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Chemical Bonds
Bettelheim, Brown, Campbell and Farrell
Chapter 3
Ionization Energy
• Energy required to remove outermost electron (most loosely held)
Ionization energy
Noble Gas Configurations
He 1s2
Ne
Ar
Kr
Xe
[He]2s2 2p6
[Ne]3s2 3p6
[Ar]4s2 4p6
[Kr]5s2 5p6
Noblegas
Noble gasnotation
He
8A
Ne
Ar
Kr
Xe131.3
Rn(222)
BoilingPoint(°C)
MeltingPoint(°C)Element
HeliumNeonArgonKryptonXenon
-272-249-189-157-112
-269-246-186-152-107
Radon -71 -62
2
4.00310
20.1818
39.9536
83.8054
86
Noble gas configuration s2p6 very stable
The Octet Rule
• Octet rule:Octet rule: Group 1A-7A elements to achieve an outer shell of eight valence electrons
• Anion: Anion: Negative ion formed when an atom gains electrons
• Cation:Cation: Positive ion formed when an atom loses electrons
The Octet Rule—Cations
Cation: Sodium atom loses an electron to form a sodium ion, which has the same electron configuration as neon
Na (11 electrons): 1s2 2s2 2p6 3s1
Na+ (10 electrons): 1s2 2s2 2p6
Na Na+ + e-
A sodiumatom
A sodiumion
Anelectron
The Octet Rule—Anions
Anion: Chlorine atom gains an electron to form a chloride ion, which has the same electron configuration as argon
chlorine atom (17 electrons): 1s2 2s2 2p6 3s2 3p5
chloride ion (18 electrons): 1s2 2s2 2p6 3s2 3p6
Anelectron
Cl Cl-+ e-
A chlorineatom
A chlorideion
Cation Names
• Groups 1A, 2A, and 3A – The name of the element followed by the word
“ion”
Li+H+ Hydrogen ion
Lithium ion
Sodium ionPotassium ion
Ion
K+
Na+
Mg2+ Magnesium ion
Calcium ion
Strontium ionBarium ion
Ca2+
Sr2+
Ba2+
Al3+ Aluminum ion
Name Ion Name Ion Name
Group 1A Group 2A Group 3A
Transition Metal Cations• Cations derived from transition and inner
transition elements more than one type of cation
• Stock System (IUPAC): – Use Roman numerals to show charge: – Fe2+ is Iron (II) Fe3+ is Iron (III)– Cu+ is Copper (I) Cu2+ is Copper (II)
• Old System:– Use the suffix -ous-ous to show the lower positive charge
and the suffix -ic-ic to show the higher positive charge– Fe2+ is Ferrous Fe3+ is Ferric– Cu+ is Cuprous Cu2+ is Cupric
Transition Metal Ion Names
Fe3+
Fe2+
Hg+
Hg2+
Cu2+
Cu+ Copper(I) ionCopper(II) ion
Iron(II) ionIron(III) ion
Mercury(I) ionMercury(II) ion
Cuprous ionCupric ion
Ferrous ionFerric ion
Mercurous ionMercuric ion
Cupr- from cuprum, the Latinname for copper
Hg from hydrargyrum, theLatin name for mercury
IonSystematic name
Common name
Origin of the symbol of theelement or the common name of the ion
Ferr- from ferrum, the Latinname for iron
Sn2+
Sn4+
Tin(II) ionTin(IV) ion
Stannous ionStannic ion
Sn from stannum, theLatin name for tin
Anion Names
• Add “ide” to the root name of the element
Anion
F-
Cl-
Br-
I-
O2-
S2-
Stemname
fluorchlorbromiodoxsulf
Anionname
fluoridechloridebromideiodideoxidesulfide
Polyatomic Ions– Contain two or more atoms– Common names often used (in parentheses)
NH4+
OH-
NO2-
NO3-
CH3COO-
CN-
MnO4-
CO32-
HCO3-
SO32-
HSO3-
SO42-
PO43-
HPO32-
H2PO4-
HSO4-
CrO42-
Ammonium
HydroxideNitrite
NitrateAcetate
Cyanide
Permanganate
Carbonate
Hydrogen carbonate (Bicarbonate)
IonName
SulfiteHydrogen sulfite (Bisulfite)
Sulfate
Phosphate
Hydrogen phosphate
Dihydrogen phosphate
Name
Hydrogen sulfate (Bisulfate)
Chromate
Ion
Naming Ionic Compounds
• Name the positive ion first, then the negative ion
• Number of each ion not included
NaBr Al2O3
MgSO4 K2S
(NH4)3PO4
Naming Ionic Compounds
NaBr Sodium bromide
Al2O3 Aluminum oxide
MgSO4 Magnesium sulfate
K2S Potassium sulfide
(NH4)3PO4 Ammonium phosphate
Formulas of Ionic Compounds
• The total number of positive charges must equal the total number of negative charges
– Li+ and Br- form LiBr (+1) + (-1)
– Ba2+ and I- form BaI2 (+2) + 2(-1)
– Al3+ and S2- form Al2S3 2(+3) + 3(-2)
Forming Chemical Bonds
• Ionic bond:Ionic bond: the force of electrostatic attraction between a cation and an anion– Atom loses or gains electrons to make a filled
valence shell (octet) and become an ion.
• Covalent bond:Covalent bond: a pair of electrons that are shared by two atoms– Atom shares electrons to make a filled
valence shell (octet)
Forming an Ionic Bond--NaCl
• Formation of sodium chloride, NaCl
Single-headed curved arrow used to show the transfer of the electron
Na + Cl Na+ Cl -
+Na+(1s22s22p6)
Cl(1s22s22p63s23p5)+Na(1s22s22p63s1)
Cl-(1s22s22p63s23p6)
Sodium atom Chlorine atom
Sodium ion Chloride ion
NaClformation.mov
Ionic Bonds• Force of attraction between a cation and an
anion.
