Unit 5 - Atomic Theory and Chemical Bonding
Emission of Energy by Atoms (pg 284)
Energy Levels of Hydrogen (pgs 285-287)
Hydrogen Orbitals (pgs 289-292)
Electron Arrangements (pgs 295-298)
Electron Configuration (pgs 299-303)
Types of Chemical Bonds (pgs 317-319)
Electronegativity (pgs 319-321)
Stable Electron Configurations (pgs 323-326)
Lewis Structures (pgs 328-332)
Unit 5 - Atomic Theory and Bonding
Upon completion of this unit, you should be able to do the
following:
1. Write the electron configuration for an atom or ion using
the periodic table.
2. Be able to draw an orbital notation diagram for any atom
or ion.
3. Determine the valence electrons of an atom.
4. Predict the types of bonds formed between two atoms and
describe the properties of each.
5. Use electronegativity to predict the percent ionic character
of bonds and the polarity of molecules.
6. Draw Lewis dot structures to represent how atoms share
or transfer valence electrons to become more stable.
7. Explain how multiple bonds can form between the same
two atoms.
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Atomic Theory
• The concept of atoms explains many important
observations, such as why compounds always have
the same composition (a specific compound always
contains the same types and numbers of atoms) and
how chemical reactions occur (they involve a
rearrangement of atoms).
• We learned to picture the atom as a positively
charged nucleus composed of protons and neutrons
at its center and electrons moving around the
nucleus in a space very large compared to the
nucleus.
• In this unit, we develop a picture of the electron
arrangements in atoms.
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Emission of Energy by Atoms
• When compounds are heated, they emit a color
characteristic of the cation. Li+, for example, emits a
red flame when heated. Na+ emits a yellow flame,
Cu2+ a green flame.
• The colors of the flames result from atoms releasing
energy in the form of visible light of specific
wavelengths, or colors.
• The heat from the flame causes the atom to absorb
energy. The atom becomes excited. Some of the
excess energy is released as light. The atom moves
to a lower energy state as it emits a photon of light.
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Emission of Energy by Atoms
• When atoms receive energy, they become excited.
They can release the energy by emitting light. The
emitted energy is carried away by a photon.
• The energy of the photon corresponds exactly to the
energy change of the emitting atom.
• High energy photons correspond to short wavelength
light. Low energy photons correspond to long
wavelength light.
• The photons of red light have less energy than the
photons of blue light because red light has a longer
wavelength than blue light.
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Energy Levels of Hydrogen
• When we study the photons of visible light emitted,
we see only certain colors.
• Only certain types of photons are produced.
• Because only certain photons are emitted, only
certain energy changes are occurring.
• So, hydrogen atoms must have certain discrete
energy levels.
• We say the energy levels of hydrogen are quantized,
that is, only certain values are allowed.
• Energy levels of all atoms are quantized.
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Energy Levels of Hydrogen
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The Hydrogen Orbitals
The probability map is
called an orbital. The
orbital shown in Figure
10.20 is called the 1s
orbital and describes
the ground (lowest)
state of energy for
hydrogen.
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The Hydrogen Orbitals
Hydrogen has
discrete energy
levels. They are
called principal
energy levels and
labeled with an
integer. Each
principal energy
level has sublevels.
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The Hydrogen Orbitals
Principal level 2 has 2 sublevels. They are called 2s and 2p.
Principal level 3 has 3 sublevels called 3s, 3p and 3d.
Principal level 4 has 4 sublevels called 4s, 4p, 4d and 4f. 10
The Hydrogen Orbitals
The principal levels describe size and shape. The s
orbital is spherical. Level 1 is smaller than level 2, which
is smaller than level 3.
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The Hydrogen Orbitals
The three 2p orbitals are lobed, not spherical. They are
oriented along the x, y or z axis.
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Electron Arrangements
• An atom has as many electrons as it does protons, so
all atoms beyond hydrogen have more than one
electron.
• Each electron appears to spin like a top on its axis. It
can only spin in one direction. We represent spin
with an arrow, ↑ or ↓. Electrons in the same orbital
must have opposite spins.
• This leads to the Pauli exclusion principle: an
atomic orbital can hold a maximum of two electrons
and those two electrons must have opposite spins.
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Electron Arrangements
• Hydrogen has an atomic number of 1 (Z =1) and
therefore a single electron to have a net charge of
zero. To show its electron configuration, we write
the principal energy level followed by the sublevel,
1s. The number of electron in the orbital is placed as
a superscript, 1s1.
• The electron configuration can also be shown using
an orbital diagram, or box diagram, as below.
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Electron Configuration
• Hydrogen (Z=1)
• Helium (Z=2)
• Lithium (Z=3)
• Berylium (Z=4)
• Boron (Z=5)
• Carbon (Z=6)
• Nitrogen (Z=7)
• Oxygen (Z=8)
• Fluorine (Z=9)
• Neon (Z=10)
1s1
1s2
1s2 2s1
1s2 2s2 2p1
1s2 2s2 2p6
The orbital diagram for nitrogen is below.
