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Unit
Four
Regents
Chemistry
Mr Grodski III
III.
A. Introduction - We have learned in Unit One that a compound is a substance
made up of two or more elements bonded together. This chemical bond is the result of
the valence electrons of two or more elements interacting together. Also in Unit One we
learned that potential energy is stored in chemical bonds. The energy that is released
from these bonds is called chemical energy. Remember that a compound is decomposed
by chemical changes only. Chemical changes are due to the breaking and the formation
of chemical bonds in compounds. Our bodies obtain energy from food by creating and
breaking chemical bonds in the food we eat.
-Generally two atoms that compose or form a chemical bond combine and.....
This is due to the fact that most composing
chemical reactions are exothermic in nature.
Release Energy
-Generally when two elements are being decomposed or the bonds are being
broken and the atoms of the compound are being separated, the compound........
This again is due to the fact that most
decomposing
chemical reactions are endothermic in nature.
Absorb Energy
In most chemical reactions there is a number of different compounds being decomposed
and composed. For instance in the combustion reaction:
H2 + O2 = H2O + Heat (Energy)
The Hydrogen and Oxygen are decomposed to compose water and the overall net change
in energy is exothermic.
Stability- Stability is a measure of potential energy and whether a substance is present
for a long period or for a short period. If something is stable it does not change easily
while something that is not stable is always changing. Sort of like relationships!
In chemistry stability means that the substance, whether that be an electron, atom,
or compound is at its lowest energy state. A lower energy state is stable because the
substance does not have a lot of energy to do something with. Its kind of like a
hyperactive kid that just ate a lot of sugar. This kid has a lot of energy and needs to burn
it off and we can say this kid is unstable because he can not control his crazy active
behavior. He is unstable because we do not know what he is going to do next. On the
other hand another kid is sluggish and tired. This kid is stable because he does not have
the energy to change or do anything but sleep.
Atoms are no different. Atoms that are reactive have higher energy states when
they are alone and are unstable but when bonded to another atom of their liking they
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Unit
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Regents
Chemistry
Mr Grodski III
achieve lower energy states and thus become stable. They do not have the energy to
change. Remember Mr. Stable and Mr. Unstable?!
STABLE = LOWER ENERGY STATES – Atoms bonded in compounds
UNSTABLE = HIGHER ENERGY STATES – Free reactive Atoms not bonded
Notice that when atoms in a bond are being pulled apart energy must be absorbed. Thus
energy must be added to the atoms and this extra energy raises their energy states. This
in turn makes the individual atoms unstable because they have energy to change or do
something. Well, then the two elements bonded now have energy to change and pull
apart and become unstable and have a higher energy state, when they absorb energy.
STABLE = BONDED ATOMS
UNSTABLE = FREE REACTIVE ATOM
The following energy diagram illustrates this principle:
The boy and girl are now
married and happy giving
off love. When Atoms
Bond they release their
extra energy and become
more stable (happy)!
The boy and girl have met and
are nervous on their first date.
Atoms when they become in
close proximately repel each
other initially thus adding to
the overall increase in energy
and becomes more unstable..
The boy and girl are single and are
alone and are unhappy and have a
lot of love to give but no one to
share it with.. Awww! Atoms who
are alone are “unhappy”
(unstable) and have a lot of
energy to give but no one to give
-Boy and Girl represent two free atoms that are not connected or bonded. Reading from
left to right you can see that the two elements are not bonded and have a higher energy
state than when bonded (Boy-Girl). Thus, the free atoms are not as stable as the bonded
atoms. One can also observe that the most unstable state or the state at which the two
elements have the highest energy is when they are coming together but not yet bonded .
The reason for this high energy at this point is that the electrons from each atom at first
repel each other. This initial repulsion is due in part by the negatively charged electrons
in the outermost shells in each atom repelling (negatives repell). With enough activation
energy (match striking), the two atoms can get past this initial repulsion and chemically
Bond.
