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Molecular Compounds Chapter 8. Honors Chemistry Greenville High School. Sect. 8.1 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

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molecular compounds chapter 8

Molecular CompoundsChapter 8

Honors Chemistry

Greenville High School

sect 8 1 compounds and molecules
Sect. 8.1 Compounds and Molecules
  • Compound: a substance that is made from the atoms of two or moreelements that are chemically bonded.
    • Notice: The type of bond is not important, can be ionic, covalent or metallic
      • Examples:
      • H2O, CO2, NaCl, C6H12O6
    • Non-examples:
      • I2, O2, Na, Si
compounds and molecules
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:

      • H2O, CO2, F2, H2, C6H12O6

Non-Examples:

      • NaCl, MgO, Al2O3,
3 types of chemical bonds
3 Types of Chemical Bonds
  • Ionic Bonds – a metal cation transfers valence electrons to a nonmetal anion
  • Metallic Bonds – postive cations in a sea of mobile valence electrons
  • Covalent Bonds – the bonds we will study in this chapter

Allthree types of chemical bonds are intramolecular forces :

the forces between atoms within a compound

covalent bonds
Covalent Bonds
  • Covalent Bonds – “Co-Workers”

Nonmetal + Nonmetal

  • two atoms share valence electrons to form a stable octet
      • Examples: H2O, CO2, NO2, SF6
    • Covalently bonded compounds are called

molecules

covalent bonds1
Covalent Bonds
  • Molecular Formula: shows how many atoms of each element a molecule contains.
    • Examples:
      • Diatomic Elements - O2, H2, Cl2
      • Molecules - CH4, NH3, H2O

Oxygen molecule

O2

Benzene

C6H6

molecular formulas
Molecular Formulas
  • The formula for water is written as H2O

What do the subscripts tell us?

  • Molecular formulas do not tell any information about the…..

structure!

(the arrangement of the various atoms).

covalent bonds2
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.
showing covalent bonding
Showing Covalent bonding
  • Show the bonding of F2
important covalent compounds
Important Covalent Compounds
  • 7 Diatomic Elements *Memorize*

O2

N2

F2

Cl2

Br2

I2

H2

These elements are NEVER found as individual atoms.

Ex: The oxygen gas we breathe is O2

sect 8 2 types of chemical bonds
Sect. 8.2 Types of Chemical Bonds
  • Covalent Bonds
    • 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.
sect 8 2 cont
Sect. 8.2 cont..
  • 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.
sect 8 2 cont1
Sect. 8.2 cont…
  • Electronegativity: How bad an element wants an electron
    • Using electronegativity differences to predict polarity and the bond type
      • Electronegativity Difference: (in Packet p. 13)
        • 0.0 - 0.4 = Nonpolar Covalent
        • 0.4 – 1.7 = Polar Covalent
        • > 1.7 = Ionic
slide17
Partial negative: element is partially neg.
  • Partial positive: element is partially pos.
types of chemical bonds
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

work on packet pg 1 and 2
Work on Packet pg. 1 and 2

Ex: Draw the electron dot diagram for the covalent bonds

**Remember Hydrogen needs only 2 electrons to fill the outer shell.

F2

CH4

bonds
Bonds
  • 2 valence electrons = 1 bond
  • Hydrogen can only form one single bond

WHY??

single bond
Single Bond
  • Single bond: when atoms share 1 pair of electrons (2 electrons total)

Draw lewis dot for H2O, then show bonds

tips for writing lewis dot structures for molecules with more than 2 atoms
~Tips for writing lewis dot structures for molecules with more than 2 atoms:
  • Central atom: is the 1st 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

H

O

H

Structural Formulas
  • structural formula: Showing bonds.
double bond
Double Bond

**Two atoms can share more thanone pair of valence electrons.

  • Double bond: when atoms share 2 pairs of electrons (4 electrons total)

Ex 1: Draw the lewis dot for CO2, then show structural formula

double bond cont
Double Bond cont…

Ex 2: Draw the lewis dot for H2CO, then show structural formula.

triple bond
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
How to find the # of bonds in a lewis structure
  • 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

slide28
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.) C2F4

c.) C2H6

exceptions to octet rule
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
Exceptions to the Octet Rule

1. Most covalent compounds of Beryllium: the number of valence electrons needed for Be is 4.

  • BeF2

2. Most covalent compounds of Group 13: Primarily Boron & Aluminum - the number of valence electronsneeded is 6

  • AlF3
  • BF3
exceptions to octet rule1

I – I – I

Exceptions to Octet rule

3.Sometime when Phosphorus, Sulfur, Iodine, Xenon & Krypton are the central element they can hold more than 8 electrons:

  • PCl5
  • I3
  • SF6
review on charges on bonding
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?

inter molecular forces imf
Intermolecular Forces (IMF)
  • Attractive forces betweenmolecules.
  • Much weaker than chemical bonds.
  • Intramolecular forces
  • are within a molecule. (bonds)
types of imf
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.

types of imf1

+

-

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.

types of imf2
Types of IMF
  • Hydrogen Bonding:
    • When Hydrogen bonds to Nitrogen, Oxygen or Fluorine (NOF)
    • Strongest of all intermolecular forces!
examples of intermolecular forces classify as london dipole or hbonding
Examples of intermolecular forces:Classify as London, Dipole or Hbonding.
  • NCl (nonpolar)
  • CO (polar)
  • HF (polar)
properties molecular compounds
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
Heat and electrical conductors
  • Covalent bonds: poor electrical and thermal conductivity.
    • No mobile electrons to conduct current

Review of bonds:

Covalent:

Ionic:

Metallic:

slide42

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.

slide43

VSEPR

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.
slide44

What are the ideal arrangements of electron pairs to minimize repulsions?

