Molecular Compounds Bonds Part II

9.1 Key points Describe how a covalent bond forms, including the energy change involved in the process. Use the octet rule to draw Lewis electron dot structures for simple molecules. Know how and when to incorporate double and triple bonds into the structures. Understand how a coordinate covalent bond differs from other covalent bonds. Be able to draw Lewis structures for polyatomic ions. Understand the concept of resonance. Know some common exceptions to the octet rule. Relate bond energy to the stability and reactivity of molecules.

LET’S FIRST REVIEW IONIC BONDING

In an IONIC bond, electrons are lost or gained, resulting in the formation of IONS in ionic compounds. FK

FK + _ The compound potassium fluoride consists of potassium (K + ) ions and fluoride (F - ) ions The ionic bond is the attraction between the positive K + ion and the negative F - ion

Compounds and Molecules Compound: a substance that is made from the atoms of two or more elements that are chemically bonded. Notice: The type of bond is not important, can be ionic, covalent or metallic Examples: H 2 O, CO 2, NaCl, C 6 H 12 O 6 Non-examples: I 2, O 2, Na, Si

Compounds and Molecules Molecule: a neutral group of a least two atoms held together by covalent bonds –Now the type of bond is important: Only covalent bonds **Notice it only has to be two atoms** It can have two or more atoms of the same element or two or atoms of different elements Examples: H 2 O, CO 2, F 2, H 2, C 6 H 12 O 6 Non-Examples: NaCl, MgO, Al 2 O 3,

3 Types of Chemical Bonds 1.Ionic Bonds – a metal cation transfers valence electrons to a nonmetal anion 2. Metallic Bonds – postive cations in a sea of mobile valence electrons 3.Covalent Bonds – the bonds we will study in this chapter All three types of chemical bonds are intramolecular forces : the forces between atoms within a compound

Covalent Bonds Covalent Bonds – “Co-Workers” Nonmetal + Nonmetal two atoms share valence electrons to form a stable octet Examples: H 2 O, CO 2, NO 2, SF 6 –Covalently bonded compounds are called molecules

Covalent Bonds Molecular Formula: shows how many atoms of each element a molecule contains. –Examples: Diatomic Elements - O 2, H 2, Cl 2 Molecules - CH 4, NH 3, H 2 O Benzene C6H6C6H6 Oxygen molecule O2O2

Molecular Formulas The formula for water is written as H 2 O What do the subscripts tell us? Molecular formulas do not tell any information about the….. structure! (the arrangement of the various atoms).

Covalent bonds Why do nonmetals share electrons? –Remember Nonmetals Hold on to their valence electrons Cannot give away electrons to bond. Still want to form a stable octet. By sharing valence electrons both nonmetal atoms get to count the electrons toward a stable octet.

So what are covalent bonds? In covalent bonding, atoms still want to achieve a noble gas configuration (the octet rule).

In covalent bonding, atoms still want to achieve a noble gas configuration (the octet rule). But rather than losing or gaining electrons, atoms now share an electron pair.

Showing Covalent bonding Show the bonding of Cl 2

Cl 2 Chlorine forms a covalent bond with itself

Cl How will two chlorine atoms react?

Cl The octet is achieved by each atom sharing the electron pair in the middle

O2O2 Oxygen is also one of the diatomic molecules

How will two oxygen atoms bond? OO

O O = For convenience, the double bond can be shown as two dashes. O O

Important Covalent Compounds 7 Diatomic Elements *Memorize* O2O2 N2N2 F2F2 Cl 2 Br 2 I2I2 H2H2 These elements are NEVER found as individual atoms. Ex: The oxygen gas we breathe is O 2

Types of Chemical Bonds Polar Vs Nonpolar –Nonmetals do not always equally share their electrons –Some nonmetals can have a stronger pull on the shared pair of electrons—like tug of war of e - –These 2 types of covalent bonds are called polar and nonpolar.

Covalent Bonds: Polar and Nonpolar –Polar: a covalent bond in which the bonded atoms have an unequal attraction for the shared pair of electrons –Nonpolar: a covalent bond in which the two bonding electrons are shared equally by the bonded atoms.

Electronegativity: How bad an element wants an electron –Using electronegativity differences to predict polarity and the bond type Electronegativity Difference: (in Packet p. 13) – = Nonpolar Covalent – 0.4 – 1.7 = Polar Covalent – > 1.7 = Ionic

* = Nonpolar * 0.4 – 1.7 = Polar * > 1.7 = Ionic Electronegativity

Partial negative: element is partially neg. Partial positive: element is partially pos.