Energy Added
H–H + O–O
H
H + O
O
→
H2O + Energy
→
Energy released
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Unit
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Chemistry
Mr Grodski III
B. Ionic Bonds - This type of bond represents a connection between the valence
electrons of one or more elements with another, where electrons are transferred from
one atom to another.
All Ionic Bonds occur between a Metal and a Nonmetal!!!!!!!!
-Since an Ionic bond is a transfer of electrons from one atom to another it makes sense
that it requires a nonmetal and metal. Remember that metals have low ionization
energies and low electronegativities and nonmetals have high electronegativities and high
ionization
energies. Thus metals have weak hold of electrons and nonmetals have a
strong attraction for electrons. This very important because in order to have an
ionic bond one element must be able to pluck an electron(s) from another. Well it should
seem obvious that the atom that does the plucking is the element that has the strongest
attractions for electrons, which is the nonmetal! And the atom that loses its electron or
gets plucked is the metal because it does not have a strong hold of its electrons.
For Example: Sodium Chloride, NaCl, table salt, is made of Cl and Na held
together by an ionic bond. Na is the metal, Cl is the nonmetal. If we look at
the electronegativity of Cl and Na from Table S it easy to see why an
electron in transferred and an ionic bond is established.
Cl = 3.2
Na = 0.9
*2.3*
Its easy to see that the Cl is more than 3 times the electronegativity of Na.
Because of this large difference in electronegativity its easy see that Cl will
pluck an electron from Na. Also one can see that the difference between the
two atoms electronegativity is 2.3 .
*This is important to note because if the difference between the two atoms
elements is 1.7 or greater then the bond is Ionic!
-We already know Na is a metal that likes to lose one electron because of the oxidation
state in the upper right corner of the periodic table (+1) and because it has one valence
electron to lose. Thus it likes to be Na+1. Notice if Na loses its valence electron it has
the same STABLE electron configuration of Ne (Noble Gas).
-Cl is nonmetal that likes to gain an electron, one of its oxidation states is -1. Thus it
likes to be Cl-1 . Notice when Cl gains an electron it has the same STABLE electron
configuration of Ar (Noble Gas)
The driving force behind chemical bonds is stable lower energy electron
configurations that are attained by differences of atomic radii.
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Mr Grodski III
1. Lewis Diagrams with Ionic compounds.- A Lewis structure of any compound
is simply the electron dot diagrams of the atoms in a compound where all valence
electrons are accounted for. Remember that only the valence electrons are available for
bonding.
The lewis diagram For NaCl is :
The two are bonded because of the
opposite charges of the two ions that attract for one another. Notice that a total number
of eight valence electrons are used . This is called the octet rule!
The octet rule means that each atom participating in a ionic bond must achieve eight
electrons in its outermost energy level, thus eight dots around each element. It’s easy to
see that Cl has attained eight and has achieved a charge of -1 because it plucked an
electron from Na. Na attained eight by losing an electron. Remember when Na loses an
electron it has the same electron configuration as Ne (a noble gas !) A noble gas is stable
because it has eight electrons in its outermost energy shell ( filled energy level ).
When an atom attains a stable octet of electrons it is at a lower energy state. When
two atoms bond they do so at lower energy states.
What is the lewis structure of K and I bonded:
Chemical Formula:
What is the lewis structure of Ca and Cl bonded:
Chemical Formula:
2. Writing Chemical Formulas of Ionic Compounds
H20
type of atom
subscript = number of Hydrogen atoms
Oxygen does not have a subscript
following it which implies one.
A.-Balance oxidation states- when writing chemical formulas for
ionic substances its important to balance the oxidation states to zero. For instance in
the case of Calcium Chloride....
Ca has a oxidation state of +2 (from upper right corner of the Periodic Table) and Cl has
an oxidation state of -1. Thus in order for the compound to have a zero or neutral charge
two Cl are needed for every Ca.
2 Cl = -2
1 Ca = +2
0
Thus:
CaCl2
When writing chemical formulas for ionic substances the metal always goes first!
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Chemistry
Mr Grodski III
B.-Kriss Cross Method - this method simply takes the oxidation
states of each atom and kriss crosses them as subscripts. For example:
Sc+3 O-2
Sc2O3
Notice if you did the math that this compound has a zero or neutral charge.
