Unit 4: Chemical Bonds - goldchemistry … · · 2013-05-21If we look at the orbital notation for...
Transcript of Unit 4: Chemical Bonds - goldchemistry … · · 2013-05-21If we look at the orbital notation for...
Unit 4: Chemical Bonds Chapter 7-9
Objectives
26 Identify the number of valence electrons for elements and their Lewis dot structure
27 Define the terms cation and anion including radius size and charge
28 Determine the isoelectronic electron configurations for atoms and their ions including the ionic charges
29 Identify the properties of ionic bonds
30 Predict the shape of molecules using the VSEPR theory
31 Identify the bonds between certain elements within a compound as non-polar, polar, or ionic
32 State and identify the three intermolecular forces including London dispersion forces and how they affect melting points, dipole forces, and hydrogen bond forces
33 Convert between formula and chemical name for covalently bonded molecules, binary ionic compounds, polyatomic ionic compounds, and hydrates
34 Identify the dissociation factor of compounds
26 Valence Shells
Before we discuss bonds, we need to
determine the number of electrons in the
valence shell of an ion.
This is accomplished by looking at the
electron configuration.
Valence Shells
O: 1s2 2s2 2p4
Consider oxygen
It’s outer energy level is the second energy
level.
This tells us there are 6 valance electrons.
Lewis Dot Structures
Once we know the number of valence
electrons, it is possible to give a visual
representation.
This is called a Lewis Dot structure.
It requires a dot for each valence electron
surrounding the element symbol.
Paired electrons should still be depicted.
Lewis Dot Structures
Oxygen has 6 valence electrons.
If we look at the orbital notation for just its valence
shell, we get the following:
2p ___ ___ ___
2s ___
Its Lewis Dot structure would look as follows.
O
Lewis Dot Structures
O
Notice how oxygen has two electrons that are not paired up.
This indicates that oxygen would like to gain two more electrons so it has 8 total electrons.
All elements are most stable with 8 valence electrons.
This is known as the Octet Rule.
◦ There are 5 exceptions: hydrogen, helium, lithium, beryllium, and boron.
◦ These prefer to have 2 electrons in their valence shell.
27 Cations and Anions
Atoms can gain or lose electrons.
When atoms lose electrons, they become
positive and are called cations.
When atoms gain electrons, they become
negative and are called anions.
Ionic Radii
Nuclear charge holds electrons a certain
distance from the nucleus.
As a cation is formed, there are less
electrons for the nucleus to hold.
This allows the nucleus to pull the outer
energy levels slightly closer.
Ionic Radii
As an anion is formed, there are more
electrons for the nucleus to hold.
The nucleus does not have enough charge to
hold the extra electrons as close, and as a
result, the radius increases slightly.
28 Isoelectronic Configurations
As discussed in Unit 3, each atom has an
electron configuration to show where each
electron belongs.
◦ For example: Al: 1s2 2s2 2p6 3s2 3p1
When ions are formed, the electrons are
either added to the last energy levels or
taken from the last energy levels.
◦ For example: assume we take three electrons
from aluminum.
Al+3: 1s2 2s2 2p6 Aluminum’s 3s and 3p orbitals
are now empty.
Isoelectronic Configuration
The term isoelectronic can be broken down
into:
◦ Iso: same
◦ Electronic: electrons
Therefore, the term means the same electron
configuration as another element.
In chemistry, this refers to a noble gas.
Isoelectronic Configuration
If we look at an element, the number of
electrons it holds are close to a noble gas.
This means it will tend to gain or lose
electrons until it matches that noble gas.
◦ For example: Oxygen has 8 electrons and is thus
close to neon’s 10.
◦ Neon has a configuration of 1s2 2s2 2p6 .
◦ Oxygen has a configuration of 1s2 2s2 2p4 .
◦ For oxygen to be isoelectronic with neon, it
needs two more electrons, thus oxygen tends to
gain electrons to get: O-2: 1s2 2s2 2p6
Charges
The elements marked below will always carry the charge indicated.
The elements in white can have charges that vary.
◦ These will be determined with a Roman numeral.
Transition
Metals
29 Ionic Properties
Every ionic compound will follow certain properties.
They are:
◦ Ionic compounds form crystalline structures.
◦ Ionic compounds are brittle.
◦ Ionic compounds have high melting and boiling points.
◦ Ionic compounds as solids will not conduct electricity.
◦ Ionic compounds in solution will conduct electricity.
30 Molecular Shapes
Covalently bonded molecules will create
different shapes.
These shapes are controlled by the bonds
formed and the paired electrons.
To determine the shapes, the valence shell
electron pair repulsion (VSEPR) theory is
used.
VSEPR Theory
VSEPR Theory states that the valence
electrons in a molecule will position
themselves so they are as far away from the
other electrons as possible.
