Transcript Document

Chemical Bonding
Page 2
• Chemical Bond
– attractive force between atoms or ions that binds
them together as a unit
– bonds form in order to…
• decrease potential energy (PE)
• increase stability
CHEMICAL FORMULA
IONIC
COVALENT
Formula
Unit
Molecular
Formula
NaCl
CO2
COMPOUND
2 elements
Binary
Compound
NaCl
more than 2
elements
Ternary
Compound
NaNO3
ION
1 atom
Monatomic
Ion
+
Na
2 or more atoms
Polyatomic
Ion
NO3
Chemical bonds are formed when
valence electrons are:
• transferred from one atom to another (ionic)
• shared between atoms (covalent)
• mobile within a metal (metallic)
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Ionic bonds are formed when metals
transfer their valence electrons to
nonmetals.
The oppositely charged ions attract
each other to form an ionic bond.
Sodium has one valence electron and chlorine has seven.
Sodium want to lose 1 electron and chlorine needs to gain 1.
Sodium transfers its valence electron to chlorine
Forming an Na+ and a Cl- ion – sodium chloride NaCl
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Electron-dot diagrams (Lewis
structures) can represent the
valence electron arrangement in
elements, compounds, and ions.
atom
ion
molecular
compound
ionic
compound
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Dots represent valence electrons.
Everything else (inner shell electrons
and nucleus) is called the Kernel and
is represented by the symbol.
Phosphorous has 5 valence electrons so we draw 5 dots
around the symbol for phosphorous.
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Draw the Lewis Dot Structures of
the first 18 elements.
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When metals lose electrons to
form ions, they lose all their
valence electrons. The Lewis Dot
Structure of a metal ion has no
dots. The charge indicates how
many electrons were lost.
Magnesium atom Magnesium ion
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When nonmetals gain electrons,
they fill up their valence shell
with a complete octet (except
hydrogen.) The ion is placed in
brackets with the charge outside
the brackets.
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A + metal ion is attracted to a –
nonmetal ion (opposites attract)
forming an ionic compound. We can
use Lewis dot structures to represent
ionic compounds.
The formula for magnesium fluoride is MgF2
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Two major categories of compounds
are ionic and molecular (covalent)
compounds. (5.2g)
• Ionic compounds are
formed when a metal
combines with a
nonmetal.
• Ionic compounds have
ionic bonds.
• Molecular compounds
are formed between
two nonmetals.
• Molecular compounds
have covalent bonds.
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Comparing the properties
compounds with ionic bonds and
compounds with covalent bonds.
 Properties of ionic
compounds
– Solids with high melting
and boiling points (strong
attraction between ions)
– Electrolytes: Do not
conduct electricity as solids
but do when dissolved or
molten – ions are charged
particles that are free to
move
– No individual molecules
 Properties of molecular
compounds
– Low melting and boiling
points (weak attraction
between molecules)
– Nonelectrolytes: Do not
conduct electricity as solids
or when dissolved or molten
– no charged particles (ions)
to move
– Solids are soft
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– Forms molecules
Ionic solids conduct electricity
when dissolved or molten.
Molecular solids do not.
Solution
doesn’t conduct
electricity
Solution
conducts
electricity
Ionic Solid
dissolved in water
Molecular Solid
dissolved in water
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Nomenclature
“Or How Do We Name Compounds”
Systematic Naming
• Compound is made up of two or more
elements
• Name should tell us how many and what type
of atoms
• Too many compounds to remember all the
names
Cation
– Positive ion
– Formed by losing
electrons
– Metals form cations
Anion
– Negative ion
– Has gained electrons
– Non metals form
anions
Ionic Compounds
• Made of cations and anions
• Metals and nonmetals
• Electrons lost by the cation are gained by the
anion
Ionic Compounds
Sodium is cation
Na
+
Cl
1+
Na +
1-
Cl
Chlorine is anion
Charges on Ions
Naming Ions
• Metal ion is written first in both name and formula
– It is named directly from element which formed the ion.
– Will nearly always be the positive ion or “cation”
– Transition metals can have more than one type of charge
– Indicate the charge with roman numerals in parenthesis.
Iron(II) or Iron(III)
– Exceptions:
• Silver always +1
• Cadmium and Zinc always +2
Name these
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Na 1+
Ca 2+
Al 3+
Fe 3+
Fe 2+
Pb 2+
Li 1+
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Sodium
Calcium
Aluminum
Iron (III)
Iron (II)
Lead (II)
Lithium
Write Formulas for these
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Potassium ion
Magnesium ion
Copper (II) ion
Chromium (VI) ion
Barium ion
Mercury (II) ion
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K1+
Mg2+
Cu2+
Cr6+
Ba2+
Hg2+
Naming Anions
• Anions are always the same.
