Download Unit 9: Atomic Structure, Periodicity and Chemical Bonding

UNIT 9:
ATOMIC STRUCTURE, PERIODICITY AND
CHEMICAL BONDING
By: Anthony Gates
AP Chemistry
QUANTUM MECHANICAL MODEL
• Previous models, such as the Bohr model, assumed that electrons
behaved like single particles moving along a set path.
• This could not explain certain properties of elements.
• The quantum mechanical (QM) model addresses known problems
with the classical shell model and is also consistent with atomic
electronic structures that correspond with the periodic table.
• Electrons behave like particles and like waves.
QM CONT.
• The QM model can be approximately solved using computers and serves as the
basis for software that calculates the structure and reactivity of molecules.
• Computers can do it faster, thus allowing for difficult calculations to be performed
faster.
• The QM model is based off of Schrodinger’s equation… an equation that is
unsolvable … unless you cut some corners.
BILLIARDS
ELECTRON MOVEMENT
• Heisenberg Uncertainty Principle: The more accurately we
know a particle’s position, the less accurately we know its
momentum, (and vice versa) thus causing a minimum amount
of uncertainty in a particle’s movement to always exist.
• This becomes less important as the mass of the object increases…
thus why we can predict the movement of billiard balls.
• This is affects our predictions of electron movement greatly.
• Thus…WE CANNOT ASSUME ELECTRONS MOVE IN A CONSTANT
CIRCULAR PATH!!!
THE AFFECT OF SPIN
• Think of billiards… has anyone ever heard of putting a little English on the ball?
• In billiards depending on where you hit the cue ball with your cue, you can cause a
different spin which in turn causes a different movement in the cue ball.
• Hitting a little high causes a topspin which causes the cue ball to keep moving forward
even after impact.
• Electrons have various spins as well and thus there movement (despite the lack of
collisions) is affected by these spins.
ORBITALS
• Due to the varying spins, distances from the nucleus and both attractive and repulsive
forces present the electrons’ movements cannot be predicted.
• However… we can predict their positions based on probability.
• Take this school for example…
• Classes
• Lockers
• Friends
• Clubs/activities
RADIAL DISTRIBUTION
• The probability of finding an electron.
• 1s Orbital
• How does this relate to the bohr model?
2s Orbital
QUANTUM NUMBERS
• n= principal quantum number
• l = angular momentum number
• Number of nodes
• 0 to n-1
• ml=magnetic quantum number
• Orientation in space
• -l to +l
• ms=electron spin quantum number
• -1/2 or +1/2
PAULI EXCLUSION PRINCIPLE
• No two electrons in the same orbital can have the same spin.
• Electrons in atoms have an intrinsic property known as spin that can
result in atoms having a magnetic moment. There can be at most
two electrons in any orbital and these electrons must have opposite
spin.
MR. GATES IF YOU PLEASE…
• Please explain/diagram atomic orbitals and their nodes…
ELECTRON CONFIGURATIONS
• Electron configurations provide a method for
describing the distribution of electrons in an
atom or ion.
• In multielectron atoms and ions, the electrons
can be thought of as being in “shells” and
“subshells,” indicated by the relatively close
ionization energies associated with some
groups of electrons.
• Inner electrons are called core electrons, and outer
electrons are called valence electrons.
COULOMB’S LAW
• Coulomb’s Law is the basis for describing the energy of interaction
between protons and electrons.
• Based on Coulomb’s Law, the force between two charged particles is
proportional to the magnitude of each of the two charges (q1 and q2),
and inversely proportional to the square of the distance, r, between
them.
• If the two charges are of opposite sign , the force between them is attractive; if
they are of the same sign, the force between them is repulsive.
COULOMB’S LAW
Potential Energy = (2.31 x 10-19Jnm)[(q1 x q2)/r]
Force ≈ (q1 x q2)/(r2)
COULOMB’S LAW & IONIZATION ENERGY
• Each electron in an atom has a different ionization energy, which can
be qualitatively explained through Coulomb’s Law.
• As two atoms come together there are three forces at work…
• Nucleus to nucleus repulsion
• Electron to electron repulsion
• Electron to nucleus attraction
• Electron to Electron shielding
COULOMB’S LAW & IONIZATION ENERGY
• The first ionization energy is the minimum energy needed to remove
the least tightly held electron from an atom or ion.
• In general, the ionization energy of any electron in an atom or ion is the
minimum energy needed to remove that electron from the atom or ion.
• The relative ionization energy can be estimated through qualitative
application of Coulomb’s Law.
• The farther an electron is from the nucleus, the lower its ionization energy.
• When comparing two species with the same arrangement of electron, the
higher the nuclear charge, the higher the ionization energy of an electron in a
given subshell.