The Lewis Structure for the above ionic compound is :
3. Characteristics of Ionic Compounds - Compounds that have ionic
bonds are very strong because of the electrostatic attraction between the metal positive
ion and the nonmetal negative ion. Strong Attractions always lead to higher melting
points or boiling points. (Remember in vapor pressures the liquid with the highest
intermolecular forces of attraction had the highest boiling point!)
Thus ionic solids have high melting points!
-Also ionic compounds in the solids state are in the fixed geometric patterns or crystal
lattice. In the geometric structure of the solid ionic crystal, ions form the crystal lattice
and are held in relatively fixed positions by electrostatic attraction. They are fixed
because the ions are in the most stable or low energy state! Because of the fixed or stable
state of the ions in a crystal lattice:
Ionic solids do not conduct electricity or heat!
-When these crystal lattices are heated (absorb heat) or dissolved in water (aq) the ions of
the metal and the nonmetal are non longer fixed (or stable) and the crystal lattice
becomes destroyed. When this happens the ions are free to move independently and are
no longer fixed in one position. This free movement of ions in the liquid state (l) or a
aqueous (aq) state allows for electrical conductivity! Lightning in a pool! Hair dryer in
tub of water!
Thus ionic compounds in the liquid or aqueous state will
conduct electricity!
Try these:
1.
2.
3.
4.
Write the lewis diagram for Li & F.
Write the chemical formula for Li & F
Write the lewis diagram for Al & Cl
Write the chemical formula for Al & Cl
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Chemistry
Mr Grodski III
C. Covalent Bonds- This type of bond represents an attraction for valence electrons
between nonmetals atoms that results in the sharing of electrons.
*All Covalent bonds occur from a nonmetal atoms bonded to nonmetals!*
Also table S rule - Difference in electronegativity is < 1.7 is a covalent bond!
Remember Ionic Bonds occur from a nonmetal and metal atoms while Covalent
Bonds result from Nonmetals and Nonmetals!
1. lewis structures for covalent bonds
-Remember that ionic bonds have a transfer of electrons in that each atom attained the
octet rule individually. In a covalent bond each element still obtains eight electrons
(octet) for each but does so by sharing! For example:
In a Ionic bond
the Na has attained 8 by losing one electron (has
the same electron configuration as Ne) and Cl when it gains one has 8! In an ionic bond
each element attains an octet independently!
In a covalent bond a lewis structure of HCL or Hydrochloric Acid is
Notice that Cl has attained an octet for itself if you count all the electrons that it shares
and it has surrounding it.
2 electrons for H
8 electrons for C l
( H has only 2 electrons in its only principle energy level thus its octet # is 2.
Its the only exception that you will deal with!)
-Notice that the Hydrogen and Chlorine share a pair of electrons. This pair of electrons
represents a covalent chemical bond! This compound is contains only one pair of
electrons thus represents a single bond. This compound can be written into a structural
formula of
-The single dash between the two atoms represents a single bond or a pair of electrons.
In a structural formula the extra electrons surrounding Cl are discarded because only the
electrons involved in the bond are illustrated as a dash in a structural formula. These
extra electrons are very important in determining shape of the structure, which we will
see later!
Write the lewis structure of NH3
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Regents
Chemistry
Mr Grodski III
Write the lewis structure of CBr4
Write the lewis structure of CO2
Notice CO2 has double bonds due to two pairs of electrons being shared!
When can use lewis structures to determine the number of bonds present!
For example determine the number of bonds in a molecule of H2
Determine the number of bonds in a molecule of Cl2
Determine the number of bonds in a molecule of O2
Determine the number of bonds in a molecule of N2
2. Nonpolar covalent bonds- Covalent bonds implies that electrons are
shared between two atoms in a bond. Ionic bonds are those bonds where there is an
transfer of electron(s) between two atoms. The reason for a transfer of electrons in an
ionic bond is that there is a large difference in electronegativity between the metal atom
and the nonmetal atom. The nonmetal atom because of its larger electronegativity plucks
the electron(s) from the lower electronegative metal atom and thus a transfer.
