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 Stability Bonds 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 electrons Metals have low electronegativity Non 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

Example of Ionic Bonding  Na Cl Result is: Na +1 Cl

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  NH 3 H 2 O

Nonpolar Covalent Bond  Atoms share the electrons equally  Electronegativity difference (ED) is 0.0  Most common in diatomics or ‘triatomics’  Br 2 I 2 N 2 Cl 2 H 2 O 2 F 2 and O 3

Coordinate Covalent Bond  Atoms share the electrons, but one atom donates both electrons  usually the bond is shared with a proton (H + )  exampleNH 3 + H +  NH 4 +

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 SolidsNetwork Solids  covalently bonded covalent in a 3-D network 1. soft1. hard, brittle 2. poor conductors2. poor conductors 3. low melting points3. high melting pointsexamples: NH 3, HCl, H 2 O, CH 4 C,SiC, and SiO 2 (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: H 2 O, NH 3, NaCl

Nonpolar Molecules  This represents balanced charge distribution  Examples:CO 2 & CH 4  All diatomics contain nonpolar molecules and nonpolar covalent bonds  Examples:O 2 & F 2

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 H 2 O

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, H 2 O, and NH 3  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

Summary of Bond Strength  Single <double<triple