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Chapter 11 Chemical Bonding

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Title: Chapter 11 Chemical Bonding


1
Chapter 11 Chemical Bonding
  • Forces that hold atoms together

2
The Nature of Bonding
  • There are several major types of bonds. Ionic,
    covalent and metallic bonds are the three most
    common types of bonds.
  • Covalent bonds electrons are shared between
    atoms.
  • Ionic bonds electrons are transferred between
    atoms, creating cations and anions.
  • Metallic bonds two or more metals bonded
    together.

3
The Nature of Covalent Bonding
  • There are two different types of covalent bonds,
    polar covalent and nonpolar covalent.
  • polar covalent electrons are not shared equally
    between the two bonded atoms. The electrons are
    pulled toward the more electronegative of the
    elements.
  • nonpolar covalent electrons are shared equally
    between the two bonded atoms.

4
Electronegativities
5
The formation of a bond between two hydrogen
atoms.
Source Andrey K. Geim/High Field Magnet
Laboratory/University of Nijmegen
6
Probability representations of the electron
sharing in HF. (a) What the probability map would
look like if the two electrons in the HF bond
were shared equally. (b) The actual situation,
where the shared pair spends more time close to
the fluorine atom than to the hydrogen atom.
7
The Nature of Covalent Bonding
  • Ionic bonds are formed when there is an
    electronegativity difference (DEN) greater than
    2.0.
  • Polar covalent bonds form when there is a DEN
    between 0.5 and 1.7.
  • Nonpolar covalent bonds form when there is a DEN
    between 0 and 0.49.

8
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9
The Nature of Covalent Bonding
  • If the DEN is between 1.7 and 2.0, an ionic bond
    will form if a metal is one of the elements, and
    a polar covalent bond will form if only nonmetals
    or metalloids are present.

10
The Nature of Covalent Bonding
  • What type of bond is formed between the following
    elements?
  • N and O K and F
  • Mg and Cl P and F
  • C and H

11
Bond Polarity
  • Covalent bonding between unlike atoms results in
    unequal sharing of the electrons
  • One end of the bond has larger electron density
    than the other
  • The result is bond polarity
  • The end with the larger electron density gets a
    partial negative charge
  • The end that is electron deficient gets a partial
    positive charge

12
The three possible types of bonds (a) a
covalent bond formed between identical atoms
(b) a polar covalent bond, with both ionic and
covalent components and (c) an ionic bond, with
no electron sharing.
13
Dipole Moment
  • Bond polarity results in an unequal electron
    distribution, resulting in areas of partial
    positive and partial negative charge
  • Any molecule that has a center of positive charge
    and a center of negative charge in different
    points is said to have a dipole moment (two
    different poles of charge).

14
Dipole Moment
  • If a molecule has more than one polar covalent
    bond, the areas of partial negative and positive
    charge for each bond will partially add to or
    cancel out each other
  • The end result will be a molecule with one center
    of positive charge and one center of negative
    charge
  • The dipole moment effects the attractive forces
    between molecules and therefore the physical
    properties of the substance

15
(a) The charge distribution in the water
molecule. (b) The water molecule behaves
as if it had a positive end and a negative end,
as indicated by the arrow.
16
(a) Polar water molecules are strongly attracted
to positive ions by their negative ends. (b) They
are also strongly attracted to negative ions by
their positive ends.
17
Polar water molecules are strongly attracted to
each other.
18
Electron Configuration in Ionic Bonding
  • Metals tend to lose their valence electrons,
    leaving a complete octet in their next-lowest
    energy level.
  • Sodium (1 valence electron) loses 1 electron
    and becomes Na1.
  • Na (Ne3s1) ? 1e- Na1(Ne)
  • Calcium (2 valence electrons) loses 2 electrons
    and becomes Ca2.
  • Ca (Ar4s2) ? 2e- Ca2(Ar)

19
Electron Configuration in Ionic Bonding
  • Nonmetals tend to gain or share valence electrons
    to complete an octet in their highest energy
    level.
  • Oxygen (6 valence electrons) gains two
    electrons to become O-2 .
  • O (He2s22p4) 2e- ? O-2 (He 2s22p6)
  • Phosphorus (5 valence electrons) gains three
    electrons to become P-3.
  • P (Ne3s23p3) 3e- ? P-3 (Ne 3s23p6)

20
Formation and Properties of Ionic Compounds
  • Ionic bonds forces of attraction that bind
    cations and anions together.
  • Ionic compound consists of electrically neutral
    group of ions joined by electrostatic forces.
  • Example Sodium chloride

21
Formation and Properties of Ionic Compounds
  • At room temperature, most ionic compounds are
    crystalline solids, where ions are arranged in
    various 3-D patterns.
  • Because of the large attractive forces of the
    ions to each other the compounds become very
    stable and have high melting points.

