Monday, January 3, 2011
Dear guests,
This Blog Post will be all about Chemical Bonding & Molecular Geometry. So before we begin with our lesson, let us first go through some basic terminologies that we will encounter in this tutorial.
Valence electrons - electrons in the outermost energy level which are involved in chemical bonding
TYPES OF CHEMICAL BOND
Ionic Bond - a bond formed between two or more atoms by transfer of one or more electrons.
Covalent Bond - a bond formed between two or more atoms by sharing of electrons
Types of Covalent Bonding According to Polarity
1. Polar Covalent Bond - a covalent bond formed by unequal sharing of electrons
2. Non-polar Covalent Bond - a covalent bond formed by equal sharing of electrons.
Metallic Bond - a bond formed between metal atoms in which bonding electrons are relatively free to move throughout a three-dimensional structure.
Lewis Electron Dot Formula - consists of chemical symbol of an element surrounded by a number of dots representing the number of valence electrons.
Octet Rule - states that atoms are stable if their outermost shell is completely filled with 8 electrons
Hybridization of Atomic Orbitals
Hybridization - the combination of atomic orbitals to produce a new set of hybrid orbitals; enables an atom to bond with a number of atoms greater than its number of unpaired valence electrons.
TYPES OF HYBRIDIZATION
1. sp3 hybridization 2. sp2 hybridization 3. sp hybridization
Bond Polarity - useful in describing the types of bond
ELECTRONEGATIVITY DIFFERENCES
Ionic Bond 1.9 and above
Polar Covalent Bond 0.1 to 1.8
Non-Polar Covalent Bond 0
Now that you know about the basics, let's go on with the topic. ;)
INTRODUCTION
Though the periodic table has only 118 or so elements, there are obviously more substances in nature than 118 pure elements. This is because atoms can react with one another to form new substances called compounds. It is formed when two or more atoms chemically bond together, the resulting compound is unique both chemically and physically from its parent atoms.
Here's an example of a chemical compound formed from two elements:
When chemically bonded together, these two substances form the compound sodium chloride, a compound so safe that we eat it every day - common table salt!
Well, there are three types of Chemical Bond and we are gonna discuss them here too so let's start with the first one which is the Ionic Bond.
Ionic Bond
Those molecules that consist of charged ions with opposite charges are called IONIC. These ionic compounds are generally solids with high melting points and conduct electrical current. Ionic compounds are generally formed from metal and a non-metal elements. The formation of an IONIC BOND is the result of the transfer of one or more electrons from a metal onto a non-metal.
For example, during the reaction of sodium with chlorine:
sodium (on the left) loses its one valence electron to chlorine (on the right), resulting in a positively charged sodium ion(left) and a negatively charged chlorine ion (right):
Notice that when sodium loses its one valence electron it gets smaller in size, while chlorine grows larger when it gains an additional valence electron. This is typical of the relative sizes of iions to atoms. Positive ions tend to be smaller than their parent atoms while negative ions tend to be larger than their parent. After the reaction takes place, the charged Na+and Cl- ions are held together by electrostatic forces, thus forming an ionic bond. Ionic compounds share many features in common:
- Ionic bonds form between metals and nonmetals.
- In naming simple ionic compounds, the metal is always first, the nonmetal second (e.g., sodium chloride).
- Ionic compounds dissolve easily in water and other polar solvents.
- In solution, ionic compounds easily conduct electricity.
- Ionic compounds tend to form crystalline solids with high melting temperatures.
Okay. Maybe to make things clearer...
I hope you get it now. Well, just to be sure, here's a music video to help you understand it more. Just for fun's sake. :>
LYRICS:
Na na na na
Na na na na
Na na na na
Anions and cations are always falling in love.
Opposite charges are what lead them to each other.
Electrostatic forces are what cause these attractions.
They also hold all the ionic bonds together.
Oh, oh, oh, oh, ohhhh, oh, oh
Anions will dance, show off what they got.
Oh, oh, oh, oh, ohhhh, oh, oh
Cations will accept, and be more than just friends.
Can't deny, can't deny,
that these ions form perfect bonds.
Can't deny, can't deny,
that these ions form perfect bonds.
p-p-p-perfect bonds
perfect, perfect bonds
p-p-p-perfect bonds
perfect, perfect bonds
Did you know cations are really positive?
Just think of cats and you'll remember this statement.
Anions are negative, well, duh, obviously!