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Electron Configuration
• Sodium (Z=11)
1s2 2s2 2p6 3s1 or [Ne] 3s1
• Magnesium (Z=12)
• Aluminum (Z=13)
• Silicon (Z=14)
• Phosphorous (Z=15)
• Sulfur (Z=16)
• Chlorine (Z=17)
• Argon (Z=18)
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Electron Configuration
• Valence electrons are the electrons in the outermost
(highest) principal energy level of an atom. These
are the electrons involved in bonding of atoms to
each other.
• Also note that the atoms of elements in the same
group have the same number of electrons in a given
type of orbital, except that the orbitals are in
different principal energy levels. Elements with the
same valence electron arrangement show very
similar chemical behavior.
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Electron Configuration
• The order of filling orbitals changes for Z=19.
Experiments show that the chemical properties of
potassium are very similar to lithium and sodium.
We predict that the 4s orbital will fill before the 3d
orbital. This means that principal energy level 4
begins to fill before level 3 is full.
• Potassium (Z=19)
• Potassium (Z=19)
• Calcium (Z=20)
• Scandium (Z=21)
[Ar] 3s2 3p6 3d 1
[Ar] 3s2 3p6 4s1
[Ar] 3s2 3p6 4s2
[Ar] 3s2 3p6 4s2 3d1
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Electron Configuration
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Electron Configuration
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Electron Configuration
Practice problems:
Determine the electron configuration for:
1. C (Z= 6)
2. Al (Z=13)
3. Cl (Z=17)
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Types of Chemical Bonds
• A bond is a force that holds two or more atoms
together and makes them function as a unit.
• In water, the fundamental unit is the H-O-H
molecule, which is held together by the two O-H
bonds.
Types of Chemical Bonds
• Ionic compounds are formed when an atom
that loses an electron relatively easily reacts
with an atom that accepts an electron. This
occurs when a metal reacts with a non-metal.
The resulting bonds are called ionic bonds.
• In an ionic bond, electrons are transferred.
Types of Chemical Bonds
• Consider diatomic hydrogen H – H .
• When two hydrogen atoms are brought close
together, the electrons are equally attracted
to both nuclei.
• When two similar atoms form a bond, the
electrons are equally attracted to the nuclei of
the two atoms. This is called a covalent bond.
• In a covalent bond, the electrons are shared
by nuclei.
Types of Chemical Bonds
• Ionic bonding and covalent bonding are
extremes. Between the extremes are cases
where atoms are not so different that
electrons are transferred, but different
enough that unequal sharing of the electrons
results. These bonds are called polar covalent
bonds.
• In HF, the fluorine atom has a stronger
attraction for the shared electrons than the
hydrogen atom does.
Types of Chemical Bonds
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Electronegativity
• The unequal sharing of electrons between two
atoms is described by a property called
electronegativity, the relative ability of an
atom in a molecule to attract shared electrons
to itself.
• The higher the atoms electronegativity value,
the closer the shared electrons tend to be to
that atom when it forms a bond.
• Fluorine has the highest electronegativity
value at 4.0. Cesium and Francium have the
lowest electronegativity value at 0.7
Electronegativity
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Electronegativity
• The polarity of a bond depends on the
difference between the electronegativity values
of the atoms forming the bonds.
• If the atoms have similar electronegativities, the
electrons are shared almost equally and the
bond shows little polarity.
• If the atoms have very different
electronegativities, a very polar bond is formed.
Electronegativity
• In extreme cases, one or more electrons are
actually transferred and ions and an ionic bond
are formed.
• Consider NaCl , for example. When a Group 1
metal reacts with a Group 17 element, ions are
formed and an ionic substance results.
Electronegativity
Page 321, example 11.1
Using the electronegativity values given in Figure
11.3, arrange the following bonds in order of
increasing polarity:
H-H, O-H, Cl-H, S-H, F-H
Stable Electron Configurations
• Representative metals form ions by losing
enough electrons to attain the configuration
of the previous noble gas that occurs before
the metal. For example, sodium will lose one
electron to attain the configuration of neon.
• Nonmetals form ions by gaining enough
electrons to attain the configuration of the
next noble gas. For example, chlorine will add
one electron to attain the configuration of
argon.
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Stable Electron Configurations
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Stable Electron Configurations
• In observing millions of stable compounds,
chemists have observed that in almost all
stable compounds, all of the atoms have
achieved a noble gas configuration.
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Stable Electron Configurations
• When a non-metal and a Group 1, 2 or 3 metal
react to form a binary ionic bond, the ions
form so that the non-metal completes the
valence-electron configuration of the next
noble gas and the metal empties the valence
orbitals to achieve the configuration of the
previous noble gas.
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Stable Electron Configurations
• When two non-metals react to form a
covalent bond, they share electrons in a way
that completes the valence-electron
configuration of both atoms.
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Lewis Structures
• Bonding involves just the valence electrons of
atoms. (See slide 18.) The Lewis structure is a
representation of a molecule that shows how
the valence electrons are arranged among the
atoms in the molecule. Dots are used to
represent the valence electrons.
Examples:
K
Br
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Lewis Structures
• Duet rule – hydrogen
• Octet rule – elements are surrounded by 8
electrons. Electrons that are shared form
bonds.
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Lewis Structures
Practice problems:
Draw the Lewis structure for:
1. HF
2. NH3
3. CH4
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