Valence Electrons
• Electrons in outermost (highest E) shell
• Determine chemical properties of an atom
• Use to predict how many bonds an atom
will make
• Group # = number of valence electrons
• Can use electron-dot (Lewis-dot)
structures to show valence electrons
Electron-Dot Structures of
First 20 Elements
Octet Rule
• All noble gases have 8 valence electrons (except
He, which has 2)
• Noble gases do not combine with other elements
because they have a full valence electron shell
(very stable)
• Octet rule = tendency for atoms to gain, lose or
share electrons to obtain a noble gas electron
configuration (usually 8 valence electrons)
• Elements other than the noble gases thus
combine with other elements to form chemical
compounds
Ions
• Ions are atoms, or groups of atoms, with a different
number of electrons than protons
• Thus, ions have an overall charge, (+) or (-)
• An atom can gain electrons to become an anion (-)
• Or can lose electrons to become a cation (+)
• Ionization energy is E required to remove an
electron (takes more E to remove 2nd electron…)
• Ionic charges of most ions from main group
elements can be predicted from group #
Ionic Compounds
• Cation + anion = ionic compound
• Atoms in ionic compounds are held together by
ionic bonds (strong attraction of opposite charges)
• There is no sharing of electrons, one atom takes
electrons from the other
• Charges always balance in ionic compounds (net
charge is zero)
• It may take more than one atom of either ion to
balance the charge
• Generally, metals form cations and nonmetals
form anions
Sodium Gives an Electron to Chlorine to
Form Sodium Chloride
Magnesium Gives One Electron to Each of
Two Chlorines to Form Magnesium Chloride
Naming Ionic Compounds
• Name of metal cation is just name of metal
• Name of anion: replace end of name with ide
• Name of cation is followed by name of anion
Example: LiBr = lithium bromide
• Subscripts are not included in names because they can be
assumed (MgCl2 = magnesium chloride)
• Some transition metals have variable charges (can’t
predict from group #), so amount of charge must be
represented in name with a roman numeral:
• Fe2+ = iron(II), Fe3+ = iron(III), FeCl2 = iron(II) chloride
To Write an Ionic Formula from Name
•
Write cation from first part of name and
anion from second part of name
•
Balance charges
•
Write formula with correct subscripts
Example: Aluminum sulfide
1. Al3+ and S2- are the ions
2. 2 x (3+) = 6+ and 3 x (2-) = 6- (balances
charge)
3. Al2S3 is correct formula
Covalent Bonds
•
•
•
Nonmetals form covalent bonds with each other
Covalent bonds involve sharing pairs of electrons
Example: H2 = H:H or H-H (from H. + .H) (two
hydrogen atoms form a hydrogen molecule)
• Covalent bonds generally follow octet rule (by
sharing electrons they gain a full valence shell)
• Exceptions to octet rule:
- Atoms with 3 valence electrons (B, Al, Ga…) can only
make 3 bonds, so only have 6 valence electrons
- Some atoms (P, S, Cl, Br, I) can have expanded octets
(10, 12 or 14 valence electrons)
- H can make only 1 bond, but two electrons gives it a full
outer shell (can only have 2 electrons in 1s orbital)
Predicting Number of Covalent Bonds from
Group Number
Multiple Covalent Bonds
• Nonmetals can also share 2 or 3 pairs of electrons to
form double or triple bonds
CO2 is O=C=O and HCN is H-CN
• This occurs when necessary to form full octets
• C, O, N, S and P are most common atoms to form
multiple bonds
• H can’t form them, why?
(It has only one valence electron to share)
• Also, note that H is a nonmetal (even though it’s in
Group 1) and does not form ionic compounds as H+)
Multiple Bonds in N2
• In nitrogen, octets are achieved by sharing
three pairs of electrons
• When three pairs of electrons are shared, the
multiple bond is called a triple bond
octets


