VI – Chemical Bonds
Chemistry Student Notes
Chemistry Unit VI – Chemical Bonds
PRE-TEST QUESTIONS
1. Which element is the second-most electronegative of all?
2. Which element is the least electronegative of all?
3. How does an atom become a cation?
4. What is the charge of a cation?
5. What is the generic electron configuration for noble gases?
I. Ions
A. Valence Electrons
1. Elements in the Periodic Table are placed in the same group because they have similar
______________ ______________.
2. Since Mendeleev first created the table, it has been found that the reason they have
similar chemical properties is due to the fact that they all have the same number of
valence electrons in their outermost energy level.
3. ______________ ______________ are the electrons found in the ______________
occupied energy level of an atom.
a. To find the number of valence electrons in an atom of a main group element, simply
look at its ______________ ______________
Group Outer e¯ configuration # Valence e¯ s
1
2
13
14
15
16
17
18
4. Why is this important?
a. Valence electrons are usually the only electrons involved in chemical __________.
5. Electron dot diagrams, or Lewis dot structures, are diagrams that show valence
electrons as dots.
Electron Dot Diagrams for Some Group A Elements
1
2
13
14
15
16
17
18
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Chemistry Student Notes
B. The Octet Rule
1. The Octet Rule, as described by Gilbert Lewis, states that atoms react with one another
by ______________ or ______________ a number of electrons in order to acquire the
stable electron structure of a noble gas, which is usually ___________ valence electrons.
2. This configuration means that an atom fills to ______________.
3. Atoms of metals tend to lose their valence electrons, leaving a complete octet in the next
lowest energy level. Atoms of some nonmetals tend to gain electrons, or even share
electrons with another nonmetal, to achieve a complete ______________.
a. Draw…
4. There are a few exceptions to the octet rule:
a. Hydrogen and Helium only need ______________ valence electrons.
b. Boron can have only ______________ valence electrons.
c. Phosphorous and Sulfur can have expanded octets of _____ and ______, respectively.
C. Formation of Cations
1. A cation is a ______________ ion due to
the ______________ of electrons.
2. When metals form cations, their ion’s
name is the same as the neutral atom.
a. e.g., Na (sodium) forms __________,
a ______________ ion.
3. Very different, though. Sodium metal
reacts explosively with water, but sodium
ions are fairly unreactive, and simply
dissolve in water, like in NaCl, for
instance.
4. In forming cations, atoms ___________
electrons in order to acquire electron
configurations like the noble gases.
a. The number of ______________ remains the ______________, however, so the
element’s identity does not change.
5. Look at sodium: Na → Na+ + e¯
a. It goes from 1s22s22p63s1 1s22s22p6
b. Now, it has the same electron configuration as ______________ (a noble gas)!
6. Look at calcium: Ca → Ca2+ + 2e¯
a. It goes from 1s22s22p63s23p64s2 1s22s22p63s23p6
b. Now, it has the same electron configuration as ______________ (a noble gas)!
D. Transition metal cations
1. Put simply: Transition metals are funky (their charges vary).
2. As they fill their outermost d sublevel, they often violate the ______________ rule and
form various ions.
3. Silver, being 1s22s22p63s23p63d104s24p64d105s1 cannot look like Xenon, but if it loses the
1 in 5s, it does grain a pseudo noble-gas electron configuration that is stable.
a. Thus, silver forms a +1 ion: ______________.
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VI – Chemical Bonds
Chemistry Student Notes
4. Transition metals are much more complicated than this in general, as many form
______________ than one type of ion. That will be discussed more later…
E. Formation of Anions
1. An anion is a ______________ ion because it has ______________ electrons.
2. The names of anions are not the same as the element.
a. Their ending is changed to –ide.
b. e.g., chlorine (Cl) forms ______________, called a ______________ ion.
3. Because they have relatively full valence shells, atoms of nonmetallic elements attain
noble-gas electron configurations more easily by gaining electrons than by losing them.
4. Fluorine goes from 1s22s22p5 1s22s22p6
a. F + e¯ F¯
b. Named ______________.
c. Looks like ______________ (a noble gas)!
5. Oxygen goes from 1s22s22p4 1s22s22p6
a. O + 2e¯ O2b. Named ______________.
c. Looks like ______________ (a noble gas)!
