Molecular compounds chapter 8
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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

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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


Molecular compounds chapter 8

  • 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


Molecular compounds chapter 8

  • 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!


Types of intermolecular forces

Types of 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:


Draw lewis dot diagrams for polyatomic ions p 6 in packet

Draw Lewis dot diagrams for polyatomic ions: p.6 in packet

  • SO42-

    2. PO43-


Molecular compounds chapter 8

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.


Molecular compounds chapter 8

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.


Molecular compounds chapter 8

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: _________


Molecular compounds chapter 8

2. H2O

  • Steric # 4 Molecular Shape:_____________

  • # of lone pairs 2 Bond angle:________________

    3. CO2

  • Steric # 2 Molecular Shape:______________

  • # of lone pairs 0 Bond angle: ______________


Molecular compounds chapter 8

  • 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?


Molecular compounds chapter 8

  • 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)


Molecular compounds chapter 8

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:


Molecular compounds chapter 8

Example: BeH2

H : Be : H

  • Steric # _____

  • # of lone pairs

  • Bond angle _________

  • Molecular Geometry __________


Molecular compounds chapter 8

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?


Molecular compounds chapter 8

Example: CO2

  • Is the molecule polar? WHY?


Molecular compounds chapter 8

Example: HCN

  • Steric # _____

  • # of lone pairs

  • Bond angle _________

  • Molecular Geometry __________

  • Is the molecule polar? WHY?


Molecular compounds chapter 8

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:


Molecular compounds chapter 8

Example: BF3

  • Steric # _____

  • # of lone pairs ______

  • Bond angle _________

  • Molecular Geometry __________

  • Is the molecule polar? _______


Molecular compounds chapter 8

  • AX2E noation

  • Steric # is 3

  • # of lone pairs is 1

  • Geometry is bent

  • Bond angle is 120

Example is GeF2


Molecular compounds chapter 8

  • Steric # _____

  • # of lone pairs ______

  • Bond angle _________

  • Molecular Geometry __________

Is this molecule polar? ____


Molecular compounds chapter 8

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:


Molecular compounds chapter 8

  • Steric # _____

  • # of lone pairs ______

  • Bond angle _________

  • Molecular Geometry __________

Is the molecule POLAR? _________


Molecular compounds chapter 8

  • AX3E notation

  • Steric # is 4

  • #of lone pairs is 1

  • Geometry is trigonal pyramidal

  • Bond angle is 107

Example NH3

The Lewis structure is:


Molecular compounds chapter 8

  • NH3

  • Steric # _____

  • # of lone pairs _____

  • Bond angle _________

  • Molecular Geometry __________

Is the molecule POLAR? _________


Molecular compounds chapter 8

  • AX2E2 notation

  • Steric # is 4

  • #of lone pairs is 2

  • Geometry is bent

  • Bond angle is 105

Example H2O.

The Lewis structure is:


Molecular compounds chapter 8

  • H2O

  • Steric # _____

  • # of lone pairs _____

  • Bond angle _________

  • Molecular Geometry __________

Is the molecule POLAR? _________


Molecular compounds chapter 8

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.


Molecular compounds chapter 8

  • PF5

  • Steric # _____

  • # of lone pairs _____

  • Bond angle _________

  • Molecular Geometry __________

Is the molecule POLAR? _________


Molecular compounds chapter 8

SIX regions of high electron density

  • AX6 notation

  • Steric # is 6

  • No lone pairs

  • Geometry is octahedral

  • Bond angle is 90

Example SF6 molecules.


Molecular compounds chapter 8

  • 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 DispersionMolecular Formula

Dipole DipoleFormula Unit

Hydrogen Bonding Lone Pair

Octet RuleChemical Bonds

ElectronegativityDouble Bond

PolarMolecule

NonpolarIntramolecular forces

SharingBetween

TransferSea of electrons

GainingCation


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