To determine the shapes of molecules using
this theory, the number of bonds and pairs
of electrons must be determined.
VSEPR Theory
Recall Lewis Dot structures from the
previous unit.
◦ Each dot represents a valence electron.
Oxygen has six valence electrons so its
Lewis Dot structure is as follows:
O These paired electrons are
called lone pairs because they
belong to only one atom.
VSEPR Theory
The Lewis Dot structures can be used to
determine the bonds created.
Take water which is H2O.
O H O H
H
H When a bond is formed, the
electrons are shown with a
single line between the
atoms.
VSEPR Theory
Notice that the water molecule still contains two
sets of lone pair electrons.
These electrons will force the hydrogens to create
a bent shape.
O H
H
VSEPR Theory
There are five shapes the basic covalent
molecules will create.
Each can be determined by looking at the
central atom of the molecule.
The two components to look at are the
number of atoms bound and the number of
lone pairs on the central atom.
VSEPR Theory
Shape Lone Pairs of
electrons on
Central Atom
Atoms bonded
to the Central
Atom
Example
Linear 0 2 CO2
Bent 1 or 2 2 H2O or HNO
Trigonal Planar 0 3 BF3
Trigonal Pyramidal 1 3 NH3
Tetrahedral 0 4 CH4
In addition, any molecule that has only two atoms will be linear.
i.e.: oxygen gas, O2
31 Intramolecular Bonds
When a bond is created, it is often stated
that the atoms share their electrons.
While this is true to some degree, the
sharing is not always equal.
Each atom has its own electronegativity.
◦ The tendency of an atom to attract a bonded
electron to itself.
◦ The greater an atom’s electronegativity, the more
time the electrons will spend near that atom.
Bond Strength
This pulling of the electrons towards one
atom can create partial charges.
Each bond can be classified as either
nonpolar covalent, polar covalent, or ionic.
To determine how atoms share, compare
their electronegativity values.
◦ The greater the difference, the more ionic
character will be present in the molecule.
Intramolecular Bonds
Nonpolar Polar Covalent Ionic
Covalent
0 0.5 2.1
Difference in Electronegativity
As we look at the difference in electronegativity, we can use the chart
above to determine the type of bond.
Take water for example: H2O
◦ The bonds formed are between hydrogen and oxygen.
◦ They have the following electronegativities: 2.2 for H and 3.4 for O
◦ The difference between the two is 1.2 and thus the bond is polar covalent.
◦ Because oxygen has the larger electronegativity, the electrons spend more time near
oxygen. This creates a partial negative charge on the oxygen atom and a partial
positive charge around the hydrogen atom.
32 Intermolecular Forces
While intramolecular forces occur inside
a molecule, intermolecular forces occur
between molecules.
Three intermolecular forces exist:
◦ London Forces
◦ Dipole Forces
◦ Hydrogen bonding
Intermolecular Force Strength
Each of the intermolecular forces hold
molecules together.
However, certain forces are stronger.
◦ London is the weakest.
◦ Dipoles use the partial charges as an attractive
force making them stronger than London.
◦ H-Bonding is the strongest because of the
partial charges and the use of hydrogen.
Determining Intermolecular Forces
To determine the intermolecular force on a
molecule, it is necessary to know whether it
is polar or not.
◦ If polar, the molecule will have a partial positive
and partial negative side.
◦ This can be determined using the
electronegativities and the Lewis Dot structure.
Determining Intermolecular Forces
Water has a Lewis Dot structure as shown.
Oxygen has an electronegativity of 3.4 while hydrogen has an electronegativity of 2.2.
◦ This would mean that oxygen is partial negative and the hydrogens are each partial positive.
Because the molecule has a postive and negative side, it is considered polar.
If the molecule would have had the same charge on the entire outside, it would be considered non polar.
Intermolecular Forces
Force Strength Polar/Nonpolar Unique
Characterestics
H-Bonding Strongest Polar Must contain H
and either O, N, F
Dipole Forces
Medium Polar
London Forces
Weak Nonpolar
Dipole forces and H-Bonding are the only forces that are polar, but H-bonding
has element requirements. If a molecule is polar but does not contain one of
the elements listed above, it must be a dipole.
33 Writing Binary Formulas and
Names
Binary compounds refer to compounds that
contain 2 elements.
When writing the name of a binary
compound, list the first element exactly as it
appears on the Periodic Table.
For the second element, drop the ending of
the element’s name and add –ide.
Binary Names-Examples
NaCl
sodium chloride
(chlorine drops the –ine)
CaBr2
calcium bromide
(bromine drops the –ine)
Writing Binary Formulas and
Names
Writing the formulas from the names
requires the use of the charges.
It is important to balance the positive charge
with the negative charge.
This is done by adding subscripts to the
elements.
Binary Formulas - Examples
calcium phosphide
Calcium has a +2 charge.