SHIELDING
• Core electrons are generally closer to the nucleus than valence
electrons, and they are considered to “shield” the valence electrons
from the full electrostatic attraction of the nucleus.
• This phenomenon can be used in conjunction with Coulomb’s Law to
explain/rationalize/predict relative ionization energies.
• Differences in electron-electron repulsion are responsible for the differences in
energy between electrons in different orbitals in the same shell.
ATOMIC RADII
• Atomic Radius increases down a group. This is due to the orbital sizes
increasing in successive principal quantum levels (numbers).
• Atomic Radius decreases to the right of the periodic table due to
increased attraction between the electrons and the nucleus at a
relatively similar distance.
ELECTRONEGATIVITY
• Electronegativity is the ability of an atom in a molecules to
attract shared electrons to it.
• Electronegativity values for the representative elements
increase going from the left to right across a period and
decrease going down a group.
• These trends can be understood qualitatively through the
electronic structure of the atoms, the shell model, and Coulomb’s
Law.
BELL RINGER!!
• Turn to the person next to you and use Coulomb’s Law and the shell
model to describe why ionization energy increases as you remove
additional electrons beyond the first.
• Use Coulomb’s Law and the shell model to justify the concept of
electron shielding.
PERIODIC TABLE
• The structure of the periodic table is a consequence of the pattern of
the electron configurations and the presence of shells (and
subshells) of electrons in atoms.
• Ignoring a few exceptions, the electron configuration of an atom can
be deduced from the element’s position on the periodic table
• http://www.ptable.com/#Orbital
PERIODIC TRENDS
• For many atomic properties, trends within the periodic table (and relative
values for different atoms and ions) can be qualitatively understood and
explained using Coulomb’s Law, the shell model, and the concept of
shielding/effective nuclear charge.
Zeff = (Atomic #) – (# of shielding e-)
• These properties include:
• First ionization energy
• Atomic and ionic radii
• Electronegativity
• Typical ionic charges
PERIODICITY
• It is useful to understand the trends in the periodic table when building
molecules since replacing an element with an element in the same
group may lead to similar properties within the molecule.
• Ex. Since SiO2 can be ceramic, SnO2 may be as well.
LIGHT
• Early physicist believed light acted like a wave and thus did
not have mass. Later this was proven wrong with what is
called the dual nature of light.
• Dual nature of light: light acts as both a wave and a particle
• This is due to Einstein proposing that light is made up of a
stream of tiny particles called photons
• E=mc2
LIGHT AND ENERGY
• The energy of a photon is related to the frequency of the
electromagnetic wave through Planck’s equation (E=hv).
• Planck’s constant (h) = 6.626 x 10-34 Js
• When a photon is absorbed (or emitted) by a molecule, the energy of the
molecule is increased (or decreased) by an amount equal to the energy of the
photon.
E = hv
c = λv
Where c is the
speed of light
3.0x108m/s and λ is
the wavelength.
PHOTOELECTRON SPECTROSCOPY
• Photoelectron spectroscopy (PES) is the study of electrons emitted by
an atom as a result of shining a light upon it.
• In the photoelectric effect, incident light ejects electrons from a material. This
requires the photon to have sufficient energy to eject the electron.
• Photoelectron spectroscopy determines the energy needed to eject
electrons from the material. Measurement of these energies
provides a method to deduce the shell structure of an atom. The
intensity of the photoelectron signal at a given energy is a measure
of the number of electrons in that energy level.
PES
• The electronic structure of atoms with multiple electrons can be
inferred from evidence provided by PES.
• http://www.chem.arizona.edu/chemt/Flash/photoelectron.html
• Different types of molecular motion lead to absorption or emission of photons in
different spectral regions.
• Infrared radiation is association with transitions in molecular vibrations and so can be
used to detect the presence of different types of bonds.
• Ultraviolet/visible radiation is associated with transitions in electronic energy levels
and so can be used to probe electronic structure.
HOMEWORK
• P. 321-324
• # 31, 35, 67, 79, 85, 87
LEWIS DOT STRUCTURE REVIEW
• Lewis diagrams can be constructed according to a wellestablished set of principles.
• Atoms must achieve noble gas configurations.
LEWIS DOT STRUCTURE REVIEW
• In cases where more than one equivalent Lewis structure
can be constructed, resonance must be included as a
refinement to the Lewis structure approach in order to
provide qualitatively accurate predictions of molecular
structure and properties (in some cases).
FORMAL CHARGES
• Formal Charge: the difference between the number of
valence electrons on the free atom and the number of the
valence electrons assigned to the atom in the molecule.
• Formal charge can be used as a criterion for determining which
of several possible valid Lewis diagrams provides the best model
for predicting molecular structure and properties.
