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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


Molecular compounds chapter 8


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 more than 2 atoms:

O

H

Structural Formulas

  • structural formula: Showing bonds.


Double bond
Double Bond more than 2 atoms:

**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… more than 2 atoms:

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


Triple bond
Triple Bond more than 2 atoms:

~ 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 more than 2 atoms:

  • 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. more than 2 atoms:

    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 more than 2 atoms:

  • 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 more than 2 atoms:

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 more than 2 atoms:

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: more than 2 atoms:

  • 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
Inter more than 2 atoms:molecular Forces (IMF)

  • Attractive forces betweenmolecules.

  • Much weaker than chemical bonds.

  • Intramolecular forces

  • are within a molecule. (bonds)


Types of imf
Types of IMF more than 2 atoms:

  • 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

more than 2 atoms:+

-

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 more than 2 atoms:

  • 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: more than 2 atoms:Classify as London, Dipole or Hbonding.

  • NCl (nonpolar)

  • CO (polar)

  • HF (polar)


Properties molecular compounds
Properties more than 2 atoms: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 more than 2 atoms:

  • Covalent bonds: poor electrical and thermal conductivity.

    • No mobile electrons to conduct current

      Review of bonds:

      Covalent:

      Ionic:

      Metallic:



Molecular compounds chapter 8

Molecular Geometry more than 2 atoms:

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 more than 2 atoms:

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 more than 2 atoms: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: more than 2 atoms: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. H angle.2O

  • 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? angle.

  • 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

Water

(asymmetrical)

Xenon tetrafluoride (symmetrical)

Xenon difluoride (symmetrical)


Molecular compounds chapter 8

Two regions of high electron density angle.

  • 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: BeH angle.2

H : Be : H

  • Steric # _____

  • # of lone pairs

  • Bond angle _________

  • Molecular Geometry __________


Molecular compounds chapter 8

Example: CO angle.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?


Molecular compounds chapter 8

Example: CO angle.2

  • Is the molecule polar? WHY?


Molecular compounds chapter 8

Example: HCN angle.

  • Steric # _____

  • # of lone pairs

  • Bond angle _________

  • Molecular Geometry __________

  • Is the molecule polar? WHY?


Molecular compounds chapter 8

Three regions of high electron density angle.

  • 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: BF angle.3

  • Steric # _____

  • # of lone pairs ______

  • Bond angle _________

  • Molecular Geometry __________

  • Is the molecule polar? _______


Molecular compounds chapter 8

  • AX angle.2E noation

  • Steric # is 3

  • # of lone pairs is 1

  • Geometry is bent

  • Bond angle is 120

Example is GeF2


Molecular compounds chapter 8

  • Steric # _____ angle.

  • # of lone pairs ______

  • Bond angle _________

  • Molecular Geometry __________

Is this molecule polar? ____


Molecular compounds chapter 8

Four regions of high electron density angle.

  • 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 # _____ angle.

  • # of lone pairs ______

  • Bond angle _________

  • Molecular Geometry __________

Is the molecule POLAR? _________


Molecular compounds chapter 8

  • AX angle.3E 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

  • NH angle.3

  • Steric # _____

  • # of lone pairs _____

  • Bond angle _________

  • Molecular Geometry __________

Is the molecule POLAR? _________


Molecular compounds chapter 8

  • AX angle.2E2 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

  • H angle.2O

  • Steric # _____

  • # of lone pairs _____

  • Bond angle _________

  • Molecular Geometry __________

Is the molecule POLAR? _________


Molecular compounds chapter 8

FIVE regions of high electron density angle.

  • 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

  • PF angle.5

  • Steric # _____

  • # of lone pairs _____

  • Bond angle _________

  • Molecular Geometry __________

Is the molecule POLAR? _________


Molecular compounds chapter 8

SIX regions of high electron density angle.

  • AX6 notation

  • Steric # is 6

  • No lone pairs

  • Geometry is octahedral

  • Bond angle is 90

Example SF6 molecules.


Molecular compounds chapter 8

  • SF angle.6

  • 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 angle.

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