• Depends on electronegativity– measure of an atom’s attraction for shared pair
of electrons in chemical bond with another atom)
- -
Electronegativity increases
Covalent Bonds
• Result of one or more pairs of electrons that are shared by two atoms– Each atom has full valence shell (octet)
• In H2, each hydrogen contributes one electron to the single bond
+H H
the single line represents a shared pair of electrons
.. H H
Molecular Compounds
• Molecular compound:Molecular compound: only covalent bonds
• Naming molecular compounds– the less electronegative element is named first (it
is generally written first in the formula)– prefixes “di-”, tri-”, etc. are used to show the
number of atoms of each element; the prefix “mono-” is generally omitted
• Exception: carbon monoxide• NO is nitrogen oxide (nitric oxide)
• SF2 is sulfur difluoride
• N2O is dinitrogen oxide (laughing gas)
Electronegativity
F has highest value Noble gases have 0 value
Ionic bonds form when electronegativity difference ≥ 1.9
Polarity of Bonds
• Nonpolar:Nonpolar: Electrons are shared equally
• Polar:Polar: Electrons are NOT shared equally
Type of Bond
less than 0.5
0.5 to 1.9
greater than 1.9
nonpolar covalent
polar covalent
ionic
two nonmetals or anonmetal and a metalloid
ElectronegativityDifference Between
Bonded Atoms
a metal and a nonmetal
Most likely to be Formed Between
Polarity of Covalent Bonds
H-Cl
BondDifference in Electronegativity Type of Bond
3.5 - 2.1 = 1.43.0 - 2.1 = 0.94.0 - 0.9 = 3.12.5 - 1.2 = 1.3
polar covalentpolar covalentionicpolar covalent
2.5 - 2.5 = 0.0 nonpolar covalent
3.0 - 2.1 = 0.9 polar covalentO-HN-HNa-FC-MgC-S
Polarity of Covalent Bonds
More electron density shown by δδ-- or the head of a crossed arrow
Less electron density shown by δδ++ or the tail of a crossed arrow
Polarity of Molecules• Polar molecule
– has polar bonds, and– Has partial positive and partial negative
charges in different parts of molecule, i.e., is a dipole (has two poles)
• Carbon dioxide, CO2, has two polar bonds but, because of its geometry, the pulls balance out so it is a nonpolar molecule
O C O-- +
Carbon dioxide(a nonpolar molecule)
::
::
Polarity of Molecules
• Water, H2O, has two polar bonds and, because of its geometry, is a polar molecule
OH H
-
+center of partial positivecharge is midway betweenthe two hydrogen atoms
Water(a polar molecule)
Polarity of Molecules
• Both dichloromethane, CH2Cl2, and formaldehyde, CH2O, have polar bonds and are polar molecules
C
H HFormaldehyde
O+-
Dichloromethane
C
H H
Cl Cl+--
Lewis Structures
• Used to decide on the arrangement of atoms in the molecule
• Bonding (shared) electrons are shown as bonds (lines)
• Nonbonding electrons are represented as a pair of Lewis dots
Drawing Lewis Structures
1. Determine the number of valence electrons in the molecule
2. Decide on the arrangement of atoms in the molecule
3. Connect the atoms by single bonds
4. Show bonding electrons as a single line; show nonbonding electrons as a pair of Lewis dots
5. In a single bondsingle bond, atoms share one pair of electrons; in a double bonddouble bond, they share two pairs, and in a triple bondtriple bond they share three pairs.
Exceptions to the Octet Rule
• H and He have a maximum of 2 electrons (duet)
• Period 2 elements have a maximum of 8 electrons (use 2s and 2p orbitals)
• Atoms of period 3 elements may have more than 8 electrons
Lewis Structures• Examples:Examples: (the number of valence
electrons is given in parentheses after the molecular formula
Carbonic acidFormaldehydeAcetyleneEthylene
Hydrogen chlorideMethaneAmmoniaWater
H
H N H C H H ClH
HC C
HC C HH
HC
HHO
H
H2O (8) NH3 (8) CH4 (8) HCl (8)
C2H4 (12) C2H2 (10) CH2O (12) H2CO3 (24)
H
HHO
H
O OC HH
O
Lewis Structures
• ExamplesNH3
CH3OH
CH3COOH
.
Valence-Shell Electron-Pair (VSEPR) Model
• Because like charges repel each other, the various regions of electron density around an atom spread so that they are as far away from each other as possible– CH4: measured H-C-H bond angles
are 109.5°
H C HH
H
Summary of Molecular Geometry
# Electron Lone Bond Angle Shape regions pairs
2 0 180o Linear
3 0 120o Planar 2 1 Angular (bent)
4 0 109.5o Tetrahedral 3 1 Trigonal pyramidal 2 2 Angular (bent)
HH
NH
CH H
HC C
H
O C
HC
HHO
H
CH C HO
HH
O109.5°4
PredictedBond
AnglesExamples
(Shape of the molecule)
2 180°
120°3
PredictedDistributionof ElectronDensity
tetrahedral
trigonal planar
linear
Methane(Tetrahedral)
Ammonia(Pyramidal)
Water(Bent)
Ethylene(Planar)
Formaldehyde(Planar)
Carbon dioxide(Linear)
Acetylene(Linear)
H
H
: :
:
: :
: :
:
:
Regions ofElectron DensityAround Central
Atom
VSEPR Model
– the measured H-N-H bond angles are 107.3°– the unshared pair is not shown on this model
– the measured H-O-H bond angle is 104.5°
H N HH
H O Htetrahedron.mov