UNIT 2: STRUCTURE & PROPERTIES OF MATTER
Big Ideas

the chemical elements are fundamental building materials of matter, and all matter can be
understood in terms of arrangements of atoms

chemical and physical properties of materials can be explained by the structure and arrangement
of atoms, ions, or molecules and the forces between them

technological devices that are based on the principles of atomic and molecular structures can have
societal benefits and costs.
The Quantum Mechanical Model of the Atom
Success Criteria:
 I can explain Bohr’s model of the atom and the conclusions drawn from studying an element’s
line spectrum
 I can explain the major refinements to Bohr’s model and the currently accepted quantum
mechanical model.

it is important to note that the model of the atom is not an exact description of
the atom, but rather a theoretical construct that fits a set of experimental data

in the early 1900s, Niels Bohr reasoned that electrons move around the nucleus
in specific circular orbits or energy levels

he came up with his theory by analyzing the line spectrum of hydrogen and
studying the work of Planck and Einstein

Max Planck proposed that energy can only be released or absorbed by atoms
in discrete chunks or quanta; light itself behaves as quantized energy particles
called photons

Bohr concluded that electrons could only exist in a series of “allowed” orbits;
electrons could jump between levels by absorbing or releasing photons carrying
an amount of energy equal to the difference in energy levels

therefore, the line spectra of elements that Bohr observed were caused
by photons emitted from electrons that dropped from higher energy
levels to lower energy levels
 large drops (ex. from 2 to 1) emit so much energy that lines
appear in the ultraviolet end of the spectrum
 drops of medium energy (ex. 3 to 2 or 4 to 2) emit energy that
is in the visible spectrum
 small drops (ex. from 6 to 5 or 5 to 4) emit a small amount of
energy and the lines appear in the infrared end of the spectrum
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
in 1924, Louis de Broglie proposed that matter has wave-like properties just
as energy has particle-like properties

large objects have wavelengths so small, that there is virtually no observable
effect on the object’s motion; for microscopic objects however (i.e. an
electron), the effect of wavelength on motion becomes significant

in 1926, Werner Heisenberg proposed his uncertainty principle; it stated that due to the dual
nature of matter it is impossible to know both the position and momentum of an electron at once

shortly after, Erwin Schrödinger combined the ideas of matter waves,
photons, and the uncertainty principle to form the quantum mechanical
model of the atom

Schrödinger used wave equations to define the probability of finding an
electron; each solution represents a wave function or orbital (these
equations were later developed by the German physicist Max Born.)

orbitals can be visualized by plotting three-dimensional electron probability
distribution graphs; the contour representations enclose about 90% of the total
electron density for the orbital

these do not show the actual paths of electrons but where there is a high
probability of finding them; although probabilities decrease as you move
away from the nucleus, they never reach zero

every electron in an atom has a unique set of four quantum numbers that describe the energy
level, orbital, orientation and spin of an electron

Principle quantum number (n)
 it is a positive integer that specifies the energy level of an atomic orbital and its relative size
 energy levels get closer together as n increases and each level can hold 2n2 electrons in it

Secondary or angular quantum number (l)
 describes an orbital’s shape (i.e. s, p d, f etc.)
 all orbitals within a sublevel are degenerate meaning they have the same energy

Magnetic quantum number (ml)
 distinguishes the orientation of one orbital from another within the same sublevel (i.e. 1 s, 3
p, 5 d, 7 f etc.)
 an orbital may have a maximum of two electrons only

Spin quantum number (ms)
 describes the direction by which an electron is spinning about its axis
 the Pauli exclusion principle proposes that only two electrons of
opposite spin can occupy the same orbital
Homework:
Read: 6.1 The Wave Nature of Light, 6.2 Quantized Energy and Photons, 6.3 Line
Spectra and the Bohr Model, 6.4 The Wave Behavior of Matter, 6.5 Quantum Mechanics
and Atomic Orbitals, 6.6 Representations of Orbitals
Answer: 6.15, 6.33, 6.49, 6.51, 6.53, 6.55
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Electron Configurations and Orbital Diagrams
Success Criteria:
 I can write long-form and short-form electron configurations for any atom or ion.
 I can apply Hund’s rule and the Pauli exclusion principle when drawing orbital diagrams.

