Lewis Dot Structures of Covalent Compounds. Atoms are made up of protons, neutrons, and electrons. The protons and neutrons are located at the center of the atom, the nucleus. These electrons can be divided into core electrons and valence electrons. The valence electrons are
Download Policy: Content on the Website is provided to you AS IS for your information and personal use and may not be sold / licensed / shared on other websites without getting consent from its author.While downloading, if for some reason you are not able to download a presentation, the publisher may have deleted the file from their server.
Atoms are made up of protons, neutrons, and electrons. The
protons and neutrons are located at the center of the atom,
the nucleus. These electrons can be divided into core
electrons and valence electrons. The valence electrons are
the outermost electrons and are the ones involved in
Electrons occupy most of
the volume of an atom
They arrange themselves
in ’shells’ at varying
distances from the
Protons and neutrons
are located in the
nucleus (center) of the
These are the outermost electrons and the ones
In chemical reactions
The number of valence electrons varies by element. For the
Main Group elements, the number of valence electrons is
equal to the Group Number that the elements belong to.
For example, Sodium (Na) belongs to Group 1A and therefore
has 1 valence electron.
The Periodic Table
For example, Bromine (Br) belongs to Group VIIA and
therefore has 7 valence electrons. We can represent the
valence electrons of an atom using a Lewis dot symbol, in
which the element symbol is surrounded by dots representing
the valence electrons.
For example, Oxygen has six valence electrons, so its Lewis
dot symbol is:
Note the six dots representing the six valence electrons
For example, neon has eight valence electrons, so its Lewis
dot symbol is:
For example, carbon has four valence electrons, so its Lewis
dot symbol is :
How many valence electrons does Potassium (K) have?
How many valence electrons does Antimony (Sb) have?
How many valence electrons does Phosphorus (P) have?
How many valence electrons does Magnesium (Mg) have?
The Noble Gas elements in Group VIIIA have either two valence electrons (He) or eight valence electrons (Ne, Ar, Kr, Xe, and Rn). These elements are extremely stable because they have full valence shells- two electrons for He in the first row and eight electrons in each of the later rows. This is the basis for the Octet Rule - elements tend to react in a way to attain the electron configuration of Group VIIIA
Metallic elements at the left side of the Periodic Table tend to
lose one or more electrons to form positive ions, such as Na+
and Mg2+, each of which has the electron configurationof the
Noble Gas that preceds it.
Nonmetals at the right side of the Periodic Table tend to either
gain electrons to form negative ions such as F-, O2-, and N3- or
to share electrons in covalent bonds. This learning objective
describes how this is done
When nonmetallic elements react with other nonmetallic
elements, they share electrons in order to obtain eight valence
Each fluorine atom has seven valence electrons. They each require one more electron to satisfy the Octet Rule.
The left fluorine atom now has a total of eight electrons and the right fluorine atom now has a total of eight electrons around it.
When nonmetallic elements react with other nonmetallic elements, they share electrons in order to obtain eight valence electrons.
The two electrons that form the covalent bond are often
Represented by a single line. The F2 molecule can be
represented using a line and dots to show the bonding pair
and the six lone pairs, respectively. This is called a Lewis dot
Some atoms have to share more than one electron in order
to satisfy the Octet Rule.
Each oxygen atom has six valence electrons. They each
require two more electrons to satisfy the Octet Rule.
The four electrons shared by the oxygen atoms form a
The double bond is represented by two single lines. Each line
in the Lewis dot structure represents two electrons
The element hydrogen is an exception to the Octet Rule. It
only needs two electrons, rather than eight, to be stable.
The hydrogen atom has one valence electron. It requires one
more electron to be stable. The fluorine atom has seven
valence electrons. It requires one more to satisfy the Octet
The hydrogen atom now has a total of two electrons around
it and is stable.
The fluorine atom now has a total of eight electrons around
it and is stable.
Add together the number of valence electrons for each atom
in the molecule. For example, CF4
Carbon has four valence electrons and each fluorine has
seven valence electrons = 4 + 4(7)
Write out the elements of the molecule so that the least
electronegative elements is in the center surrounded by the
other elements. For example, CF4
Place a covalent bond between the central atom and the outside atoms. Remember each covalent bond contains two
There are 24 valence electrons remaining. Add electrons to
the outer atoms as lose pairs to satisfy the Octet Rule.
Nitrogen has 5 valence electrons and each hydrogen has 1
The total number of valence electrons = 5 + 3 (1) = 8
Hydrogen is always an outer atom and is never at the centre
of a molecule
Add the bonding electrons. This uses 6 of the 8 valence
The 2 remaining valence electrons are not added to the outer
atoms, because each H has its maximum of 2 valence
Add the bonding electrons. This uses 6 of the 8
The 2 remaining valence electrons are not added to
the outer atoms, because each H has its maximum
of 2 valence electrons.
