OFB Chapter 3 lecture notes - Georgia Institute of...
Transcript of OFB Chapter 3 lecture notes - Georgia Institute of...
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1/12/2004 OFP Chapter 3 1
Chemical Periodicity and the Formation of Simple Compounds
OFB Chapter 3
3-1 Groups of Elements
3-2 The Periodic Table
3-3 Ions and Ionic Compounds
3-4 Covalent Bonding and Lewis Structures
3-5 Drawing Lewis Structures
3-6 Naming Compounds in which Covalent Bonding Occurs
3-7 the Shapes of Molecules
3-8 Elements Forming More than One Ion
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1/12/2004 OFP Chapter 3 2
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1/12/2004 OFP Chapter 3 3
Modern periodic law
The chemical and physical properties of the elements are periodic functions of their atomic numbers.
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1/12/2004 OFP Chapter 3 4
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1/12/2004 OFP Chapter 3 5
Elements are classified as metal, non-metals, or semi-metals, and also fall into groups based on similarities in chemical and physical properties.
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1/12/2004 OFP Chapter 3 6
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1/12/2004 OFP Chapter 3 7
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1/12/2004 OFP Chapter 3 8
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1/12/2004 OFP Chapter 3 9
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1/12/2004 OFP Chapter 3 10
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1/12/2004 OFP Chapter 3 11
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1/12/2004 OFP Chapter 3 12
The Electronegativity (the power of an atom when in chemical combination to attract electrons to itself) is a periodic property.
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1/12/2004 OFP Chapter 3 13
Lewis Structures: Ionic compounds
Whenever possible, the valence electrons in a compound are distributed in such a way that each main-group element in a molecule (except hydrogen) is surrounded by eight electrons (an octet of electrons). Hydrogen should have two electrons in such a structure.
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1/12/2004 OFP Chapter 3 14
Lewis Structures:
Na · Na + · Ca · Ca 2+
Sodium atom
Sodium ion
Calcium ion
Calcium atom
Loss of a valence electron
Gain of a valence electron
Combination to form the compound NaCl
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1/12/2004 OFP Chapter 3 15
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1/12/2004 OFP Chapter 3 16
Lewis Structures: Covalent Compounds
Whenever possible, the valence electrons in a compound are distributed in such a way that each main-group element in a molecule (except hydrogen) is surrounded by eight electrons (an octet of electrons). Hydrogen should have two electrons in such a structure.
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1/12/2004 OFP Chapter 3 17
Ammonia
:NH3
Water
H2O:
Methane
CH4
Ammonia and Water have a pair of “lone electrons” called an “lone pairs”
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1/12/2004 OFP Chapter 3 18
Ammonia (NH3)
H· ·H
H·
·N···.
H:N:H·· ··H
lone pair
H―N―H··
|
H
Nitrogen (Group V)
5 valence
+ 2 in 1s = 7 electrons
Lewis Structures
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1/12/2004 OFP Chapter 3 19
methane
ethane
ethylene
acetylene
carbon monoxideC HH
H
H
C O
C C
H
H
H
H
CC
H
H
H
H
H
H CC HH
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1/12/2004 OFP Chapter 3 20
Drawing Lewis Structures
1. Count up the total number of valence electrons available (symbolized by A) by first adding the group numbers of all the atoms present. If the species is a negative ion, add the absolute value of the total charge; if it is a positive ion, subtract it.
2. Calculate the total number of electrons needed (N) for each atom to have its ownnoble-gas set of electrons around it (two for hydrogen, eight for the elements from carbon on in the periodic table).
3. Subtract the number in step 1 from the number in step 2. This is the number of shared (or bonding) electrons present (S).
4. Assign two bonding electrons (as one shared pair) to each connection between two atoms in the molecule or ion.
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1/12/2004 OFP Chapter 3 21
Drawing Lewis Structures5. If any of the electrons earmarked for sharing remain, assign them in pairs by making some of the bonds double or triple bonds. In some cases, there may be more than one way to do this. In general, double bonds form only between atoms of carbon, nitrogen, oxygen, and sulfur.
6. Assign the remaining electrons as lone pairs to the atoms, giving octets to all atoms except hydrogen.
7. Determine the formal charge on each atom, and write it next to that atom. Check that the formal charges add to give a correct total charge on the molecule or molecular ion.