Types of Chemical Bonds Examples: Determine the electronegativity difference, the bond type and indicate partial positive and partial negative charges. a.) H and I H= ___ I=___, Δ = ____ Bond type=_______________ H - I b.) K and Br K=____ Br=_____, Δ = _____ Bond type=_______________ K - Br

Ex: Draw the electron dot diagram for the covalent bonds **Remember Hydrogen needs only 2 electrons to fill the outer shell. F 2 CH 4

Bonds 2 valence electrons = 1 bond Hydrogen can only form one single bond WHY??

Single Bond Single bond: when atoms share 1 pair of electrons (2 electrons total) Draw lewis dot for H 2 O, then show bonds

~Tips for writing lewis dot structures for molecules with more than 2 atoms: Central atom: is the 1 st element in the compound or molecule (except H) 1. **The central atom ALWAYS goes in the middle!!! *** 2. Rearrange dots so that every element has 8 valence electrons (H and He only need 2 val)

Structural Formulas structural formula: Showing bonds. HHO

Double Bond **Two atoms can share more than one pair of valence electrons. Double bond: when atoms share 2 pairs of electrons (4 electrons total) Ex 1: Draw the lewis dot for CO 2, then show structural formula

Double Bond cont… Ex 2: Draw the lewis dot for H 2 CO, then show structural formula.

Triple Bond ~ Triple bond : when atoms share 3 pairs of electrons (6 electrons total) Draw the lewis dot for HCN and show structural formula.

How to find the # of bonds in a lewis structure 1.Find the total # of valence electrons. 2. Use the formula to find the number of bonds. # of val e - needed (all have 8 or 2 e - ) - # of val e - available = ____/2 to find the # of bonds

1.Find the total # of valence electrons. 2. Use the formula to find the number of bonds. # of val e - needed (all have 8 or 2 e - ) - # of val e - available = ____/2 to find the # of bonds Ex: Find the number of bonds for each molecule or compound and write the lewis dot and structural formula: a.) CO b.) C 2 F 4 c.) C 2 H 6

Exceptions to Octet rule For some molecules, it is impossible to satisfy the octet rule Yet the stable molecules do exist Two types of exceptions: –Atoms that cannot hold 8 valence electrons Hydrogen, helium, beryllium, boron, aluminum –Atoms that can hold more than 8 valence electrons Phosphorus, sulfur, iodine, xenon, krypton

Exceptions to the Octet Rule 1. Most covalent compounds of Beryllium: the number of valence electrons needed for Be is Most covalent compounds of Group 13: Primarily Boron & Aluminum - the number of valence electrons needed is 6 AlF 3 BF 3 BeF 2

3. Sometime when Phosphorus, Sulfur, Iodine, Xenon & Krypton are the central element they can hold more than 8 electrons: Exceptions to Octet rule PCl 5 SF 6 I 3 I – I – I

Review on charges on bonding: Ionic Bonds: –Have a full positive or full negative charge. –Ionic bonds do NOT have partial charges. Why? Polar Covalent Bonds: –Have partial positive or partial negative charges. Why? Nonpolar Covalent Bonds: – Have NO partial positive or partial neg. charge. Why?

Intermolecular Forces (IMF) Attractive forces between molecules. Much weaker than chemical bonds. Intra molecular forces are within a molecule. (bonds)

Types of IMF London Dispersion Forces: –Occurs between nonpolar molecules (diatomics) –Caused by motion of electrons ( “e - sloshing” ), they create a temporary dipole (slight charge) – Weakest of all forces. View animation online.animation

Types of IMF Dipole-Dipole Forces: –Occurs between polar molecules –Where one side is partial positive and one is partial negative. –Stronger than London Dispersion forces. + + - - View animation online.animation

Types of IMF Hydrogen Bonding: –When Hydrogen bonds to Nitrogen, Oxygen or Fluorine (NOF) – Strongest of all intermolecular forces!

Types of intermolecular forces:

Examples of intermolecular forces: Classify as London, Dipole or Hbonding. NCl (nonpolar) CO (polar) HF (polar)

Properties Molecular Compounds Low melting points and boiling points. –The IMF between molecular compounds are weaker than ionic or metallic compounds –This means that only a small amount of energy is required break the bonds Strongest Bonds  Weakest Bonds

Heat and electrical conductors Covalent bonds: poor electrical and thermal conductivity. –No mobile electrons to conduct current Review of bonds: Covalent: Ionic: Metallic:

Draw Lewis dot diagrams for polyatomic ions: p.6 in packet 1.SO PO 4 3-

Molecular Geometry Lewis structures fail to indicate three- dimensional shapes of molecules. The shape of a molecule controls some of its chemical and physical properties.

