Download CH 5: The Atom

CH 5: The Atom
Renee Y. Becker
CHM 1025
Valencia Community College
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Dalton Model of the Atom
• John Dalton proposed that all matter is made up
of tiny particles.
• These particles are molecules or atoms.
• Molecules can be broken down into atoms by
chemical processes.
• Atoms cannot be broken down by chemical or
physical processes.
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Dalton’s Model
• According to the law of definite composition, the
mass ratio of carbon to oxygen in carbon dioxide
is always the same. Carbon dioxide is composed
of 1 carbon atom and 2 oxygen atoms.
• Similarly, 2 atoms of hydrogen and 1 atom of
oxygen combine to give water.
• Dalton proposed that 2 hydrogen atoms could
substitute for each oxygen atom in carbon
dioxide to make methane with 1 carbon atom
and 4 hydrogen atoms. Indeed, methane is
CH4!
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Dalton’s Theory
A Summary of Dalton’s Atomic Theory:
1. An element is composed of tiny, indivisible,
indestructible particles called atoms.
2. All atoms of an element are identical and
have the same properties.
3. Atoms of different elements combine to form
compounds.
4. Compounds contain atoms in small whole
number ratios.
5. Atoms can combine in more than one ratio
to form different compounds.
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Dalton’s Atomic Theory
• The first two parts of Dalton’s theory were later
proven incorrect.
– We will see this later.
• Proposals 3, 4, and 5 are still accepted today.
• Dalton’s theory was an important step in the
further development of atomic theory.
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Subatomic Particles
• About 50 years after Dalton’s proposal, evidence
was seen that atoms were divisible.
• Two subatomic particles were discovered.
– negatively charged electrons, e–
– positively charge protons, p+
• An electron has a relative charge of -1, and a
proton has a relative charge of +1.
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J. Thomson proposed a
subatomic model of the
atom in 1903.
Thomson’s Model of the Atom
Thomson proposed that the
electrons were distributed
evenly throughout a
homogeneous sphere of
positive charge.
This was called the “plum
pudding” mode of the
atom.
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Mass of Subatomic Particles
• Originally, Thomson could only calculate the
mass-to-charge ratio of a proton and an
electron.
• Robert Millikan determined the charge of an
electron in 1911.
• Thomson calculated the masses of a proton and
electron:
– an electron has a mass of 9.11 × 10-28 g
– a proton has a mass of 1.67 × 10-24 g
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Types of Radiation
• There are three types of radiation:
– alpha (a), beta (b), & gamma (g)
• Alpha rays are composed of helium atoms
stripped of their electrons (helium nuclei).
• Beta rays are composed of electrons.
• Gamma rays are high-energy electromagnetic
radiation.
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Rutherford’s Gold Foil Experiment
• Rutherford’s student
fired alpha particles at
thin gold foils. If the
“plum pudding” model of
the atom was correct, αparticles should pass
through undeflected.
• However, some of the
alpha particles were
deflected backwards.
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Explanation of Scattering
• Most of the alpha particles passed through the foil
because an atom is largely empty space.
• At the center of an atom is the atomic nucleus,
which contains the atom’s protons.
• The α-particles that bounced backwards did so
after striking the dense nucleus.
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Explanation of Scattering
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Rutherford's Model of the Atom
• Rutherford proposed a
new model of the atom:
– The negatively charged
electrons are
distributed around a
positively charged
nucleus.
• An atom has a diameter of
about 1 × 10-8 cm and the
nucleus has a diameter of
about 1 × 10-13 cm.
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Subatomic Particles
• Based on the heaviness of the nucleus,
Rutherford predicted that it must contain neutral
particles in addition to protons.
• Neutrons, n0, were discovered about 30 years
later. A neutron is about the size of a proton
without any charge.
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Atomic Notation
• Each element has a characteristic number of protons in the
nucleus. This is the atomic number, Z.
• The total number of protons and neutrons in the nucleus of
an atom is the mass number, A.
• We use atomic notation to display the number of protons
and neutrons in the nucleus of an atom:
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Periodic Table
• We can use the periodic table to obtain the atomic
number and atomic mass of an element.
• The periodic table shows the atomic number,
symbol, and atomic mass for each element.
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Atomic Notation
• The atomic number is:
– ALWAYS the # of protons
– Usually the # of electrons
• Unless the atom has a charge
• If the atom has a negative charge it has an extra
electron (example Cl- has 18 electrons)
• If the atom has a positive charge it has lost an
electron (example Na+ has 10 electrons)
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Atomic Notation
• The Mass # is:
– The sum of the protons and neutrons
• To find the # of neutrons
# neutrons = Mass # - Atomic #
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Using Atomic Notation
• An example:
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Si
• The element is silicon (symbol Si).
• The atomic number is 14: silicon has 14 protons.
• The mass number is 29: the atom of silicon has 29
protons + neutrons.
• The number of neutrons is:
29 – 14 =
15 neutrons.
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Example 1
• How many electrons, protons, and neutrons are
in the following elements of the periodic table?
Element
Atomic
Notation
# electrons
# protons
# neutrons
Nickel (Ni)
Nitrogen (N)
Chlorine (Cl)
Sodium (Na)
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Isotopes
• All atoms of the same element have the same
number of protons.
• Most elements occur naturally with varying
numbers of neutrons.
• Atoms of the same element that have a different
number of neutrons in the nucleus are called
isotopes.
• Isotopes have the same atomic number but
different mass numbers.
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Isotopes
• We often refer to an isotope by stating the
name of the element followed by the mass
number.
– cobalt-60 is
– carbon-14 is
60
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Co
14
C
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Wave Nature of Light
• Light travels through space as a wave, similar to
an ocean wave.
– Wavelength is the distance light travels in
one cycle.