CALCULATING FORMAL CHARGES
Formal Charge = (# valence e- on free atom) – (# of valence e- assigned
to atom in molecule)
BOND POLARITY
• Two or more valence electrons shared between atoms of
identical electronegativity constitute a nonpolar covalent
bond.
• Two or more valence electrons shared between atoms of
unequal electronegativity constitute a polar covalent bond.
POLAR COVALENT BONDS
• The difference in electronegativity for the two atoms
involved in a polar covalent bond is not zero.
• The atom with a higher electronegativity will develop a
partial negative charge relative to the other atom in the
bond.
• For diatomic molecules the partial negative charge on the more
electronegative atom is equal in magnitude to the partial positive
charge on the less electronegative atom.
POLAR COVALENT BOND CONT.
• Greater differences in electronegativity lead to greater
partial charges, and consequently greater bond dipoles.
• Bond Dipoles: a polarity within a bond; when a bond has a
center of positive charge and a center of negative charge.
• Typically this is shown via an arrow pointing towards the atom
with the partial negative charge.
• The sum of partial charges in any molecule or ion must be
equal to the overall charge on the species.
DIPOLE MOMENTS
• Dipole moment: occurs when a
molecule has a center of positive
charge and a center of negative
charge.
• If the sum of the bond dipoles do not
cancel each other out, the molecule
is said to have a dipole moment.
• All polar molecules have dipole
moments.
VSEPR
• Valence Shell Electron Pair
Repulsion
• The VSEPR model uses the
Coulombic repulsion
between electrons as a basis
for predicting the
arrangement of electron
pairs around a central atom.
VSEPR CONT.
• The combination of Lewis diagrams with VSEPR model provides a
powerful model for predicting structural properties of many covalently
bonded molecules and polyatomic ions, including the following.
• Molecular Geometry
• Bond Angles
• Relative Bond Energies Based on Bond Order
• Relative Bond Lengths (multiple bonds, effects of atomic radius)
• Presence of a dipole moment
LEWIS STRUCTURE LIMITATIONS
• As with any model, there are limitations to the use of the Lewis
structure model, particularly in cases with an odd number of valence
electrons.
• Recognizing that Lewis diagrams have limitations is of
significance.
• Students don’t need to know the exceptions themselves, but simply that there are
exceptions to the octet rule
• Boron, PCl5, etc.
HOMEWORK
• Pg. 383-386
• # 25a, 25b, 25d, 67, 75, 81, 82
GRAPHING BOND FORMATION
• The formation of a nonpolar covalent bond can be represented
graphically as a plot of potential energy vs. distance for the
interaction of two identical atoms.
• The relative strengths of attractive and repulsive forces as a function of
distance determine the shape of the graph.
• The bond length is the distance between the bonded atoms’ nuclei, and
is the distance of minimum potential energy where the attractive and
repulsive forces are balanced.
GRAPHING CONTINUED
• The bond energy is the energy
required for the dissociation of
the bond. This is the net
energy of stabilization of the
bond compared to the two
separated atoms.
• Typically, bond energy is given
on a per mole basis.
LATTICE ENERGY
• Lattice Energy: the change in energy that takes place when separated
gaseous ions are packed together to form an ionic solid.
• Energy required to bring molecules together to form crystals.
• Based on Coulomb’s Law
IONIC CRYSTALS
• The cations and anions in an ionic crystal are arranged in a
systematic, periodic 3-D array that maximizes the attractive forces
among cations and anions while minimizing repulsive forces.
• Coulomb’s Law describes the force of attraction between the cations
and anions in an ionic crystal.
• Because the force is proportional to the charge on each ion, large charges lead
to stronger interactions.
• Because the force is inversely proportional to the square of the distance between
the centers of the ions (nuclei), smaller ions lead to stronger interactions.
HYBRIDIZATION
• Organic chemists commonly use the terms “hybridization” and
“hybrid orbital” to describe the arrangement of electrons around the
central atom.
• When there is a bond angle of 180◦, the central atom is said to be sp hybridized;
• for 120◦, the central atom is sp2 hybridized;
• …and for 109◦, the central atoms is sp3 hybridized.
• Students should be aware of this terminology, and be able to use it.
• Students do not need to know the hybridization of molecules with expanded octets
(more than four pairs of electrons on the center atom). Students are responsible for
the shape of the molecule.
SIGMA VS. PI
• Bond formation is associated with overlap between atomic
orbitals. In multiple bonds, such overlap leads to the
formation of both sigma and pi bonds.
• The overlap is stronger in sigma than pi bonds, which is reflected in
sigma bonds having larger bond energy than pi bonds.
• The presence of a pi bond also prevents the rotation of the bond,
and leads to structural isomers.
• Structural isomers: where the molecules contains the same atoms, but one or
more bonds differ.
OVERLAP CONT.
