chemical bonding and nomenclature chemical bonding and nomenclature by paul surko new dimensions...
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Chemical Bonding Chemical Bonding and Nomenclatureand Nomenclature
By Paul SurkoNew Dimensions High SchoolPoinciana, FL
s
8
I want you to meet a friend of mine?Bonding, the way atoms are attracted to each other to form molecules, determines nearly all of the chemical properties we see. And, as we shall see, the number “8” is very important to chemical bonding.
5.1 What are Molecules?5.1 What are Molecules?Molecules are a
combination of atoms bonded together. Bonding determines the chemical
properties of the molecule (compound).
5.5 Ionic Bonding-Being Like 5.5 Ionic Bonding-Being Like the Noble Gasesthe Noble Gases
All atoms want to have the same number of electrons as the Noble Gases. The Noble Gases have very stable electron configurations. In order to achieve the same electron configuration as the Noble Gases metal atoms will 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+3
Cl becomes Cl- O becomes O-2
Mg becomes Mg+2Na becomes Na+
The positive and negative ions are attracted to each other electrostatically.
Opposites Attract!Opposites Attract!
Putting Ions TogetherPutting Ions Together
Na+ + Cl- = NaCl
Ca+2 + O-2= CaO Na+ + O-2 = Na2O
Al+3 + S-2 = Al2S
3Ca+2 + N-3 = Ca
3N
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
5.2 The Covalent Bond5.2 The Covalent BondAtoms can form molecules by sharing Atoms can form molecules by sharing
electrons in the covalent bond. This is done electrons in the covalent bond. This is done only among non-metal atoms.only among non-metal atoms.
5.3 Dot Structures-Octet Rule5.3 Dot Structures-Octet Rule(All atoms want 8 electrons around them.)(All atoms want 8 electrons around them.)
Valence electrons are those in the outermost orbitals. They are the ones that can form bonds.
Lewis came up with a way to draw valence electrons so that the bonding could be determined.
Rules to Write Dot StructuresRules to Write Dot Structures1.1. Write a skeleton molecule with the lone atom in the middle Write a skeleton molecule with the lone atom in the middle (Hydrogen can never be in the middle)(Hydrogen can never be in the middle)2.2. Find the number of electrons needed (N) Find the number of electrons needed (N) (8 x number of atoms, 2 x number of H atoms)(8 x number of atoms, 2 x number of H atoms)3.3. Find the number of electrons you have (valence eFind the number of electrons you have (valence e--'s) (H)'s) (H)4.4. Subtract to find the number of bonding electrons (N-H=B) Subtract to find the number of bonding electrons (N-H=B) 5.5. Subtract again to find the number of non-bonding Subtract again to find the number of non-bonding electrons (H-B=NB)electrons (H-B=NB)6.6. Insert minimum number of bonding electrons in the Insert minimum number of bonding electrons in the skeleton between atoms only. Add more bonding if needed skeleton between atoms only. Add more bonding if needed until you have B bonding electrons.until you have B bonding electrons.7.7. Insert needed non-bonding electrons around (not Insert needed non-bonding electrons around (not between) atoms so that all atoms have 8 electrons around between) atoms so that all atoms have 8 electrons around them. The total should be the same as NB in 5 above.them. The total should be the same as NB in 5 above.
Let's Try it!Let's Try it!1.1.SS
2.2.NN
3.3.HH
4.4.BB
5.5.NBNB
6.6.EE ..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 ●●
Let's Try it!Let's Try it!1.1.SS
2.2.NN
3.3.HH
4.4.BB
5.5.NBNB
6.6.EE ..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
Let's Try it!Let's Try it!1.1.SS
2.2.NN
3.3.HH
4.4.BB
5.5.NBNB
6.6.EE .. ..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 ●● ●●
Let's Try it!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
5.6 Polarity-Unequal Sharing 5.6 Polarity-Unequal Sharing of Electronsof Electrons
Even though all atoms want the same number of electrons as the Noble Gases, some want to get or give them more than others. The magnitude of this attraction for electrons is called “Electronegativity”. The more electronegative an atom is, the more it wants the electrons.
Some atoms want to gain electrons so bad, they take them altogether to form negative ions. Some want to lose them so bad that they become positive ions.
Examples of Polar and Non-Examples of Polar and Non-Polar CompoundsPolar Compounds
H2O Water is a bent molecule. The lone pair of electrons
from the Lewis structure distorts its shape and it becomes a very polar molecule.
NaCl Since Na is a metal it gives up its electron to form Na+ and Cl takes the electron completely to form Cl-.
HCl The Chlorine wants the electrons more than the Hydrogen. Thus we have +δHCl-δ.
Cl2 (Cl—Cl) The Chlorine molecules want the electrons
equally so they form a non-polar molecule with NO partial or full charges.
CO2 Carbon Dioxide is a linear molecule. It has no lone pairs
of electrons from the Lewis structure. The two oxygen atoms pull equally and make it a non-polar molecule.
..:O:H ●●
H
.. ..O::C::O●● ●●
5.7 Nomenclature5.7 NomenclatureNaming of CompoundsNaming of Compounds
Binary Compounds have two types of atoms (not diatomic which has Binary Compounds have two types of atoms (not diatomic which has only two atoms). 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
5.7 Nomenclature5.7 NomenclatureNaming of CompoundsNaming of Compounds
Binary Compounds have two types of atoms (not diatomic which has Binary Compounds have two types of atoms (not diatomic which has only two atoms). only two atoms).
Metals (Transition Metals) and Non-MetalsOlder System
Metal (Latin) _______ + ous or ic + Non-Metal ________ide Ferrous Bromine
Ferrous Bromide FeBr2
Compare with Ferric Bromide FeBr3
Non-Metals and Non-MetalsUse 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
Let’s Practice!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
Polyatomic IonsPolyatomic Ions(partial list from page 195 (193 2(partial list from page 195 (193 2ndnd edition)) edition))
Ammonium……………...Ammonium……………... Nitrate……………………Nitrate…………………… Permanganate…………. . Permanganate…………. . Chlorate…………………Chlorate………………… Hydroxide……………….Hydroxide………………. Cyanide………………….Cyanide…………………. Sulfate…………………...Sulfate…………………... Carbonate……………….Carbonate………………. Chromate………………..Chromate……………….. Acetate…………………..Acetate………………….. Phosphate……………….Phosphate……………….
NHNH44++
NONO33--
MnOMnO44--
ClOClO33--
OHOH--
CNCN--
SOSO4 4 2 -2 -
COCO332-2-
CrOCrO442-2-
CC22HH33OO22--
POPO443-3-
Acids Acids (with H in front)(with H in front)Binary acids (without oxygen in formula)
Hydro _________ ic Acid
HCl Hydrochloric acid HBr Hydrobromic acid
Oxy acids (with oxygen in formula)
-ate goes to –ic and –ite goes to -ous
HNO3 Nitric acid HNO2 Nitrous acid
H2SO4 Sulfuric acid H2SO3 Sulfurous acid
H3PO4 Phosphoric acid H3PO3 Phosphorous acid
Lets Practice!Lets Practice!
HFNa2CO3
H2CO3
KMnO4
HClO4
H2S
NaOH
CuSO4
PbCrO4
H2O
NH3
Hydrooxic acid (no……just water)
Nitrogen trihydride (no..just ammonia)
Copper (II) sulfate or Cupric sulfate
Lead (II) chromate or Plubous chromate
Sodium hydroxide
Hyrdogen sulfuric acidPerchloric acid
Potassium permanganate
Sodium carbonate
Hydroflouric acid
Carbonic acid