Chapter 5
Chemical Bonds
Some Definitions
• Valence electrons
• Octet Rule
• Ion
– Cation
– Anion
• Electronegativity
• Polarity
Ions on the Periodic Table
CATION
+1 +2
CATION Non-Ionic
ANION
+3 ? -3 -2 -1 ?
Electron-dot structures
•
•
•
•
Also called Lewis Dot Structures
Represent an element’s valence electrons
Representative of the valence electrons for
the entire group
Features:
1. Each electron is represented by a dot
surrounding the atomic symbol of an element
2. Dots are only allowed in 4 positions (top,
bottom, left, right), no in-betweens
Electron-dot structures
•
What is the electron-dot structure for
sodium?
•
What is the electron-dot structure for
oxygen?
•
What is the electron-dot structure for neon?
Practice
• Write the electron-dot symbols for a sodium
metal atom and a chlorine atom, and their
ions.
Na
+
Na
Cl
Cl
Ionic Bonding
Sodium Chloride
Ionic Bonding
Atoms shuffle electrons, then are attracted to
each other because of net charge
=
IONIC BOND
Formulas of Ionic Bonding
• Textual way of describing what is physically happening
• Ions come first, then attraction of charges = BALANCE
Writing Ionic Formulas
Write the formula for calcium fluoride.
• Write the symbols for the ions, and their
charges.
• Cross over the charge numbers so they
become subscripts.
• Rewrite the formula, dropping the charge
superscripts.
2+
Ca 1
FCaF
2
2
1-
Naming Ionic Compounds
•
Between 2 Elements
1. Name of metal ion
2. Name of non-metal ion, –ide suffix
Those Pesky Transition Metals
•
•
Chromium Cr2+
chromium(II)
3+
Cr
chromium(III)
Variable valence =+ we cannot
predict ionic
Copper
Cu
copper(I)
charge from the Cu
group
number
2+
copper(II)
iron(II)
Roman Iron
numeral Fe
in3+2+parenthesis
is used to
Fe
iron (III)
indicateManganese
ion
Mn2+
manganese(II)
Mn3+
manganese(III)
Nickel
Ni2+
nickel(II)
Ni3+
nickel(III)
Zinc
Zn2+
zinc(II)
Naming Ionic Compounds
Write the name for ZnCl2
1. Cross over the subscripts so they become
charges
•
Remember that the metal (positive ion) is
always first.
2. Write the name of each element and add –
ide ending to the non-metal
3. Be sure to include a roman numeral in
parenthesis if the metal is a transition
element
Practice
• React lithium and sulfur together. What will
happen?
Li
+
Li
Stable?
S
-
2-
S
+
Li
2Li+ + S2- = Li2S
Bond, Chemical Bond
•
Ionic Bonding
1. Electron transfer from metal to non-metal
2. The more electronegative atom wins!
•
Covalent Bonding
1. Electron sharing between non-metals
2. Atoms of similar electronegativity
Covalent Bonding
• 2 Atoms share valence electrons to create
stable bonds
• No transference of electrons
• Between atoms of similar electronegativity
• Non-metals
• Atoms seeking to fill an octet: achieve the
next highest noble gas configuration
Naming Covalent Compounds
• First non-metal = elemental name
• Second non-metal = elemental name with
–ide ending
– E.g. oxide, sulfide, bromide
• Subscripts = use prefixes to name multiple
atoms
• Exception: for oxygen, drop the “a” at the
end of the prefixes
– E.g. tetroxide, hexoxide, pentoxide
Example
• Name the following covalent compound
N2O3
dinitrogen trioxide
C Cl4
* carbon tetrachloride
* By convention, the mono- on the first
atom is not generally included
Example
• Write the formula for the following
compound:
tetraphosphorus hexoxide
4 P’s
6 O’s
P4O6
Electron-dot Formulas
1. Determine the arrangement of atoms
•
For “central” atom(s), it will generally be single,
surrounded by multiple other atoms
2. Determine the total number of valence electrons
3. Attach the central atom to each bonded atom by
an electron pair
•
Represented by a dash
4. Arrange the remaining electrons around the outer
elements first, then the central atom
5. If octets are not complete, form a multiple bond
Practice
• Draw the electron-dot formula for nitrogen
tribromide
No multiple bonds necessary!
