Chemical Bonds
Mutual attraction that binds atoms
together to form compounds
Types of Chemical Bonds
• Ionic bond = bond that results from attraction
between oppositely charged ions (transfer of
electrons)
• Covalent bond = bond resulting from the
sharing of electrons
Metallic Bonds
= metal atoms in a delocalized cloud of electrons
• Remember: metals like to give up electrons,
so no atom “wants” the free e-
Bonding happens on a spectrum
• Bonds are rarely purely ionic or purely
covalent
• Remember Electronegativity (the ability of
atoms to attract electrons)
– Comparing the electronegativities of atoms
involved in bond, can determine whether bond is
ionic or covalent
see page 161 PT for electronegativity values
Electronegativity difference
above 1.7 = ionic
• Ex. Cs and F
• Cs = 0.7
• F = 3.3
Electronegativity difference of
1.7 or less = covalent
• Polar-covalent (0.3-1.7),
– ex. H2O: H-2.1, O-3.5, diff = 1.4
• Non-polar covalent (under 0.3)
– Ex. Bonds between atoms of the same element
are always non-polar (purely) covalent, ex. H2, O2,
N2
Covalent Bonding
Forms molecular compounds that
consist of molecules
Molecular Compounds
• Molecule = a neutral group of atoms that are
held together by covalent bonds
• Chemical formula (molecular formula) =
indicates the number of atoms of each element in
a compound
–
–
–
–
–
H2O
O2
CO2
HCl
C6H12O6
Diatomic molecule = made
up of 2 atoms of the same
element
Diatomic Molecules
(Hydrogen and the magic 7)
• Some elements always exist as diatomic
molecules
Empirical Formula
• The “reduced” form of a compound
• Has the lowest ratio
• Molecular formula = C6H12O6
• Empirical formula = CH2O
Potential Energy and Bonding
• Nature favors covalent bonding
– Atoms have lower PE when they are bonded
– As atoms near each other, their charged particles
interact
• Nucleus is attracted to e• e- repel each other
• 2 nuclei repel each other
+
-
-
+ +
PE↓
PE↑
PE↑
Characteristics of Covalent Bonds
• Bond length = avg. distance btw 2 bonded
How are they
atoms
related?
• Bond energy = E required to break the bond
Covalent Bond = overlap of
orbitals where shared
electrons live
Octet Rule
• Chemical compounds tend to form so that
each atom, by gaining, losing, or sharing
electrons, has an octet (8) electrons in its
highest occupied energy level (valence)
• F: ↑↓ ↑↓ ↑↓ ↑↓ ↑__
1s 2s
2p
• F: ↑↓ ↑↓ ↑↓ ↑↓ __↓
Bonding
electron pair in
overlapping
orbitals
Example
1s
H: ↑__
Bonding electron pair in
overlapping orbitals
1s 2s
2p
3s
3p
Cl: ↑↓ ↑↓ ↑↓ ↑↓ ↑↓ ↑↓ ↑↓ ↑↓ __↓
Exceptions to the Octet Rule
• Hydrogen (1 valence) and Helium (2 valence)
experience stability with only 2 electrons in the
1s-orbital (it is their highest and ONLY energy
level)
• Boron (3 valence) tends to form stable
compounds with 6 valence e– Ex. BF3
• Other elements can be surrounded by more than
8 e- when they bond with highly electronegative
elements (like halogens)
– This bonding will involve d orbitals In addition to s and
p
– Ex. PF5 and SF6
Electron Dot Notation
# valence e-
electron-dot notation
example
1
X
Na
2
X
Mg
3
X
B
4
X
C
5
X
N
6
X
O
7
X
F
8
X
Ne
Lewis Structures = formulas in which atomic
symbols represent nuclei and inner shell electrons; only
valence electrons are drawn
• Use electron-dot notation to represent
molecules
– Examples:
Structural Formulas indicate
the kind, number, arrangement,
and bonds, but not the
unshared pairs of electrons
Types of Covalent Bonds
• Single bonds = result from the sharing of 1
electron pair
• Double bonds = result from the sharing of 2
electron pairs
• Triple bonds = result from the sharing of 3
electron pairs
Bond Length and Bond Energy
• REMEMBER: bond energy is energy required
to break the bond
• Shorter bonds have higher energies and are
harder to break
Shorter length
single bonds
double bonds
triple bonds
Higher energy
Molecular Geometry
= the 3-dimensional arrangement of a
molecule’s atoms in space
Geometry and bond polarity will determine
molecular polarity and intermolecular forces.
Forces of attraction between molecules
Intramolecular Forces
• Forces between atoms or ions that keep
compounds together
– Ionic attraction
– Sharing electrons
Intermolecular Forces
1. Dipole-Dipole forces = attractive forces
between polar molecules
Hydrogen-bonds (H-bonds) = type of D-D force
btw H and certain other elements
H-N, H-O, or H-F
Intermolecular Forces (cont)
2. Ion-Dipole forces = attraction btw an ion and
a polar molecule
Intermolecular Forces (cont)
3. London Dispersion forces = attractive forces
resulting from temporary dipoles induced by
ions
Naming Covalent Compounds
• Use the prefixes based on the number of each
type of atom and suffix –ide for 2nd element
1 = mono2 = di3 =tri4 = tetra5 = penta6 = hexa7 = hepta8 = octa9 = nona10 = deca
CO = carbon monoxide
CO2 = carbon dioxide
BF3 = boron trifluoride
CCl4 = carbon tetrachloride
N2O5 = dinitrogen pentoxide
B3F6 = triboron hexafluoride
Only used in naming organic
compounds
Naming Covalent Compounds
• Exceptions: covalent compounds containing
hydrogen, e.g.
