chemical bonding and nomenclature adapted from paul surko
TRANSCRIPT
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Chemical Bonding and Nomenclature
Adapted from Paul Surko
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Bonding, the way atoms are attracted to each other to form compounds, determines nearly all of the chemical properties we see. The number “8” is very important to chemical bonding.
What is Bonding????????
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What are Compounds?Compounds are a
combination of atoms bonded together. Bonding determines the chemical
properties of the compound.
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Ionic Bonding-Great with 8All atoms want 8 valence electrons. Metals give up electrons to form positive ions (cations) and non-metal atoms will receive or take additional electrons to become negative ions (anions). IONS are charged particles.
N becomes N-3
Al becomes Al+3Cl becomes Cl-
O becomes O-2
Mg becomes Mg+2
Na becomes Na+
The positive and negative ions are attracted to each other electrostatically.
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Properties of Ionic Compounds
Made of cations and anions Exist in crystalline structures (solids) at STP High MPs and BPs Conduct electricity in aqueous and molten
states
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Opposites Attract!
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Putting Ions TogetherNa+ + Cl- = NaCl
Ca+2 + O-2= CaO Na+ + O-2 = Na2O
Al+3 + S-2 = Al2S
3 Ca+2 + N-3 = Ca3N
2
Ca+2 + Cl- = CaCl2
You try these!
Mg+2 + F- =
NH4
+ + PO4
-3 =
K+ + Cl- =
Al+3 + I- =
Sr+2 + P-3 =
Li+ + Br- =
Sr3P
2
AlI3
MgF2
(NH4)
3PO
4
KCl
LiBr
Not NH43
PO4
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NomenclatureNaming of Ionic Compounds with TMs
Binary Compounds have two types of atoms (not diatomic which has only two atoms).
Metals (Groups I, II, and III) and Non-Metals
Metal _________ + Non-Metal _________ideSodium Chlorine
Sodium Chloride NaCl
Metals (Transition Metals) and Non-Metals
Metal ______ +Roman Numeral (__) + Non-Metal ________ide Iron III Bromine
Iron (III) Bromide FeBr3
Compare with Iron (II) Bromide FeBr2
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Metals (Transition Metals) and Non-MetalsOlder System
Ferrous Bromine
Ferrous Bromide FeBr2
Compare with Ferric Bromide FeBr3
Metal (Latin) _______ + ous or ic + Non-Metal ________ide
Nomenclature--Naming of Ionic Compounds with TM
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Let’s Practice!Name the following.
CaF2
K2S
CoI2
SnF2
SnF4
OF2
CuI2
CuI
SO2
SrS
LiBr
Strontium SulfideLithium Bromide
Copper (I) Iodide or Cuprous Iodide
Sulfur dioxide
Copper (II) Iodide or Cupric Iodide
Oxygen diflourideTin (IV) Flouride or Stannic Flouride
Tin (II) Flouride or Stannous Flouride
Cobalt (II) Iodide or Cobaltous IodidePotassium Sulfide
Calcium Flouride
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Polyatomic Ions(partial list from page 195 (193 2nd edition))
Ammonium……………... Nitrate…………………… Permanganate…………. . Chlorate………………… Hydroxide………………. Cyanide…………………. Sulfate…………………... Carbonate………………. Chromate……………….. Acetate………………….. Phosphate……………….
NH4+
NO3-
MnO4-
ClO3-
OH-
CN-
SO4 2 -
CO32-
CrO42-
C2H3O2-
PO43-
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Lets Practice!
Na2CO3
KMnO4
NaOH
CuSO4
PbCrO4
NH3ammonia
Copper (II) sulfate or Cupric sulfate
Lead (II) chromate or Plubous chromate
Sodium hydroxide
Potassium permanganate
Sodium carbonate
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The Covalent BondAtoms can form molecules by sharing
electrons in the covalent bond. This is done only among non-metal atoms.
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Properties of Covalent Compounds
Generally exist as liquids and gases at STP NO crystalline structure Low MPs and BPs Do NOT conduct electricity
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Dot Structures-Octet Rule(All atoms want 8 electrons around them.)
Lewis came up with a way to draw valence electrons so that the bonding could be
determined.
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Rules to Write Dot Structures1. Write a skeleton molecule with the lone atom in the middle (Hydrogen can never be in the middle)2. Find the number of electrons needed (N) (8 x number of atoms, 2 x number of H atoms)3. Find the number of electrons you have (valence e-'s) (H)4. Subtract to find the number of bonding electrons (N-H=B) 5. Subtract again to find the number of non-bonding electrons (H-B=NB)6. Insert minimum number of bonding electrons in the skeleton between atoms only. Add more bonding if needed until you have B bonding electrons.7. Insert needed non-bonding electrons around (not between) atoms so that all atoms have 8 electrons around them. The total should be the same as NB in 5 above.
