Unit 6: Chemical Bonding Guided Notes
I.
Ch. 6.1 Introduction to Chemical Bonding
A.
Warm Up:
1.
What is the total number of electrons in all the
phosphorus?
2.
s-orbitals of a neutral atom of
The electron configuration of a neutral calcium atom is shown below.
1s2 2s2 2p6 3s2 3p6 4s2
How many valence electrons are in the atom?
B.
Introduction to Chemical Bonding
1.
Most atoms do not exist freely as isolated atoms in nature (except for _____________
_________________).
2.
C.
a)
lower energy state than as free atoms
b)
filled _____________ _______________ ______________
Chemical Bond – a link (glue) that holds two atoms together.
1.
2.
has.
D.
Most atoms tend to bond to other atoms to form more stable arrangements.
What part of an atom is involved in chemical bonding?
a)
Electrons in the highest energy level.
b)
These electrons are called __________________ electrons.
The periodic table can be used to determine the number of valence electrons an atom
a)
s-block - number of valence electrons equal to group number.
b)
p-block - number of valence electrons equal to group number minus 10
c)
d-block - usually 2 valence electrons - Why?
Electron-dot notation is used to represent an atom’s valence electrons.
1.
Write the symbol of the atom.
2.
Use dots to represent valence electrons. Put one dot on each side of symbol before
doubling (RLTB).
a)
3.
Example: Sulfur (S)
The maximum number of valence electrons an atom can have is ______________.
E.
Octet Rule
1.
Noble gas atoms are unreactive because their electron configurations are especially
stable.
a)
This stability results from the fact that the noble-gas atoms’ outer s and p
orbitals are completely filled by a total of eight electrons. (except He)
2.
Octet rule: chemical compounds tend to form so that each atom, by gaining, losing, or
sharing electrons, has eight electrons in its highest energy level (except H).
F.
Types of Chemical Bonds
1.
The two of the most common types include:
a)
Ionic bonding – involves the _______________ of valence electrons from one
atom to another.
b)
Covalent bonding – involves ________________ of valence electrons between
two atoms.
II.
Ch. 6.2 Ions and Ionic Bonding
A.
Warm Up:
1.
How many moles are contained in 6.0 g of carbon?
2.
When two chemicals are mixed together in a beaker that is sitting in an ice water bath,
the ice in the water melts as the chemicals react. Is the reaction exothermic or endothermic?
Explain your reasoning.
B.
Ions and Ionic Bonding
1.
One way for atoms to reach stability is by either _____________ or ______________
electrons.
C.
2.
When this occurs an ______________is formed.
3.
Ion –
4.
Occurs when the number of protons is _______________to the number of electrons.
Cations
1.
Whenever atoms lose electrons they form ________________.
2.
Cations are ________________________charged atoms.
3.
_________________tend to form cations.
a)
Example: Sodium
D.
Anions
1.
Whenever atoms gain electrons they form ____________________.
2.
Anions are _____________________charged atoms.
3.
____________________tend to form anions.
a)
E.
Example: Chlorine (Cl)
We can use the periodic table to predict the type of ion an atom will form.
Group Number
1
2
13
15
16
17
F.
Type of Ion formed
Ionic Bonding
1.
Ionic bonding involves the _________________of electrons from one atom to another.
2.
We can use electron dots to show how ionic bonds are formed.
a)
Example: Na and Cl
3.
_____________________ charges serve as the glue that holds the atoms together in the
bond.
4.
Ionic bonds are extremely ________________chemical bonds.
5.
Ionic bonds usually involve a combination of a ________________ with a
__________________.
G.
Use electron dots to demonstrate how calcium (Ca) and bromine (Br) form a stable
compound by ionic bonding. Give the chemical formula for the compound formed.
H.
Let’s check your understanding
1.
Describe whether each statement describes an ANION or CATION.
2.
3.
a)
Ions usually formed by metals ____________________________
b)
Ion formed when an atom loses an electron ____________________________
c)
Negatively charged ion ____________________________
d)
Ion that contains more protons than electrons____________________________
e)
Type of ion sulfur forms to become stable ____________________________
Using the periodic table, predict the ion charge each of the following atoms will form:
a)
phosphorus (P)
(d) aluminum (Al)
b)
calcium (Ca)
(e) lithium (Li)
c)
iodine (I)
(f) oxygen (O)
Provide the following information.
Magnesium Atom
Magnesium Ion
a)
Number of Protons____________
Number of Protons______________
b)
Number of Electrons____________ Number of Electrons______________
4.
Use electron-dot structures to demonstrate the formation of ionic compounds involving
the following elements and give the formula for the stable compound formed.
a)
magnesium (Mg) and sulfur (S)
b)
aluminum (Al) and chlorine (Cl)
c)
barium (Ba) and nitrogen (N)
III.
Ch. 6.3 Covalent Bonding
A.
Warm Up:
1.
Use electron-dot structures to demonstrate the formation of ionic compounds involving
the following elements. Give the formula for the stable compound formed.