I did mention before opposites attract,
so don't get bored so easily!
Oh, oh, oh, oh, ohhhh, oh, oh
Anions will dance, show off what they got.
Oh, oh, oh, oh, ohhhh, oh, oh
Cations will accept, and be more than just friends.
Can't deny, can't deny,
that these ions form perfect bonds.
Can't deny, can't deny,
that these ions form perfect bonds.
p-p-p-perfect bonds
perfect, perfect bonds
p-p-p-perfect bonds
perfect, perfect bonds
Well, wasn't that FUN? :} Haha.
If you want to learn more about Ionic Bonds, OPEN THIS LINK IN A NEW WINDOW. :)
NEXT. Let's learn all about Covalent Bonds.
Covalent Bonds
The second major type of atomic bonding occurs when atoms shareelectrons. As opposed to ionic bonding in which a complete transfer of electrons occurs, covalent bonding occurs when two (or more) elementsshare electrons. Covalent bonding occurs because the atoms in thecompound have a similar tendency for electrons (generally to gain electrons). This most commonly occurs when two nonmetals bond together. Because both of the nonmetals will want to gain electrons, the elements involved will share electrons in an effort to fill their valence shells. A good example of a covalent bond is that which occurs between two hydrogen atoms. Atoms of hydrogen (H) have one valence electron in their first electron shell. Since the capacity of this shell is two electrons, each hydrogen atom will "want" to pick up a second electron. In an effort to pick up a second electron, hydrogen atoms will react with nearby hydrogen (H) atoms to form the compound H2. Because the hydrogen compound is a combination of equally matched atoms, the atoms will share each other's single electron, forming one covalent bond. In this way, both atoms share the stability of a full valence shell.
HERE ARE MORE VIDEOS TO HELP YOU UNDERSTAND THE COVALENT BOND:
Polar and nonpolar covalent bonding
There are, in fact, two subtypes of covalent bonds. The H2 molecule is a good example of the first type of covalent bond, the nonpolar bond. Because both atoms in the H2 molecule have an equal attraction (or affinity) for electrons, the bonding electrons are equally shared by the two atoms, and a nonpolar covalent bond is formed. Whenever two atoms of the same element bond together, a nonpolar bond is formed.
A polar bond is formed when electrons are unequally shared between twoatoms. Polar covalent bonding occurs because one atom has a stronger affinity for electrons than the other (yet not enough to pull the electrons away completely and form an ion). In a polar covalent bond, the bonding electrons will spend a greater amount of time around the atom that has the stronger affinity for electrons. A good example of a polar covalent bond is the hydrogen-oxygen bond in the water molecule.
Water molecules contain two hydrogen atoms(pictured in red) bonded to one oxygen atom (blue). Oxygen, with six valence electrons, needs two additional electrons to complete its valence shell. Each hydrogen contains one electron. Thus oxygen shares the electrons from two hydrogen atoms to complete its own valence shell, and in return shares two of its own electrons with each hydrogen, completing the H valence shells.
The primary difference between the H-O bond in water and the H-H bond is the degree of electron sharing. The large oxygen atom has a stronger affinity for electrons than the small hydrogen atoms. Because oxygen has a stronger pull on the bonding electrons, it preoccupies their time, and this leads to unequal sharing and the formation of a polar covalent bond.
For you to have some covalent fun, enjoy this song about covalent bonds. :)
Metallic Bond
Another chemical bonding mechanism is the metallic bond.In the metallic bond, an atom
achieves a more stable configuration by sharing the electrons in its outer shell with many
other atoms. Metallic bonds prevail in elements in which the valence electrons are not tightly
bound with the nucleus, namely metals, thus the name metallic bonding.
In this type of bond, each atom in a metal crystal contributes all the electrons in its valence
shell to all other atoms in the crystal.
Another way of looking at this mechanism is to imagine that the valence electrons are not
closely associated with individual atoms, but instead move around amongst the atoms within
the crystal. Therefore, the individual atoms can "slip" over one another yet remain firmly
held together by the electrostatic forces exerted by the electrons. This is why most metals
can be hammered into thin sheets (malleable) or drawn into thin wires (ductile). When an
electrical potential difference is applied, the electrons move freely between atoms, and a
current flows.