– Frequency is the number of wave cycles
completed each second.
• Light travels at a constant speed: 3.00 × 108 m/s
(given the symbol c).
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Wavelength vs. Frequency
• The longer the wavelength of light, the lower the
frequency.
• The shorter the wavelength of light, the higher the
frequency.
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Radiant Energy Spectrum
• The complete radiant energy spectrum is an
uninterrupted band, or continuous spectrum.
• The radiant energy spectrum includes most types of
radiation, most of which are invisible to the human eye.
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Visible Spectrum
• Light usually refers to radiant energy that is
visible to the human eye.
• The visible spectrum is the range of
wavelengths between 400 and 700 nm.
• Radiant energy that has a wavelength lower
than 400 nm and greater than 700 nm cannot be
seen by the human eye.
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The Wave/Particle Nature of Light
• In 1900, Max Planck proposed that radiant
energy is not continuous, but is emitted in small
bundles. This is the quantum concept.
• Radiant energy has both a wave nature and a
particle nature.
• An individual
unit of light
energy is
a photon.
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Bohr Model of the Atom
• Niels Bohr speculated that electrons orbit about
the nucleus in fixed energy levels.
• Electrons are found only in specific energy
levels, and nowhere else.
• The electron energy
levels are quantized.
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Emission Line Spectra
• When an electrical voltage is passed across a gas in
a sealed tube, a series of narrow lines is seen.
• These lines are the emission line spectrum. The
emission line spectrum for hydrogen gas shows
three lines: 434 nm, 486 nm, and 656 nm.
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Evidence for Energy Levels
• Bohr realized that this was the evidence he
needed to prove his theory.
• The electric charge temporarily excites an
electron to a higher orbit. When the electron
drops back down, a photon is given off.
• The red line is the
least energetic and
corresponds to an
electron dropping
from energy level 3
to energy level 2.
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“Atomic Fingerprints”
• The emission line spectrum of each element is
unique.
• We can use the line spectrum to identify elements
using their “atomic fingerprint.”
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“Neon Lights”
• Most “neon signs” don’t actually contain neon gas.
• True neon signs are red in color.
• Each noble gas has its own emission spectrum, and
signs made with each have a different color.
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Energy Levels and Sublevels
• It was later shown that electrons occupy energy
sublevels within each level.
• These sublevels are given the designations s, p,
d, and f.
– These designations are in reference to the
sharp, principal, diffuse, and fine lines in
emission spectra.
• The number of sublevels in each level is the
same as the number of the main level.
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Energy Levels and Sublevels
• The first energy level has 1 sublevel:
– 1s
• The second energy level has 2 sublevels:
– 2s and 2p
• The third energy level has 3 sublevels:
– 3s, 3p, and 3d
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Electron Occupancy in Sublevels
• The maximum number of electrons in each of the energy
sublevels depends on the sublevel:
– The s sublevel holds a maximum of 2 electrons.
– The p sublevel holds a maximum of 6 electrons.
– The d sublevel holds a maximum of 10 electrons.
– The f sublevel holds a maximum of 14 electrons.
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Electrons per Energy Level
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Electron Configurations
• Electrons are arranged about the nucleus in a
regular manner. The first electrons fill the
energy sublevel closest to the nucleus.
• Electrons continue filling each sublevel until it is
full, and then start filling the next closest
sublevel.
• A partial list of sublevels in order of increasing
energy is:
1s < 2s < 2p < 3s < 3p < 4s < 3d < 4p < 5s <
4d …
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Filling Diagram for Sublevels
• For now, use
this Figure to
predict the order
of sublevel
filling.
Increasing Energy
• The order does
not strictly follow
1, 2, 3, etc.
Core
[He]
[Ne]
[Ar]
[Kr]
[Xe]
[Rn]
1s
2s
3s
4s
5s
6s
7s
2p
3p
4p
5p
6p
7p
3d
4d 4f
5d 5f
6d
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Electron Configurations
• The electron configuration of an atom is a
shorthand method of writing the location of
electrons by sublevel.
• The sublevel is written followed by a superscript
with the number of electrons in the sublevel.
– If the 2p sublevel contains 2 electrons, it is written
2p2.
• The electron sublevels are arranged according to
increasing energy.
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Writing Electron Configurations
• First, determine how many electrons are in the atom.
Bromine has 35 electrons.
• Arrange the energy sublevels according to increasing
energy:
– 1s 2s 2p 3s 3p 4s 3d …
• Fill each sublevel with electrons until you have used all the
electrons in the atom:
– Fe: 1s2 2s2 2p6 3s2 3p6 4s2 3d 10 4p5
• The sum of the superscripts equals the atomic number of
bromine (35).
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Example 2
Write the electron configuration for the following
a) Cl
e) N
b) Cl-
f) Mg
c) C
g) Mg2+
d) P
h) Fe
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Electrons
• Two main types of electrons
– Core – all electrons that are not in outermost
shell
– Valence – electrons in the outermost shell
• Most important electrons
• Where reactions happen
• These are the electrons that could be taken
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Example 3
• Write the electron configuration for fluorine
• Label the valence and core electrons
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Quantum Mechanical Model
• An orbital is the region of space where there is a
high probability of finding an atom.
• In the quantum mechanical atom, orbitals are
arranged according to their size and shape.
• The higher the energy of an orbital, the larger its
size.
• s-orbitals have
a spherical
shape
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Shapes of p-Orbitals
• Recall that there are three different p sublevels.
• p-orbitals have a dumbbell shape.
• Each of the p-orbitals has the same shape, but each
is oriented along a different axis in space.
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Location of Electrons in an Orbital
• The orbitals are the region of space in which the electrons
are most likely to be found.
• An analogy for an electron in a p-orbital is a fly trapped in two
bottles held end-to-end.
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