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 thepositively charged nuclei

• Single, double, and triple covalent bonds involveone, two and three pairs of electronsrespectively.

• Bond length decreases and bond strengthincreases as the number of shared electronsincrease

• Bond polarity results from the difference inelectronegativities of the boned atoms

• A covalent bond is formed by electrostaticattraction between a pair of electrons and thepositively charged nuclei

• Single, double, and triple covalent bonds involveone, two and three pairs of electronsrespectively.

• Bond length decreases and bond strengthincreases as the number of shared electronsincrease

• Bond polarity results from the difference inelectronegativities of the boned atoms

Applications and skills

• Deduction of the polar nature of a covalentbond from electronegativity values

• Valence electrons – electrons in the outermostoccupied energy level of an atom

• Lewis structures – a symbolic representation ofthe arrangement of the valence electrons

Li Be B C

N O F Ne

• Chemical reactivity is determined by valenceelectrons

• Valence electrons – electrons in the outermostoccupied energy level of an atom

• Lewis structures – a symbolic representation ofthe arrangement of the valence electrons

Li Be B C

N O F Ne

• Chemical reactivity is determined by valenceelectrons

• Stable octet – when electrons have acompletely filled outermost energy level

– Atoms want to gain or lose or share electrons inorder to achieve a stable octet

• Stable octet – when electrons have acompletely filled outermost energy level

– Atoms want to gain or lose or share electrons inorder to achieve a stable octet

A covalent bond forms by electronsharing

• Covalent bond – the electrostatic attractionbetween a pair of electrons and positivelycharged nuclei– shared electrons are attracted to the nuclei of both

atoms– Usually occurs between non-metals

• Molecule – a group of atoms held together bycovalent bonds

• Diatomic – molecule containing two atoms• Triatomic – molecule containing three atoms

• Covalent bond – the electrostatic attractionbetween a pair of electrons and positivelycharged nuclei– shared electrons are attracted to the nuclei of both

atoms– Usually occurs between non-metals

• Molecule – a group of atoms held together bycovalent 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 andformaldehyde

Figure 2.5c-d

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

• Non-bonding pairs or lone pairs – electronsnot involved in a bond

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

• Non-bonding pairs or lone pairs – electronsnot involved in a bond

Atoms can share more than one pair ofelectrons to form multiple bonds

• Single bond – 2 electrons, 1 pair

• Double bond – 4 electrons, 2 pairs

• Triple bond – 6 electrons, 3 pairs

• 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 dotdiagram if that is what is asked for

Bond length

• Bond length – the distance between the twobonded 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 length – the distance between the twobonded 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 bondenthalpy (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

• Bond Strength – described in terms of bondenthalpy (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 andionic bonds

Ionic Bonding Covalent bonding

Formed between a cation and anion Usually formed between non-metals

Formed by atom either losing or gainingelectrons in order to attain a nobel gasconfiguration

Formed from atoms sharing electronswith each other to attain a nobel gaselectron configuration

Electrostatic attraction betweenoppositely charged ions

Electrostatic attraction between a sharedpair of electrons and the positivelycharged nuclei

Electrostatic attraction betweenoppositely charged ions

Electrostatic attraction between a sharedpair of electrons and the positivelycharged 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 stateDo not conduct in solid state

Do not conduct electricity because noions are present to cary harge

Electronegativity Difference ∆χp

• Ionic ∆χp > 1.8

• Pure covalent (non polar) ∆χp = 0

• Polar covalent 0 < ∆χp ≤ 1.8

• Ionic ∆χp > 1.8

• Pure covalent (non polar) ∆χp = 0

• Polar covalent 0 < ∆χp ≤ 1.8

Polar bonds result from unequalsharing 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

• 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 twoseparated opposite charges

– More electronegative is partially negative– Use the symbol δ (delta) to represent partial charge

• δ - or δ+

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 twoseparated opposite charges

– More electronegative is partially negative– Use the symbol δ (delta) to represent partial charge

• δ - or δ+

• Gizmo• Read 4.2