– H2S (hydrogen sulfide)
– CH4 (methane)
– NH3 (ammonia)
Organic Compounds
• Compounds containing carbon
– Hydrocarbons = carbon chain + hydrogen
Propane
– Carbohydrates = carbon + hydrogen + oxygen
Glucose
Naming Hydrocarbons
Consist of chains of
carbon atoms
Name = prefix + -ane
# Carbon atoms
prefix
1
Meth-
2
Eth-
3
Prop-
4
But-
5
Pent-
6
Hex-
7
Hept-
8
Oct-
9
Non-
10
Dec-
Drawing Hydrocarbons
Ionic Compounds
= Compounds consisting of a positive ion and a
negative ion held together by the attraction of
opposite electrical charges
The overall charge of an ionic compound = 0
The + and - ions cancel each other
Formula Units
• Molecules are the basic unit of covalent
compounds (a molecule of water, H2O)
• The smallest unit of an ionic compound is the
formula unit (a formula unit of sodium
chloride, NaCl)
Crystal Lattice
= an orderly arrangement of ions
Attractive forces
btw opposite ions
btw nuclei and e- of adjacent ions
Repulsive forces
btw like ions
btw e- and ebtw nucleus and nucleus
Lattice Energy and Bond Strength
To compare bond strength in ionic compounds,
we compare lattice energy
Related to the energy in the bond =
how tightly the ions in a crystal are
held together
Remember: higher energy bonds are
harder to break
Molecular (covalent) vs. Ionic
• Covalent and ionic bonds are strong
attractions that hold the atoms in a compound
together
Molecules in Covalent
Compounds
Held together by intermolecular
forces
Formula Units in Ionic
Compounds
Held together by attractive forces in
a crystalline lattice
Molecular v. Ionic Properties
Molecular
• Lower melting points
(liquids and gases)
Ionic
• Higher melting points
(solids)
• Hard and brittle
• Conduct electricity as
liquids or when
dissolved in water
• Many are soluble in
water
• Some aren’t. WHY?
Types of Ions
• Monatomic = ions that
consist of one element
Cl- = chloride
Br- = bromide
Na+
Mg2+
Al3+
• Polyatomic = ions that
consist of more than
one type of atom
(CO3)2- = carbonate
(PO4)3- = phosphate
(C2H3O2)1- = acetate
(ClO3)1- chlorate
(NO3)1- = nitrate
Naming Binary (2 elements)
Ionic Compounds
Cation (metal) comes 1st
– Element name
Anion (nonmetal) comes 2nd
– Element root + -ide
Naming Ionic Compounds that Contain
Polyatomic Ions
• Cation (metal)
– Element name
• Polyatomic ion
– Ion name
Stock System for Transition Metals
• Some transition metals (Group 3-12) can form
more than one type of cation
– Ex. Copper can form +1, +2, or +3
• When writing the name, use a Roman numeral
to indicate the charge of the transition metal
ion
– Ex. Copper (I) chloride, copper (II) chloride, etc.
Writing Formulas for Ionic Compounds
1. Symbol for cation + symbol/formula for
cation
2. Write the charges as super scripts
3. Cross the numbers down to subscripts
Naming Acids
• Binary Acids = consist of 2 elements
Hydrogen + a halogen
• HF = hydrofluoric acid
• HCl = hydrochloric acid
• HBr = hydrobromic acid
• HI = hydroiodic acid
Naming Acids
• Oxyacids = consist of H, O, and another
nonmetal
H2SO4 = sulfuric acid
H2SO3 = sulfurous acid
H2CO3 = carbonic acid
HNO3 = nitric acid
HNO2 = nitrous acid
HC2H3O2 = acetic acid
HClO4 = perchlorous acid
HClO3 = chloric acid
HClO2 = chlorous acid
HClO = hypochlorous acid
Assigning Oxidation Numbers
• The oxidation # for any uncombined element or
diatomic molecule is ZERO Ex. Zn, Cu, O2, N2, Cl2
• The oxidation # for any monatomic ion is its
CHARGE
Ex. Ba+2, K+1
• H is usually +1, unless it’s combined with a metal
to form a metal hydride then it’s -1 HCl (+1), NaH (-1)
• O is usually -2, unless it’s a peroxide, then it’s -1
MgO (-2), H2O2 (-1)
Assigning Oxidation Numbers
• In binary covalent compounds (nonmetal +
nonmetal) the positive one is first and the
negative one is second
• The sum of the oxidation numbers for all
atoms in a neutral compound is ZERO
• The sum of the oxidation numbers in a
polyatomic ion is equal to the CHARGE of the
polyatomic ion
Practice
• H2
• CaCl2
• KClO4