6. All of the halogens need ______________ more electron to complete their octet.
a. They form ions collectively called _________ _______ and have a charge of ______.
II.
Ionic Bonds and Ionic Compounds
A. Formation of Ionic Compounds
1. Ionic compounds are composed of ______________ and ______________.
2. Although they are composed of charged ions, ionic compounds are electrically
______________ overall, because the ______________ charge ______________.
B. Ionic Bonds
1. Cation and anions, having opposite charges, are ___________________ to one another.
2. Ionic bonds are the electrostatic attraction that binds oppositely charged ions together.
a. They are formed by the ______________ of electrons from one atom to another
(usually from a metal to a nonmetal), which results in positively and negatively
charged ions which are then attracted to one another.
b. Draw sodium reacting with chlorine….
3. Ionic compounds do not form basic units: instead they are continuous ______________
of positive and negative ions alternating.
C. Chemical Formulae
1. A chemical formula shows the ______________ and ______________ of atoms in the
smallest representative unit of a substance.
2. For ions, this is represented as the smallest ______________________ ratio of cations to
anions called a ______________ ______________.
a. NaCl does not mean there is a bunch of NaCl units. Instead, it means that there is
____ Na+ ion for every _____ Cl¯ ion.
i. No charge is written in the formula because it is ______________ overall.
b. MgCl2 means that there is ____ Mg2+ ion for every _____ Cl¯ ions.
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Chemistry Student Notes
i. Again, no charge is written in the formula because it is ______________ overall.
D. Properties of Ionic Compounds
1. Most ionic compounds are crystalline ______________ at room temperature.
2. The coordination number of an ion is the number of ions of opposite charge that
______________ the ion in a crystal.
a. NaCl, for instance, consists of each Na+ ion surrounded by ____ other Cl¯ ions and
each Cl ¯ ion is surrounded by ____ Na+ ions. Its coordination number is _____.
b. CsCl has a coordination number of ____ because Cs is a bigger atom than Na.
3. Ionic compounds generally have ______________ melting points.
a. The crystal lattice structure of ionic compounds make them very ______________.
e.g., NaCl melts at about ______________.
4. Ionic compounds can conduct an electric current when melted or dissolved in water.
a. Either way, the ions have to be able to ______________ around.
III.
Bonding in Metals
A. Metallic Bonds and Metallic Properties
1. Metals can be thought of as large ______________ that are loosely attracted to their
______________ ______________.
2. These valence electrons can be modeled as a “______________ of ________________.”
a. They are mobile and drift from atom to atom, belong to all metal cations at one time.
3. Metallic bonds consist of the attraction of the free-floating valence electrons for the
positively-charged metal ions.
4. Draw the metallic bonding below…
5. The “sea of electrons” helps explain many ______________ of metals.
a. The mobile electrons allow metal to be ______________ conductors of electricity.
i. After all, electricity is the “______________ of electrons.”
b. Because metallic bonding is the same in every direction, metals do not shatter when
hit, they bend and flatten: this is called ______________.
i. Ionic compounds, however, ______________ when hit because the ions line up
and repel each other.
c. Metals are also capable of being pulled into wires: ______________.
B. Alloys
1. Alloys are mixtures composed of two or more elements, at least one of which is a metal.
2. The properties of alloys are often superior to those of the component elements.
a. i.e., you get the ____________________ without the _____________________.
3. Some common alloys:
a. Brass –
b. Bronze –
c. Sterling silver –
d. Steel – ______________ and ______________; often combined with other elements
for specific properties of the steel.
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Chemistry Student Notes
IV. Molecular Compounds
A. Ionic Vs. Covalent Bonding
1. When metals react with nonmetals to form a compound, they have a ______________
difference in electronegativity values.
a. This difference is usually either ______ or _______ or higher.
b. As a result of this huge difference, electrons are completely transferred from the
______________ to the ______________, resulting in positive and negative ions that
are attracted to one another.
c. Draw…
2. When two nonmetals form a compound, their electronegativities are very similar in value.
a. This difference is usually less than ________.
b. Because they’re so similar, they do not completely give up or gain the electrons.
c. Instead, they ______________ them!
d. Draw…
3. When atoms share their electrons in a chemical bond, it is called a _____________
______________.
B. Molecules
1. A molecule is a ______________ group of atoms joined together by covalent bonds.
a. Many elements are found in nature in the form of molecules.
b. Molecules are single, basic units of matter, whether of an element or a compound.
