Chapter 11 Chemical Bonds: The Formation of Compounds from Atoms - PowerPoint PPT Presentation

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Chapter 11 Chemical Bonds: The Formation of Compounds from Atoms

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Chapter 11 Chemical Bonds: The Formation of Compounds from Atoms Objectives: Describe the trends in the periodic table Know how to draw Lewis Structures of atoms – PowerPoint PPT presentation

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Title: Chapter 11 Chemical Bonds: The Formation of Compounds from Atoms


1
Chapter 11Chemical BondsThe Formation of
Compounds from Atoms
  • Objectives
  • Describe the trends in the periodic table
  • Know how to draw Lewis Structures of atoms
  • Understand and predict the formation of ionic
    bonds
  • Understand and predict covalent bonds
  • Describe electronegativity
  • Know how to draw complex lewis structures of
    compounds
  • Understand the formation of compounds containing
    polyatomic ions
  • Describe molecular shape, including the VSEPR
    model

2
Periodic Trends in Atomic Properties
  • Periodic table designed to show trends
  • Use trends to predict properties and reactions
    between elements
  • Trends include
  • Metals, nonmetals, metalloids
  • Atomic radius
  • Ionization energy
  • Electronegativity

3
Metals, Nonmetals and Metalloids
  • Metals
  • Lustrous, malleable, good conductors of heat and
    electricity
  • Left-hand side of table
  • Most elements are metals
  • Tend to lose electrons and form positive ions

4
Metals, Nonmetals and Metalloids
  • Nonmetals
  • Nonlustrous, brittle, poor conductors
  • (Hydrogen displays nonmetallic properties under
    normal conditions but is UNIQUE element)

5
Metals, Nonmetals and Metalloids
  • Metalloids
  • Found along border between metals and nonmetals
  • Metal Nonmetal
  • Usually electrons are transferred from metal to
    nonmetal

6
Atomic Radius
  • Increases down each group
  • Decreases from left to right across a period
  • Increase in positive charge stronger pull on
    electrons gradual decrease in atomic radius

7
Atomic Radius
8
Ionization Energy
  • The energy required to remove an electron from
    the atom

9
Ionization Energy
  • Ionization energy in Group A elements
  • Ionization energy
  • Nonmetals tend to gain electrons (rather than
    give them up)

10
Ionization Energy
11
Lewis Structures
  • Diagram that shows valence electrons
  • Dots number of s and p electrons
  • Paired dots
  • Simple way of showing electrons

12
Lewis Structures
  • When drawing
  • 3, 6, 9
  • Just like orbital filling diagram
  • Examples draw Lewis Structures of B, N, F, Ne

13
Lewis Structures
F
B
  • N

Ne
14
The Ionic Bond
  • Ionic bond
  • Attraction between electrostatic charges is a

15
The Ionic Bond
16
The Ionic Bond
  • NOT A MOLECULE
  • Bond not just between

17
The Ionic Bond
  • Typically metal nonmetal

18
Predicting Formulas of Ionic Compounds
  • In almost all stable chemical compounds of
    representative elements, each atom attains a
    noble gas electron configuration. This concept
    forms the basis for our understanding of chemical
    bonding.

19
Predicting Formulas of Ionic Compounds
  • How many electrons must be gained or lost to
    achieve noble gas configuration?

20
Predicting Formulas of Ionic Compounds
  • Elements in a family usually form compounds with
    the same atomic ratios

21
Predicting Formulas of Ionic Compounds
  • The formula for sodium oxide is Predict the
    formula for
  • Sodium sulfide
  • Sodium Ne3s1 must
  • Sulfur Ne3s23p4 must
  • Soformula must

22
Predicting Formulas of Ionic Compounds
  • Rubidium Oxide
  • Rubidium Kr5s1 must
  • Oxygen He2s22p4
  • Soformula must be
  • This makes sense b/c rubidium is in same family
    as sodium

23
The Covalent Bond
  • A pair of electrons
  • Most common type of bond
  • Electron orbital expands to include both nuclei

24
The Covalent Bond
25
The Covalent Bond
  • Atoms may share more than one pair of electrons
  • Double bond
  • Triple bond
  • Multiple bonds are
  • Covalent bonding between identical atoms means
    electrons are
  • Covalent bonding between different atoms leads to

26
Electronegativity
  • The attractive force that an atom of an element
    has for shared electrons
  • Atoms have different electronegativities
  • Electrons will spend more time near atom with
    stronger (larger) electronegativity
  • Soone atom assumes a
  • The other assumes a