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LEWIS ELECTRON DOT FORMULA & THE OCTET RULE
Octet Rule
The octet rule is one of the more poorly named rules in chemistry, since it turns out to be violated almost as much as it is followed. It comes from the realization that since atoms want to have noble gas structure and some noble gases have eight electrons in their valence shell, atoms should have eight electrons around them.
The atoms that commonly follow the octet rule are
- Carbon: C
- Nitogen: N
- Oxygen: O
- The halogens, F, Cl, Br, I
However, there are almost as many exceptions to the rule as there are atoms that follow it. Ones you will possibly see:
- Hydrogen: H will have 2 electrons to gain a noble gas shell.
- Phosphorus: P often has 10 electrons
- Sulfur: S often has 12 electrons
- Atoms in groups 1, 2, and 3 want 2, 4, and 6 electrons respectively. I.e., Na (group 1) wants 2, Be (Group 2) wants 4 and Al (Group 3) wants 6.
- Any molecule with an odd number of electrons: at least one atom must therefore have an odd number of electrons (The molecule is known as a free radical.)
Steps in writing Lewis structures of Covalent Compounds
1. Determine the total number of valence electrons in compound by adding the valence electrons
in each atom
2. The atom with the highest covalency number is considered as the central atom. Write the
skeletal structure of the compound, using chemical symbols and placing thecentral atom in the
middle of the other atoms.
3. Bond the other atoms to the central atom by a single bond. Take into account the number of
electrons used in the bonding. Each bond is equal to two valence electrons.
4. Distribute the remaining valence electrons to the attached atoms first and then to the central
atom last.
5. Check if the octet rule is followed by each atom
6. If there is a deficiency in the octet rule, form a multiple bond.
The Lewis Dot Structure Of Nitrogen Trifluoride
To build a picture of Nitrogen Trifluoride, we need to start with the electron dot diagrams for the elements involved in the molecule. In this case those are Nitrogen and Fluorine. There are many ways to fit Nitrogen and Fluorine together in theory so we use the chemical formula to tell us the number of each type of atom involved. In this case the formula for Nitrogen Trifluoride is NF3, meaning there is one Nitrogen atom and three Fluorine atoms in each NF3 molecule, as seen below.
Once we have the electron dot diagrams for these elements, it is then a simple matter of putting the puzzle together. We need to fit the four atoms into a single molecule and have no gaps left over. This means that when we look at each atom in the finished molecule it must have eight electrons surrounding it, regardless of how many other atoms it is sharing them with.
Each Fluorine atom has only one gap in its outer shell, so it can bond in only one place. Nitrogen has three gaps in its outer shell so it makes sense to place the Nitrogen in the middle of the molecule and attach the Fluorine atoms to it. Note that the eight electrons in the outer shell can be moved around to suit the build of the molecule, though they must be evenly spread across the four quarters of the atom. Building NF3 looks like this:
When we join the pairs of electrons up, the molecule becomes complete. If we count the number of electrons around each atom we can see that each atom is surrounded by eight electrons. This means that its outer electron shell is full and the molecule is (relatively) stable.
When drawing molecules, it is traditional to replace each pair of electrons with a single solid line joining the atoms together. In the case of the electron pairs that do not join to another atom, the line just stops in space. These are called lone pairs. They are often left out of Lewis Dot Structure drawings, as shown below, but they do have an effect on the structure of the molecule.
Here's another video to make things more easier. :)
And now that we know how to make an Electron Dot Structure, Let's learn about Molecular Geometry!
MOLECULAR GEOMETRY (&VESPR)
Molecular geometry or molecular structure is the three-dimensional arrangement of atoms within a molecule. It is important to be able to predict and understand the molecular structure of a molecule because many of the properties of a substance are determined by its geometry.
Molecular geometry or molecular structure is the three-dimensional arrangement of atoms within a molecule. It is important to be able to predict and understand the molecular structure of a molecule because many of the properties of a substance are determined by its geometry.
At this point we are ready to explore the three dimensional structure of simple molecular (covalent) compounds and polyatomic ions. We will use a model called the Valence Shell Electron-Pair Repulsion (VSEPR) model that is based on the repulsive behavior of electron-pairs. This model is fairly powerful in its predictive capacity. To use the model we will have to memorize a collection of information.