-But in a covalent (sharing) bond the atoms involved are all nonmetals, which have
similar electronegativities. When two nonmetal atoms that are bonded have the same
electronegativities then the electrons are “equally shared” and are NONPOLAR! Think
of it as a tug of war between the two atoms. If each atom in the bond has the same
electronegativity, which is the same attraction for electrons, then the each atom will pull
on the electrons equally and no one wins and no poles are produced.
-All nonmetal atoms that are bonded to themselves are Nonpolar Covalent Bonds.
These molecules represent all the bonds in the HOFBrINCl ‘s.
( H2 , O2 , F2 , Br2 , I2 , N2 , Cl2 )
For instance I2 , one of the HOFBrINCl’s equally shares electrons, its lewis structure:
-Notice each Iodine atom has 8 electrons and shares a single bond between them.
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Chemistry
Mr Grodski III
3. Polar covalent bond- In this covalent bond the electrons that the two
nonmetals atoms are sharing unequally. The reason that they are sharing electrons
unequally is that the two atoms have different electronegativities.
-A covalent bond is a bond between two nonmetals. If those two nonmetals are the same
element as in the case of HOFBrINCl’s then its a nonpolar covalent bond. If the two
nonmetals are different elements than they must have different electronegativities and
thus its a polar covalent bond!
-In keeping with the tug of war example, with a polar covalent bond the nonmetal with
the strongest attraction for electrons or the highest electronegativity will pull the shared
electrons closer to its atom and win! If the shared electrons are closer to one atom than
another than one side of the bond will have a partial negative charge and the other side a
partial positive charge, thus two poles..and polar!
-for example
has a polar covalent bond. It is a covalent bond because H and
Cl are both nonmetals and if you subtract the two electronegativities obtained from table
S - 3.2 - 2.2 = 1.0 1.0 is less than 1.7 thus it must be covalent. The electronegativities are
different thus it must be a polar covalent bond.
-If we were to draw a lewis structure for the HCL showing the pull of electrons due to the
differences in electronegativities it would look something like this:
-Notice that the Chlorine atom which is the most electronegative of the two atoms pulls
the electrons closer to itself and has a partial negative charge. Hydrogen has its electrons
pulled farther away and has a partial positive charge. THIS IS NOT LIKE A IONIC
BOND BECAUSE THE ELECTRON ALTHOUGH UNEQUALLY SHARED ARE
NOT BEING PLUCKED JUST PULLED UNEQUALLY!
- The arrow below the lewis structure above demonstrates a Dipole Moment, which is
the direction of the charge. Charge always flows from positive to negative. All
polar bonds have a dipole moment.
4. Coordinate covalent bonds- This type of covalent bond is no different
than any other covalent bond except in this case one of the two atoms involved in
the bond donates a complete pair of electrons to be shared and the other atom
brings none.
This is an example of a Ammonium
Ion
-For example NH4+ :
shows an example of a coordinate covalent bond
between the top H and the Nitrogen.
represent electrons from the hydrogen.
represent electrons from the nitrogen. Notice Nitrogen has five valence electrons and
Hydrogen has one valence electron. The top H and the N is a coordinate covalent bond
because both electrons (a bond represents a pair) are from the Nitrogen and the hydrogen
did not bring any electrons!
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Chemistry
Mr Grodski III
5. Polar vs. Nonpolar Molecules-
To Be Polar or Not to be Polar That is The Question!
To Be Polar:
1. Molecule must have polar bonds!(not HOFBrINCl’s)
2. Must be asymmetrical!! (no symmetry)
Is CCL4 , H2 , NH3, H20 Polar or Nonpolar molecules?
Knowing our two rules: TO BE POLAR (Molecule)
1. Must have Polar Bonds!
2. Must be asymmetrical!
If one of these two are false for a molecule then it must be Nonpolar!
A nonpolar molecule is a molecule that can have
or
Nonpolar Bonds
Polar Bonds
The reason that CH4 is nonpolar is due to the fact that the second rule is not
true. It does have symmetry!
