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Chemical Bonding

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A mutual electrical attraction between the nuclei and valence ... 7 valance electrons. X. X. X. X. Covalent Bonding. Shared electron pairs and unshared pairs: ... – PowerPoint PPT presentation

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Title: Chemical Bonding


1
Chapter 6
  • Chemical Bonding

2
Chapter Sections
  • Introduction to chemical bonding
  • Covalent bonding molecular compounds
  • Ionic bonding ionic compounds
  • Metallic bonding
  • Molecular geometry

3
Section 1
  • Introduction to chemical bonding

4
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

5
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.

6
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

7
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.

8
Degree of Ionic/Covalent Character in Chemical
Bonds
100 50 5 0
Ionic Polar-Covalent Nonpolar-Covalent
9
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)

10
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.

11
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

12
Ionic/Covalent Character Due to Electronegativity
Differences
100 50 5 0
3.3 1.7 0.3 0
Ionic Polar-Covalent Nonpolar-Covalent
13
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

14
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.

15
Assignment
  • Page 177 3, 4, 5
  • Page 209 6

16
Section 2
  • Covalent Bonding Molecular Compounds

17
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

18
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

19
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

20
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!

21
Covalent Bonding
  • The electron clouds of the bonded atoms are
    overlapped and form a bond length.

22
Covalent Bonding
  • Energy is released when these atoms join together
    with a bond.
  • Energy must be added to separate these atoms
    called bond energies.
  • Bond energy is expressed in kilojoules per mole.

23
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.

24
Examples of electron dot notations
  • 1 valence electron
  • 3 valence electrons
  • 5 valence electrons
  • 7 valance electrons

X
X
X
X
25
Covalent Bonding
  • Shared electron pairs and unshared pairs
  • ClCl Shared pair
  • Unshared pairs

26
Covalent Bonding
  • These electron dot representations are called
    Lewis structures.
  • Dots represent the valence electrons

27
  • Lewis structures

28
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.

29
Covalent Bonding
  • Lewis structure for ammonia (NH3)

30
Covalent Bonding
  • Practice
  • Draw Lewis structure for methane CH4
  • Ammonia NH3
  • Hydrogen Sulfide H2S
  • Phosphorus trifluoride PF3

31
Covalent Bonding
  • Some atoms can form multiple bonds especially
    C, O, N.
  • Double bonds are bonds that share 2 pair of
    electrons
  • CC means CC
  • Triple bonds share 3 pair
  • CC means CC

32
Covalent Bonding
  • Resonance
  • Some substances cannot be drawn correctly with
    Lewis structure diagrams
  • Some electrons share time with other atoms ex.
    Ozone O3

33
Covalent Bonding
  • Electrons in ozone may be represented as O
    OO
  • Other times it may be represented as OOO
  • Actually these structures are shared electrons
    resonate (go back forth) between them

34
Covalent Bonding
  • Assignment
  • p. 189 4 a e

35
Section 3
  • Ionic Bonding and Ionic Compounds

36
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
37
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

38
Section 3 Ionic Bonding Compounds
  • Crystals in ionic compounds exist in orderly
    arrangements known as a crystal lattice.

39
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.

40
Section 3 Ionic Bonding Compounds
  • The ratio of ions depends on the charges.
  • What would result when F- combines with Ca2?
  • CaF2

41
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 ? -

42
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

43
Polyatomic Ions
  • A group of atoms covalently bonded together but
    with a charge.
  • Sulfate SO42-
  • Carbonate CO32-
  • Nitrate NO3-
  • Ammonium NH4

44
Section 4
  • Metallic Bonding

45
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-
46
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

47
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.

48
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)

49
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)

50
Section 5
  • Molecular Geometry

51
Molecular geometry
  • A molecules properties depend on bonding of
    atoms, but also the molecular geometry.

52
Molecular geometry
  • Is the three dimensional arrangement of a
    molecules atoms in space.

53
VSEPR Theory
  • Valence Shell Electron Pair Repulsion
  • Electrons around a nucleus repel each other to be
    as far away from each other as possible.

54
VSEPR Theory
  • AB2 forms linear molecule as with beryllium
  • However, water (H2O) is bent due to electrons
    repulsion!

55
VSEPR Theory
  • AB3 forms trigonal planar molecule-ex. ammonia
  • AB4 forms tetrahedral molecule ex. methane
  • See page 200 for other shapes

56
Hybridization
  • Explains how atoms 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.

57
Hybridization
  • Methane (CH4) is an example of hybridization
  • Carbons normal configuration is 2s22p2
  • In methane all the electrons in the 2nd energy
    level become equal in energy and is referred to
    as sp3

58
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

59
Intermolecular Forces
  • Strongest IM forces exist in polar molecules.
  • Polar molecules act as tiny dipoles (equal
    opposite charges separated by short distances)

60
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.

61
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

62
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.

63
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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