• Change the element ending to -- ide
• F1- Fluorine to Fluoride
Name These
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Cl1N3Br 1O2I1Sr2+
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Chloride
Nitride
Bromide
Oxide
Iodide
Strontium
Write These
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Sulfide ion
Iodide ion
Phosphide ion
Strontium ion
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S2I1P3Sr2+
Polyatomic Ions
• Tightly bound groups of atoms acting as a
single ion.
• Names given in table in book. (pg 123)
• Most are anions that contain oxygen. Names
end in –ate (one more O), or –ite (one less O).
• SO32- = sulfite; SO42- = sulfate
• Exceptions: Ammonium cation NH4+, Cyanide
CN-, and hydroxide OH-
Naming Binary Ionic Compounds
• 2 elements involved
• Ionic – metal (cation) and a non-metal (anion)
• Naming is easy with representative elements
in A groups
• NaCl = Na+ Cl- = sodium chloride
• MgBr2 = Mg2+Br- = magnesium bromide
Naming Binary Ionic Compounds
• The problem comes with the transition
metals.
• Need to figure out their charges
• All ionic compounds will have a neutral charge
– Same number of + and – charges
• Use the anion to determine the charge on the
positive ion.
Naming Binary Ionic Compounds
• Try naming these
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KCl
Na3N
CrN
ScP
PbO
PbO2
Na2Se
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Potassium chloride
Sodium nitride
Chromium (III) nitride
Scandium (III) phosphide
Lead (II) oxide
Lead (IV) oxide
Sodium selenide
Tertiary Ionic Compounds
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Will have polyatomic ions
At least 3 elements
Use blue sheet
Name these ions
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NaNO3
CaSO4
CuSO3
(NH4)2O
•Sodium nitrate
•Calcium sulfate
•Copper (II) sulfite
•Ammonium oxide
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LiCN
Fe(OH)3
(NH4)2CO3
NiPO4
• Lithium cyanide
• Iron (III) hydroxide
• Ammonium
carbonate
• Nickel (III)
phosphate
Polyatomic ions are groups of
atoms covalently bonded
together that have a negative or
positive charge.
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Polyatomic ions are held
together by covalent bonds but
form ionic bonds with other ions.
H
Covalent
bonds
H N H
+
Cl
Ionic bond
-
H
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Writing Formulas
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Charges have to add up to zero.
Get charges on pieces from Periodic Table
Cations from element name on table
Anions from table change ending to –ide, or
use name of polyatomic ion
• Balance the charges
• Put polyatomics in parenthesis
Writing Formulas
• Write formula for calcium chloride
– Calcium is Ca2+
– Chloride is Cl1– Ca+2Cl-1 would have a +1 charge
– Need another Cl1– Ca+2Cl2-1
=
CaCl2
Writing Formulas
• Crisscross method
Calcium chloride
2+
Ca
1Cl
CaCl2
No need to write the
one
Iron (III) sulfide
3+
2S
Fe
Fe 2 S3
Fe2S3
Write Formulas for These
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Lithium sulfide
Tin (II) oxide
Tin (IV) oxide
Magnesium fluoride
Copper (II) sulfate
Iron (III) phosphide
Iron (III) sulfide
Ammonium chloride
Ammonium sulfide
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Li2S
SnO
SnO2
MgF2
CuSO4
FeP
Fe2S3
(NH4)Cl
(NH4)2S
Things to Look For
• If cations have ( ), the roman numeral is their
charge.
• If anions end in –ide they probably are off the
periodic table (monoatomic)
• If anion ends in –ate or –ite it is a polyatomic
ion
Molecular Compounds
Writing Names and Formulas
Covalent Bonding / Compounds
• Compounds in which the electronegativity
difference is less than 2.0
• Between a nonmetal and nonmetal
• Can’t be held together because of opposite
charges
• Can’t use charges to figure out how many of
each atom
Covalent Bonding
• Smallest piece of a covalently bonded
compound is a molecule
• Electrons are shared between atoms in bond
Carbon Dioxide
Water
H2O
CO2
Ammonia
NH3
In a multiple covalent bond, more
than one pair of electrons are
shared between two atoms.
(5.2e)
•Diatomic oxygen has a double bond O=O (2 shared pairs)
because oxygen needs 2 electrons to fill its valence shell
•Diatomic nitrogen has a triple bond NN (3 shared pairs)
because nitrogen needs 3 electrons to fill its valence shell
•Carbon dioxide has two double bonds
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Regents Question: 08/02 #17
Which molecule contains a triple covalent
bond?
(1) H 2
(2) N 2
(3) O 2
(4) Cl 2
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Molecular polarity can be determined
by the shape of the molecule and the
distribution of charge.
