Chapter 8Covalent Bonding General Chemistry I T.ARA
Chemical Bonding • Now that we know something about electron configurations, we can take a closer look at the ways atoms form bonds. • There are two main types of chemical bonds: • Ionic Bonds • Covalent Bonds
A. Ionic Bonding • Ionic bonding involves the transfer of valence electrons from one atom (usually a metal) to another atom (a nonmetal) such that each atom gains a noble gas configuration. • The ionic bond is the electrostatic attraction between the cation (+) and the anion (-) that result from the electron transfer. • The two bonded atoms do not share electrons.
A. Ionic Bonding: Sodium Chloride • A sodium atom will readily lose a valence electron: Na [Ne]3s1 Na+ [Ne] + e- • A chlorine atom will readily acceptan electron: Cl [Ne]3s23p5 + e- Cl- [Ar] • The sodium cation & the chloride anion are held together by an electrostatic (coulombic) attraction– opposite charges attract. Na+ + Cl- NaCl (ionic bond)
A. Ionic Bonding: Sodium Chloride Using Lewis symbols:
1. Polyatomic Ions • Polyatomic Ion:a group of covalently bonded atoms with an overall positive or negative charge • The atoms within a polyatomic ion are covalently bonded together, but polyatomic ions form ionic bonds with other ions. eg. NO3- is a polyatomic anion.
B. Covalent Bonding • In a covalent bondthe electrons are the “glue” that holds the atoms together. • A covalent bond is formed when two atoms share one or more pairs of electrons. • Each atom will form enough covalent bonds to achieve a noble gas configuration. G.N. Lewis
B. Covalent Bonding: Hydrogen H• + H• H : H • The two electrons are shared evenly between the two hydrogen atoms. • It is as if each atom has two electrons – the noble gas configuration of [He]. H : H
B. Covalent Bonding: Hydrogen • The actual electron probability cloud(2) for the two electrons in the H-H bond looks like this.
1. Lewis Structures • A Lewis structure is a way of drawing a molecule that shows all valence electrons as dots or lines that represent covalent bonds. eg. The Lewis structure for H2 can be drawn in two ways: H : H or H–H a) A single line represents two covalently shared electrons – also known as a single bond.
1. Lewis Structures • The Lewis Structure for F2: • Two electrons are shared between the two F atoms (one single covalent bond). • Each F atom also has three unshared electron pairs. These non-bonding electron pairs are called lone pairs.
2. The Octet Rule • Note that by sharing electrons, it is as if each F atom has eight electrons - the noble gas configuration of [Ne]. • The Octet Rule:Main group elements with more than two valence electrons gain, lose, or share electrons to achieve a noble gas configuration characterized by eight valence electrons.
2. The Octet Rule • As the previous examples illustrated, the number of covalent bonds an atom must form to achieve an octet is equal to eight minus it group number.
a) Multiple Covalent Bonds • It is not always possible for atoms to gain a full octet by sharing single electron pairs with other atoms. • In other words, it is not always possible to construct a valid Lewis structure using only single bonds. eg. N2 In this structure, each nitrogen atom would have only six valence electrons – two short of an octet.
a) Multiple Covalent Bonds • Two atoms can share more than one electron pair to gain a full octet. • Double Bond: when 2 electron pairs (4 electrons) are shared between 2 atoms • Triple Bond:when 3 electron pairs (6 electrons) are shared between two atoms • Double & triple bonds are referred to as multiple covalent bonds.
a) Multiple Covalent Bonds Ethane (C2H6), Ethylene (C2H4) & Acetylene (C2H2)
a) Multiple Covalent Bonds 4 total bonds 3 total bonds 2 total bonds 1 bond When necessary, atoms will form any combination of single, double and triple covalent bonds to gain a full octet.
Draw a valid Lewis structure for the cyanide ion (CN-). Lewis Structures of Ions: Add one additional valence electron for every negative charge & subtract one valence electron for every positive charge.
b) Formal Charge • As you just saw, even when the octet rule is obeyed, some compounds have an overall charge. • This means that the compound contains one or more charged atoms. Formal charge: the charge a bonding atom would have if its bonding electrons were shared equally Formal Charge = Atomic Group # – # lone pair electrons – ½ (# bonding electrons)
Draw all possible Lewis structures for N2O (O-N-N) & assign formal charges.
Draw all possible Lewis structures for N2O (O-N-N) & assign formal charges.
