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Electrons in Atoms Gallium and Germanium: Discovered in 1875 & 1886 Mendeleev arranged the elements in order of increasing mass. In the 1860 s, the proton was not ... – PowerPoint PPT presentation

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Title: Electrons%20in%20Atoms

Electrons in Atoms
  • Why do ions have the charges they have? Like
    Al3 or Fe2 or Fe3 or O2-
  • Why does an atom become an ion in the first
  • Why are the BrINClHOF elements the only ones that
    make molecules with themselves? Why dont any
    other elements do that?
  • How do fireworks make colors when they explode?
  • How do fluorescent lights work?
  • How do we know how hot the sun is?
  • How is there life on this planet?
  • The electrons in atoms can answer all of these
    questions and so many others.
  • It all happens because of electrons..

Daltons Model of the Atom1803
  • Atoms are tiny, indestructible spheres
  • No internal structure

Thomsons Model1897
  • Referred to as the plum-pudding model.
  • The whole atom is a sphere of positive charge,
    with little negative electrons embedded in it.

Rutherfords Model1911
  • Small, dense core of positive charge.
  • Electrons circle the nucleus in fixed orbits.

Rutherfords Model
  • Electrons revolve around the nucleus like planets
    around the sun (fixed orbits).
  • This model failed to explain some properties of

Niels Bohrs Model1913
  • Electrons orbit the nucleus in specific orbits a
    fixed distance away.

Neils Bohrs Model (1913)
  • They orbit at a particular energy level. They
    can move to a higher level, but they need energy
    to do so.
  • A quantum of energy is the required amount to
    move an e- to a higher level. Exactly this
    amount, no in-between.

Neils Bohrs Model (1913)
  • This model of the atom had shortcomings. It
    failed to explain some phenomenon in nature.
  • So, a better version was still out there waiting
    to be discovered..

  • To understand the electronic structure of atoms,
    one must understand the nature of electromagnetic
  • The distance between corresponding points on
    adjacent waves is the wavelength (?).

  • The number of waves passing a given point per
    unit of time is the frequency (?).
  • For waves traveling at the same velocity, the
    longer the wavelength, the smaller the frequency.

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Electromagnetic Radiation
  • All electromagnetic radiation travels at the same
    velocity the speed of light (c),
  • 3.00 ? 108 m/s.
  • Therefore,
  • c ??
  • This all suggests light is a wave.

The Nature of Energy
  • The wave nature of light does not explain how an
    object can glow when its temperature increases.
  • Max Planck explained it by assuming that energy
    comes in packets called quanta.
  • Equantum h?
  • Einstein used this assumption to explain the
    photoelectric effect.

Photoelectric Effect
  • The emission of electrons from a metal when light
    is shined upon the metal.
  • Depending on the metal used, only light of a
    certain wavelength (color) would cause an
    electron to be emitted.

Photoelectric Effect
  • Einstein concluded that energy is proportional to
  • Ephoton h?
  • where h is Plancks constant, 6.626 ? 10-34 J-s.
  • This suggests light as a particle.

Nature of Energy
  • Therefore, if one knows the wavelength of light,
    one can calculate the energy in one photon, or
    particle, of that light
  • c ??
  • E h?

The Nature of Energy
  • Another mystery in the early 20th century
    involved the emission spectra observed from
    energy emitted by atoms and molecules.

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The Nature of Energy
  • For atoms and molecules one does not observe a
    continuous spectrum, as one gets from a white
    light source.
  • Only a line spectrum of discrete wavelengths is

The Nature of Energy
  • Niels Bohr adopted Plancks idea of quanta and
    explained these phenomena in this way
  • Electrons in an atom can only occupy certain
    orbits (corresponding to certain energies).
  • Electrons in permitted orbits have specific,
    allowed energies these energies will not be
    radiated from the atom.
  • Energy is only absorbed or emitted in such a way
    as to move an electron from one allowed energy
    state to another the energy is defined by
  • E h?

The Nature of Energy
  • The energy absorbed or emitted from the process
    of electron promotion or demotion can be
    calculated by the equation
  • ?E -Rh(1/nf2 - 1/ni2)
  • where RH is the Rydberg constant, 2.18 ? 10-18 J,
    and ni and nf are the initial and final energy
    levels of the electron.

The Wave Nature of Matter
  • Louis de Broglie posited that if light can behave
    with material properties (photons), matter should
    exhibit wave properties.
  • He demonstrated that the relationship between
    mass and wavelength was
  • ? h/mv
  • In other words, if light waves can act like
    particles, then things can move like waves.

Heisenbergs Uncertainty Principle
  • Heisenberg showed that the more precisely the
    momentum of a particle is known, the less
    precisely is its position known
  • (?x) (?mv) ? h/4p
  • For regular-sized objects, the uncertainty is
    practically zero, but in many cases, our
    uncertainty of the whereabouts of an electron is
    greater than the size of the atom itself!

