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Chapter 6. Chemical Bonding. Chapter Sections:. Introduction to chemical bonding Covalent bonding & molecular compounds Ionic bonding & ionic compounds Metallic bonding Molecular geometry. Section 1:. Introduction to chemical bonding. Introduction to chemical bonding.

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Chapter 6

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Chapter 6

Chemical Bonding


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Chapter Sections:

  • Introduction to chemical bonding

  • Covalent bonding & molecular compounds

  • Ionic bonding & ionic compounds

  • Metallic bonding

  • Molecular geometry


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Section 1:

  • Introduction to chemical bonding


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Introduction to chemical bonding

  • What is a chemical bond???

    A mutual electrical attraction between the nuclei and valence electrons of different atoms that binds the atoms together


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Introduction to chemical bonding

  • Why do atoms bond?

    They are working to achieve more stable arrangements where the bonded atoms will have lower potential energy than they do when existing as individual atoms.


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Introduction to chemical bonding

  • Types of Chemical Bonding:

    1. Ionic – an electrical attraction that forms between cations (+) and anions (-)

    2. Covalent – are formed when electrons are shared between atoms

    3. Metallic – formed by many atoms sharing many electrons


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Introduction to chemical bonding

  • However….

    • Bonds are never purely covalent or purely ionic.

    • The degree of ionic-ness or covalent-ness depends on property of electronegativity.


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Degree of Ionic/Covalent Character in Chemical Bonds

100%

50%

5%

0%

Ionic

Polar-Covalent

Nonpolar-Covalent


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Introduction to chemical bonding

  • Recall what electronegativity is:

    The degree of attraction that an atom has to electrons that are within a bonded compound.

    (see page 161)


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Introduction to chemical bonding

  • To determine the degree of ionic-ness or covalent-ness you must take each of the electronegativities for the elements in the compound and subtract them.


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Introduction to chemical bonding

  • If difference is 0-0.3 = nonpolar covalent

  • If difference is 0.3 – 1.7 = polar covalent

  • 1.7 and above = Ionic


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Ionic/Covalent Character Due to Electronegativity Differences

3.3

1.7

0.3

0

100%

50%

5%

0%

Ionic

Polar-Covalent

Nonpolar-Covalent


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Introduction to chemical bonding

2.5 - 2.1 = 0.4

Polar Covalent

2.5 - 0.7 = 1.8

Ionic

2.5 – 3.0 = 0.5

Polar Covalent

  • Sulfur + Hydrogen

  • Sulfur + Cesium

  • Sulfur + Chlorine


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Introduction to chemical bonding

  • In general however…

    If bonding elements are on opposite sides of the periodic table then they tend to be ionic.

    If elements are close together, then they tend to be covalent.


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Assignment:

  • Page 177 #3, 4, & 5

  • Page 209 #6


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Section 2:

  • Covalent Bonding & Molecular Compounds


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Covalent Bonding

  • What is a molecule?

    A neutral group of atoms that are held together by covalent bonds.

  • May be different atoms such as H2O or C6H12O6

  • May be the same atoms such as O2


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Covalent Bonding

  • Molecular compounds are made of molecules ….. Not ions!

  • We represent molecular compounds by chemical formulas that show numbers of atoms of each kind of element in the compound. CH4 - methane


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Covalent Bonding

  • Diatomic molecules are those elements that exist in pairs of like atoms that are bonded together.

  • There are 7 diatomic molecules:

    H2 N2 O2 F2 Cl2 I2 Br2


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Covalent Bonding

Formation of a covalent bond:

  • When atoms are far apart they do not attract – potential energy is zero.

  • As they come closer the electrons are attracted to protons but electrons and electrons repel – but e- to p attraction is stronger!


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Covalent Bonding

  • The electron clouds of the bonded atoms are overlapped and form a “bond length.”


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Covalent Bonding

  • Energy is released when these atoms join together with a bond.

  • Energy must be added to separatethese atoms – called bond energies.

  • Bond energy is expressed in kilojoules per mole.


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Covalent Bonding

  • Octet Rule – Atoms will either gain, lose, or share electrons so that their outer energy levels will contain eight electrons (H is an exception since it can only have 2 in the outer level).

  • These electrons that are being gained, lost, or shared are represented by using the electron dot diagrams.


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Examples of electron dot notations

  • 1 valence electron

  • 3 valence electrons

  • 5 valence electrons

  • 7 valance electrons

X

X

X

X


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Covalent Bonding

  • Shared electron pairs and unshared pairs:

    Cl:ClShared pair

    Unshared pairs


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Covalent Bonding

  • These electron dot representations are called Lewis structures.

  • Dots represent the valence electrons


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Lewis structures


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Covalent Bonding

  • Lewis structures can also be represented using structural formulas.

  • Dashes indicate bonds of shared electrons (unshared e- are not shown

    Cl - Cl

  • One pair (2 e-) is shared here.


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Covalent Bonding

  • Lewis structure for ammonia (NH3)


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Covalent Bonding

  • Practice:

    • Draw Lewis structure for methane CH4

    • Ammonia NH3

    • Hydrogen Sulfide H2S

    • Phosphorus trifluoride PF3


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Covalent Bonding

  • Some atoms can form multiple bonds – especially C, O, & N.