B.
a)
magnesium (Mg) and sulfur (S)
b)
aluminum (Al) and chlorine (Cl)
c)
barium (Ba) and nitrogen (N)
Covalent Bonding
1.
Covalent Bond: Valence electrons ________________ between two atoms.
2.
Molecule – a group of _____________________ bonded atoms.
a)
Molecules usually consist of a nonmetal bonded to a nonmetal or a metalloid
bonded to nonmetal.
b)
Example: H2
(1)
s-orbital of one hydrogen atom overlaps with the s-orbital of an other
hydrogen atom.
3.
A dash is used to represent a shared pair of electrons.
a)
4.
C.
Example: Cl2
Single covalent bond – share ________ __________ of electrons between two atoms.
Drawing Molecules
1.
We can use electron dot notation to help show how the atoms in a molecule are held
together.
2.
These drawings are known as _______________ __________________.
3.
All atoms seek to fill their outer energy shell, most follow the Octet rule, where they will
have 8 electrons on their outer level.
a)
Hydrogen only wants 2, as it’s in the 1s energy level.
DRAWING LEWIS STRUCTURES NOTES
A Lewis structure is a diagram that shows how the atoms in a molecule are bonded together.
Use the following rules to help you determine the Lewis structure of a molecule.
Rule #1: Determine the total number of ____________ electrons in the molecule.
Example: Draw the Lewis structure for NCl3.
Nitrogen
_____ atom x
Chlorine
_____ atoms x
_____ valence electrons
_____ valence electrons
=
_____
=
_____
_______ total valence electrons
Rule #2: If the molecule has more than two atoms, then determine the __________ atom in
the molecule. The central atom is usually the atoms with the ____________ valence
electrons (except for ____________ because it can only form one covalent bond)
Example: __________ is the central atom in this example.
Rule #3: Connect all other atoms to the central atom using a __________. A dash represents
a __________ pair of electrons between two atoms.
Rule #4: Make outer atoms stable by using _____________.
Rule #5: If any electrons left over, give them to the _______________ atom. Check to make
sure all atoms are stable.
Another example: Draw the Lewis structure for H2O
IV.
Ch. 6.4 Multiple Covalent Bonding and Polyatomic Ions
A.
Warm Up:
1.
The half-life of gold-198 is about 3 days. How many half-lives have elapsed when 12.5%
of a gold-198 sample remains in a container?
2.
When an electron moves from its ground state to an excited state, the electron moves
from a _______ energy level to a _______ energy level and energy is _____________.
B.
Multiple Covalent Bonds
1.
Sometimes it is possible to form more than one covalent bond between two atoms.
2.
Example: O2
a)
Double covalent bond – share _________ pairs of electrons between the same
two atoms.
3.
Example: N2
a)
Triple covalent bond – share __________ pairs of electrons between the same
two atoms.
4.
C.
Double and triple covalent bonds are referred to as multiple covalent bonds.
LEWIS STRUCTURES WITH MULTIPLE COVALENT BONDS
1.
Example:
CH2O
D.
If all electrons have been used and the central atom is not stable, then consider a
multiple bond.
E.
Convert two dots to a dash.
1.
F.
V.
Example:
CO2
Polyatomic Ions
1.
Group of covalently bonded atoms that has a charge. (a charged molecule)
2.
Example: SO4-2 (sulfate ion)
Ch. 6.5 Polar and Nonpolar Bonds
A.
Warm Up:
1.
Would sulfuric acid (H2SO4) be considered an element, compound, or mixture? Explain
the reasoning of your choice.
2.
B.
How many atoms are contained in 97.6 g of platinum (Pt)?
Polar and Nonpolar Covalent Bonds
1.
Covalent bonding involves sharing electrons between atoms.
2.
Nonpolar covalent bond – __________ sharing of electrons between two atoms.
3.
a)
Both atoms have same attraction for shared pair.
b)
Example: H – H
Polar covalent bond – ______________ sharing of electrons between atoms.
a)
One atom has greater attraction for shared pair. (Electronegativity Tug – of –
War)
b)
4.
Example: H – Cl
This creates partial (d) charges on each atom in the bond.
C.
The atom that has a _______________ _____________ for shared electrons takes on a
partial negative charge. The atom with a ________________ _______________ takes on a
partial positive charge.
1.
How can you determine which atom has greater attraction for electrons?
2.
Can use the _________________ in electronegativities between two atoms to predict
the type of bond formed between the atoms.
3.
Table of values on page 161 in book
D.
Use electronegativities to determine the type of bond that will form. If the bond is
polar, decide which atom is partially negative and which is slightly positive.
1.
C and S
2.
Na and O
3.
P and Cl
Bond Type
Nonpolar
Covalent
Polar Covalent
Ionic
Electronegativity
Difference
VI.
Ch. 6.6 H. Chem: VSEPR Theory and Molecular Geometry
A.
Warm Up:
1.
