1 Oxidation Numbers: Rules 1)The oxidation number of the atoms in any free, uncombined element, is...
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Transcript of 1 Oxidation Numbers: Rules 1)The oxidation number of the atoms in any free, uncombined element, is...
![Page 1: 1 Oxidation Numbers: Rules 1)The oxidation number of the atoms in any free, uncombined element, is zero 2)The sum of the oxidation numbers of all atoms.](https://reader035.fdocuments.in/reader035/viewer/2022062407/56649d365503460f94a0e2ba/html5/thumbnails/1.jpg)
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Oxidation Numbers: Rules
1) The oxidation number of the atoms in any free, uncombined element, is zero
2) The sum of the oxidation numbers of all atoms in a compound is zero
3) The sum of the oxidation numbers of all atoms in an ion is equal to the charge of the ion
4) The oxidation number of fluorine in all its compounds is –1
5) The oxidation number of other halogens in their compounds is usually –1
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Oxidation Numbers: Rules6) The oxidation number of hydrogen is +1
when it is combined with more electronegative elements (most nonmetals) and –1 when it is combined with more electropositive elements (metals)
7) The oxidation number of Group 1A elements is always +1 and the oxidation number of Group 2A elements is always +2
8) The oxidation number of oxygen in most compounds is –2
9) Oxidation numbers for other elements are usually determined by the number of electrons they need to gain or lose to attain the electron configuration of a noble gas
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Ionic Bonding
Na – e– Na+
Cl + e– Cl–
Na + Cl Na+ + Cl–
Na+ cations and Cl– anions are electrostatically attracted to each other resulting in an extended ionic lattice
We say that Na+ and Cl- ions are held together by ionic bonding
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F2 Molecule
This bond is called a nonpolar covalent bond It is characterized by the symmetrical
charge distribution
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HF Molecule
F is more electronegative than H In this molecule the electron pair
will be shifted towards the F atom
This bond is called a polar covalent bond The charge distribution is not symmetrical
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Electron Density Distribution
Blue – low electron density (more positive) Red – high electron density (more negative)
H F
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Polar Bonds
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Polar Molecules Polar molecules can be attracted
by magnetic and electric fields We sometimes represent these
molecules as dipoles The direction of the dipole is from
the positive to the negative pole Each dipole is characterized by a
dipole moment The larger the difference in the
electronegativities of the bonded elements, the higher the dipole moment of the molecule
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The Continuous Range of Bonding Types
Covalent and ionic bonding represent two extremes:
In pure nonpolar covalent bonds electrons are equally shared by the atoms
In pure electrostatic ionic bonds electrons are completely transferred from one atom to the other
Most compounds fall somewhere between these two extremes
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All bonds have some ionic and some covalent character For example, HI is about 17%
ionic and 83% covalent
As the electronegativity difference increases, the bond becomes more polar less covalent more ionic
The Continuous Range of Bonding Types
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Example 1
Which of these bonds is more polar: NO CCl NaH NaBr
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Example 2
Which of these bonds is less covalent: AlI AlCl AlF AlBr
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Example 3
Which of these bonds has the highest dipole moment: CB CC CN CO CF
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The Octet Rule In most of their compounds,
the representative elements achieve noble gas configurations
Lewis dot formulas are based on the octet rule
Electrons which are shared among two atoms are called bonding electrons
Unshared electrons are called lone pairs or nonbonding electrons
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H2O Molecule
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NH3 Molecule
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NH4+ Ion
Lewis formulas can also be drawn for polyatomic ions
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CO2 Molecule
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N2 Molecule
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Covalent Bonding
Covalent bonds are formed when atoms share electrons If the atoms share 2 electrons a
single covalent bond is formed
If the atoms share 4 electrons a double covalent bond is formed
If the atoms share 6 electrons a triple covalent bond is formed
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The Octet RuleS = N - A
S = total number of electrons shared in bonds
N = total number of electrons needed to achieve a noble gas configuration 8 for representative elements 2 for H atoms
A = total number of electrons available in valence shells of the atoms A is equal to the periodic group
number for each element
A-S = number of electrons in lone pairs
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Examples F2
H2O
CH4
CO2
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Examples N2
CO
C2H2
HCN
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For ions we must adjust the number of electrons available, A: Add one e- to A for each negative charge Subtract one e- from A for each positive charge
Examples
NH4+
BF4–
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Example: CO32-
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Resonance
There are three possible structures for CO32-
The double bond can be placed in one of three places
These are called equivalent resonance structures
The real structure of the CO32- anion is an
average of these three resonance structures
COO
O
COO
O
COO
O
2- 2- 2-
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Resonance
There are no single or double bonds in CO32-
All three bonds are equivalent
They are intermediate between the single and double bond
C
OO
O
2-
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Resonance: Other Examples SO3
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Resonance: Other Examples NO3
–
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Resonance: Other Examples SO4
2–
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Exceptions to the Octet Rule
In those cases where the octet rule does not apply, the substituents attached to the central atom nearly always attain noble gas configurations
The central atom does not have a noble gas configuration but may have fewer than 8 or more than 8 electrons
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Examples
BBr3 AsF5
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Assignments & Reminders
Go through the lecture notes
Read Chapter 7 completely, except for Sections 7-7 & 7-8
Read Sections 4-5 & 4-6 of Chapter 4
Homework #4 due by Oct. 16 @ 3 p.m.
Review Session @ 5:15 p.m. on Sunday