an electron configuration is a shorthand notation that represents the number and arrangement of
electrons in its orbitals

there are an infinite number of electron configurations for an atom; however, chemists use the
ground state configuration (ex. F - 1s22s22p5)

an atom’s chemical reactivity depends mainly on its valence electrons; these electrons can be
highlighted by using a shorthand configuration

this notation places the nearest noble gas from the preceding period in square brackets followed
by the outer energy levels (ex. F - [He]2s22p5)

an orbital diagram shows an atom’s orbitals and
their relative energies, as well as the electrons that
occupy them; higher orbitals are closer together in
energy than lower orbitals

when filling orbital diagrams, three rules need to be
followed:
 the aufbau principle: it states that each
additional electron that is added will occupy
the lowest available energy level
 Hund’s rule: when electrons are added to
orbitals of the same energy level, each
orbital receives one electron before any
pairing occurs
 the Pauli exclusion principle: electrons are shown with either a spin up or spin down

some elements have configurations that are different from the expected because filled or halffilled d orbitals are more stable
ex.
Group 6 - Cr - [Ar]3d54s1,
Group 11 - Cu - [Ar]3d104s1
Homework:
Read: 6.7 Many Electron Atoms, 6.8 Electron
Configurations, 6.9 Electron Configurations
and the Periodic Table
Answer: 6.63, 6.65, 6.67, 6.69, 6.71, 6.73
Review: Chemical Bonding
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
bonds form to give the atoms greater
stability and the lowest possible
energy; bonding is a continuum
(bonding character changes with
changing differences in
electronegativity - ∆EN)

in general, bonds with a ∆EN between
0 and 0.5 are mostly covalent, bonds
with a ∆EN of 0.5 to 1.7 are polar
covalent, and bonds with a ∆EN of 1.7
and 3.3 are mostly ionic
Characteristics
Bonding
Ionic
A force of attraction between
oppositely charged ions
Electronegativity
Occurs between elements that have a
large difference in electronegativity
(i.e. > 1.7)
High
Crystalline, hard, brittle
Solid
High
High
Melting and boiling points
Physical properties
State at room temperature
Conductivity in solution
Solubility in water
Covalent
The sharing of pairs of electrons
accomplished by the formation of
a new orbital from overlapping
orbitals
Occurs between elements that
have a small difference (i.e. <1.7)
in electronegativity
Low
Soft
Liquid or gas
Low
Low

a polar covalent bond is a bond with an unequally shared pair of electrons resulting in a dipole;
one end of the molecule has a partial positive charge and the other a partial negative charge

a coordinate covalent bond occurs when one atom contributes both electrons (ex. H3O+)

many molecules, especially organic ones, require
resonance structures to represent their bonding; this
occurs when more than one Lewis structure is possible (ex. SO2)

resonance is caused by delocalized electrons that are shared within hybrid orbitals

resonance bonds are a hybrids of both single and double bonds or double and triple bonds;
therefore, bond lengths are in-between the lengths of these bonds

“symmetrical” molecules are non-polar no matter how polar the individual bonds are; all basic
VSEPR arrangements are non-polar if all sites are occupied by identical atoms (ex. BeF2, BF3,
CCl4, PF5, SF6)

“asymmetrical” molecules are polar molecules if: the bonds are polar (ex. HCl), the central atom
has lone pairs (ex. NH3), the central atom is attached to more than one type of atom (ex. CH3Cl)
Valence-Shell Electron-Pair Repulsion (VSEPR) Theory
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Success Criteria:
 I can apply the VSEPR framework to determine the bonding pairs and lone pairs for any molecule
or polyatomic ion.
 I can draw these molecules and ions and determine their geometry, shape and bond angles.