The table below contains several columns. We already have a concept of bonding pair of electrons and non-bonding pairs of electrons. Bonding pairs of electrons are those electrons shared by the central atom and any atom to which it is bonded. Non-bonding pairs of electrons are those pairs of electrons on an individual atom that are not shared with another atom. In the table below the term bonding groups (second from the left column) is used in the column for the bonding pair of electrons. Groups is a more generic term. Group is used when a central atom has two terminal atoms bonded by single bonds and a terminal atom bonded with two pairs of electrons (a double bond). In this case there are three groups of electrons around the central atom and the molecualr geometry of the molecule is defined accordingly. The term electron-pair geometry is the name of the geometry of the electron-pairs on the central atom, whether they are bonding or non-bonding. Molecular geometry is the name of the geometry used to describe the shape of a molecule. The electron-pair geometry provides a guide to the bond angles of between a terminal-central-terminal atom in a compound. The molecular geometry is the shape of the molecule. So when asked to describe the shape of a molecule we must respond with a molecular geometry. If asked for the electron-pair geometry on the central atom we must respond with the electron-pair geometry. Notice that there are several examples with the same electron-pair geometry, but different molecular geometries. You should note that to determine the shape (molecular geometry) of a molecule you must write the Lewis structure and determine the number of bonding groups of electrons and the number of non-bonding pairs of electrons on the central atom, then use the associated name for that shape.
The table below summarizes the molecular and electron-pair geometries for different combinations of bonding groups and nonbonding pairs of electrons on the central atom.
| Geometry | Type | # of Electron Pairs | Ideal Bond Angle | Examples |
| linear | AB2 | 2 | 180° | BeCl2 |
| trigonal planar | AB3 | 3 | 120° | BF3 |
| tetrahedral | AB4 | 4 | 109.5° | CH4 |
| trigonal bipyramidal | AB5 | 5 | 90°, 120° | PCl5 |
| octohedral | AB6 | 6 | 90° | SF6 |
| bent | AB2E | 3 | 120° (119°) | SO2 |
| trigonal pyramidal | AB3E | 4 | 109.5° (107.5°) | NH3 |
| bent | AB2E2 | 4 | 109.5° (104.5°) | H2O |
| seesaw | AB4E | 5 | 180°,120° (173.1°,101.6°) | SF4 |
| T-shape | AB3E2 | 5 | 90°,180° (87.5°,<180°) | ClF3 |
| linear | AB2E3 | 5 | 180° | XeF2 |
| square pyramidal | AB5E | 6 | 90° (84.8°) | BrF5 |
| square planar | AB4E2 | 6 | 90° | XeF4 |
Here are some of the molecular geometry representations.
Lets consider the Lewis structure for CCl4. We can draw the Lewis structure on a sheet of paper. The most convenient way is shown here:
Notice that there are two kinds of electron groups in this structure. Bonding electrons, which are shared by a pair of atoms and nonbonding electrons, which belong to a particular atom but do not participate in bonding. In CCl4 the central carbon atom has four bonding groups of electrons. Each chlorine atom has three nonbonding pairs of electrons.
The arrangement of the atoms is correct in my structure. That is the carbon is the central atom and the four chlorine atoms are terminal. But this drawing does not tell us about the shape of the molecule. Lets look at what I mean. Here we have a ball and stick model of CCl4.
This is a nice representation of a two dimensional, flat structure. The Cl-C-Cl bond angles appear to be 90 degrees. However, the actual bond angles in this molecule are 109.5 degrees. What does this do to our geometry?
Lets rotate this molecule to see what has happened. We see the actual molecular geometry is not flat, but is tetrahedral. This ball and stick model does not adequately represent why the molecule has to have this 3-dimensional arrangement. The shape we see is the only possible shape for a central carbon atom with four bonds. This geometry is a direct result of the repulsion experienced by the four groups of bonding electrons.
The shape of this molecule is a result of the electrons in the four bonds positioning themselves so as to minimize the repulsive effects. This was seen in the 'balloon' example we used in class. When four balloons of the same size are tied together the natural arrangement is as a tetrahedron. With bond angles of 109.5 degrees. As indicated in Table I, any compound containing a central atom with four bonding groups (pairs) of electrons around it will exhibit this particular geometry.
By recognizing that electrons repel each other it is possible to arrive at geometries which result from valence electrons taking up positions as far as possible from each other. The position of the atoms is dictated by the position of the bonding groups of electrons.
These ideas have been explored and have resulted in a theory for molecular geometry known as Valence Shell Electron-Pair Repulsion Theory.
what we could have been, 4:53 AM.