  • We need to identify the number of regions of high electron density, called the steric number,on the central atom.
  • Regions of high electron density include:
    • Single bonds
    • Double bonds
    • Triple bonds
    • Unshared (lone) pairs of electrons

**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
Examples: Draw the Lewis Dot Structure and fill in the following:

1. CH4

  • Steric # ____
  • # of lone pairs _____

2. H2O

  • Steric # ____
  • # of lone pairs _____

3. CO2

  • Steric # ____
  • # of lone pairs _____
examples use table to determine molecular shape and bond angle
Examples: Use table to determine molecular shape and bond angle.

1. CH4

  • Steric # 4 Molecular Shape: __________
  • # of lone pairs 0 Bond angle: _________
slide48
2. H2O
  • Steric # 4 Molecular Shape:_____________
  • # of lone pairs 2 Bond angle:________________

3. CO2

  • Steric # 2 Molecular Shape:______________
  • # of lone pairs 0 Bond angle: ______________
slide49

How does Molecular Geometry affect Polarity?

  • 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?
slide50

How does Molecular Geometry affect Polarity cont..

  • 3. One lone pair on the central atom-
  • Polar? Nonpolar?
  • Two or more lone pairs on the central atom
  • Polar? Nonpolar?

Water

(asymmetrical)

Xenon tetrafluoride (symmetrical)

Xenon difluoride (symmetrical)

slide52

Two regions of high electron density

  • AX2 notation
  • Steric # is 2
  • No lone pairs
  • Geometry is linear
  • Bond Angle is 180

Look at the example of the BeF2(g) molecule.

The Lewis Structure is:

slide53

Example: BeH2

H : Be : H

  • Steric # _____
  • # of lone pairs
  • Bond angle _________
  • Molecular Geometry __________
slide54

Example: CO2

  • 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?
slide55

Example: CO2

  • Is the molecule polar? WHY?
slide56

Example: HCN

  • Steric # _____
  • # of lone pairs
  • Bond angle _________
  • Molecular Geometry __________
  • Is the molecule polar? WHY?
slide57

Three regions of high electron density

  • AX3 notation
  • Steric # is 3
  • No lone pairs
  • Geometry is trigonal planar
  • Bond Angle is 120

Example of BF3 molecules.

The Lewis Structure is:

slide58

Example: BF3

  • Steric # _____
  • # of lone pairs ______
  • Bond angle _________
  • Molecular Geometry __________
  • Is the molecule polar? _______
slide59

AX2E noation

  • Steric # is 3
  • # of lone pairs is 1
  • Geometry is bent
  • Bond angle is 120

Example is GeF2

slide60

Steric # _____

  • # of lone pairs ______
  • Bond angle _________
  • Molecular Geometry __________

Is this molecule polar? ____

slide61

Four regions of high electron density

  • AX4 notation
  • Steric number is 4
  • No lone pairs
  • Geometry is tetrahedral
  • Bond angle is 109.5

Look at the example of CH4 molecules.

The Lewis Structure is:

slide62

Steric # _____

  • # of lone pairs ______
  • Bond angle _________
  • Molecular Geometry __________

Is the molecule POLAR? _________

slide63

AX3E notation

  • Steric # is 4
  • #of lone pairs is 1
  • Geometry is trigonal pyramidal
  • Bond angle is 107

Example NH3

The Lewis structure is:

slide64

NH3

  • Steric # _____
  • # of lone pairs _____
  • Bond angle _________
  • Molecular Geometry __________

Is the molecule POLAR? _________

slide65

AX2E2 notation

  • Steric # is 4
  • #of lone pairs is 2
  • Geometry is bent
  • Bond angle is 105

Example H2O.

The Lewis structure is:

slide66

H2O

  • Steric # _____
  • # of lone pairs _____
  • Bond angle _________
  • Molecular Geometry __________

Is the molecule POLAR? _________

slide67

FIVE regions of high electron density

  • AX5 notation
  • Steric Number 5
  • No lone pairs
  • Geometry is trigonal bipyramidal
  • Bond angle is 90/120

Example of PF5 molecules.

slide68

PF5

  • Steric # _____
  • # of lone pairs _____
  • Bond angle _________
  • Molecular Geometry __________

Is the molecule POLAR? _________

slide69

SIX regions of high electron density

  • AX6 notation
  • Steric # is 6
  • No lone pairs
  • Geometry is octahedral
  • Bond angle is 90

Example SF6 molecules.

slide70

SF6

  • Steric # _____
  • # of lone pairs ______
  • Bond angle _________
  • Molecular Geometry __________

Is the molecule POLAR? _________

word bank ch 8 packet p 10
Word Bank Ch 8 Packet p. 10

London Dispersion Molecular Formula

Dipole Dipole Formula Unit

Hydrogen Bonding Lone Pair

Octet Rule Chemical Bonds

Electronegativity Double Bond

Polar Molecule

Nonpolar Intramolecular forces

Sharing Between

Transfer Sea of electrons

Gaining Cation

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