2. A diatomic molecule is a molecule consisting of only two atoms.
a. There are seven _________________ that exist as diatomic molecules that you need
to know:
i. Hydrogen (H2)
ii. Oxygen (O2)
iii. Nitrogen (N2)
iv. Fluorine (F2)
v. Chlorine (Cl2)
vi. Bromine (Br2)
vii. Iodine (I2)
3. Molecular compounds tend to have relatively ______________ melting and boiling
points than ionic compounds.
a. Thus, many molecular compounds are ______________ or ______________ at room
temperature.
C. Molecular Formulae
1. A molecular formula is the chemical formula of a molecular compound.
a. Shows how many atoms of each element a molecule contains.
2. Subscripts are written ______________ the symbol they modify and below. Ones are not
written and are assumed.
a. H2O means there are ___ hydrogen atoms bound to ____ oxygen atom.
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VI – Chemical Bonds
Chemistry Student Notes
b. Note that whereas an ionic formula unit is the ratio of cations to anions, a molecule is
an actual basic unit with a definite number of atoms in each molecule.
3. The shortcomings of a molecular formula, though, are that they do not tell you about a
molecule’s structure.
a. They do not indicate a molecule’s ______________.
b. They do not show which atoms are bonded to ______________.
V.
The Nature of Covalent Bonding
A. Introduction to Covalent Bonding
1. In forming covalent bonds, electron sharing usually occurs so that the atoms of each
element attain the electron configuration of ______________ ______________.
2. Lewis structures, or Lewis diagrams, help represent the covalent bonds by a pair of dots
between atoms to represent a shared pair of electrons.
a. Or, the pair of shared electrons can be represented by a dash ( — ).
B. Single Covalent Bonds
1. Two atoms held together by sharing a single pair of electrons are joined by a single
covalent bond.
2. Hydrogen has ____ valence electron.
a. In order to have the same configuration as Helium, a noble gas, it needs ___ more.
b. Two hydrogen atoms share their ____ valence electron so that each hydrogen atom
has two valence electrons.
c. Thus, each hydrogen atom looks like helium, with ____ valence electrons.
d. Draw Hydrogen (H2)…
3. The halogens have ______________ valence electrons and need only one more to
achieve an octet. Thus, they form ______________ covalent bonds to have their octet.
a. Draw Fluorine (F2)…
4. Notice that the fluorine atoms above only share one pair of electrons. The other pairs of
electrons that are not involved in bonding are called an ______________
______________, also known as a __________ ____________ or a nonbonding pair.
C. Double Covalent Bonds
1. Sometimes, atoms need more than one electron to achieve an octet and become
chemically stable.
2. A double covalent bond involves ______________ shared pairs of electrons.
3. Oxygen has ____ valence electrons. When it bonds with another oxygen atom to form an
oxygen molecule (O2), it forms a double covalent bond.
a. Draw…
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Chemistry Student Notes
D. Triple Covalent Bond
1. A triple covalent bond involves ______________ shared pairs of electrons.
2. Nitrogen has ____ valence electrons. When it bonds with another nitrogen atom to form a
nitrogen molecule (N2), it forms a triple covalent bond.
3. Draw…
E. Lewis Structures
1. A Lewis Structure, or Lewis diagram, shows which atoms are bonded to which atoms in
a molecule.
2. The bonding electrons are represented either as a pair of dots or as a dash.
3. Lone pairs are represented as a pair of dots.
4. How to draw Lewis structures:
a. Count the number of valence electrons of each atom in the formula.
b. Determine the ______________ atom
i. Usually the ______________ electronegative element.
ii. NEVER ______________.
iii. If ______________ is present, it’s a good choice.
c. Bond each atom to the central atom with a PAIR of electrons.
d. Continue doing this until all atoms are bonded.
e. Once all atoms have been used and are present in the Lewis Structure, then make sure
all atoms have an ______________.
f. If you do not have enough electrons to complete octets, make double bonds.
5. A couple tips and tricks before getting started:
a. Carbon will form as many as ______________ single bonds, 1 triple and 1 single, 2
doubles, etc. where they add up to four bonds total.
b. Halogens and Hydrogen will never form more than ______________ bond.