Symmetry simply implies that the molecule is shaped in a particular way that
cancels out the dipole moment! Remember that the dipole moment is the direction of the
charge and charge always flows from the positive pole to the negative pole of the Dipole
(polar molecule). Remember that a polar molecule has two poles due to differences in
Electronegativity. For example:
Electrons are being
pulled toward Br
thus Br becomes -
Electrons are being
pulled away from I,
thus I becomes +
Electronegativities from table S:
2.9
2.7
Because of the differences in the electronegativities in these two atoms there is unequal
sharing of electrons in the bond. Br has a stronger attraction for electrons than I which is
demonstrated by a higher electronegative value of 2.9. If Br has a stronger attraction for
electrons than it pulls the electrons closer toward itself and Br gains a partial negative
charge. I has its electrons pulled away from it by the more electronegative Br and gains
a partial positive charge. If the BrI molecule has two ends with different charges than its
polar (or a Dipole). The dipole moment goes from right to left. It is not canceled out here
thus BrI is polar.
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Chemistry
Mr Grodski III
This molecule has polar bonds but is symmetrical
thus it is nonpolar Again symmetry means that
the dipole moments cancell out. In this case the
direction of the charge of all the atoms attached
is toward the Carbon thus the dipole moments
cancells out! If the dipole moments all went away
from the carbon it would also cancell out.
A molecule with polar bonds and is nonpolar must be a molecule that must
have symmetry( - dipole moments cancel out!)
No polar bonds no dice!
Again symmetry is dependent upon dipole moments canceling out and not
the
overall shape of molecule.
This molecule is also a Tetrahedral but
it is a Polar molecule because it does
not have symmetry, which means that
the dipole moments do not cancel out!
Cl is more electronegative than C thus
the direction of charge is different!
This molecule is a Tetrahedral and is
nonpolar because it is has symmetry
which means that dipole moments
cancell out!
Therefore just because a molecule is Linear, Trigonal planar, or
Tetrahedral does not mean they are automatically nonpolar because of their
symmetrical shapes! All dipole moments must cancel out in each molecule
in order for a molecule to be non polar.
If they have different atoms attached to them as shown in the above CH3Cl
molecule they will not have symmetry because there dipole moments or the
direction of each charge will not cancel out!Bent and Pyramidal shapes are always
polar because of there asymmetrical shapes regardless if the atoms attached are all the
same or if they are different!
Shapes of molecules are determined by the number of valence electrons, bonding and nonbonding
electrons. Since valence electrons are all the same in a group or family of elements ( column ) then we can
predict for shapes of molecules based on this information.
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Unit
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Chemistry
Mr Grodski III
D. Polyatomic IonsA single atom with a charge is a monatomic atom. Example: Na+ . A compound of two
or more covalently bonded atoms with a charge is called a polyatomic ion. Polyatomic
Ions are very stable and behave like a monoatomic particle and the bonds that hold them
are stronger than those of the metal it ionically bonds with. Therefore, during reactions,
the polyatomic ion usually remains intact as it passes from the reactants to the products
AgNO3
+
NaCl
AgCl
+
NaNO3
-Notice the Nitrate (NO3) stays intact through reactions. Notice that the nitrate ion from
table E is -1 and Ag has a +1 oxidation state (from upper righthand of Ag in periodic
table) and AgNO3 is neutral. Check NaCl you will see that it also will be neutral.
Table E lists all of the polyatomic ions that you will be using in this course. Polyatomic
ions are most often have negative oxidation states THUS ionically bond with Metals.
Although the bonds which keep the atoms in a polyatomic ion are covalent bonds, the
polyatomic ions possess a charge to satisfy the octet rule for the covalent bonds in the
ion.
The above polyatomic ion is has a negative charge of one because it needs one electron
from a metal in order to satisfy the octet on the Left oxygen. The triangle represents the
electron from the metal. Notice all octets are satisfied. When Na loses an electron it
becomes positively charged and ionically bonds with NO3 and the lewis structure looks
like this:
** Polyatomic Ions have both Covalent Bonds and Ionic bonds**
When naming polyatomic compounds use the Kriss Cross method and makes sure the
charges balance out to zero! Remember all polyatomic ions are listed in Table E!