27
Electronegativity
  • Electronegativity trends and periodic table
  • See table 11.5 page 237
  • Generally increases from left to right
  • Decreases down a group
  • Highest is fluorine (4.0)
  • Lowest is francium (0.7)

28
Electronegativity
29
Electronegativity
  • Polarity is determined by difference in
    electronegativity
  • Nonpolar covalent
  • Polar covalent
  • Ionic compound

30
Electronegativity
31
Electronegativity
  • If the electronegativity difference is greater
    than 1.7-1.9 then the bond will be more ionic
    than covalent
  • Above 2.0
  • Below 1.5
  • See Continuum on page 239

32
Electronegativity
  • Polar bonds form between two atoms
  • Molecules can also be polar or nonpolar
  • Dipole
  • Polar
  • Nonpolar

33
Lewis Structures of Compounds
  • Convenient way of showing ionic or covalent bonds
  • Usually the single atom in a formula is the
    central atom

34
The Ionic Bond
  • LEWIS STRUCTURES of ionic bonds

35
The Covalent Bond
  • LEWIS STRUCTURES of covalent bonds
  • Use dashes instead of dots

36
The Covalent Bond






37
Lewis Structures of Compounds
  1. Obtain the total number of valence electrons
  2. Add the valance electrons of all atoms
  3. Ionic add one electron for each negative charge
    and subtract one electron for each positive charge

38
Lewis Structures of Compounds
  • Write the skeletal arrangement of the atoms and
    connect with a single covalent bond
  • Subtract two electrons for each single bond
  • This gives you the net number of electrons
    available for completing the structure

39
Lewis Structures of Compounds
  1. Distribute pairs of electrons around each atom to
    give each atom a noble gas structure
  2. If there are not enough electrons then try to
    form double and triple bonds

40
Lewis Structures of Compounds
  • Write the Lewis Structure for methane CH4
  • Total number of valence electrons is eight
  • Draw skeletal structure
  • Dashes equal two electrons being shared
  • Subtract the eight electrons shown as dashes
  • Check that all atoms have a noble gas structure

41
Lewis Structures of Compounds
  • Methane, CH4

H
H
H
C
H
42
Lewis Structures of Compounds
  • Carbon Dioxide, CO2
  • Total valence electrons 16

O C O
Not Enough! Must try double bonds
43
Complex Lewis Structures
  • Some molecules and polyatomic ions have strange
    behaviors
  • No single Lewis structure is consistent
  • If multiple structures are possible the molecule
    shows resonance
  • Resonance structures show all possibilities

44
Complex Lewis Structures
  • Carbonate ion, CO32-

2-
2-
2-
Carbon only has 6 electrons try double bonds
more than one location..form resonant structures
O C O O
O C O O
O C O O
45
Compounds ContainingPolyatomic Ions
  • Polyatomic ion stable group of atoms that has a
    positive or negative charge
  • Behaves as a single unit in many chemical
    reactions
  • Sodium carbonate (Na2CO3)
  • Carbonate ion (co3) has covalent bonds
  • Sodium atoms are ionically bonded to carbonate ion

46
Compounds ContainingPolyatomic Ions
  • Easier to dissociate ionic bond than break
    covalent bond
  • More in chapters 6 and 7

47
Molecular Shape
  • Three-dimensional shape of molecule important
  • Explains
  • Helpful to know how to predict the geometric
    shape of molecules
  • Linear?
  • V-shaped?
  • Trigonal planar?
  • Tetrahedral?

48
The VSEPR Model
  • Valence Shell Electron Pair Repulsion Model
  • Make predictions about shape
  • Electron pairs will

49
The VSEPR Model
  • Linear Structure
  • 180o apart

50
The VSEPR Model
  • Trigonal Planar
  • 120o apart

51
The VSEPR Model
  • Tetrahedral structure
  • 109.50 apart
  • When drawing
  • Wedged line to show atom protruding from page
    dashed line to show atom receding from page

52
The VSEPR Model
  • Pyramidal shape
  • Four pairs of electrons on central atom BUT only
    three shared
  • Electrons are tetrahedral but actual shape is
    more of a pyramid

53
The VSEPR Model
  • Electron pairs determine shape BUT name for shape
    is determined by position of atoms

54
The VSEPR Model
  • V-shaped or bent
  • Four electron pairs but only two shared
  • Electron arrangement is
  • But, molecule is
  • Water
  • Helps explain some properties

55
The VSEPR Model
  • Predict the shape for .
  • Draw the Lewis Structure
  • Count the electron pairs and determine the
    arrangement that will minimize repulsions
  • Determine the positions of the atoms and name the
    structure
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