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Chapter 9 Chemical Bonding PowerPoint PPT Presentation


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Chapter 9 Chemical Bonding. Bonding of Atoms. Bonding Extremes. In chapter 4 you learned that there were two types of bonds. Ionic Bonds – Electrons are transferred from one atom to another forming ions. Covalent (molecular) Bonds – Electrons are shared between atoms

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Chapter 9 Chemical Bonding

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Chapter 9 chemical bonding l.jpg

Chapter 9 Chemical Bonding

Bonding of Atoms


Bonding extremes l.jpg

Bonding Extremes

  • In chapter 4 you learned that there were two types of bonds.

  • Ionic Bonds – Electrons are transferred from one atom to another forming ions.

  • Covalent (molecular) Bonds – Electrons are shared between atoms

  • This is simple, but not all bonds are one type or the other. Instead, there is a continuum from one type to the other.

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Examples of the extremes l.jpg

Examples of the Extremes

  • Na + Cl [Na]+ + [ Cl ]-

  • This is an example of the ionic extreme.

  • C + O + O O C O

  • This is an example of the covalent extreme.

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A tug of war between atoms l.jpg

A tug of war between atoms

  • Picture two atoms, each one is represented by a person below. Both pull on the other’s valence electrons. Neither is strong enough to pull the electrons totally away, but one can pull a bit harder and the electrons move toward the harder pull.

d +

d -

e-

e-

The electrons are shared, but not equally, the resulting molecule has a somewhat negative end and a somewhat positive end.

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Electronegativity

  • The relative attraction of an atom for the electrons in a bond is called electronegativity

  • Electronegativity is a periodic property and each element has been assigned a quantitative number as shown in table 9.2 on page 304.

  • The difference in the electronegativity values between any two elements will determine the bonding type that will occur between them.

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Some examples l.jpg

Take a look at Hydrogen and Chlorine

The value for H is 2.1

The value for Cl is 3.0

The electronegativity difference (DEN) is .9

Using figure 9.1 on page 303, the bond type is read as polar covalent since .9 is between 0.5 and 2.0

Look at Hydrogen and Fluorine

H is still 2.1

F is 4.0

DEN = 4-2.1 or 1.9

It is also polar, but to a greater degree than is the HCl.

Some Examples

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For you to try l.jpg

Lithium and Bromine

Cesium and Fluorine

Fluorine and Fluorine

Sodium and Chlorine

Potassium Chlorine

Oxygen and Oxygen

Silicon and Carbon

Phosphorus and Bromine

DEN = 1.8 is classified as ionic, but has less ionic character than K and Cl

DEN = 3.3 Ionic

DEN = 0 Pure Covalent

DEN = 2.1 Ionic

DEN = 2.2 Ionic

DEN = 0 Pure Covalent

DEN = .7 Polar

DEN = .7 Polar

For you to try

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Periodic trends in electronegativity l.jpg

Periodic Trends in Electronegativity

Electronegativity Increases

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Variation in % Ionic Character

100%

50%

Percentage Ionic Character to the bond

  • 0 0.5 1.0 1.5 2.0 2.5 3.0

Electronegativity difference between atoms

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From the graph l.jpg

From the Graph

  • Once an electronegativity difference of 1.7 is reached, the bond has better than 50% ionic character.

  • The bond might be labeled correctly as polar covalent, but for the most part the rules you learned about metals and nonmetals forming ionic compounds and two nonmetals forming covalent compounds remains a good rule of thumb for deciding which bonding type is present.

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Polar bonding l.jpg

Polar Bonding

  • We’ve looked at the extremes which are a bond with a DEN so large that it is completely ionic or with a DEN that is zero where the bond is completely covalent.

  • Now we will look closer at the unequal sharing that occurs in polar bonds and what it means.

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Unequal sharing l.jpg

Unequal Sharing

  • Hydrogen and Chlorine are 2 nonmetals and by our earlier rule they will form a covalent bond.