Distinguish between the Bohr model and the quantum mechanical model of the atom.
2.
Carbon-14 is a radioactive isotope that decays by releasing a beta particle. Write the
nuclear equation that represents this decay.
B.
Molecular Geometry and VSEPR Theory
1.
Whenever the atoms of a molecule bond together, the molecule takes on a particular
_____________.
2.
________________ _______________ refers to the three-dimensional arrangement of
a molecule’s atoms in space.
3.
_______________ _______________ is the model used to explain the shapes of
molecules.
C.
VSEPR –
D.
The main idea of this theory is that electron pairs on the central atom of a molecule will
position themselves in such a way as to ______________________ repulsion.
E.
There are two types of electron pairs:
1.
2.
F.
G.
Example – CH4
1.
_____ electron pairs
2.
_____ bonding / _____ nonbonding pairs
3.
molecular geometry = ________________________
Example – NH3
1.
_____ electrons pairs
2.
_____bonding / _____ nonbonding
3.
molecular geometry = __________________
H.
I.
Example – H2O
1.
_____ electron pairs
2.
_____ bonding / _____ nonbonding
3.
molecular geometry = ___________________
VSEPR and Multiple Covalent Bonds
1.
For the VSEPR model, multiple bonds are counted as ______________________ when
determining the shape of the molecule.
2.
3.
4.
J.
Example – CH2O
a)
_____ electron pairs
b)
_____ bonding / _____ nonbonding
c)
molecular geometry = __________________
Example – NOCl
a)
_____ electron pairs
b)
_____ bonding / _____ nonbonding
c)
molecular geometry = ______________
Example – CO2
a)
_____ electron pairs
b)
_____ bonding / _____ nonbonding
c)
molecular geometry = _________________
COMMON MOLECULAR GEOMETRIES
Bonding Pairs
Nonbonding Pairs
Shape
Example
2
0
Linear
CO2
3
0
Trigonal Planar
CH2O
2
1
Bent
NOCl
4
0
Tetrahedral
CH4
3
1
Trigonal Pyramidal
NH3
2
2
Bent
H 2O
K.
Your Turn:
1.
Predict the shape of the following molecules
a)
PBr3
b)
TeF2
c)
SiO2
VII. H. Chem: Ch. 6.7 Dipoles and Molecular Polarity
A.
B.
Warm Up:
1.
Determine the molecular geometry of PCl3.
2.
Are the P–Cl bonds in PCl3 polar or nonpolar? Explain the reasoning of your choice.
Dipoles and Molecular Polarity
1.
We have discussed that covalent bonds can be classified as either _______________ or
_________________.
2.
_______________ bonds are created when there is an _______________ sharing that
occurs between two atoms.
C.
Molecules as a whole can also be classified as either ______________ or
_________________.
1.
This affects many of the properties of the substance such as ___________________ and
________________.
D.
Polar Molecules
E.
___________ molecules are created when the center of partial positive charge
________ ______ coincide with the center of partial negative charge.
1.
Example: HCl
2.
Polar molecules are often called _________________.
a)
A ___________ is created when the center of partial positive charge and the
center of partial negative charge are separated by a ____________________.
F.
For molecules that contain more than one bond, polar molecules are created when the
polar bonds are _______________________ arranged.
1.
G.
Example: H2O
a)
Molecule has a ________________ shape.
b)
The molecule has two distinct regions of __________________.
c)
Therefore, water is considered a ________________ molecule.
Nonpolar Molecules
1.
Molecules with polar _____________ do not always create polar
____________________.
2.
a)
Example: CCl4
b)
Shape is _____________________.
This time the centers of positive charge and negative charge _________________.
a)
This has the net effect of __________________ each other.
b)
CCl4 is a __________________ molecule.
3.
Whenever polar bonds are _________________ arranged, they produce
___________________ molecules.
a)
H.
Another example: CO2
Molecular Polarity and Molecular Geometry
1.
The _____________ of the molecule determines how the bonds are ________________
in the molecule.
2.
The following geometries give asymmetrical arrangements and produce
______________ molecules:
a)
3.
The following geometries give symmetrical arrangements and produce
__________________ molecules:
a)
I.
CHCl3 is commonly known as chloroform. Would this molecule be considered polar or
nonpolar? Explain the reasoning of your choice.
J.
Just remember “SNAP”
1.
SNAP
2.
Molecular polarity is determined by the ______________ and
_________________________ in the molecule. To make your life easier, just look the atoms.
3.
If the atoms in the molecule are symmetrical, the charges are balanced by each other.
The molecules are considered to be __________________.
K.
4.
However, if the molecule is asymmetrical, it is considered to be ________________.
5.
So if it is lopsided, it is ______________. If it is balanced, it is __________________.
Intermolecular Forces
1.
The negative region in one polar molecule _______________ the positive region in
adjacent molecules. So the molecules all attract each other from opposite sides.
2.
The forces of attraction between molecules are known as __________________
____________.