Valence Shell Electron Pair Repulsion Theory - predicts the shapes of a number of molecules and polyatomic ions. Electron pairs move to create the most stable arrangement. -The repulsions between electron pairs causes molecular shapes to adjust so that the electron pairs stay as FAR APART as possible. VSEPR

1)We need to identify the number of regions of high electron density, called the steric number, on the central atom. 2)Regions of high electron density include: Single bonds Double bonds Triple bonds Unshared (lone) pairs of electrons What are the ideal arrangements of electron pairs to minimize repulsions? **Double and triple** bonds only count as ONE region of high electron density just like a single bond or a lone pair.

Examples: Draw the Lewis Dot Structure and fill in the following: 1. CH 4 –Steric # ____ –# of lone pairs _____ 2. H 2 O –Steric # ____ –# of lone pairs _____ 3. CO 2 –Steric # ____ –# of lone pairs _____

Examples: Use table to determine molecular shape and bond angle. 1. CH 4 –Steric # 4 Molecular Shape: __________ –# of lone pairs 0 Bond angle: _________

2. H 2 O –Steric # 4 Molecular Shape:_____________ –# of lone pairs 2 Bond angle:________________ 3. CO 2 –Steric # 2 Molecular Shape:______________ –# of lone pairs 0 Bond angle: ______________

How does Molecular Geometry affect Polarity? 1.One polar bond on central atom Molecule polar? Molecule nonpolar? 2. More than one polar bond on the central atom will cancel out polarities if they have equal electronegativities. Molecule polar? Molecule nonpolar?

How does Molecular Geometry affect Polarity cont.. 3. One lone pair on the central atom- Polar? Nonpolar? 4.Two or more lone pairs on the central atom Polar? Nonpolar? Water (asymmetrical) Xenon difluoride (symmetrical) Xenon tetrafluoride (symmetrical)

Two regions of high electron density AX 2 notation Steric # is 2 No lone pairs Geometry is linear Bond Angle is 180  Look at the example of the BeF 2(g) molecule. The Lewis Structure is:

Steric # _____ # of lone pairs Bond angle _________ Molecular Geometry __________ H : Be : H Example: BeH 2

Steric # _____ # of lone pairs ____ Bond angle _________ Molecular Geometry __________ Is the molecule polar? Electronegativity Difference between Carbon & Oxygen is.89 So the bonds are polar But is the molecule? Example: CO 2

Is the molecule polar? WHY? Example: CO 2

Steric # _____ # of lone pairs Bond angle _________ Molecular Geometry __________ Is the molecule polar? WHY? Example: HCN

Three regions of high electron density Example of BF 3 molecules. The Lewis Structure is: AX 3 notation Steric # is 3 No lone pairs Geometry is trigonal planar Bond Angle is 120 

Steric # _____ # of lone pairs ______ Bond angle _________ Molecular Geometry __________ Is the molecule polar? _______ Example: BF 3

Example is GeF 2 AX 2 E noation Steric # is 3 # of lone pairs is 1 Geometry is bent Bond angle is 120 

Is this molecule polar? ____ Steric # _____ # of lone pairs ______ Bond angle _________ Molecular Geometry __________

Four regions of high electron density AX 4 notation Steric number is 4 No lone pairs Geometry is tetrahedral Bond angle is  Look at the example of CH 4 molecules. The Lewis Structure is:

Is the molecule POLAR? _________ Steric # _____ # of lone pairs ______ Bond angle _________ Molecular Geometry __________

Example NH 3 The Lewis structure is: AX 3 E notation Steric # is 4 #of lone pairs is 1 Geometry is trigonal pyramidal Bond angle is 107 

Is the molecule POLAR? _________ NH 3 Steric # _____ # of lone pairs _____ Bond angle _________ Molecular Geometry __________

AX 2 E 2 notation Steric # is 4 #of lone pairs is 2 Geometry is bent Bond angle is 105  Example H 2 O. The Lewis structure is:

Is the molecule POLAR? _________ H 2 O Steric # _____ # of lone pairs _____ Bond angle _________ Molecular Geometry __________

FIVE regions of high electron density Example of PF 5 molecules. AX 5 notation Steric Number 5 No lone pairs Geometry is trigonal bipyramidal Bond angle is 90  /120 

Is the molecule POLAR? _________ PF 5 Steric # _____ # of lone pairs _____ Bond angle _________ Molecular Geometry __________

SIX regions of high electron density Example SF 6 molecules. AX 6 notation Steric # is 6 No lone pairs Geometry is octahedral Bond angle is 90 

Is the molecule POLAR? _________ SF 6 Steric # _____ # of lone pairs ______ Bond angle _________ Molecular Geometry __________

London DispersionMolecular Formula Dipole DipoleFormula Unit Hydrogen Bonding Lone Pair Octet RuleChemical Bonds ElectronegativityDouble Bond PolarMolecule NonpolarIntramolecular forces SharingBetween TransferSea of electrons GainingCation