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Tuesday, January 11, 2011
CHEMISTRY
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6 comments:
some notes on covalent bonds
said...
Covalent bonds: formed from sharing of two electrons, usually one donated from each of the two bonding atoms. There are 3 kinds of covalent bonds based on number of pairs of electrons shared between the two atoms:
single covalent bond - 1 shared pair double covalent bond - 2 shared pairs triple covalent bond - 3 shared pairs Bond energy: single < double < triple Bond length: single < double < triple Lewis Structures: The simplest covalent bond is between two hydrogen atoms. Electron-dot (or Lewis) symbol for hydrogen is H.. The two electrons from two hydrogen atoms pair up to make H:H. The pair of dots can be replaced by a dash to represent a bond, as in H-H. Hydrogen as the element exists as these "diatomic molecules" and its molecular formula is shown as H2.
The halogens also exist as diatomic molecules. Fluorine, chlorine, bromine and iodine all have the same type of Lewis structure, with seven outer electrons (2 on the right side, 2 on the top, 2 on the left side, and 1 on the bottom, so two F atoms share their unpaired electrons with each other to form the single covalent bond, F:F or F-F, to get F2. The rest of the pairs of electrons around each halogen are called "unshared pairs". The number of paired and unpaired electrons around an atom play an important part in determining the shape of a molecule, which in turn contribute to chemical and physical properties of the molecule.
Oxygen atoms, with six outer level electrons each ( 2 on the right side, 2 on the top, and 1 each on left side and bottom), share two pairs of electrons and thus have a double bond between them. This is shown by O::O or O=O for O2.
Nitrogen atoms, with only five outer electrons has 2 on the right side, and 1 each on the top, left side and bottom. Two nitrogen atoms share three electrons each, to form three pairs between them. This bonding is written as :N:::N:, N=N for N2
Things get a bit more complicated when atoms of different kinds of elements bond. One of the simplest bonds is between hydrogen and chlorine. The single unpaired electron on H pairs up with the unpaired electron on the chlorine and the result is H:Cl or H-Cl, or HCl.
Writing/Drawing Lewis Structures for simple molecules.
Count the number of valence electrons in each of the atoms in the molecule. Count the total number of electrons that are needed to give each atom in the molecule an octet in the valence level, or two electrons for hydrogen. Find the difference between the number of electrons available for sharing (#1) and the total number needed to complete the octet for each (#2). This gives the number of electrons that will be shared among the bonding atoms. Divide the number of shared electrons by 2 to give the number of electron pairs, thus the number of bonds between all of the atoms in the molecules. A double or triple bond count as two or three bonds because of the number of electron pairs they represent. Examples: H2O CO2 CH2O C2H4
There are several rules to remember when writing Lewis structures:
H is always terminal, only one single bond per hydrogen Oxygen almost always has two shared pairs and two unshared pairs of electrons. The two shared pairs may be in two separate single bonds or in one double bond. Carbon atoms will always have 4 bonds, either as 4 separate single bonds, or a combination of single, double and/or triple bonds so that there are 4 pairs of electrons around each carbon atom in the molecule. Carbon atoms may single, double, or triple bond to each other, single or double bond to oxygen atoms, or single, double, or triple bond to nitrogen atoms. Halogen atoms form single bonds only, and usually form just one bond. An example of an exception is the chlorate, ClO31-, ion.
What is electronegativity? Electronegativity is a measure of the attraction of an atom for electrons in a covalent bond.
Fluorine, the most reactive non-metal, is assigned the highest value since it has the greatest attraction for the electron being shared by the other element. Oxygen is also highly electronegative and has a strong attraction for electrons. Metals have low electronegativities since they have weak attraction for any shared electrons.
When two unlike atoms are convalently bonded, the shared electrons will be more strongly attracted to the atom of greater electronegativity. Such a bond is said to be polar. A polar bond results in the unequal sharing of the electrons in the bond.
The presence or absence of polar bonds within molecule plays a very important part in determining chemical and physical properities of those molecules. Some of these properties are melting points, boiling points, viscosity and solubilities in solvents.
Predicting Bond Types The difference in electronegativities of two elements can be used to predict the nature of the chemical bond. Bond type can be described as belonging to one of three classes: nonpolar covalent
polar covalent
ionic
When differences are 1.7 or greater, the bond is usually ionic. Less than 1.7, the bond is usually covalent, and unless the difference is less than 0.5 the bond has some degree of polarity. Differences of less than 0.5 are considered to be nonpolar.
The most commonly used electronegativity scale is Pauling's. It is the one we will usually use and side 2 of the Sargent-Welch Periodic Table gives the value for each element. See key for location within the element's block.
At the top of side 2 is a table of electronegativity differences and percent ionic character which you are to use to determine the type of bond between two atoms.
Note the electronegativities of elements from left to right across periods and down groups. What is the trend for each of these? (Notice that Pauling didn't assign a value to the Noble gases)
Across a period: the electronegativities generally increase from left to right across a period with the Group VII element having the highest value for the period. Down a group: the electronegativities generally decrease from top to bottom down a group. Francium is the element with the lowest electronegativity.
Hi there! This post couldn't be written any better! Reading through this post reminds me of my previous roommate! He constantly kept preaching about this. I'll forward this post to him. Pretty sure he's going to have a good read. Thank you for sharing!
6 comments:
Covalent bonds:
formed from sharing of two electrons, usually one donated from each of the two bonding atoms.