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Regents
Chemistry
Mr Grodski III
E. Metallic Bonding - Metallic bonding occurs between two metals. A metal
bond consists of an arrangement of positive ions that are located at the crystal lattice
sites and are immersed in a “sea” of mobile electrons. These mobile electrons can be
considered as belonging to the whole crystal rather than to individual atoms. Remember
most metals are transitional elements that have an abundant amount of electrons in their d
orbitals that creates this “sea” and many of metals physical and chemical characteristics.
F. Characteristics of Different solids
1. Ionic Compounds – (Ionic Compounds- binary & ternary Salts)
∼Metals bonded to nonmetals∼
1. high melting points
2. low electrical conductivity in solid state
3. high electrical conductivity in liquid or aqueous state
4. mostly soluble in water (dissolves): called electrolyes.
5. All ionic substances are called salts
2. Molecular solids- (covalently bonded- often referred to as Organic
because they predominate in living tissue)
∼nonmetals bonded to nonmetals∼
1. Soft
2. Good electrical insulators (rubber)
3. poor heat conductors
4. low melting points
5. mostly insoluble in water: does not dissolve in water!
3. Metallic solids- (transitional metals bonded to themselves)
∼metals bonded to metals∼
1. good conductors of electricity and heat
2. great strength
3. malleability and ductility
4. luster & colored solutions when forming complex ions in water
4. Network Solids- ( Covalently bonded crystals) - All of our crystals on
earth are ionic crystals that result from Ionic Compounds (ternary and binary salts)
except for a few crystals that are covalently bonded. These substances are called
Network Solids because they have a network of arrangement much like the crystal lattice
of ionic and metallic compounds. There are only three in existence and they include:
diamond (C), silicon carbide (SiC), and silicon dioxide (SiO2).
Characteristics include:
1. Hard
2. Poor conductors of heat and electricity
3. High melting points
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G. Naming and Writing Formulas of Chemical Compounds
1. Formulas:
a. Empirical Formulas - Simplest atomic ratio in which elements
combine to form a compound. For example H2O2 ( Hydrogen peroxide) has a empirical
formula of HO. One part Hydrogen to one part Oxygen.
b. Molecular Formulas - Covalent structures are comprised of
molecules. These molecules sometimes consist of atoms combined chemically in
multiples of their simplest ratio but many time they include molecules that have multiples
of their smallest ratio. For example H2O2 is the molecular formula for Hydrogen peroxide
yet its empirical formula is HO. The molecular formula is often referred to as the “true”
formula.
2. Naming:
a. Binary compounds (two elements!)
1. Binary salts (Ionic compounds) - In a binary compound
composed of a metal and a nonmetal, the metallic element is usually named normally and
written first. The nonmetal ends in ide
For example: NaCl is named sodium chloride
CaCl2:
LiI :
2. Binary compounds of two nonmetals (covalent)
The less electronegative element is usually named and written first normally and the
other nonmetal ends in ide . Prefixes and Suffixes are used to indicate the number of
atoms of each of the nonmetals.
mono- (1), example, CO : carbon monoxide
di- (2),
example, CO2 : carbon dioxide
tri- (3),
example, SO3 : sulfer trioxide
tetra- (4),
example, CCl4 : carbon tetrachloride
penta- (5),
example, PCl5 : phosphorus pentachloride
N2O : dinitrogen monoxide
b. Ternary compounds (Salts- Ionic) - contain three different
kinds of atoms. Most compounds that contain three or more different kinds of atoms
contain polyatomic ions. The names of these compounds with the exception of the acids,
which we will learn, later includes just the normal name of the metal and the name of the
polyatomic ion. NaNO3 is called sodium nitrate and there is no use of prefixes or
suffixes in this case.