  • Double bonds are bonds that share 2 pair of electrons

    C=C means C::C

  • Triple bonds share 3 pair

    C≡C means C:::C


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Covalent Bonding

  • Resonance:

  • Some substances cannot be drawn correctly with Lewis structure diagrams

  • Some electrons share time with other atoms – ex. Ozone – O3


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Covalent Bonding

  • Electrons in ozone may be represented as: O = O–O

  • Other times it may be represented as O–O=O

  • Actually these structures are shared – electrons “resonate” (go back & forth) between them


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Covalent Bonding

  • Assignment:

    p. 189 #4 a – e


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Section 3:

  • Ionic Bonding and Ionic Compounds


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Section 3: Ionic Bonding & Compounds

  • Ionic compounds are formed of positive and negative ions

  • When combined these charges equal zero

    Ex: Na = 1+

    Cl = 1-

0 charge


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Section 3: Ionic Bonding & Compounds

  • Ionic substances are usually solids

  • Ionic solids are generally crystalline in shape

  • An ionic compound is a 3-D network of + and – ions that are attracted to each other


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Section 3: Ionic Bonding & Compounds

  • Crystals in ionic compounds exist in orderly arrangements known as a crystal lattice.


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Section 3: Ionic Bonding & Compounds

  • Ionic substances are not referred to as “molecules”

  • Ionic substances are referred to as “formula units”

  • A formula unit is the simplest ratio of the ions that are bonded together.


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Section 3: Ionic Bonding & Compounds

  • The ratio of ions depends on the charges.

  • What would result when F-combines with Ca2+?

    • CaF2


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Section 3: Ionic Bonding & Compounds

  • When ions are written using electron dot structures the dots are written and symbols for their charges.

  • Na.  Na+

  • Cl  -


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Compared to molecular compounds, ionic compounds:

  • Have very strong attractions

  • Are hard, but brittle

  • Have higher melting points and boiling points

  • When dissolved or in the molten state they will conduct electricity


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Polyatomic Ions:

  • A group of atoms covalently bonded together but with a charge.

  • SulfateSO42-

  • CarbonateCO32-

  • NitrateNO3-

  • AmmoniumNH4+


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Section 4:

  • Metallic Bonding


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Metallic Bonding

  • Metals are excellent electrical conductors in the solid state.

  • This is due to highly mobile valence electrons that travel from atom to atom.

e-


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Metallic Bonding

  • Generally metals have either 1 or 2 s electrons

  • p orbitals are vacant

  • Many are filling in the d level

  • Electrons become delocalized and move between atoms


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Metallic Bonding

  • A metallic bond is the mutual sharing of many electrons among many atoms.

  • Electrons travel in what is known as the zone of conduction.


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Metallic Properties

  • High electrical conductivity

  • High thermal conductivity

  • High luster

  • Malleable (can be hammered or pressed into shape)

  • Ductile (capable of being drawn or extruded through small openings to produce a wire)


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Metallic Bond Strength

  • Varies with nuclear charge and number of electrons shared.

  • High bond strengths result in high heats of vaporization (when metals are changed into gaseous phase)


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Section 5:

  • Molecular Geometry


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Molecular geometry…

  • A molecule’s properties depend on bonding of atoms, but also the molecular geometry.


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Molecular geometry…

  • Is the three dimensional arrangement of a molecule’s atoms in space.


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VSEPR Theory

  • Valence Shell Electron Pair Repulsion

  • Electrons around a nucleus repel each other to be as far away from each other as possible.


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VSEPR Theory

  • AB2 forms linear molecule as with beryllium

    However, water (H2O) is bent due to electrons repulsion!


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VSEPR Theory

  • AB3 forms trigonal planar molecule-ex. ammonia

  • AB4 forms tetrahedral molecule ex. methane

  • See page 200 for other shapes


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Hybridization

  • Explains how atom’s orbitals become rearranged to form covalent bonds.

  • Hybridization is the mixing of 2 or more orbitals of similar energies on the same atom to produce new orbitals of equal energies.


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Hybridization

  • Methane (CH4) is an example of hybridization:

    • Carbon’s normal configuration is 2s22p2

    • In methane all the electrons in the 2nd energy level become equal in energy and is referred to as sp3


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Intermolecular Forces:

  • What happens to liquid molecules when they are heated?

  • As energy is added particles overcome their attraction to each other.

  • IM Forces are the forces of attraction between molecules – not within the molecule.

  • IM forces vary in strength but are weaker than bonds that join atoms


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Intermolecular Forces:

  • Strongest IM forces exist in polar molecules.

  • Polar molecules act as tiny “dipoles” (equal & opposite charges separated by short distances)


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Intermolecular Forces:

  • Dipole – dipole forces attract between molecules such as between two water molecules.

  • Positive H region is attracted to negative O region of a different molecule.


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Intermolecular Forces:

  • Another IM force is Hydrogen bonding.

  • Is a strong type of dipole-dipole force

  • Explains high boiling points of H-containing substances such as water and ammonia


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Intermolecular Forces:

  • In hydrogen bonding, a hydrogen atom is attracted to an unshared pair of electrons of an electronegative atom in a nearby molecule.

  • The double helix of DNA is held together by hydrogen bonding.


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Intermolecular Forces:

London Dispersion forces:

  • Are very weak bonds

  • Occur due to the fact that since electrons are in constant motion that briefly there are moments where electrons are unevenly distributed and thus the molecule briefly has a charged area.


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