There are 3 kinds of covalent bonds based on number of pairs of electrons shared between the two atoms:
single covalent bond - 1 shared pair
double covalent bond - 2 shared pairs
triple covalent bond - 3 shared pairs
Bond energy: single < double < triple
Bond length: single < double < triple
Lewis Structures: The simplest covalent bond is between two hydrogen atoms. Electron-dot (or Lewis) symbol for hydrogen is H.. The two electrons from two hydrogen atoms pair up to make H:H. The pair of dots can be replaced by a dash to represent a bond, as in H-H. Hydrogen as the element exists as these "diatomic molecules" and its molecular formula is shown as H2.
The halogens also exist as diatomic molecules. Fluorine, chlorine, bromine and iodine all have the same type of Lewis structure, with seven outer electrons (2 on the right side, 2 on the top, 2 on the left side, and 1 on the bottom, so two F atoms share their unpaired electrons with each other to form the single covalent bond, F:F or F-F, to get F2. The rest of the pairs of electrons around each halogen are called "unshared pairs". The number of paired and unpaired electrons around an atom play an important part in determining the shape of a molecule, which in turn contribute to chemical and physical properties of the molecule.
Oxygen atoms, with six outer level electrons each ( 2 on the right side, 2 on the top, and 1 each on left side and bottom), share two pairs of electrons and thus have a double bond between them. This is shown by O::O or O=O for O2.
Nitrogen atoms, with only five outer electrons has 2 on the right side, and 1 each on the top, left side and bottom. Two nitrogen atoms share three electrons each, to form three pairs between them. This bonding is written as :N:::N:, N=N for N2
Things get a bit more complicated when atoms of different kinds of elements bond. One of the simplest bonds is between hydrogen and chlorine. The single unpaired electron on H pairs up with the unpaired electron on the chlorine and the result is H:Cl or H-Cl, or HCl.
Writing/Drawing Lewis Structures for simple molecules.
Count the number of valence electrons in each of the atoms in the molecule.
Count the total number of electrons that are needed to give each atom in the molecule an octet in the valence level, or two electrons for hydrogen.
Find the difference between the number of electrons available for sharing (#1) and the total number needed to complete the octet for each (#2). This gives the number of electrons that will be shared among the bonding atoms.
Divide the number of shared electrons by 2 to give the number of electron pairs, thus the number of bonds between all of the atoms in the molecules. A double or triple bond count as two or three bonds because of the number of electron pairs they represent.
Examples:
H2O
CO2
CH2O
C2H4
There are several rules to remember when writing Lewis structures:
H is always terminal, only one single bond per hydrogen
Oxygen almost always has two shared pairs and two unshared pairs of electrons. The two shared pairs may be in two separate single bonds or in one double bond.
Carbon atoms will always have 4 bonds, either as 4 separate single bonds, or a combination of single, double and/or triple bonds so that there are 4 pairs of electrons around each carbon atom in the molecule. Carbon atoms may single, double, or triple bond to each other, single or double bond to oxygen atoms, or single, double, or triple bond to nitrogen atoms. Halogen atoms form single bonds only, and usually form just one bond. An example of an exception is the chlorate, ClO31-, ion.
What is electronegativity?
Electronegativity is a measure of the attraction of an atom for electrons in a covalent bond.
Fluorine, the most reactive non-metal, is assigned the highest value since it has the greatest attraction for the electron being shared by the other element. Oxygen is also highly electronegative and has a strong attraction for electrons.
Metals have low electronegativities since they have weak attraction for any shared electrons.
When two unlike atoms are convalently bonded, the shared electrons will be more strongly attracted to the atom of greater electronegativity. Such a bond is said to be polar. A polar bond results in the unequal sharing of the electrons in the bond.
The presence or absence of polar bonds within molecule plays a very important part in determining chemical and physical properities of those molecules. Some of these properties are melting points, boiling points, viscosity and solubilities in solvents.
Predicting Bond Types
The difference in electronegativities of two elements can be used to predict the nature of the chemical bond.
Bond type can be described as belonging to one of three classes:
nonpolar covalent
polar covalent
ionic
When differences are 1.7 or greater, the bond is usually ionic.
Less than 1.7, the bond is usually covalent, and unless the difference is less than 0.5 the bond has some degree of polarity.
Differences of less than 0.5 are considered to be nonpolar.
The most commonly used electronegativity scale is Pauling's. It is the one we will usually use and side 2 of the Sargent-Welch Periodic Table gives the value for each element.
See key for location within the element's block.
At the top of side 2 is a table of electronegativity differences and percent ionic character which you are to use to determine the type of bond between two atoms.
Note the electronegativities of elements from left to right across periods and down groups.
What is the trend for each of these? (Notice that Pauling didn't assign a value to the Noble gases)
Across a period: the electronegativities generally increase from left to right across a period with the Group VII element having the highest value for the period.
Down a group: the electronegativities generally decrease from top to bottom down a group. Francium is the element with the lowest electronegativity.
Hey Chumi cud u plz post ur Bekius notes plz.
thanx so much they r such a big help.
Attn gr8 ppl of the 12th grade I majorly need bikiyus notes so to those kind soles who wish to help please post notes!!!!
Thank you in advance
whoever wrote the first two comments is pretty brilliant....
i could barely take the test
Hi there! This post couldn't be written any better! Reading through this post reminds me of my previous roommate! He constantly kept preaching about this. I'll forward this post to him.
Pretty sure he's going to have a good read. Thank you for sharing!
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