N

+ 

N



N:::N

triple bond
To Write Dot-Structures from Covalent Formulas
1. Find central atom (usually forms most bonds)
2. Find total # of valence electrons (add valence
electrons from each atom in compound)
3. Use a pair of electrons to attach each atom to
the central atom
4. Subtract # of bonded pairs from total and place
remaining valence electrons as nonbonded pairs
5. If not all atoms have a complete octet, move
nonbonded pairs to complete (form multiple
bonds)
Naming Covalent Compounds
• More than one covalent compound can often be
formed from the same elements (multiple bonds)
• Prefixes in the name tell how many of that type
of atom are in the compound (subscripts in
formula)
• Prefixes are optional for single atoms
• CO = carbon monoxide (or carbon oxide)
• CO2 = carbon dioxide (or monocarbon dioxide)
• Naming of organic carbon compounds has a
whole unique set of rules (covered in chem 102)
Polyatomic Ions
• A polyatomic ion is a covalently bonded group
of atoms with an overall net charge (number of
electrons is different from number of protons)
• Most contain covalent bonds between oxygen
and another nonmetal (NO3-, HSO3-, PO43-…)
• They usually have a charge of -1, -2 or -3
• The only common positively charged one is
NH4+
Compounds Containing Polyatomic Ions
Naming Polyatomic Ions
• The ending ate is used for the most common ion for
that combination of atoms
• Related ions with one less O end in ite
PO43- = phosphate
PO33- = phosphite
• When H is added, add hydrogen to the beginning of
the name and reduce negative charge by 1
HPO42- = hydrogen phosphate
HPO32- = hydrogen phosphite
Writing Formulas with Polyatomic Ions
• All ions exist in pairs with ions of opposite charge
• Net charge of a compound or a solution is zero
• Bonds between polyatomic ions and other ions (mono- or
polyatomic) are ionic bonds
• Bonds within polyatomic ions are covalent, so the atoms
stay together as a group, even when dissolved in water
• Use same rules for writing formulas for ionic compounds
containing polyatomic ions as those for monoatomic ions
(write ions, balance charge, write formula)
• Example:
magnesium sulfate = Mg2+ + SO42- = MgSO4
Shapes of Molecules
• Molecules have 3-D shapes
• Shape is important to chemical reactivity
• Enzymes (proteins) in your cells bind to molecules
with specific shapes
• Molecular shape can be predicted using VSEPR
(valence-shell electron pair repulsion) theory
• VSEPR = electron groups (bonds or non-bonded
pairs) are placed as far apart as possible around
central atom to minimize repulsion between
negative charges
• Shape is determined by number of electron groups
How to Predict Molecular Shape
•
Write electron-dot structure
•
Determine number of electron groups and shape
2 groups = linear
3 groups = trigonal planar
4 groups = tetrahedral
•
If there are lone pairs, remove to determine final shape
3 groups (one lone pair) = bent
4 groups (one lone pair) = pyramidal
4 groups (two lone pairs) = bent
Bond Polarity
• The electrons in covalent bonds are not always
shared equally
• Some atoms attract the electrons more
• The ability of an atom to attract shared electrons
through a covalent bond is called electronegativity
• Higher electronegativity = stronger pull on
electrons
• Fluorine is the most electronegative element (4.0),
and values for all other elements are based on F
• When electrons in a covalent bond are not equally
shared, the bond is called polar covalent
Predicting Bond Type from Electronegativity Difference
• Bond type is a continuum, from ionic to nonpolar
covalent, with polar covalent in the middle
• In general:
- A metal plus a nonmetal = ionic bond
- For a nonmetal plus a nonmetal, an electronegativity
difference of:
< 0.5 = nonpolar covalent bond
≥ 0.5 = polar covalent bond
• Use the symbols + and - to indicate partial charge (
= delta) on atoms in polar covalent bonds
• + = electron poor and - = electron rich
Polar and Nonpolar Molecules
• Molecules can be overall polar or nonpolar
• Look first at the shape:
1. If symmetrical, then nonpolar
2. If non-symmetrical, then:
- If has polar bonds, then polar
- If has non-bonded electon pairs, then polar
- If no polar bonds and no lone pairs, then
nonpolar
• A polar molecule contains polar bonds
• The separation of positive and negative
charge is called a dipole
• In a polar molecule, dipoles do not cancel
+ H–Cl
dipole
••
H– N–H
H
dipoles do not cancel
• A nonpolar molecule contains nonpolar
bonds
Cl–Cl
H–H
or a symmetrical arrangement of polar
bonds
O=C=O
Cl
Cl–C–Cl
Cl
dipoles cancel