Chapter 6

Section 1

Sharing Electrons

• Ionic bond: Electrons transfer from one atom to another to form ions

• Covalent bond: Sharing of electrons between atoms

– Example: water

– O2 + 2H2 2H2O

Molecular Orbitals

• 2 Hydrogen atoms approach each other

– The nucleus of each atom attracts its own electron and the electron of the other hydrogen atom

– The nucleus repel each other and the electron clouds repel too

– There isn’t enough attraction to take electrons away so they share

– Form a single electron cloud around by Hydrogen nuclei

Molecular Orbitals

• After they share electrons we have H2

• Both Hydrogen atoms are now stable because of the shared electrons

• Covalent bond: Forms when two or more valence electrons are shared between atoms

Molecular Orbitals

• Single Bond: – Two atoms share one pair of electrons

• Molecular orbital:– Space around the two nuclei where shared

electrons are moving

Covalent Bonds

• Nuclei bonded together are not a fixed distance from each other– They vibrate back and forth coming closer

and then stretching further apart

• Bond length:– The average distance between two bonded

atoms

Covalent Bonds

• Once the nuclei are bonded together are they permanently bonded together?– No!– If the bond energy is reached the bond can break and

the two nuclei are no longer bonded

• Bond energy:– Energy required to break a bond between two atoms

and separate them– Typically stronger bonds the bond length is short– Generally the highest bond energy comes when

atoms are bonded to H and F

Electronegativity and Bonding

• Electronegativity:– Tendency of an atom to attract bonding

electrons to itself when it bonds with another atom

– Table 6-2 Electronegativity values– Values generally increase as you go left to

right across a period– Down a group values decrease

Electronegativity and Bonding

• Nonpolar covalent bond:– Bonding electrons are shared equally– H bonded to H– Electronegativities of two atoms are equal

• Polar covalent bonds: – 2 atoms form a covalent bond but one atom attracts

electrons more strongly than the other– Attraction is not strong enough to transfer the electron

to the other atom though

Electronegativity and Bonding

• If the electronegativities are greatly different than an ionic bond forms

• Figure 6-7 page 199

Determining Ionic or Covalent

• The difference between the electronegativities is what is important– Difference greater than 2.1 the bond is ionic

– Difference between 0 and .5 the bond is nonpolar covalent

– Difference between 0.5 and 2.1 the bond is polar covalent

Examples

• Is the compound AlF3 ionic or covalent

• Is the compound AlCl3 ionic or covalent

Polar Molecules

• Polar:– Suggests that the ends have opposite

charges such as a battery or magnet

– Example: HF • F gets the – charge• H gets the + charge

• Dipole:– One end has a partial positive charge and the

other has a partial negative

Dipole Moment

• As the polarity of the molecule increases its bond strength also increases

• A larger dipole moment indicates a higher degree of polarity