Mg3(PO4)2
Write the chemical formula for Potassium Chlorite:
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4. Stock System - This is newer method in naming chemical compounds that
utilizes a Roman Numeral to identify a particular oxidation state of an element that has
multiple oxidation states. For example Nitrogen has 8 oxidation states. They include -3, 2, -1, +1, +2, +3, +4, and +5. These oxidation states can be found at the upper right hand
corner of each element in the Periodic Table. Oxidation states are the ions or charges
that a particular element can have (losing or gaining electrons or both). In the case of
Nitrogen it can either lose electrons or gain electrons to obtain positive or negative ions.
The Stock System eliminates the need to to use prefixes or suffuxes. It does this by
using Roman Numerals to identify the Positive Ions of a particular element. For example
Nitrogen with its multiple oxidation states will be written as a compound with oxygen
under both formats.
Formula
Name of Compound
Stock System Name
N2 O
NO
N2O3
NO2
N2O5
dinitrogen monoxide
nitrogen monoxide
dinitrogen trioxide
nitrogen dioxide
dinitrogen pentoxide
nitrogen(I) oxide
nitrogen(II) oxide
nitrogen(III) oxide
nitrogen(V) oxide
If you check the charges of both elements you would see that they would balance to
zero!
H. Molecular Attraction - Molecular attractions account for the
intermolecular attractions between molecules in liqiuds and gases. These are not
chemical bonds but can be very significant in strength. Remember that the strength
of the intermolecular attractions affects vapor pressure and boiling points. Strong
intermolecular attractions equal lower vapor pressure and higher boiling points!
1. Dipole - Dipole Attractions- This type of molecular attraction utilizes
the attraction between two dipoles (two polar molecules). The positive end of one
dipole will attract the other end of another dipole. For example in the case of HCl :
These attractions can be very strong. The strength of a dipole - dipole interaction is
determined by the difference in electronegativities obtained from table K. The larger the
difference in electronegativity the larger greater the dipole force.
Dipole - Dipole forces of attraction ocurr from two polar molecules!!!
2. Hydrogen Bonding - The strongest type of Molecular Attractions are
Hydrogen Bonding or H-Bonding. This type of attraction is the a special type of dipole dipole attraction. Its still requires two polar molecules but one of the molecules must
have a hydrogen covalently bonded to a highly electronegative atom. In this way the
hydrogen has its pair of sharing electrons pulled away from it to make it very polar. Just
like dipole - dipole the strongest H-Bonding occurs when the difference in
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electronegativity between the H and the element its bonded to is greatest. Since F is the
most electronegative atom than the strongest H - Bonding takes place in molecules in HF.
-Water has a relatively high boiling point because of H - Bonding!
H -Bond
3. Van der Waals Force (London Forces)- his type of intermolecular
attraction occurs between nonpolar molecules. This type of attractions are the weakest
because of the absence of electrostatic forces that polar molecules have. Van er Waals
forces appear to be due to chance distribution of electrons resulting in momentary dipole
attractions. That is even a nonpolar molecule can have have a momentary dipole due to
electrons moving to one side of the molecule. If more electrons move to one side of a
molecule it can give that side a momentary negative charge which can attract a
momentary positive charge from another molecule.
The strength of a Van der Waals force is stronger in larger molecules and or heavier, denser atoms
because they have more electrons that can produce larger momentary charges, which are stronger.
The strength of a van der Waals force also increases as two nonpolar molecules come closer to
one another and weaker as they move apart.
4. Molecule Ion Attraction- This type of attraction is one where a
molecule like water is attracted to ion. In aqueous solutions, where a chemical is
dissolved in water the water molecule surrounds the ion and dissolves it. Dissolving of a
particle by water is called hydration. When NaCl(s) is mixed with water, the crystal
lattice of the ionic compound is broken down and free ions are released into the solvent
(water). Ions have charges and the negative charge from Cl is attracted to the positive H
end of the polar water molecule. The positive charge of the Na is attracted to the
negative O in water.
When something dissolves in water the water molecules surround the ions as shown. if
something is not soluble than the solvent (water in this case cannot surround it). Notice
that a polar molecule could only dissolve an ionic solid because of its dipoles.
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