Br N Br
Br
Practice
• Draw the electron-dot formula for carbon
dioxide.
Multiple bonds necessary!
O C O
Diatomic Molecules
• Several non-metals are naturally diatomic:
more stable electron configuration
• Oxygen and nitrogen form multiple bonds
to achieve stability
N
N
N N
Polyatomic Ions
•
A group of atoms with a net electrical
charge




Most are non-metals covalently bonded to
oxygen
Net charge: -1, -2, -3 shared among the atoms
Two notable exceptions are +1:
1. Ammonium (NH4+)
2. Hydronium (H30+)
Formed when a larger molecule splits
unevenly
Names of Polyatomic Ions
• Memorize the common –ate ions
•
•
•
•
•
•
Nitrate
Chlorate
Bromate
Carbonate
Sulfate
Phosphate
NO3ClO3BrO3CO32SO42PO43-
Remember these exceptions!
hydroxide OHammonium NH4+
hydronium H3O+
– Add an O = per- prefix on –ate ion (e.g. perchlorate, ClO4-)
– Lose an O = -ate changes to –ite (e.g. chlorite, ClO2-)
– Lose another O = hypo- prefix on –ite ion (e.g.
hypochlorite, ClO-)
• All forms (-ate, -ite, per- hypo-) of an ion have
the same charge!
Naming Polyatomic Ions
• Recognize the polyatomic ion as a UNIT
• Name the metal ion first, then the polyatomic
ion
• No extra prefixes or suffixes
CaCO3
Calcium Carbonate
VSEPR Models
•
Based on several assumptions:
1. Atoms in a molecule are bound together by at least
one pair of electrons: a bonding pair
2. Some atoms may possess unbound, or lone pair,
electrons
3. Arrangements of electrons in the molecule will
minimize electron interactions
4. Lone pairs occupy more space than bonded pairs
5. Double bonds occupy more space than single
bonds
6. Multiple bonds are considered as single bonds
when determining molecule shape
VSEPR Models
•
Electron group geometry vs. Molecular
geometry (shape)
–
–
•
•
Electron group geometry = placement of electrons
around the central atom
Molecular geometry = 3-dimensional arrangement
of the molecule
Electron group geometry determines shape
Shape only refers to atoms and bonded
electron groups, not lone pairs
Points to Remember
1. ONLY the number of electron groups that
surround the central atom are considered in
determining molecular shape
2. 1 electron group may consist of a triple bond, a
double bond, a single bond, or a lone pair
3. A lone pair will give a different molecular shape
than a single, double, or triple-bonded atom at
the same position
VSEPR Models
Electron
Pairs
Bonded
Pairs
Lone Pairs
Molecular
Geometry
Examples
2
2
0
Linear
CO2
O =C =O
CO2
Source: http://www.molecules.org/VSEPR_table.html
VSEPR Models
Electron
Pairs
Bonded
Pairs
Lone Pairs
Molecular
Geometry
Example
3
2
1
Bent 120°
NO2-
N
O
NO2
Source: http://www.molecules.org/VSEPR_table.html
-
O
VSEPR Models
Electron
Pairs
Bonded
Pairs
Lone Pairs
Molecular
Geometry
Example
3
3
0
Triangular
(Planar)
NO3CO32-
O
N
NO3Source: http://www.molecules.org/VSEPR_table.html
O
O
VSEPR Models
Electron
Pairs
Bonded
Pairs
Lone Pairs
Molecular
Geometry
Example
4
2
2
Bent
109.5°
H2O
H
H
O
H2O
Source: http://www.molecules.org/VSEPR_table.html
VSEPR Models
Electron
Pairs
Bonded
Pairs
Lone
Pairs
Molecular
Geometry
Example
4
3
1
Pyramidal
NH3
H
NH3
Source: http://www.molecules.org/VSEPR_table.html
N H
H
VSEPR Models
Electron
Pairs
Bonded
Pairs
Lone
Pairs
Molecular
Geometry
Example
4
4
0
Tetrahedral
CH4
H
H
CH4
Source: http://www.molecules.org/VSEPR_table.html
C H
H
Electronegative Atom
Monster
Electronegativity
increases
Predicting Bond Type
Source: http://www.chem.ufl.edu/~chm2040/Notes/Chapter_11/writing.html
Polarity
• Is a BOND polar or non-polar?