Phosphorus has a -3 charge
To balance their charges, we need to have multiple atoms.
If we add another calcium, we will have an overall charge of +4 to -3.
Let’s add another phosphorus to give us a overall charge of +4 to -6.
Since we are only off by 2, adding another calcium will balance the charges at +6 to -6.
Ca3P2
Writing Formulas
We got to the answer on the last slide by using logical method of adding one atom at a time.
This can also be done by looking for the least common multiple.
Since we had charges of 2 and 3, the least common multiple is 6.
◦ (2 x 3 = 6)
◦ (3 x 2 = 6)
Therefore, the atom with the charge of 2 requires 3 atoms and the atom with the charge of 3 requires 2.
Ca3P2
Writing Formulas
magnesium bromide
MgBr2
Magnesium has a +2 charge.
Bromine has a -1 charge.
The least common multiple is two.
◦ 1 x 2 = 2
◦ 2 x 1 = 2
Transition Metals
If you recall from the slide on charges,
transition metals did not have a defined
charge.
Their charges vary and thus a Roman
numeral is used to determine their charge.
This Roman numeral is always listed directly
after the metal.
Transition Metals
Fe2O3
For the formula above, we know Fe is iron and O is oxide (oxygen as an ion)
Oxygen has a -2 charge and since there are 3, this compound has a overall -6 charge.
Since we have to have a +6 charge as well, we have to consider a number times 2 to give 6. (? x 2 =6)
In this case, the answer would be three.
Therefore, the name of the compound is iron (iii) oxide.
Transition Metals
Manganese (ii) chloride
In this compound, Manganese has a +2
charge and chlorine has a -1 charge.
Therefore, the least common multiple is 2.
MnCl2
Polyatomic Ions
Some elements will combine covalently
(Unit 6) and still have a charge.
As long as they have a charge, they can
create ionic compounds.
They are treated as though they are single
entities.
Naming the Polyatomics
From a formula, naming the polyatomic ionic
compounds requires element to be named
and the polyatomic ion.
For instance:
CaSO4
◦ Ca represents calcium
◦ SO4 represents sulfate
The name of this compound is calcium sulfate.
Writing formulas
Writing formulas from the name works the
same as binary compounds.
◦ Determine the charge on each.
◦ Find the least common multiple.
◦ Add the proper subscripts.
If more than one polyatomic ion is required,
add parenthesis around the ion before
adding the subscript.
Writing formulas-Example
Iron (iii) nitrate
Iron (iii) refers to Fe+3
Nitrate refers to NO3-1
Therefore, the least common multiple is 3
and three nitrates are required.
Fe(NO3)3
The parenthesis tells us that there are 3 N and 9 O.
Hydrates
Hydrates are unique ionic compounds that
attract water.
Each hydrate is surrounded by a certain
number of water molecules.
These water molecules need to be identified
in both the formula and the name.
Hydrates
To indicate a hydrate, a dot is used to indicate a weak bond.
The number of water molecules are indicated with a numeric prefix.
The ionic part of the compound is named as previously described.
For example: copper (ii) sulfate pentahydrate
CuSO4 • 5H2O
Molecular Nomenclature
When naming covalent molecules, first
identify each element.
◦ If there is more than one of the first element,
add the appropriate prefix to the front of its
name.
The second element should always includes
its prefix and its ending should change
to –ide.
Molecular Nomenclature
CO2
To name this compound, first identify each element.
Since there is only one of the first element, no prefix is needed.
There are two of the second element so the prefix di- will be added. Notice the ending of oxygen has already be changed.
carbon oxide di
Molecular Nomenclature
Writing the formulas will simply work in the
opposite direction.
Identify and write the symbol for each
element.
Use the prefixes to determine the subscript
of each element.
Molecular Nomenclature
Tetraphosphorus decoxide
First, record the symbols for each element.
Tetra- indicates four so the subscript on phosphorus will be four.
Deca- means ten so the subscript on oxide will be ten
P O P4O10
34 Dissociation Factors
Dissociation factors describe how many
pieces an ionic compound can divide into.
This is calculated by adding the subscripts of
each ion.
Be careful because the one’s are omitted
when writing formulas.
Dissociation Factors
Assume we have calcium chloride:
CaCl2
If this molecule breaks apart, we will have 1
calcium ion and 2 chloride ions.
This means the dissociation factor is 3.
Dissociation Factors
The same idea applies to polyatomics:
Ca(NO3)2
If this molecule breaks apart, we will have 1
calcium ion and 2 nitrate ions.
This means the dissociation factor is 3.
The polyatomic ions do not break apart.
This concludes the tutorial on
measurements.
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Prefixes
mono 1
di 2
tri 3
tetra 4
penta 5
hexa 6
hepta 7
octa 8
nona 9
deca 10
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Valence Shell: Outer most energy level of an
atom