• Possible shapes
– Linear
– Bent
– Pyramidal
– Tetrahedral
(X2 HX CO2)
(H2O)
(NH3)
(CH4 CCl4)
A polar molecule is called a dipole. It has a positive
side and a negative side – uneven charge distribution.
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Symmetrical (nonpolar)
molecules include CO2 , CH4 ,
and diatomic elements. ..
Symmetrical molecules are not dipoles.
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Asymmetrical (polar)
molecules include HCl, NH3 ,
and H2 O. (5.2l)
The negative side of the molecule is the side that has the
atom with the higher electronegativity.
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Differences between ionic and covalent
bonding:
Ionic bonding
•
electron is “stolen”
•
high electronegativity
difference
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between metal & nonmetal
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Formation of crystal
structure
think proportions of
atoms in
formula unit NaCl 1:1
Na
+
Cl
+
Na +
Cl
Molecules are easier to name and work
with
• Ionic compounds use charges to determine
how many of each.
– Have to figure out charges
– Have to figure out numbers
• Molecular compound’s name tells you the
number of atoms.
Naming
• The second part of all names end with -ide
• Prefixes are used to indicate number of each
atom
Prefixes
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1
2
3
4
5
6
7
8
monoditritetrapentahexaheptaocta-
• 9 nona• 10 deca-
Naming Continued
• To write the name…write two words
Prefix-name Prefix-name –ide
• One exception is we don’t write mono- if
there is only one of the first element.
• No double vowels when writing names
– (oa oo)
Name These
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N2O
NO2
Cl2O7
CBr4
CO2
BaCl2
H2O
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Dinitrogen monoxide
Nitrogen dioxide
Dichlorine heptoxide
Carbon tetrabromide
Carbon dioxide
Barium chloride
Dihydrogen monoxide
Write Formulas for These
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Diphosphorous pentoxide
Tetraiodine monoxide
Sulfur hexaflouride
Nitrogen trioxide
Carbon tetrahydride
Phosphorous trifluoride
Aluminum chloride
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P2O5
I4O
SF6
NO3
CH4
PFl3
AlCl3
Lewis Dot Structure
(AKA Electron Dot Structure)
1. Write the symbol for each atom and show
each of their valence electrons as dots
(ignore all electrons below valence shell)
Cl
Cl2
Cl
Cl
Cl Cl
2. The number of electrons
before you combine the
atoms will equal number
you have after.
Comparison of Bonding Types
ionic
covalent
ions
molecules
nonconductive
molten salts
conductive
transfer of
electrons
high mp
DEN > 1.7
valence
electrons
sharing of
electrons
low mp
DEN < 1.7
The bonds holding metals
together in their crystal lattice
are called metallic bonds.
• All metals have metallic bonds
• “Positive ions immersed in a sea of mobile
electrons”
– Bonds are between Kernels, leaving the valence
electrons free to move from atom to atom
– Mobile electrons give metals the ability to
conduct electricity
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Intermolecular Forces
• Weaker than covalent bonds
• Weak intermolecular forces – lower boiling point
The stronger the intermolecular
forces, the higher the boiling
points and melting points.
Strongest
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Ionic Solids
Molecules with Hydrogen bonds
Polar molecules
Nonpolar molecules
Weakest
For nonpolar molecules, the greater the mass, the greater
the force of attraction.
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Hydrogen Bonds
• Hydrogen bonds are considered to be dipoledipole type interactions
• Hydrogen bonds vary from about 4 kJ/mol to
25 kJ/mol (so they are still weaker than typical
covalent bonds.
• But they are stronger than dipole-dipole and
or dispersion forces.
Hydrogen Bonds
Hydrogen
Bonds
ion-dipole forces
• Attractive forces between neutral molecules
and charged (ionic) compounds
Ion-dipole forces
(Ion-Molecule attraction)
•are important in solutions of ionic substances in polar solvents
•(e.g. a salt in aqueous solvent)
Van der Waals Forces
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Weak bonds
Liquefy gases
Bonds that combine gas molecules to form liquid
Ex. CO2 – liquid in toy car
- liquid nitrogen
• Molecules must be close to each other
• Larger atoms have stronger Van-der Waals forces
Dipole-dipole Forces
• Polar molecules attract one another when the partial positive
charge on one molecule is near the partial negative charge on
the other molecule
• The polar molecules must be in close proximity for the dipoledipole forces to be significant
• Dipole-dipole forces are characteristically weaker than iondipole forces
• Dipole-dipole forces increase with an increase in the polarity
of the molecule
London Dispersion Forces
• Nonpolar molecules would not seem to have any basis for
attractive interactions
• However, gases of nonpolar molecules can be liquefied
indicating that if the kinetic energy is reduced, some type of
attractive force can predominate.
• Fritz London (1930) suggested that the motion of electrons
within an atom or non-polar molecule can result in a transient
dipole moment
London Forces