Bonding
Energy and Chemical Bonds
Chemical Bond: A force of attraction between atoms in a compound
All elements bond for one reason: to acquire an electron configuration of a noble gas (8 valence electrons =stable octet or 2 valence for He)
BARF
Bonds and Energy changes: As bonds are formed, energy is
released. This is an exothermic process.
As bonds are broken, energy is absorbed. This is an endothermic process.
Bonds and StabilityBonds form to make
elements/compounds stable.The greater the release of
energy, the stronger the bond.
Electronegativity (review)The measure of attraction an
atom has for electronsMetals have low
electronegativityNon metals have high
electronegativity
Ionic Bonding
Results from the complete transfer of electrons between metals (that lose electrons) and nonmetals(that gain electrons) Metals form positive ions called cations (memory hint: has a “t” in the name—
looks like a plus sign for positive) Non metals form negative ions called
anions
Electronegativity Difference
If the difference between the electronegativities (higher minus lower) is 1.7 or more, the bond is ionic.
Exception: the bond is always ionic for a metal hydride: (group 1 and 2 + hydrogen is always ionic)
See examples
Covalent Bonding
Results from the sharing of electrons between two atoms
Polarity: Unequal sharing of electrons Each atom attracts electrons
by different amounts (like a tug of war)
Polar Covalent Bond
one atom has a slightly higher affinity for electrons
electronegativity difference (ED) is between 0.01 and 1.69
example of polar covalent bonds NH3 H2O
Nonpolar Covalent Bond
Atoms share the electrons equally Electronegativity difference (ED) is
0.0 Most common in diatomics or
‘triatomics’ Br2 I2 N2 Cl2 H2 O2 F2 and O3
Coordinate Covalent Bond
Atoms share the electrons, but one atom donates both electrons
usually the bond is shared with a proton (H+)
example NH3 + H+ NH4+
Metallic Bond These bonds are only found in metals Metals do not have a strong attraction
for electrons. The electrons are loosely held, so therefore, their electronegativity is very low.
These are often described as “positive ions surrounded by a sea of mobile electrons ”
the positive ions form a strong attraction for the electrons surrounding them causing strong bonds
Ionic Solids
contains ionic bonds Properties
crystalline in structure (regular geometric pattern)
relatively high melting and boiling points in the solid form, a poor conductor of
electricity in the liquid or aqueous form, a good
conductor of electricity ionic solids dissolve in water
Solids with Covalent Bonds
Molecular Solids Network Solids covalently bonded covalent in a 3-D network1. soft 1. hard, brittle2. poor conductors2. poor conductors3. low melting points 3. high melting pointsexamples: examples: NH3, HCl, H2O, CH4 C,SiC, and SiO2
(diamond, silicon carbide, silicon dioxide)
Metallic Solids
Contains metallic bonds Good conductors of heat and
electricity in due to“sea of mobile electrons”
examples: Cu, Ag, Au
Types of Molecules, Symmetry and Polarity
Polar Molecules (dipoles) This represents unbalanced charge
distribution along the bond Examples: H2O, NH3, NaCl
Nonpolar Molecules This represents balanced charge
distribution Examples:CO2 & CH4
All diatomics contain nonpolar molecules and nonpolar covalent bonds
Examples:O2 & F2
Molecular Attraction
These are forces between molecules, not to be confused with attractive forces between atoms which are bonds.
There are four types of molecular attractions:
Dipole – Dipole attraction
Neighboring polar molecules orient themselves so that oppositely charge regions line up.
Examples: HCl and H2O
Hydrogen “bonds”
This happens when hydrogen is bonded to a small highly electronegative element.
Only happens with F, O and N 3 substances are HF, H2O, and NH3
these attractive forces are SO strong they have been called bonds
hydrogen bonds are the reason that water has such a relatively high boiling point; this also gives insects the ability to “walk on water”
van der Waal’s Forces
weak intermolecular forces of attraction between individual molecules as molecular mass increases, van der
Waal’s forces increase as the distance between molecules
increases, van der Waal’s forces decrease.
The stronger the van der Waal’s forces, the higher the melting and boiling point.
Molecule-Ion Attraction
Ions are attracted to the negative and positive ends of water molecules or other polar solvents
Example: NaCl in water—sodium has a positive charge and is attracted to the negative or oxygen end of the water molecules
Multiple Covalent Bonds
Double covalent bond 2 pairs of electrons are shared Stronger than a single bond Shorter than a single bond More stable than a single bond
Triple Covalent Bond 3 pairs of electrons are shared stronger than a single or double
bond shorter than a single or double bond more stable than a single or double
bond