Lewis structures are 2-D structures that do not tell any information about molecular shape; in
1957, two Canadian scientists developed the VSEPR model for predicting the shape of molecules

VSEPR theory proposes that molecular shape depends on the repulsions between all bonding
pairs (BP) and non-bonding pairs (LP) in the valence energy level of an atom

thus, electron pairs reside in orbitals which are as far apart as possible; this results in a minimum
total energy for the molecule

repulsion is greatest between two lone pairs; repulsion is the least between bonded pairs;
repulsion between a lone pair and bonded pair is in between

therefore, bond angles decrease as the number of lone pair electrons increases; also, multiple
bonds exert a greater repulsive force than single bonds

when all electron groups are bonding pairs, a molecule will have one of the five basic geometric
arrangements

variations in these arrangements will occur if one or more of the electron groups includes a lone
pair (E); note that double and triple bonds are considered one bonding pair
1) linear (AX2)
2 BP
180o
BeF2
CO2
2) trigonal planar (AX3)
angular

3 BP
2 BP, 1LP
120o
BF3
SnCl2
SO3
3) tetrahedral (AX4)
trigonal pyramidal

angular

4 BP
3 BP, 1LP
2 BP, 2LP
109.5o CH4
NH3
H2O
4) trigonal bipyramidal (AX5)
seesaw

T-shaped

linear

5 BP
4 BP, 1LP
3 BP, 2LP
2 BP, 3LP
90o
120o
PCl5
SF4
ClF3
XeF2
5) octahedral (AX6)
square pyramidal

square planar

6 BP
5 BP, 1LP
4 BP, 2LP
90o
SF6
IF5
XeF4
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Homework:
Read: 9.1 Molecular Shapes, 9.2 VSEPR Model, 9.3 Molecular Shape and
Molecular Polarity
Answer: 9.11, 9.13, 9.19, 9.21, 9.27, 9.29, 9.31, 9.33, 9.35, 9.41, 9.43
Valence Bond Theory & Hybridization
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Success Criteria:
 I can determine the hybridization scheme for any molecule or polyatomic ion.
 I can draw simple orbital diagrams to show hybridization.

VSEPR theory is generally accurate for predicting molecular geometry but does not explain why
and how covalent bonds are formed

the marriage of Lewis’s notion of bonding pairs and the idea of atomic orbitals leads to a model
of chemical bonding called valence-bond theory

simple molecules can be explained by the overlap of atomic orbitals to form covalent
bonds; this provides a way for a pair of electrons to be shared between two nuclei
ex.
ex.
the bond in H2 is caused by the overlap of two 1s orbitals
the bond in HCl is caused by the overlap of a 1s orbital from hydrogen and a 3p
orbital from chlorine

for many complex molecules, hybridization theory is needed to relate the formula, Lewis
structure, geometry, shape, and bonding

for these molecules, the s, p and d orbitals often mix to form new hybrid orbitals that are
stronger; these have new shapes and properties

not all hybrid orbitals of the central atom must be used for bonding, therefore lone pairs can
participate in hybridization

in many cases, electrons are bumped up or promoted to higher energy levels (ex. in most sp3d and
sp3d2 hybrids); also, filled or empty orbitals can overlap to form a coordinate covalent bond

there are 5 types of hybrid orbitals corresponding to each of the basic geometric arrangements
 linear – one s and one p orbital hybridize to form two sp orbitals
 trigonal planar – one s and two p orbitals hybridize to form three sp2 orbitals
 tetrahedral – one s orbital and three p orbitals hybridize to form four sp3 orbitals
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 trigonal bipyramidal – one s orbital, three p orbitals, and one d orbital hybridize to form
five sp3d orbitals
 octahedral – one s orbital, three p orbitals, and two d orbitals hybridize to form six
sp3d2orbitals

a sigma (σ) bond occurs when electron density is heavily concentrated between the nuclei of two
atoms

σ bonds are formed by an overlap of two s orbitals (ex. in H2), two p orbitals end to end (ex. in
F2), or between sp orbitals (ex. in HF); they are responsible for single bonds

bonds can also be formed by the sideways overlap of two p
orbitals; these are called pi (π) bond

pi bonds occurs when electron density is concentrated in
two separate regions that lie on opposite sides of the nuclei

a double bond consists of one σ bond and one π bond and a triple bond consists of one σ bond
and two π bonds; π bonds usually remain in unhybridized p orbitals and are weaker than σ bonds

in molecules that have multiple bonds and more than one resonance structure (ex. benzene, the pi
bonds are delocalized and spread among several atoms
Homework:
Read: 9.4 Covalent Bonding and Orbital Overlap, 9.5 Hybrid Orbitals, 9.6
Multiple Bonds
Answer: 9.47, 9.51, 9.53, 9.55, 9.59, 9.63, 9.65
Review: Intermolecular Forces
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
these are forces that act between atoms, molecules or ions; they were studied extensively by
Johannes van der Waals and can be categorized into various groups
Dipole-dipole forces