• In systems such as benzene, where atomic p-orbitals overlap strongly
with more than one other p-orbital, extended pi bonding exists, which
is delocalized across more than two nuclei.
• Such descriptions provide an alternative description to resonance in Lewis
structures.
• A useful example of delocalized pi bonding is molecular solids that conduct
electricity. The discovery of such materials at the end of the 1970’s overturned a
long-standing assumption in chemistry that molecular solids will always be
insulators.
COMPARING ATOMIC MODELS
• Molecular orbital theory describes covalent bonding in a
manner that can capture a wider array of systems and
phenomena than the Lewis of VSEPR models.
• Molecular orbital diagrams, showing the correlation
between atomic and molecular orbitals, are useful
qualitative tools related to molecular orbital theory.
Use the details of modern atomic theory to explain each of the following
experimental observations.
(a) Within a family such as the alkali metals, the ionic radius increases as
the atomic number increases.
(b) The radius of the chlorine atom is smaller than the radius of the
chloride ion, Cl-. (Radii : Cl atom = 0.99Å; Cl- ion = 1.81 Å)
(c) The first ionization energy of aluminum is lower than the first
ionization energy of magnesium. (First ionization energies: 12Mg = 7.6
ev; 13Al = 6.0 ev)
(d) For magnesium, the difference between the second and third
ionization energies is much larger than the difference between the first
and second ionization energies. (Ionization energies for Mg: 1st = 7.6
ev; 2nd = 14 ev; 3rd = 80 ev)
Answer:
(a) The radii of the alkali metal ions increase with increasing atomic number because –
(1) the outer principal quantum number (or shell or energy level) is larger. OR
(2) There is an increase in shielding. (3) The number of orbitals increases.
(b) The chloride ion is larger than the chlorine atom because - (any of these)
(1) the electron-electron repulsion increases.
(2) the electron-proton ratio increases.
(3) the effective nuclear charge decreases.
(4) shielding increases.
(c) The first ionization energy for Mg is greater than that for Al because - (either of these)
(1) the 3p orbital (Al) represents more energy than the 3s orbital (Mg) represents.
(2) the 3p electron in an Al atom is better shielded from its nucleus than a 3s electron in a
Mg atom.
(3) [half credit] a 3p electron is easier to remove than a 3s electron.
(d) In a Mg atom, the first two electrons lost are removed from the 3s orbital whereas the 3rd
electron comes from a 2p orbital; a 2p orbital is much lower in energy than the 3s is; so
more energy is needed to remove a 2p electron.
Explain each of the following observations using principles of atomic
structure and/or bonding.
(a) Potassium has a lower first-ionization energy than lithium.
(b) The ionic radius of N3- is larger than that of O2-.
(c) A calcium atom is larger than a zinc atom.
(d) Boron has a lower first-ionization energy than beryllium.
Answer:
(a) potassium’s valence electron is 4s while lithium’s is 2s . potassium’s electron is
shielded by more electrons than lithium and is therefore more easily removed at a
lower energy.
(b) The addition of electrons to a neutral atom produces an anion that is significantly
larger than its parent atom. Even though both ions are isoelectronic, there is a greater
nuclear positive charge in the oxide ion causing its electrons to be more tightly pulled
toward the nucleus.
(c) Even though a zinc atom contains 10 more electrons than calcium, these are all 3d
electrons, filling an inner shell, not adding another larger one. There is a corresponding
increase of 10 more protons for the zinc and this increase in nuclear charge pulls the
electrons in more tightly and reducing its size.
(d) Boron’s last electron is 2p and it receives the benefit of effective shielding by its
completed 2s electrons. Thus it is easier to remove this electron.
1
1
1
In the SO2 molecule, both of the bonds between sulfur and oxygen
have the same length. Explain this observation, supporting your
explanation by drawing in the box below a Lewis electron-dot
diagram (or diagrams) for the SO2 molecule.
On the basis of your Lewis electron-dot diagram(s) in part (c), identify
the hybridization of the sulfur atom in the SO2 molecule.
• SO2 is a resonance structure that switches
between the two forms and “evens out” the
bond length
• sp2
Structures of the pyridine molecule and the benzene molecule are
shown below. Pyridine is soluble in water, whereas benzene is not
soluble in water. Account for the difference in solubility. You must
discuss both of the substances in your answer.
water is polar and can form hydrogen bonds since it has a hydrogen
attached to an oxygen; the lone pair of electrons on the nitrogen
creates a slightly polar nitrogen and it can hydrogen bond to the
hydrogen in the water. A C-H bond, as those in benzene, is nonpolar and can not hydrogen bond with water. Since there is little
attraction between water and benzene (a non-polar molecule) and
“like dissolves like”, benzene will not dissolve in water but the
polar pyridine will.