Placethe remaining 2 Valence
electrons on the central
This is the
Check all atoms in the molecule to ensure that each has 8
electrons(2 for hydrogen). If an atom has fewer than 8
electrons, create double or triple bonds. (Note: Double
bonds only exist between C,N,O and S atoms)
Hydrogen : 1 bond = 2 electrons (stable)
Carbon : 4 bonds = 8 electrons (stable)
= 8 electrons (stable)
Apply Rules 1-5 to the molecule
Rule 1: Count the valency electrons
Rule 2: Place the least electronegative element at the centre, except for H, which is always an outer atom
Rule 3: Add covalent bonds between the centre and the outer atoms
Rule 4: Add lone pairs to the outer atoms
Rule 5: Add lone pairs to the centre atom
Carbon has 4 valence electrons, each hydrogen has 1 valence
electron, and oxygen has 6 valence electrons.
Total number of valence electrons : 4 + 2(1) + 6 = 12
Carbon is at the centre of the molecule because it is less
electronegative than oxygen. Hydrogen is always an outer
atom and is never at the centre of the molecule.
Add the bonding electrons.
This uses 6 of the 12 valence
Add the remaining 6 lectrons to
the outer atom. Hydrogen does
not need any more electrons, but
Oxygen needs 6 to complete its
Rule 5 There are no valence electrons left to add to the centre
Oxygen shares one of its lone pairs with C and O and give the desired 8 electron total
This is the
Exceptions to the Octet Rule
The Octet Rule applies to Groups IVA through VIIA in the
second row of the Periodic Table, but there are exceptions to the rule among some other elements. The following two cases are an example
Boron has 3 valence electrons and each Fluorine has 7 valence electrons
Total number of electrons = 3 + 3 (7) = 24
Boron is at the centre of
the molecule because it is
less electronegative than
Add the bonding electrons.
This uses 6 of the 24 valence
Add the remaining electrons
to the outer atoms. Each
Fluorine has the required 8
This uses the remaining
electrons leaving none to add
to the Boron central atom
Check the number of electrons around each atom. Each
Fluorine atom has 8 electrons, but the Boron Atom has only
6. This is an exception to the Octet Rule. A B=F bond is not
an option, because double bonds exist only between C,N,O,
and S atoms
This is the Lewis
dot structure BF3
Phosphorus has 5 valence
electrons and each fluorine
has 7 valence electrons
Total number of electrons
Phosporus is at the centre
because it is less
electronegative than fluorine
Add the bonding electrons. This uses 6 of the 24 valence
Add the remaining electrons to the outer atoms. Each Fluorine requires 6 more electrons
This uses the remaining
electrons leaving none to
the central P atom
Check the number of electrons
around each atom. Each
Fluorine atom has 8 electrons,
but the phoshorus atom has 10.
This is an exception to the
Check the number of electrons
around each atom. Each fluorine
atom has 8 electrons, but the phoshorus
atom has 10 . This is an exception to the
How Elements Form Compounds
Some atoms lose or gain electrons to become stable charged particles called ions
When atoms loses electrons, they form positively charged ions called cations
When atoms gain electrons, they form negatively charged ions called anions.
Sodium chloride is a relatively harmless compound because the sodium and chlorine atoms have stable ions .
The compound formed is called an ionic compound because it is made up of positive and negative ions that have resulted from the transfer of from a metal to a nonmetal.
The positive and negative ions are attracted to each other because they have opposite charges.
When ionic compounds are placed in water, the ions separate and are surrounded by water molecules. They are electrolytes.
They are also conductive
The location of the alkalis metals (dark green), the alkaline earth metals (light green), and the halogens (red) in the PT
There are over 100 elements in the PT
Thousands of different compounds are formed when these elements combine.
How can we name these compounds?
How can we write formulas to represent them?
We have seen from past discussions that
The PT and a knowledge of the electronic structure could be used to predict ionic charge of elements
Ionic charges (or valences) of some elements in the PT
Naming Ionic Compounds
The name of the metal first, followed by name of the of the nonmetal.
The ending of the name of the nonmetal changes and ends with “ide”
Names and Ionic Charges of some nonmetals
Names and Formulas for Atoms with More Than One Ionic Charge
Some metals are able to form more than one kind of ion.
For example, the element copper forms two completely different compounds when it reacts with chlorine
One of the compound is white: the other is yellow
Ionic charge on the copper in the white compound is 1+ . Its chemical formula is CuCl
The ionic charge on the copper in the yellow compound is 2+, its formula is CuCl2
What are these compounds?
When a compound containing this ion is dissolved in water, the positive metal ion and the nitrate ion separate from each other but the nitrate ion itself stays together as a unit surrounded by water molecules
An example is
The nitrate ion
Writing Formulas for Polyatomic Compounds
The ionic charges of polyatomic ions makes it possible for them to form ionic compounds
Common Polyatomic ions and Their Ionic charges
When a polyatomic ion such as nitrate or sulfate combines with other elements
We follow the same rules for writing formulas
What is the formula for the ionc compound formed by sodium and a sulfate ion?