3. Resonance Structures • When more than one Lewis structure can be drawn for a molecule, the structures are called resonance structures. • Each resonance structure contributes to the overall structure of the molecule. • Individual resonance structures DO NOT ACTUALLY EXIST– we use resonance structures conceptually to help us understand molecular structures!!! • The actual structure is a weighted average of the resonance structures – called a resonance hybrid.
3. Resonance Structures eg. Resonance Structures of Ozone • Each bond is in between a single & a double bond. • Each terminal O has a partial negative charge.
3. Resonance Structures • Resonance structures are connected by double-headed arrows ( ). • Each resonance structure must have the same overall charge as the molecule. • Lower energy resonance structures contribute more to the overall structure of the molecule. • Resonance structures are lowest in energy when: • All atoms with full octets • The minimum # of formal charges • Negative charges on electronegative atoms (more on this in a minute)
3. Resonance Structures Which N2O resonance structure(s) are lowest in energy? A B C B & C both have only two formal charges – lower in energy than A. High in Energy: Too many formal charges!! C is lowest in energy because the negative charge is on O (more electronegative than N) – more on this in a minute!
4. Bond Lengths & Bond Strengths • How does the type of bond (bond order) affect the properties of the bond? • Bonds get shorter as the number of electrons shared between the two atoms increases. (Bonds get shorter as the bond order increases.) Bond Length: Single Bond (1) > Double Bond (2) > Triple Bond (3)
4. Bond Lengths & Bond Strengths • The bond order also affects the strength of the bond (the bond enthalpy). • As the bond order increases, the bond gets stronger (harder to break).
C. Ionic vs. Covalent: Bond Polarity • Now we know the difference between ionic and covalent bonds. • Given a specific compound, how do you know which type of bond to expect?
1. The Continuum • In reality, all bonds have some covalent character & some ionic character. • In other words, all bonds fit somewhere on a continuum between covalent & ionic. • The real questions are: • How evenly are the electrons being shared between the two atoms? • How can you predict how evenly the electrons will be shared?
1. The Continuum Covalent Polar Ionic Covalent - electrons shared - electrons shared - electrons not equally unequally shared at all - No separation - some separation - complete separation of charge of charge of charge - held together by - held together by - held together by shared electrons shared electrons electrostatic attraction
2. Electronegativity • In 1932, Linus Pauling proposed the idea of electronegativity to explain the ways atoms share electrons in bonds. Electronegativity: the ability of an atom in a covalent bond to attract shared electrons to itself; the “electron-pulling” power of an atom
2. Electronegativity The higher the EN, the more tightly an atom holds its electrons.
a) Bond Polarity • The polarity of a bond describes how evenly the electrons are shared between the two atoms. • The more polar a bond, the less evenly the electrons are shared. • A polar bond is indicated by using + and - to represent the partial chargeson the atoms. + = less electronegative atom - = more electronegative atom
a) Bond Polarity • The greater the difference in electronegativity between the two atoms, the more polar the bond between them will be.
a) Bond Polarity • “Approximate” Guidelines: • If EN ≤ 0.5, the electrons are shared fairly evenly & the bond is nonpolar covalent. • If EN 1.9, the electrons are localized on the more electronegative atom & the bond is ionic. • If 1.9 > EN > 0.5, the electrons are shared unequally & the bond is polar covalent.
The Continuum H2 LiF HCl MgO CsI N2 Covalent Polar Covalent Ionic EN 0 0.5 1.0 1.5 2.0 2.5 3.0
D. Exceptions to the Octet Rule • Most structures containing main group elements follow the octet rule, but there are exceptions in certain compounds: • Some atoms have fewer than eight electrons (less than an octet). • Some atoms have more than eight electrons (more than an octet).
1. Less than an Octet • Elements in group 3A have three valence electrons. • Using those three electrons to form three bonds gives the central atom only six valence electrons (less than an octet). • eg. BH3
1. Less than an Octet • Because boron has less than an octet, BH3 is very reactive – it will react to form an octet. • The nitrogen in ammonia uses its lone pair electrons to form a bond with boron – both B and N now have an octet.
2. More than an Octet – Expanded Valence • Elements in the third period & lower can form stable compounds in which the central atom has more than 8 electrons – an expanded valence. • Unlike elements in the first two periods (like N & O), third row elements (like P & S) can use their empty 3d orbitals to accommodate extra electrons.
2. More than an Octet – Expanded Valence • eg. PCl5, SBr6, ClF3