6. Let’s try some!
a. Hydrofluoric acid (HF)
b. Methane (CH4)
c. Ammonia (NH3)
d. Water (H2O)
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e. Propane (C3H8)
f. Ethanol (C2H5OH)
VI. Miscellany
A. Coordinate Covalent Bonds
1. A coordinate covalent bond happens when one atom contributes ______________ of
the bonding electrons.
2. In a structural formula, you can show the coordinate covalent bonds as ______________
that point from the atom that donates the pair of electrons to the atom that receives them.
3. For example, Carbon monoxide…
B. Polyatomic Ions
1. A polyatomic ion is a tightly bound ______________ of atoms that has a positive or
negative charge overall and behaves as a unit.
2. Polyatomic ions are covalently bonded, but since they have an overall charge, they act as
an ______________.
a. Technically, they are both __________________ and ______________.
3. Some common polyatomic ions include:
a. ______________, NH4+
b. ______________, OH¯
c. ______________, NO3¯
d. ______________, SO42e. ______________, PO43C. Bond Dissociation energies
1. The bond dissociation energy is the energy required to break the bond between
______________ covalently bonded atoms.
2. A large bond dissociation energy corresponds to a ______________ covalent bond.
a. Double bonds are harder to break than single bonds.
b. Triple bonds are harder to break than double bonds.
D. Resonance
1. A resonance structure is when a molecule can have ______________ than one Lewis
structure represent its molecule.
a. This usually occurs when the atom is symmetrical and has single and double bonds.
2. Each resonance structure is drawn, and a double-headed arrow is used to show that both
structures are representative of that compound.
3. The actual bonding of the single and double bonds in a resonance structure is actually a
hybrid, or mixture, of the extremes represented by the resonance forms.
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Chemistry Student Notes
4. Ozone (O3)
E. Exceptions to the Octet Rule
1. The octet rule cannot always be satisfied in all molecules, such as molecules whose total
number of valence electrons is an odd number. There are also molecules in which an
atom has fewer, or more, than a complete octet of valence electrons.
2. Nitrogen dioxide (NO2)
a. Has a total of _____ valence electrons.
b. It is impossible to draw a dot structure for NO2 that satisfies the octet rule for all
atoms. And yet, it exists as a stable molecule.
3. Boron trifluoride (BF3)
a. Boron is deficient by ______________ electrons and has only ____ electrons.
4. Some atoms, like Phosphorous and Sulfur will expand their octets to include ten or
twelve electrons.
VII. Bonding Theories
A. Molecular Orbitals
1. Just as an atomic orbital belongs to a particular atom, a molecular orbital belongs to a
molecule as a whole.
a. Molecular orbitals are formed and filled when atoms bond in a molecule.
2. Sigma bonds
a. A sigma bond (σ) is a bond formed when two atomic orbitals ______________ to
form a molecular orbital that is symmetrical around the axis connecting the two
atomic nuclei.
b. Both s orbitals (as in H2) and p orbitals (as in F2) can form sigma bonds.
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Chemistry Student Notes
c. Draw the sigma bond for s orbitals and for p orbitals…
3. Pi Bonds
a. In the sigma bond of the fluorine molecule, the p atomic orbitals overlap end-to-end.
b. A pi bond (π) is a covalent bond in which the bonding electrons are most likely to be
found in sausage-shaped regions above and below the bond axis of the bonded atoms.
c. Pi bonds are weaker than sigma bonds because pi bonds overlap less.
d. Draw a pi orbital…
e. There are up to two pi orbitals possible per bond.
i. A ______________ bond consists of a sigma bond plus a pi bond.
ii. A ______________ bond consists of a sigma bond plus two pi bonds.
B. VSEPR Theory
1. Molecules do not exist in the two-dimensional plane like we draw on paper. They are 3dimensional and this affects the way they interact and react.
2. VSEPR theory, or valence-shell electron-pair repulsion theory, explains the threedimensional shape of molecules.
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Chemistry Student Notes
3. VSEPR theory states that electron pairs ______________ one another and will orient
themselves into a shape in order to orient themselves as far apart as possible.
4. ______________ pairs of electrons are also important in predicting the shapes of
molecules, since they will repel the electrons involved in bonding as well.