• Is a MOLECULE polar or non-polar?
• Be careful! Polar bonds can make a nonpolar molecule!
Bond =
non-polar
Cl — Cl
Molecule =
non-polar
Polarity
• Is a BOND polar or non-polar?
• Is a MOLECULE polar or non-polar?
• Be careful! Polar bonds can make a nonpolar molecule!
Opposing dipoles cancel!
Bonds =
polar
O=C=O
Molecule =
non-polar
Polarity
• Is a BOND polar or non-polar?
• Is a MOLECULE polar or non-polar?
• Be careful! Polar bonds can make a nonpolar molecule!
Dipoles do not cancel!
Bonds =
polar
H
O
Molecule =
polar
H
Between Molecules: Hydrogen Bonds
• Hydrogen attached to highly electronegative
atoms F, O, N
– Strong dipoles
• Attraction of opposite dipoles
http://images.google.com/imgres?imgurl=http://polymer.bu.edu/~fstarr/net.jpg&imgrefurl=http://polymer.bu.edu/~fstarr/water.ht ml&h=248&w=350&sz=17&tbnid=4GT8aTZv8MkJ:&tbnh=82&tbnw
=116&hl=en&start=49&prev=/images%3Fq%3Dhydrogen%2Bbond%26start%3D40%26svnum%3D10%26hl%3Den%26lr%3D%26rls%3DGGLC,GGLC:1969 -53,GGLC:en%26sa%3DN
How Plants Work
Source: http://www.agridept.gov.lk/Techinformations/Hponics/images/P_24.jpg
Source:http://upload.wikimedia.org/wikipedia/en/thumb/c/c8/Coastredwood.jpg/250px-Coastredwood.jpg
Hydrogen Bonding at our Core
Source:http://colossus.chem.umass.edu/chandler/ch112/dna_hbond.gif
Between Molecules: Dipole-Dipole
• Different nonmetals bonded
together
• Partial charges:
Dipoles
• Attraction of
opposite dipoles
http://images.google.com/imgres?imgurl=http://library.tedankara.k12.tr/chemistry/vol2/dipole-dipole%2520forces/z105.jpg&imgrefurl=http://library.tedankara.k12.tr/chemistry/vol2/dipoledipole%2520forces/z105.htm&h=570&w=409&sz=65&tbnid=ljWYJbTg9PEJ:&tbnh=131&tbnw=93&hl=en&start=2&prev=/images%3Fq%3Ddipoledipole%26svnum%3D10%26hl%3Den%26lr%3D%26rls%3DGGLC,GGLC:1969-53,GGLC:en
Between Molecules: Dispersion
• Non-polar
molecules
• Temporary dipoles
because of nonuniform electron
cloud (not
dispersed evenly)
• Weak bonds
Source: http://touregypt.net/vdc/slide.86.jpg
To Sum Up the Forces…
Practice
• For your molecule, determine:
– Name
– Electron-dot structure
– 3-dimensional structure
– VSEPR shape
– Intramolecular force
– Intermolecular force