Formal charge = group number - lone-pair electrons - ½ (number of electrons in bonding pairs)
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1/12/2004 OFP Chapter 3 22
Drawing Lewis Structures
1. Count up the total number of valence electrons available (symbolized by A) by first adding the group numbers of all the atoms present. If the species is a negative ion, add the absolute value of the total charge; if it is a positive ion, subtract it.
C HH
H
H
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1/12/2004 OFP Chapter 3 23
Drawing Lewis Structures
2. Calculate the total number of electrons needed (N) for each atom to have its ownnoble-gas set of electrons around it (two for hydrogen, eight for the elements from carbon on in the periodic table).
C HH
H
H
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1/12/2004 OFP Chapter 3 24
Drawing Lewis Structures
3. Subtract the number in step 1 from the number in step 2. This is the number of shared (or bonding) electrons present (S).
C HH
H
H
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1/12/2004 OFP Chapter 3 25
Drawing Lewis Structures
4. Assign two bonding electrons (as one shared pair) to each connection between two atoms in the molecule or ion.
C HH
H
H
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1/12/2004 OFP Chapter 3 26
Drawing Lewis Structures
1. Count up the total number of valence electrons available (symbolized by A) by first adding the group numbers of all the atoms present. If the species is a negative ion, add the absolute value of the total charge; if it is a positive ion, subtract it.
C C
H
H
H
H
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1/12/2004 OFP Chapter 3 27
Drawing Lewis Structures
2. Calculate the total number of electrons needed (N) for each atom to have its ownnoble-gas set of electrons around it (two for hydrogen, eight for the elements from carbon on in the periodic table).
C C
H
H
H
H
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1/12/2004 OFP Chapter 3 28
Drawing Lewis Structures
3. Subtract the number in step 1 from the number in step 2. This is the number of shared (or bonding) electrons present (S).
C C
H
H
H
H
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1/12/2004 OFP Chapter 3 29
Drawing Lewis Structures
4. Assign two bonding electrons (as one shared pair) to each connection between two atoms in the molecule or ion.
C C
H
H
H
H
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1/12/2004 OFP Chapter 3 30
Drawing Lewis Structures5. If any of the electrons earmarked for sharing remain, assign them in pairs by making some of the bonds double or triple bonds. In some cases, there may be more than one way to do this. In general, double bonds form only between atoms of carbon, nitrogen, oxygen, and sulfur.
C C
H
H
H
H
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1/12/2004 OFP Chapter 3 31
Drawing Lewis Structures
1. Count up the total number of valence electrons available (symbolized by A) by first adding the group numbers of all the atoms present. If the species is a negative ion, add the absolute value of the total charge; if it is a positive ion, subtract it.
C O
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1/12/2004 OFP Chapter 3 32
Drawing Lewis Structures
2. Calculate the total number of electrons needed (N) for each atom to have its ownnoble-gas set of electrons around it (two for hydrogen, eight for the elements from carbon on in the periodic table).
C O
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1/12/2004 OFP Chapter 3 33
Drawing Lewis Structures
3. Subtract the number in step 1 from the number in step 2. This is the number of shared (or bonding) electrons present (S).
C O
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1/12/2004 OFP Chapter 3 34
Drawing Lewis Structures
4. Assign two bonding electrons (as one shared pair) to each connection between two atoms in the molecule or ion.
C O
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1/12/2004 OFP Chapter 3 35
Drawing Lewis Structures
5. If any of the electrons earmarked for sharing remain, assign them in pairs by making some of the bonds double or triple bonds. In some cases, there may be more than one way to do this. In general, double bonds form only between atoms of carbon, nitrogen, oxygen, and sulfur.
C O
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1/12/2004 OFP Chapter 3 36
Drawing Lewis Structures
6. Assign the remaining electrons as lone pairs to the atoms, giving octets to all atoms except hydrogen.
C O
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1/12/2004 OFP Chapter 3 37
Drawing Lewis Structures
7. Determine the formal charge on each atom, and write it next to that atom. Check that the formal charges add to give a correct total charge on the molecule or molecular ion.
Steps 1.-6.