5. To determine VSEPR geometry, first figure out which atom is the central atom.
a. Then figure out how many atoms are bonded to it (double and triple bonds count as
single bonds for this purpose).
b. Determine how many lone pairs are attached.
c. Use the reference sheet to determine the shape.
6. Some common examples of the shapes you need to know:
a. HCl
b. CO2
c. BF3
d. CH4
d. NH3
e. H2O
7. Why is this important?
a. The shape affects how molecules interact with one another and how they will react, as
well as hydrogen bonding and other intermolecular forces.
C. Hybrid Orbitals
1. In hybridization, several atomic orbitals mix to form the same total number of
equivalent hybrid orbitals.
2. Methane hybridizes by taking the three orbitals in 2p and the one orbital in 2s and
hybridizing them to make 4 orbitals of equal energy called an _______.
3. Ethene hybridizes by taking 2 of the p orbitals and the one s orbital to form 3 orbitals of
equal energy called an ______.
a. It has a ______________ bond!
4. Ethyne (acetylene) hybridizes by taking just 1 of the p orbitals and the one s orbital to
form 2 orbitals of equal energy called an _____.
a. It has a ______________ bond!
VIII. Polar Bonds and Molecules
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VI – Chemical Bonds
Chemistry Student Notes
A. Bond Polarity
1. Covalent bonds usually occur between two atoms of different elements, which also have
different electronegativity values.
2. Electrons are ______________ between the atoms like in tug-of-war.
3. A polar covalent bond, known also as a polar bond, is a covalent bond in which the
electrons are shared unequally.
a. The more electronegative atom attracts the electrons more strongly and gains a
slightly ______________ charge. The less electronegative atom has a slightly
______________ charge.
b. The lower-case Greek letter ______________ (δ) is used to show the locations of the
partial charges (they are less than the full charge that exists on a proton or electron).
c. Dipole arrows are also used. They’re drawn along each bond, pointing towards the
__________________ pole.
d. HCl is a polar molecule, as are water (H2O) and ammonia (NH3), due to their shapes.
4. A nonpolar covalent bond is a covalent bond in which the electrons are shared equally
between the two atoms.
a. Molecules of ______________ (H2, N2, O2, etc.) are nonpolar covalent.
b. Carbon dioxide is also nonpolar covalent. (Draw)
B. Polar Molecules
1. Electronegativity differences can be used to predict the type of bond formed.
Electronegativity Difference Most Probable Type of Bond
Example
Range
2. A molecule that has two poles is called a dipolar molecule, or ______________.
3. When polar molecules are placed between two oppositely charged plates, they tend to
become oriented with respect to the positive and negative plates.
C. Intermolecular Forces
1. Intermolecular attractions are weaker than either ionic or covalent bonds, but are still
strong enough to explain several properties of various compounds.
2. Van der Waals Forces
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Chemistry Student Notes
a. The two weakest attractions between molecules are collectively called van der
Waals forces.
b. Consists of dipole interactions and dispersion forces.
3. London dispersion forces (the weakest of these three)
a. As electrons constantly move around in the atom, at some random moment in time,
they all ______________ ______________ on one side of the atom, so that that side
gets a small ______________ charge for a brief moment in time.
b. This can induce a dipole in other, nearby, atoms and attract the atoms together.
c. This is how dry ice (solid CO2) and liquid nitrogen (N2) – both ______________ –
can exist in a phase other than ______________!
d. This is the only intermolecular force that affects molecules and atoms of everything –
even ______________ molecules and even ______________ ______________, too!
4. Dipole interactions (a van der Waals force)
a. Dipole interactions occur when ______________ molecules are attracted to one
another.
b. Kind of like ions, but much weaker.
5. Hydrogen bonding (the strongest of these three forces)
a. Hydrogen bonds are attractive forces in which a hydrogen covalently bonded to a
very electronegative atom is also weakly bonded to an unshared electron pair of
another electronegative atom.
i. Note: “electronegative atom” here refers specifically to ____, _____, _____
b. Extremely important force in ______________.
c. Explains why ice is less dense than liquid water, why water has a high boiling point,
surface tension, why water droplets are ______________, etc.
D. Network Solids
1. Network solids are solids in which all of the atoms are covalently bonded to each other.
2. They have extremely high melting points because there are so many covalent bonds that
have to be broken.
3. ______________ is an excellent example of this.
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