Quantum Mechanics
  • Erwin Schrödinger developed a mathematical
    treatment into which both the wave and particle
    nature of matter could be incorporated.
  • It is known as quantum mechanics.
  • Amazingly accurate in describing electrons and
    microscopic behaviors, but also exceedingly

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Richard Feynman on Quantum Mechanics (1965)
  • There was a time when the newspapers said that
    only twelve men understood the theory of
    relativity. But after people read the paper a
    lot of people understood the theory of
    relativity. On the other hand I think I can
    safely say that nobody understands quantum

Schrodingers Wave Equations
  • The wave equation is designated with a lower case
    Greek psi (?).
  • The square of the wave equation, ?2, gives a
    probability density map of where an electron has
    a certain statistical likelihood of being at any
    given instant in time.

Some Folks are Really Into It.
Quantum Numbers
  • Solving the wave equation gives a set of wave
    functions, or orbitals, and their corresponding
  • Each orbital describes a spatial distribution of
    electron density.
  • An orbital is described by a set of three quantum

Principal Quantum Number, n
  • The principal quantum number, n, describes the
    energy level on which the orbital resides.
  • The values of n are integers 1.
  • n also describes the relative size of the
    orbital, 2 larger than 1, and so on.

Angular Momentum Quantum Number (l)
  • This quantum number defines the shape of the
  • Allowed values of l are integers ranging from 0
    to n - 1.
  • We use letter designations to communicate the
    different values of l and, therefore, the shapes
    and types of orbitals. This is where s, p, d f
    come into play

Value of l 0 1 2 3
Type of orbital s p d f
Magnetic Quantum Number (ml)
  • The magnetic quantum number describes the
    three-dimensional orientation of the orbital.
  • Allowed values of ml are integers ranging from -l
    to l
  • -l ml l.
  • Therefore, on any given energy level, there can
    be up to 1 s orbital, 3 p orbitals, 5 d orbitals,
    7 f orbitals, etc.

Magnetic Quantum Number (ml)
  • Orbitals with the same value of n form a shell.
  • Different orbital types within a shell are

s Orbitals
  • The value of l for s orbitals is 0.
  • They are spherical in shape.
  • The radius of the sphere increases with the value
    of n.

s Orbitals
  • Observing a graph of probabilities of finding an
    electron versus distance from the nucleus, we see
    that s orbitals possess n-1 nodes, or regions
    where there is 0 probability of finding an

p Orbitals
  • The value of l for p orbitals is 1.
  • They have two lobes with a node between them.

d Orbitals
  • The value of l for a d orbital is 2.
  • Four of the five d orbitals have 4 lobes the
    other resembles a p orbital with a doughnut
    around the center.

Energies of Orbitals
  • For a one-electron hydrogen atom, orbitals on the
    same energy level have the same energy.
  • That is, they are degenerate.

Energies of Orbitals
  • As the number of electrons increases, though, so
    does the repulsion between them.
  • Therefore, in many-electron atoms, orbitals on
    the same energy level are no longer degenerate.

Spin Quantum Number, ms
  • In the 1920s, it was discovered that two
    electrons in the same orbital do not have exactly
    the same energy.
  • The spin of an electron describes its magnetic
    field, which affects its energy.

Spin Quantum Number, ms
  • This led to a fourth quantum number, the spin
    quantum number, ms.
  • The spin quantum number has only 2 allowed
    values 1/2 and -1/2.

Pauli Exclusion Principal
  • No two electrons in the same atom can have
    exactly the same energy.
  • Therefore, no two electrons in the same atom can
    have identical sets of quantum numbers.

Electron Configurations
  • This shows the distribution of all electrons in
    an atom.
  • Each component consists of
  • A number denoting the energy level,
  • A letter denoting the type of orbital,
  • A superscript denoting the number of electrons in
    those orbitals.

Orbital Diagrams
  • Each box in the diagram represents one orbital.
  • Half-arrows represent the electrons.
  • The direction of the arrow represents the
    relative spin of the electron.

Hunds Rule
  • For degenerate orbitals, the lowest energy is
    attained when the number of electrons with the
    same spin is maximized.

Periodic Table
  • We fill orbitals in increasing order of energy.
  • Different blocks on the periodic table (shaded in
    different colors in this chart) correspond to
    different types of orbitals.

Atomic Orbitals
  • Region of space where there is a high probability
    of finding an electron.
  • Principal Quantum (n) --denotes the energy level
    of electrons (1,2,3,4,etc.)
  • Also denotes the of sublevels at that energy
    level (s,p,d,f)
  • Sublevels describe the shapes and sizes of
    orbitals where e- may be found.

Shapes of Orbitals
  • s-spherical, with nucleus at the center
  • p-dumbbell, or figure-8, with nucleus at the
  • das shown
  • fas shown
  • As you increase energy levels, the shape of each
    remains the same, but size gets larger.

