CH 5: The Atom

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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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Daltons 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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Daltons Theory

• A Summary of Daltons Atomic Theory
• An element is composed of tiny, indivisible,
  indestructible particles called atoms.
• All atoms of an element are identical and have
  the same properties.
• Atoms of different elements combine to form
  compounds.
• Compounds contain atoms in small whole number
  ratios.
• Atoms can combine in more than one ratio to form
  different compounds.

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Daltons Atomic Theory

• The first two parts of Daltons theory were later
  proven incorrect.
• We will see this later.
• Proposals 3, 4, and 5 are still accepted today.
• Daltons theory was an important step in the
  further development of atomic theory.

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Subatomic Particles

• About 50 years after Daltons 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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Thomsons Model of the Atom

• J. Thomson proposed a subatomic model of the atom
  in 1903.
• 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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Rutherfords Gold Foil Experiment

• Rutherfords student fired alpha particles at
  thin gold foils. If the plum pudding model of
  the atom was correct, a-particles 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 atoms protons.
• The a-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 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

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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 dont 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 lt 2s lt 2p lt 3s lt 3p lt 4s lt 3d lt 4p lt 5s lt 4d

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Filling Diagram for Sublevels

• The order does not strictly follow 1, 2, 3, etc.
• For now, use this Figure to predict the order of
  sublevel filling.

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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

• Cl
• Cl-
• C
• P

1. N
2. Mg
3. Mg2
4. 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.