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Transcript
Atomic Structure and the Composition of Matter
• The atom is a basic building block of minerals.
• Matter is a special form of energy; it has mass and
occupies space. Neither matter nor energy may be created
or destroyed - they may only be converted from one form
to the other.
• Energy is the ability to do work and it occurs in a number
of forms, including:
– Potential; Kinetic; Electrical; Heat; Chemical; Nuclear
– Radiant (the only form in which there is an absence of
matter)
• Atoms are the smallest division of matter that retain the
characteristics of the elements.
Elements and the Periodic Table
• Approximately 300 different kinds of atoms that are
capable of independent, prolonged existence. These
are called nuclides.
• If nuclides are grouped by chemical characteristics,
about 100 sets result and these are referred to as
elements.
• The modern Periodic Table was devised in 1869 by
Julius Meyer and Dmitri Mendeleev. It organizes
the elements into groups and families with similar
chemical and physical properties.
Periods
Groups or Families
Periodic Table
Crustal Abundance
Crustal volume:<1% of Earth
Mantle volume: 83% of Earth
Core volume: 16% of Earth
Crustal mass:<1% of Earth
Mantle mass: 68% of Earth
Core mass: 31% of Earth
What element is most abundant for the entire Earth?
Atomic Particles: Basics
• Atoms are composed of electrons and two large
nuclear particles called protons and neutrons.
• Protons and neutrons are approximately equal in
mass and are ~1800 times more massive than the
electron. Both nuclear particles are composed of
quarks, smaller fundamental particles.
• Protons have unit positive charge (+1), while
electrons have unit negative charge (-1). Neutrons
carry no charge.
• Atoms are electrically neutral and thus the
number of electrons must equal the number of
protons.
Basic Terminology
• Atomic number (Z): The atomic number represents the
number of unit positive charges on the nucleus and is
equal to the number of protons within the nucleus, since
each proton carries unit positive charge. In electrically
neutral atoms, it also represents the number of electrons,
which carry unit negative charge.
• Mass number (A): The mass number is equal to the total
number of nucleons, which is the sum of the number of
protons and neutrons. A does not equal the total mass of
the atom; rather, it represents a whole number
approximation of the mass, as expressed in amu.
• The number of neutrons is simply defined as the A - Z.
Isotopes and Isobars
•
A specific type of atom is designated by using its chemical symbol,
which is an abbreviation of its name in German, Latin, or English, with
the A, the mass number, placed in the upper left and Z, the atomic
number, placed in the lower left corner. For example, 23Na11, has a
mass number of 23 and an atomic number of 11.
•
Isotopes are atoms of the same element that differ in mass. For
example, 87Sr and 86Sr or 238U and 235U. Isotopes have similar
chemical characteristics and are studied using a mass spectrograph
or spectrometer. Most elements have at least two naturally occurring
isotopes.
•
Isobars are nuclides that have the same mass number but different
atomic numbers. For example, 36S and 36Ar are isobars; they both
contain a total of 36 nucleons (protons plus neutrons), but the sulfur
isotope has 16 protons and 20 neutrons, while the argon isotope has 18
protons and 18 neutrons. Isobars do not have similar chemical
characteristics!
Atomic Weight
• Atomic weight is the weighted average of the atomic
masses of the naturally occurring isotopes. For example,
a natural sample of the element chlorine contains a mixture
of 75.53% 35Cl and 24.47% 37Cl. Thus the atomic weight is
obtained by multiplying the mass of each isotope (in amu)
times its fractional abundance:
• 0.7553 (34.97 amu) + 0.2447 (36.95 amu) = 35.45 amu
Atomic Models
• Bohr Model
– Electron shells
• Quantum Mechanics
– Orbitals
– Afbau Filling Order
– Quantum numbers and superposition of states
EM propagation and spectrum
λ is wavelength (m)
ν is frequency (cycles/s= s-1)
c is velocity of EM rad (ms-1)
λ = ν/ c and cλ = ν
ν = 1/λ is wavenumber (m-1)
Note that in 1900 Planck
Determined that the energy of
a photon is quantized:
E = hυ
where h is Planck’s constant
Blackbody radiation
Planck’s Law
8# ch" $5
! " d " = ch / " kT
d"
e
$1
where h is Planck’s constant
and k is Boltzman’s constant
! mT =
ch
= 2.90x10 "3 mK
4.97k
Rutherford’s Scattering Experiment
Balmer Series: Visible spectrum for Hydrogen
Bohr Atomic Model
Oxygen Atom
The Bohr model for the atom envisioned these electrons in stable orbits of specified
radius and energy, where we could exactly pinpoint the position of any individual
electron. Each energy level was permitted to have a specified number of electrons, and
was called a shell. We know now that this simple view is not correct; it is impossible
to exactly determine the position of an electron in space.
Complete emission
spectrum and energy
levels for H atom
The Quantum Mechanical View
•
Using the theory of quantum or wave mechanics we can calculate the
probabilities of various electron configurations, and thus show that
specified regions near the nucleus have higher probabilities for finding
an electron than others. Each electron does, however, have a specific
energy. Must solve wave equation for specific states!
•
The combination of the energy and probability gives rise to the current
understanding for electron distributions, which are referred to as
electron orbitals; these orbitals are referred to as s (sharp), p
(principal), d (diffuse), and f (fundamental).
•
With increasing atomic number, each new element has an additional
electron also added to it extra-nuclear cloud. From theory and
experiment, we know that these electrons are added in a systematic
fashion, with the lowest energy orbitals being filled first.
•
This process is called aufbau filling (1s -> 2s -> 2p -> 3s -> 3p -> 4s > 3d -> 4p -> 5s, etc. ).
S, P, and D orbital probability
density functions
S orbitals
P orbitals
1S orbital
D orbitals
Pauli Exclusion Principle:
no two electrons may have
the same quantum number
Hunds Rule: electrons prefer
parallel spins before pairing
up in subshell orbital
Quantum Numbers
•
•
•
•
n = 6; l = 4; m = 1
n = principal
l = orbital (0 -> (n-1))
ml = magnetic (± l)
ms = spin (±1/2)
– l = 0 -> s orbital
= 1 -> p orbital
= 2 -> d orbital
= 3 -> f orbital
From: http://www.daugerresearch.com/orbitals/
Animated Orbital Simulations
Superposition of States
| 4, 3, 3 | and | 4, 1, 0 |
| 3, 2, 2 | and | 3, 1, -1 |
From: http://www.daugerresearch.com/orbitals/
H atom quantum numbers and X-rays