Rule 1: write the symbols of the metal and of the polyatomic group
Rule 2: write the ionic charges
Crisscross rule: crisscross the ionic charges
Note that polyatomic ions do not ”reduce” . Formula cannot be simplified Na1SO2 because SO4 is a group
The formula is Na2SO4
Try this: what is the formula of lead(IV) carbonate?
There are many types of polyatomic ions, but one special group is known as the Oxyacids
Oxyacids are compounds formed when hydrogen combines with polyatomic ions that contain oxygen. Ionic charge for hydrogen in these compounds is 1+
Imagine that you find an unlabelled container of solid white crystals in the kitchen.
You are sure the crystals are either salt or sugar
A simple taste test will tell you what the crystals are.
But imagine you find the same crystals in the lab. A taste is too dangerous. What do you do?
Dissolve the crystals in water and test for conductivity.
If it conducts electricity, the compound must contain ions
Salt or sodium chloride is an ionic compound
In ionic compounds, metals with 1, 2, or 3 electrons in their outer shell lose electrons to nonmetals, which often have 5, 6, or 7 electrons in their outer shell.
If the solution does not conduct electricity, it must be a different kind of compound
Most compounds you encounter every day do not contain ions.
Rather, they contain neutral groups of atoms called molecules.
Sugar is a molecular compound. It is made up of molecules in which nonmetal atoms, such as hydrogen and oxygen share electrons to form stable arrangements.
Water and carbon dioxide are also molecular compounds, whether in in a gas, a liquid, or a solid state, the particles in ionic and molecular compounds are different as shown
Salt is an example of an ionic compound made up of ions of opposite charge. Ice (H2O) is an example of a molecular compound made up of neutral molecules
Hydrogen gas is a molecule formed when two hydrogen atoms combine. Each hydrogen atom has one electron.
For the two hydrogen atoms to become stable, both must gain an electron.
They do this by sharing a pair of electrons, one from each atom
The result is a covalent bond--- a shared pair of electrons held between two nonmetal atoms that holds the atoms together in a molecule.
Many nonmetals form molecules in this way. For example chlorine gas is a molecule that consists of two chlorine atoms held together with a covalent bond. Each chlorine atom has 7 electrons in its outer orbit and needs to gain electron to be stable
Many nonmetallic elements exist as covalently bonded molecules. Table below lists elements that form diatomic molecules.
Molecular compounds are all around us a bottle of soda contains water molecules, sucrose, glucose, or fructose
Writing formulas for Molecular Compounds
Formulas can be written using a method similar to the one used for ionic compounds.
The number of electrons that metals and nonmetals transfer to become stable ions can be a clue to the formula of an ionic compound.
Similarly, the number of electrons that a nonmetal needs to share to become stable is a clue to the number of covalent bonds it can form
The combining capacity of a nonmetal is a measure of the number of covalent bonds that it will need to form a stable molecule
Table 1: Combinig Capacities of Nonmetal Atoms
Carbon has four electrons in its outer(valence) orbit. If it lost 4 electrons, it would form a positive ion. If it gained 4 electrons, it would have the electron arrangement of neon and would form a negative ion
It turns out that carbon cannot form either ion. Instead it “gains” 4 electrons by sharing: carbon has a combining capacity of 4.
For example, when carbon shares one of its outer orbit electrons with each of four different hydrogen atoms, as shown in figure, the result is methane CH4, the major component of natural gas
As a result of forming covalent bonds through sharing electrons, the atoms end up with a stable arrangement in their orbit similar to that of a noble gas.
You can use the combining capacity to write the formulas of molecular compound s without having to consider the electronic structure
How would you write the formula for a compound formed between Carbon and Sulfur?
Rule 1: Write the symbols, with the left hand element from Table 1 with the combining capacities
Rule 2: Crisscross the combining capacities to produce subscripts
The formula is C2S4
Rule 3: Reduce the subscripts if possible
The formula C2S4 is reduced to C1S2
Rule 4: Any “1” subscript is not needed.
The correct formula is CS2
Naming Molecular Compounds
Many molecular compounds have simple names. The compound H2S is called hydrogen sulfide, much as if it is ionic. Other molecular compounds have names that are very familiar to us even though they do not follow a system
Common names have been used for centuries for water (H2O): ammonia (NH3), hydrogen peroxide (H2O2) and methane (CH4)
The names of molecular compounds often contain prefixes. These prefixes are used to count the number of atoms when the same two elements form different combinations.
For example , the gas that you exhale is carbon dioxide (CO2) while the poisonous combination of carbon and oxygen that can be formed in automobiles is carbon monoxide
The prefixes “di” and “mono” differentiate between the two molecules
Table 2: Prefixes in Molecular Compounds