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Let's Try it!1.S
2.N
3.H
4.B
5.NB
6.E ..H:O:H ●●
H O H Water H2O
2 x 2 = 4 for Hydrogen1 x 8 = 8 for Oxygen4+8=12 needed electrons
8 – 4 = 4 non-bonding electrons
2 x 1 = 2 for Hydrogen1 x 6 = 6 for Oxygen You have 8 available electrons
12 - 8 = 4 bonding electrons
8 H
12 N
4 B
4 NB
-
-
H:O:H
..H:O:H ●●
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Let's Try it!1.S
2.N
3.H
4.B
5.NB
6.E ..H:N:H ●●
HH N H Ammonia NH
3
3 x 2 = 6 for Hydrogen1 x 8 = 8 for Nitrogen6+8=14 needed electrons
8 – 6 = 2 non-bonding electrons
3 x 1 = 3 for Hydrogen1 x 5 = 5 for Nitrogen You have 8 available electrons
14 - 8 = 6 bonding electrons
8 H
14 N
6 B
2 NB
-
-
..H:N:H
..H:N:H ●●
H
HH
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Let's Try it!1.S
2.N
3.H
4.B
5.NB
6.E .. ..O::C::O●● ●●
O C O Carbon Dioxide CO2
1 x 8 = 8 for Carbon2 x 8 = 16 for Oxygen8+16=24 needed electrons
16 – 8 = 8 non-bonding electrons
1 x 4 = 4 for Carbon2 x 6 = 12 for Oxygen You have 16 available electrons
24 - 16 = 8 bonding electrons
16 H
24 N
8 B
8 NB
-
-
O::C::O
.. .. O::C::O ●● ●●
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Let's Try it!1.S
2.N
3.H
4.B
5.NB
6.E .. .. ..O::C: O:●● ●●
OO C O Carbonate CO
3-2
3 x 8 = 24 for Oxygen1 x 8 = 8 for Carbon24+8=32 needed electrons
24 – 8 = 16 non-bonding electrons
3 x 6 = 18 for Oxygen1 x 4= 4 for Carbon You have 22 + 2 more available e-'s
24 H32 N
8 B
16 NB
-
-
..O::C:O
.. .. .. O::C: O: ●● ●●
O
..:O: ..
:O:
32 - 24 = 8 bonding electrons
-2
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Nomenclature of Covalently Bonded Compounds--Molecules
Non-Metals and Non-Metals
Use Prefixes such as mono, di, tri, tetra, penta, hexa, hepta, etc.
CO2 Carbon dioxide CO Carbon monoxide
PCl3 Phosphorus trichloride CCl4 Carbon tetrachloride
N2O5 Dinitrogen pentoxide CS2 Carbon disulfide
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VSEPR Theory
Valence Shell Electron Pair Repulsion Theory—Geometric Shapes Linear, Bent Trigonal Planar, Trigonal Pyramidal Tetrahedral Trigonal Bipyramidal Octahedral
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VSEPR Theory
Why is H2O bent and CO2 linear?
O in water has lone pairs causing bending whereas the C in carbon dioxide does not
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VSEPR Theory
Lone pairs on the central atom cause crowding (INCREASED REPULSION) and result in bending
*Remember only the lone pairs on the central atom matter—the lone pairs on the external atoms do not crowd
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Polarity of Molecules
Polar Molecule: a molecule that has uneven distribution of charge—dipole moments do NOT cancel
Nonpolar Molecule: a molecule that has even distribution of charge—all of the dipoles cancel
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Polarity of Molecules Cont’d
Examples on the Board H2O
CH4
CO2
NH3
*BF Why does water exist as a liquid at STP and
carbon dioxide exists as a gas at STP?
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Intermolecular Forces
Intermolecular Forces: forces of attraction that exist between two molecules Hydrogen Bonding: an IMF results from the
attraction between hydrogen and a highly electronegative element like F, N, or O Rather strong force Responsible for water’s high surface tension, holding
together DNA, and varying BPs and MPs
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Intermolecular Forces
Dipole-Dipole Force: an IMF that exists between two polar molecules (hydrogen bonding is a special type of dipole-dipole force) Rather strong Used to predict MPs and BPs
Van der Waals Force: an IMF that exists between two nonpolar molecules Very weak Instanteous Predicts the low MPs and BPs of nonpolar molecues
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Forces between Ionic Solids
Electrostatic Forces: a force of attraction that exists between ionic compounds due to opposite charges VERY strong Responsible for high MPs and high BPs
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Hybridization Theory A theory that suggests that orbitals from
atoms will merge and create bonding orbitals of equivalent energy Sigma: bonding that occurs by overlapping
orbitals end to end Pi: bonding that occurs by overlapping orbitals
side to side
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Hybridization Theory
Sigma Bonds () Pi Bond ()
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Sigma and Pi Bonds
Single Bond: sigma only Double Bond: 1 sigma and 1 pi Triple: 1 sigma and 2 pi
How many sigma and pi bonds are found in the following: N2
C2H4
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Types of Hybridization
sp sp2
sp3
sp3d sp3d2
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Note: this theory uses both orbitals involved in bonding and orbitals holding lone pairs so CH4, NH3, and CH4 all have sp3 hybridization
Hybridization Theory
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Molecular Orbital Theory
Based on quantum mechanics Treats the electron as a moving object Relates to probability of location not exact
location
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Molecular Orbital Theory
Bonding Orbital: area of high electron probability that has lower energy than the orbitals of the separate atoms
Antibonding Orbital: area of high electron probability that has higher energy than the orbitals of the separate atoms
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Molecular Orbital Theory
Nonbonding Orbital: an orbital that does not contribute stability nor does it destabilize the molecule
Open parking space near door Open parking space far from door Occupied parking space near door
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Paramagnetism of Oxygen
MOT explains the paramagnetism of oxygen