C HH
H
H
Methane CH4
Formal charge = group number -lone-pair electrons - ½ (number of electrons in bonding pairs)
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1/12/2004 OFP Chapter 3 38
Drawing Lewis Structures
7. Determine the formal charge on each atom, and write it next to that atom. Check that the formal charges add to give a correct total charge on the molecule or molecular ion.
C HH
H
H
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1/12/2004 OFP Chapter 3 39
Drawing Lewis Structures
7. Determine the formal charge on each atom, and write it next to that atom. Check that the formal charges add to give a correct total charge on the molecule or molecular ion.
C C
H
H
H
H
Ethylene C2H4
Formal charge = group number -lone-pair electrons - ½ (number of electrons in bonding pairs)
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1/12/2004 OFP Chapter 3 40
Drawing Lewis Structures
7. Determine the formal charge on each atom, and write it next to that atom. Check that the formal charges add to give a correct total charge on the molecule or molecular ion.
C C
H
H
H
H
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1/12/2004 OFP Chapter 3 41
Drawing Lewis Structures
7. Determine the formal charge on each atom, and write it next to that atom. Check that the formal charges add to give a correct total charge on the molecule or molecular ion.
C O
Formal charge = group number - lone-pair electrons -½ (number of electrons in bonding pairs)
Carbon monoxide
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1/12/2004 OFP Chapter 3 42
3-7 the Shapes of Molecules
The VESPR theory
The Valence Shell Electron-Pair Repulsion Theory
Steric NumberSN = (number of atoms bonded to a central atom) + ( number of lone pairs on central atom)
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1/12/2004 OFP Chapter 3 43
Text Error OFB 4th Edition
Page 123
Exercise 3-12
SeOF4
The steric number for the Se atom is 5, not 4
3-7 The Shapes of Molecules
The VESPR theory
See website for other corrections in the text and solutions manual.
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1/12/2004 OFP Chapter 3 44
The VSEPR Theory
The Valence Shell Electron-Pair Repulsion Theory
Electron pairs in the valence shell of an atom repel each other on a spherical surface formed by the underlying core of the atom.
The geometry which applies to a particular arrangement is determined by the steric number(SN) of the central atom.
“Steric” means “having to do with space.” The steric number of an atom in a molecule can be determined by drawing the Lewis structure of the molecule and adding the number of atoms that are bonded to it and the number of lone pairs that it has.
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1/12/2004 OFP Chapter 3 45
Geometry and Steric Number• SN = 2 Linear 180°• SN = 3 Trigonal planar 120°• SN = 4 Tetrahedral 109.5°• SN = 5 Trigonal bipyramidal
– 90° (equatorial - axial or – 120° (equatorial – equatorial)
• SN = 6 Octahedral 90°• SN for Double and triple bonds
count the same as single bonded atoms
• SN = 5 Lone pairs occupy equatorial positions in preference to axial positions.
• When lone pairs are present, the situation is more complicated due to repulsive forces– Lone pr. vs lone pair >– Lone pr vs. bonding pair > – Bonding pr. Vs. bonding pair
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1/12/2004 OFP Chapter 3 46
Examples with no Lone Pairs on the central Atom
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The VSEPR Theory
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The VSEPR Theory
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The VSEPR Theory
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The VSEPR Theory
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The VSEPR Theory
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The VSEPR Theory
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1/12/2004 OFP Chapter 3 59
Dipole Moments• Bonded atoms share electrons
unequally, whenever they differ in Electronegativity
• E.g., HCl. The Cl atom carries a slightly negative electric charge and the H atom a slightly positive charge of equal magnitude. Aligns itself in an electric field
• Dipolar or polar molecules posses a dipole moment, µ
• To find µ, – first draw the VESPR molecular
geometry. – Next assign a dipole moment to each
bond, – then vectorially add the dipole moments
for each bond
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1/12/2004 OFP Chapter 3 61
3-8 Elements Forming more than One Ion
• Oxidation State (Oxidation Number)
– Ionic vs Covalent bonding– Not formal electric charges,
rather what the charge would be if the compound were ionic
– Range from -3 to +7• Examples
1. CrO3–
2. TlCl3–
3. Mn3N2–
4. VCl4 –
5. Mn2O7–
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Chapter 3Chemical Periodicity and the
Formation of Simple Compounds
• Examples / exercises– All (3-1 to 3-16)
• HW Problems– 8, 9, 18, 34, 45, 57, 59, 60,
69, 70