these attractions occur between polar molecules that orient
themselves so that oppositely charged ends of the molecules are
near one another (ex. HCl and HCl)

the energy required to separate polar molecules from one another
is much greater than that needed to separate non-polar molecules
of similar molar mass
Ion-dipole forces

these are forces of attraction between ions and polar molecules (ex. NaCl in H2O)
London dispersion forces

these occur between atoms that momentarily
becomes slightly polar due to the uneven
distribution of electrons

these forces occur between all particles; they are
the main forces that act between non-polar
molecules (ex. halogens and noble gases)

they are stronger in larger molecules because
they have more electrons
Hydrogen bonding

it is a particularly strong form of dipole-dipole force; it exists between an H atom and the lone
pair of the very electronegative O, N or F atom

this explains the high boiling
points of H2O, HF, and NH3
compared to other binary
hydrides

this bond is only 5% as strong
as a single covalent bond, but
multiple hydrogen bonds
creates a stronger molecule (ex.
DNA)
Bonding in Solids
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Success Criteria:
 I can determine the hybridization scheme for any molecule or polyatomic ion.
 I can draw simple orbital diagrams to show hybridization.

solids can be divided into two categories based on the arrangements of their atoms
 crystalline solids have organized arrangements and distinctive shapes (ex. quartz,
diamond)
 amorphous solids have indistinct shapes because their arrangements lack order (ex.
glass, rubber)

crystalline solids are built of identical repeating units called unit cells; these unit cells create a
three-dimensional crystal lattice

atomic solids
 these are made up of individual atoms held together solely by dispersion forces (ex. noble
gases)
 since these only experience London dispersion forces, they tend to have the lowest
melting and boiling points

molecular compounds
 these are made up of molecules (ex. solid methane or carbon dioxide)
 they tend to be soft and have low melting and boiling points

metals
 bonds are created by a dense core of cations with a delocalized
sea of free-electrons
 the free-electron model can explain the conductivity, ductility,
and malleability
 most metals have high melting and boiling points

ionic crystals
 these are formed by an array of ions arranged at regular
positions in a crystal lattice
 they tend to have very high melting and boiling points

network solid
 these are composed of atoms bonded covalently in
continuous two or three-dimensional arrays (ex. (SiO2)n)
 carbon can form various allotropes such as diamond ,
graphite, and fullerenes (C60)
 these tend to have the highest melting and boiling points of all solids
 carbon nanotubes are sheets of graphite rolled up; they are stronger
than metals of the same dimensions
Homework:
Read: 11.2 Intermolecular Forces, 12.1 Classification of Solids, 12.2
Structure of Solids, 12.3 Metallic Solids, 12.4 Metallic Bonding, 12.5
Ionic Solids, 12.6 Molecular Solids, 12.7 Covalent-Network Solids
Answer:12.11, 12.43, 12.63
Science, Technology, Society and the Environment
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
there are many technologies that are based on the principles of atomic and molecular structures
Photoelectron Spectroscopy (PES)

it is used to study the energy levels of atomic core electrons by using the photoelectric effect

a sample of element is radiated with photons from an X-ray or UV-source; absorption of the
photon causes the ejection of an electron from any energy level

a spectrum results showing the energy required to remove each electron giving information about
the number of electrons and their whereabouts
X-ray Crystallography

it is used to determine the structure of molecules (i.e. the double helical
nature of DNA was discovered this way)

X-rays are passed through a crystal and are diffracted; from the angles
and intensities of diffraction, a three-dimensional picture of the
electron density within the crystal is produced

from the electron density, the positions of the atoms in the crystal can
be determined, as well as information on the chemical bonds
Nuclear Magnetic Resonance (NMR)

it is used to investigate the properties of atoms and molecules

like electrons, nuclei possess an intrinsic and thus generate a
small magnetic field

when an external magnetic field is applied, the parallel
alignment of the nuclear magnetic field is lower in energy than
the anti-parallel

in the case of magnetic resonance imaging (MRI) machines, a person’s body is placed in a
magnetic field to create pictures with details which are useful for diagnosing diseases
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