4.2 - Covalent Bond

Covalent Bonding 4.2 pg 97

Understandings • A covalent bond is formed by electrostatic attraction between a pair of electrons and the positively charged nuclei • Single, double, and triple covalent bonds involve one, two and three pairs of electrons respectively. • Bond length decreases and bond strength increases as the number of shared electrons increase • Bond polarity results from the difference in electronegativities of the boned atoms

Applications and skills • Deduction of the polar nature of a covalent bond from electronegativity values

• Valence electrons – electrons in the outermost occupied energy level of an atom • Lewis structures – a symbolic representation of the arrangement of the valence electrons

Li

Be

B

C

N

O

F

Ne

• Chemical reactivity is determined by valence electrons

• Stable octet – when electrons have a completely filled outermost energy level – Atoms want to gain or lose or share electrons in order to achieve a stable octet

A covalent bond forms by electron sharing • Covalent bond – the electrostatic attraction between a pair of electrons and positively charged nuclei

– shared electrons are attracted to the nuclei of both atoms – Usually occurs between non-metals

• Molecule – a group of atoms held together by covalent bonds • Diatomic – molecule containing two atoms • Triatomic – molecule containing three atoms

Molecules formed by covalent bonds: Hydrogen and oxygen

Figure 2.5a-b

Molecules formed by covalent bonds: Methane and formaldehyde

Figure 2.5c-d

• Octet rule – atoms tend to form a stable arrangement of 8 valence electrons – Exception H and He

• Non-bonding pairs or lone pairs – electrons not involved in a bond

Atoms can share more than one pair of electrons to form multiple bonds • Single bond – 2 electrons, 1 pair • Double bond – 4 electrons, 2 pairs • Triple bond – 6 electrons, 3 pairs

• A pair of electrons is shown as a line or two dots • Each dot or x is an electron

Be careful that you are drawing a Lewis dot diagram if that is what is asked for

Bond length

• Bond length – the distance between the two bonded nuclei – Bond length decreases as number of bonds increases – Triple bond is shorter than double bond involving same type of atoms – Double bond shorter than single bond involving same type of atoms

• Bond Strength – described in terms of bond enthalpy (chapter 5) – A measure of the energy to break the bond

• A short bond is stronger – Takes more energy to break a shorter bond – While double bond is stronger than a single bond it is not twice as strong

Comparison of covalent bonds and ionic bonds Ionic Bonding

Covalent bonding

Formed between a cation and anion

Usually formed between non-metals

Formed by atom either losing or gaining electrons in order to attain a nobel gas configuration

Formed from atoms sharing electrons with each other to attain a nobel gas electron configuration

Electrostatic attraction between oppositely charged ions

Electrostatic attraction between a shared pair of electrons and the positively charged nuclei

Lattice structure

Molecules

Higher melting and boiling points

Lower melting points and boiling points

Low volatilites

May be volatile

Soluble in water

Typically insoluble in water

Conduct electricity in molten state Do not conduct in solid state

Do not conduct electricity because no ions are present to cary harge

Electronegativity Difference ∆χp • Ionic ∆χp > 1.8 • Pure covalent (non polar) ∆χp = 0 • Polar covalent 0 < ∆χp ≤ 1.8

Polar bonds result from unequal sharing of electrons

• Non-polar Covalent bonds – electrons evenly shared – Atoms have the same /almost the same electronegativity – (Have a difference in electronegativity of zero) – Eg. Cl2 or H-H

Polar Covalent Bond – Electrons are unevenly shared – Atoms have a significantly different electronegativities (less than 1.8) – Eg. HCl or H2O

dipole - refers to the fact that the bond has two separated opposite charges – More electronegative is partially negative – Use the symbol δ (delta) to represent partial charge • δ - or δ+

• Gizmo • Read 4.2