Download Electrons in Atoms

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 place?
• Why are the BrINClHOF elements the only ones that
make molecules with themselves? Why don’t 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……..
Dalton’s Model of the Atom
1803
• Atoms are tiny, indestructible spheres
• No internal structure
Thomson’s Model
1897
• Referred to as the “plum-pudding” model.
• The whole atom is a sphere of positive
charge, with little negative electrons
embedded in it.
Rutherford’s Model
1911
• Small, dense core of
positive charge.
• Electrons circle the
nucleus in fixed orbits.
Rutherford’s Model
• Electrons revolve around
the nucleus like planets
around the sun (fixed
orbits).
• This model failed to explain
some properties of atoms.
Niels Bohr’s Model
1913
• Electrons orbit the
nucleus in specific orbits
a fixed distance away.
Neils Bohr’s 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 inbetween.
Neils Bohr’s 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…..
Waves
• To understand the electronic structure of
atoms, one must understand the nature of
electromagnetic radiation.
• The distance between corresponding
points on adjacent waves is the
wavelength ().
Waves
• 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.
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 = hν
• 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
frequency:
Ephoton = h
where h is Planck’s
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.
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
observed.
The Nature of Energy
•
Niels Bohr adopted Planck’s idea of quanta and
explained these phenomena in this way:
1. Electrons in an atom can only occupy certain orbits
(corresponding to certain energies).
2. Electrons in permitted orbits have specific, “allowed”
energies; these energies will not be radiated from the
atom.
3. 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
What Is Light?
Particle or Wave?
• In the 1600’s, Sir Isaac Newton described
light as a stream of particles.
• By the 1900’s, experiments had shown
that light was a wave.
• In the early 1900’s, an experiment
produced results that made either
explanation incorrect…….
• In 1924, Louis de Broglie proposes
that matter (particles) moves in
waves.
• Light is a stream of particles
(photons) that move in waves.
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 mechanics.”
Light
Erwin Schrödinger’s Model
1926
• Mathematical equations
describe the motion of
electrons.
Erwin Schrödinger’s Model (1926)
• Quantum mechanical model came
from his equations.
• Same as Bohr’s model except that
electrons exist in the most
probable locations (orbitals) at
certain energy levels
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 center
• d—as shown
• f—as 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 lowestenergy orbitals first
• Pauli Exclusion Principle- 2 e- in same
orbital must have opposite spin
• Hund’s Rule- e- occupy orbitals of the
same energy so that there’s 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:
2
2
6
2
6
1
10
• 1s 2s 2p 3s 3p 4s 3d
Another…
• Chromium, Cr, also is an
exception to the Aufbau Principle
• According to Aufbau, Cr should
have this configuration:
• 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 element’s electron
configuration describe?
• All of the orbitals that the element’s
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.
Organizing Principle
• Chemists used the properties of elements to sort them
into groups.
•1829: Dobereiner published a
classification system.
•Grouped elements into
triads.
•Similar chemical
properties.
•Example: Cl, Br, and I
J.W. Dobereiner (17801849)
First Periodic Table
• In 1869, the first table
having elements
organized by their
properties was
published by a Russian
chemist and professor
named Dmitri
Mendeleev.
• He listed them in order
of atomic mass.
Gallium and
Germanium:
Discovered in
1875 & 1886
• Mendeleev arranged
the elements in order
of increasing mass.
• In the 1860’s, 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.
• 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
properties.
Electron Configurations in
Groups
Helium (He)
Neon (Ne)
Argon (Ar)
Krypton (Kr)
1s2
1s22s22p6
1s22s22p63s23p6
1s22s22p63s23p63d104s24p6
Noble Gases
Lithium (Li)
Sodium (Na)
1s22s1
1s22s22p63s1
Potassium (K)
1s22s22p63s23p64s1
Alkali Metals
Blocks of Elements
Metals
• 80% of elements
Nonmetals
Metals
Metalloids
Metals (Cont.)
• Conductors of heat
• Conductors of electric current
• High luster
• Ductile
• Malleable
• Solids @ room temp. (except
Hg)
Nonmetals
• Most are gases @ room
temp
• Poor conductors of heat
• Poor conductors of
electric current
• Solid nonmetals are brittle
Metalloids
Properties similar to those of
metals and nonmetals
Behaviors can be controlled by changing the
conditions.
– 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 atom.
• 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 properties—as shown.
As you approach the
nonmetals, metallic properties
decrease.