Electron Configurations
  • Orbitals of an atom will fill so that the atom is
    in its most stable state. There are 3 rules that
    govern this
  • Aufbau Principle- e- occupy lowest-energy
    orbitals first
  • Pauli Exclusion Principle- 2 e- in same orbital
    must have opposite spin
  • Hunds Rule- e- occupy orbitals of the same
    energy so that theres a max of same spin e-

Exceptions to Aufbau
  • If you did the configuration for Cu according to
    the three rules, it would look like this
  • 1s22s22p63s23p64s23d9
  • In actuality, it is this
  • 1s22s22p63s23p64s13d10

  • Chromium, Cr, also is an exception to the Aufbau
  • According to Aufbau, Cr should have this
  • 1s22s22p63s23p64s23d4
  • But it actually has this
  • 1s22s22p63s23p64s13d5

Why Would an Atom Do This?
  • Because a filled shell is the most stable
    arrangement, and a half-filled shell is the next
    best arrangement.

Valence Electrons
  • The electrons that exist in the outermost energy
    level of an atom are valence electrons.
  • A full shell or a half-filled shell is the most
    stable arrangement.
  • Noble gases always have a full valence, or a full
    outer shell, which is what every other element is
    trying to achieve. (Max. of 8 valence electrons)

  • What does the term orbital describe?
  • A region around the nucleus where an electron is
    most likely to be found.
  • What does an elements electron configuration
  • All of the orbitals that the elements electrons
    occupy, and how those electrons are distributed.

  • We do not need to focus on all the electrons that
    an atom has, we really only need to focus on the
    valence electrons. Why?
  • Because they are the outermost electrons, and
    they are the only electrons that can possibly
    interact with other atoms.
  • How many valence electrons does Oxygen have?
  • 6

  • Why are the alkali metals so reactive?
  • They all have an s1 electron (1 valence electron)
    that they are trying to lose.
  • Noble gases are also called inert gases. Why are
    the noble gases so unreactive?
  • Because they have a full outer shell (eight
    valence electrons) and do not need any more or
    less electrons.

First Periodic Table
  • In 1869, the first table having elements
    organized by their properties was published by a
    Russian chemist and professor named Dmitri
  • He listed them in order of atomic mass.

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Gallium and Germanium Discovered in 1875 1886
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  • Mendeleev arranged the elements in order of
    increasing mass.
  • In the 1860s, the proton was not yet discovered.
  • In 1913, British physicist Henry Moseley arranged
    the elements in order of increasing atomic number
    ( of protons).

Some Vocabulary
  • Vertical columns are called
  • groups or families.
  • Horizontal rows are called periods.
  • How many elements are in period 2?
  • 8
  • How many elements are in period 6?
  • 32
  • How many elements are in group 2?
  • 6

The Periodic Law (Cont.)
  • Elements within a column of a group have similar
  • Properties in a period change as you move across
    a period from left to right.
  • The pattern of properties within a period repeats
    as you move form one period to the next.
  • Periodic Law When elements are arranged in order
    of increasing atomic number, there is a periodic
    repetition of their physical and chemical

Electron Configurations in Groups
Helium (He) 1s2
Neon (Ne) 1s22s22p6
Argon (Ar) 1s22s22p63s23p6
Krypton (Kr) 1s22s22p63s23p63d104s24p6
Noble Gases
Lithium (Li) 1s22s1
Sodium (Na) 1s22s22p63s1
Potassium (K) 1s22s22p63s23p64s1
Alkali Metals
Blocks of Elements
  • 80 of elements

Metals (Cont.)
  • Conductors of heat
  • Conductors of electric current
  • High luster
  • Ductile
  • Malleable
  • Solids _at_ room temp. (except Hg)

  • Most are gases _at_ room temp
  • Poor conductors of heat
  • Poor conductors of electric current
  • Solid nonmetals are brittle

  • Properties similar to those of
  • metals and nonmetals
  • Behaviors can be controlled by changing the
  • Example Silicon

Classifying the Elements
  • Group 1A Elements Alkali Metals
  • Group 2A Elements Alkaline Earth Metals
  • Group 7A Elements Halogens
  • Group 8A Elements Noble Gases
  • Group B Elements Transition Metals
  • Below the Main Body- Inner Transition Metals

Periodic Trends
  • Atomic Size
  • Atomic radius one half the distance between the
    nuclei of two atoms of the same element
  • Increases from top to bottom within a group
  • Decreases from left to right across a period

Atomic Radius vs. Atomic Number
Increase within Group 1 Shielding Effect
Periodic Trends in Atomic Size
Size decreases
Size Increases
Trends in Ionization Energy
  • Ionization Energy
  • The energy required to remove an electron from an
  • First Ionization Energy
  • The energy required to remove the first electron
    from at atom.
  • First ionization energy tends to decrease from
    top to bottom within a group and increase from
    left to right across a period.

First Ionization Energy vs. Atomic Number
Why is the 1st ionization energy for the noble
gases higher?
Periodic Trends in Ionization Energy
Energy Increases
Energy Decreases
  • Electronegativity- The ability of an atom of an
    element to attract electrons when the atom is
    BONDED to another atom in a compound.

Periodic Trends
  • Metallic propertiesas shown. As you approach
    the nonmetals, metallic properties decrease.
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