
• Work marks the beginning of the development of the modern
atomic theory
• Revived Democritus’ ideas and backed them up with scientific
experimentation and data!
• Dalton’s Atomic Theory:
1. Matter is composed of extremely small particles called atoms
2. Atoms are indivisible and indestructible
3. Atoms of a given element are identical in size, mass, and chemical
properties
4. Atoms of a specific element are different from those of another
element
5. Different atoms combine in simple whole-number ratios to form
compounds
6. In a chemical reaction, atoms are separated, combined, and
rearranged
Dalton’s Atomic Theory
• Dalton’s theory successfully explained:
• The law of conservation of mass (proposed by Lavoisier)
• States that mass is conserved in any process, such as a chemical reaction.
• Conservation is the result of the separation, combination, or
rearrangement of atoms
• The law of multiple proportions (proposed by Proust)
• Certain elements can combine to form two or more different chemical
compounds
• Experimental evidence led to general acceptance of his atomic
theory!
• Not all of the theory was accurate.
• It has had to be revised with new information
Modern Atomic Theory
1. All matter is made up of a very small particles called atoms
2. Atoms of the same element have the same chemical
properties, atoms of different elements have different
properties
3. Any natural sample of an element has an average mass
characteristic of that element
4. Compounds are formed when atoms of two or more
elements unite
5. Atoms are not subdivided in physical or chemical reactions
Protons, Neutrons, and Electrons
DEFINING THE ATOM
The Atom
• Atom:
• Smallest particle of an element that retains the properties of the
element
• Size:
• Consider this:
• In 2006 the population of Earth was about 6.5 x 109 people
• A solid copper penny contains 2.9 x 1022 atoms
• This is 5 trillion times larger than the worlds population!!!!!
Electrons
• J.J. Thomson
• 1879 Cathode Ray Tube
• Provided the first evidence that atoms are made of even smaller
particles
• Cathode-Ray Tube:
•
•
•
•
Vacuum tube, all gases removed
Metal electrodes at each end
Electricity runs from the negative cathode to the positive anode
Using a magnet, the negative charged cathode ray bent towards a
positively charged plate
Electrons
• Observations of the cathode-ray behavior led to the
hypothesis that cathode rays consist of identical negatively
charged particles
• Thomson proposed the “Plum Pudding” model of the atom
Thomson’s Model
• Found the electron
• 1 unit of negative charge
• Mass 1/2000 of hydrogen atom
• Later refined by Millikan to 1/1840
• Concluded that there must be a positive charge since atom
was neutral
• Atom was like plum pudding
• A bunch of positive stuff, with electrons able to be removed
Nucleus
• Ernest Rutherford – 1908
• Discovered the nucleus
• The core of the atom and
contains protons and
neutrons
• Rutherford’s nuclear model
proposed that an atom is
mostly empty space
• This is based on studies of
the bombardment of thin
metal foils with fast-moving
positively charged particles
Rutherford’s Experiment
Rutherford’s Experiment
• His experiment demonstrated that:
• An atom is mostly empty space
• There is a small dense, positive piece at the center
• Alpha particles are deflected if they got close enough to the
positive center
Protons and Neutrons
• Protons:
• Rutherford also concluded that the nucleus also contained
positively charged particles call protons.
• Subatomic particle with a positive charge equal in magnitude to the
negative charge of an electron
• The number of protons determines the identity of an atom!!
• Neutrons:
• Discovered by James Chadwick (Rutherford’s coworker)
• Showed the nucleus also contained a neutral subatomic particle,
the neutron
• Electrically neutral subatomic particle
• The number of neutrons can vary within an element  Isotopes!
Rutherford Activity
Structure of the Atom
• There are two regions
• The nucleus:
• Protons and Neutrons
• Positive charge
• Almost all of the mass
• Electron Cloud:
• Most of the volume of an atom
• Region were electrons can be
found
Comparing Subatomic Particles
Properties of Subatomic Particles
Particle
Symbol
Relative
Charge
Relative Mass
(Mass of Proton =
1)
Actual Mass (g)
Electron
e-
1-
1/1840
9.11 x 10-28
Proton
p+
1+
1
1.67 x 10-24
Neutron
n0
0
1
1.67 x 10-24
Atomic Number
• Atoms of each element contain a unique positive charge in
their nuclei
• Therefore, the number of protons in an atom identifies that
particular atom!
• The number of protons in an atom is referred to as the atomic
number
• Because all atoms are neutral:
• The number of protons =
the number of electrons
• If you know the atomic number
of an element you now know
both the number of protons and
the number of electrons!
Atomic
Numbe
r
Mass Number
• The total number of protons and neutrons in the nucleus of an
atom
• Mass # = p + n
• Usually found at the bottom of the atomic symbol
• How to find the number of neutrons:
• Mass number – atomic number = # neutrons
• Example:
• Nitrogen:
• Mass number: 14
• Atomic number: 7
• Number of Neutrons: 7
Fill in the Gaps
Atomic #
Mass #
3
7
39
P
N
19
6
6
6
19
7
10
E
Mass Number Continued
• Although a given type of atom will usually contain a certain
number of neutrons in the nucleus, a small percentage will not
• Ex: most hydrogen atoms contain no neutrons
• A small percentage contain one neutron and a smaller
percentage two neutrons
• What do we call atoms with a different number of neutrons?
Isotopes
• In nature, most elements are found as a mixture of isotopes.
• Isotope:
• Atoms that have the same number of protons but different
numbers of neutrons
• This means that the number of neutrons in a atom can
change!
• Protons MUST stay the same
• Different isotopes of the same element have different mass
numbers (i.e. different neutrons)
• A given element has at least two isotopes, and they all have
different mass numbers
• Example:
• Gadolinium (Gd) has 6 stable isotopes
Isotopes
• Isotopes are chemically alike because they have identical
numbers of protons and electrons
• It is useful to compare the relative masses of atoms to a
standard reference isotope
• Carbon-12 is the standard reference isotope
• Carbon-12 has a mass of exactly 12 atomic mass units (amu)
Atomic Mass
• A weighted average mass of the atoms in a naturally occurring
sample of the element
• Mass of an atom in atomic mass units (amu)
• Equal to 1/12th of the mass of carbon
• Weighted average mass reflects both the mass and the
relative abundance of the isotopes as they occur in nature
• This is the number which appears on the periodic table
Atomic Mass Number
• Example:
• 99% of all carbon atoms are the isotope containing 6 neutrons,
the remaining 1% is the heavier isotope containing 7 neutrons,
which raises the aver mass of carbon from 12.000 to 12.011
Calculations
• Carbon has two major, stable isotopes.
• Carbon-12 (98.93% abundance)
• Carbon-13 (1.07% abundance)
• Carbon-14 is in trace amounts and is radioactive
• Calculate the average atomic mass of carbon
• Carbon 12: (0.9893 x 12) = 11.8716
• Carbon 13: (0.0107 x 13) = 0.1391
• AAM: 12.0107
Calculations
Rubidium has two common isotopes, 85Rb and 87Rb. If the
abundance of 85Rb is 72.2% and the abundance of 87Rb is 27.8%,
what is the average atomic mass of rubidium?
• Average atomic mass = (0.722 x 85) + (0.278 x 87)
• Average atomic mass = (61.37) + (24.186)
• Average atomic mass = 85.56 amu