  • Their DEN is .9 which puts them in the polar covalent range.

  • Since Cl is more electronegative than H, it will have a stronger attraction for the electrons and the Cl end of the molecule will have a somewhat negative charge while the H end will have a somewhat positive charge.

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Hydrogen chloride l.jpg

Hydrogen Chloride

d +

d -

Cl

The electron from H is pulled toward the Cl.

H

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Water is a polar molecule

  • Water is H2O, but let’s look at each H-O bond separately.

  • The DEN for H-O is 1.4, which makes it covalent.

  • The Oxygen is most electronegative and the electrons will be pulled in that direction

  • Let’s look at a model like the one for HCl

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The hydrogen oxygen bond l.jpg

The Hydrogen – Oxygen bond

O

d -

H

The electrons are pulled harder by the oxygen than the hydrogen atom.

d +

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A look at the lewis structure of water l.jpg

A look at the Lewis structure of Water

Oxygen has 6 valence electrons and Hydrogen has 1.

O

O

d -

H

d +

H

H

H

H

O

H

d +

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Water s polarity and it s shape l.jpg

Water’s Polarity and it’s shape

  • Water is bent because of it’s shared and unshared pairs of electrons.

These two pair of electrons are shared.

H

These two pair of electrons are unshared.

O

H

The unshared pair repel (like charges repel) and the molecule is “bent”. The angle is 105o.

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Water polarity and it s shape 2 l.jpg

Water Polarity and it’s shape 2

  • We’ll talk more about water in a later chapter, but if the 4 pair of electrons surrounding water in the H2O molecule were all equal the bond angle would be that of a tetrahedron (109.5o)

  • The nonbonding electrons require more room and “squeeze” the bonding pair closer and thus reduce the bond angle to 105o.

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Coffee filter chromatography l.jpg

Coffee Filter Chromatography

  • The mini lab on page 312 will help to illustrate some of water’s unique properties.

  • It will also give you some insight into polar and non-polar substances and how they might be expected to behave.

  • Follow the directions closely and answer the analysis questions.

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Lewis structures and molecular shape l.jpg

Lewis Structures and Molecular Shape

  • The following steps may be used to determine the correct Lewis Structure.

  • 1.) Decide which atoms are bonded

  • 2.) Count ALL valence electrons

  • 3.) Place 2 electrons in each bond.

  • 4.) Complete the octets of the atoms attached to the central atom by adding e- in pairs.

  • 5.) Place any remaining electrons on the central atom in pairs.

  • 6.) If the central atom does not have an octet, form double bonds. If necessary, form triple bonds.

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Examples of lewis structures and molecular shape l.jpg

Examples of Lewis Structures and Molecular Shape

  • SO3 - and ClO4 – Draw Lewis Structures

  • First SO3 – S 1 atom x 6 electrons = 6

    O 3 atoms x 6 electrons = 18 Total = 24 e-

There are no electrons left and S does not have 8 electrons so one pair from one of the oxygen atoms needs to be pulled in.

O

S

O

O

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Sulfur trioxide s shape l.jpg

Sulfur trioxide’s shape

The molecule will have one double bond between one Oxygen and the Sulfur and single bonds joining the other two Oxygen atoms to the sulfur.

The shape of the molecule will be somewhat triangular.

O

O

S

O

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Websites with 3D Models

  • Virtual Chemistry site

  • If you use this site at home, you’ll need to get a Chime Plugin The MDL Chime Site

  • You will be building gumdrop models and will need to consider the spacing and direction some of the parts of the molecules are directed.

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Formal charge and lewis structures l.jpg

Formal Charge and Lewis Structures

  • [Formal Charge] = [Number of e- in valence shell of the isolated atom] – [(Number of bonds to the atom) + (Number of unshared e-)]

  • Let’s look at the structures that are possible for Sulfuric acid. H2SO4.

  • There is a structure that obeys the octet rule, but it is not the structure that is found by experiment.

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