22
Sodium Chloride Crystals
23
The structure of lithium fluoride.
24
Electron Configuration in Ionic Bonding
  • Scientists have learned that all of the elements
    within each group behave similarly because they
    have the same number of valence electrons.
  • Valence electrons - of electrons in the highest
    occupied energy level of an atom.
  • The number of valence electrons is related to the
    group numbers on the periodic table.

25
Electron Configuration in Ionic Bonding
  • Group 1 elements 1 valence electron.
  • Group 2 elements 2 valence electrons.
  • Groups 3-12 elements 2 valence electrons.
  • Group 13 elements 3 valence electrons.
  • Group 14 elements 4 valence electrons.
  • Group 15 elements 5 valence electrons.
  • Group 16 elements 6 valence electrons.
  • Group 17 elements 7 valence electrons.
  • Group 18 elements 8 valence electrons.

26
Determining Valence Electrons for an Ion or a
Compound
  • 1. Multiply the number of valence electrons by
    the number of moles of each element.
  • 2. Add up all the electrons for each of the
    elements.
  • 3. If there is a charge and it is negative, add
    that number of electrons to the total.
  • 4. If there is a charge and it is positive,
    subtract that number of electrons from the total.
  • Total of electrons should always be an even
    number!

27
Determining Valence Electrons Examples
  • Determine the number of valence electrons in each
    of the following compounds and ions
  • NH41
  • CH2ClBr
  • PO4-3

28
Electron Configuration in Ionic Bonding
  • Valence electrons are the only electrons involved
    in bonding, and are the only ones written when
    drawing electron dot structures.
  • In forming compounds, atoms tend to achieve the
    electron configuration of a noble gas, having 8
    valence electrons which as known as having a
    stable octet (octet for 8 valence electrons).

29
Lewis Symbols of Atoms and Ions
  • Also known as electron dot symbols
  • Use symbol of element to represent nucleus and
    inner electrons
  • Use dots around the symbol to represent valence
    electrons
  • put one electron on each side first, then pair
  • Elements in the same group have the same Lewis
    symbol
  • Because they have the same number of valence
    electrons
  • Cations have Lewis symbols without valence
    electrons
  • Anions have Lewis symbols with 8 valence electrons

30
The Nature of Covalent Bonding
  • Structural formula chemical formulas that show
    the arrangement of atoms in molecules and
    polyatomic ions.
  • Octet rule atoms gain or lose electrons to
    acquire the stable electron configuration of a
    noble gas, usually having 8 valence electrons.

31
Lewis Structures
  • You can represent the formation of the covalent
    bond in H2 as follows
  • This uses the Lewis dot symbols for the hydrogen
    atom and represents the covalent bond by a pair
    of dots.

32
Lewis Structures
  • The shared electrons in H2 spend part of the time
    in the region around each atom.
  • In this sense, each atom in H2 has a helium
    configuration.

33
Lewis Structures
  • The formation of a bond between H and Cl to give
    an HCl molecule can be represented in a similar
    way.
  • Thus, hydrogen has two valence electrons about it
    (as in He) and Cl has eight valence electrons
    about it (as in Ar).

34
Lewis Structures
  • Formulas such as these are referred to as Lewis
    electron-dot formulas or Lewis structures.

35
The Nature of Covalent Bonding
  • Exceptions to the octet rule
  • H needs 2 electrons to be stable
  • Be needs 4 electrons to be stable
  • B needs 6 electrons to be stable

36
The Nature of Covalent Bonding
  • Steps for Drawing Lewis-dot structures
  • Determine the number of valence electrons in the
    molecule.
  • - When drawing determining valence electrons for
    an ion, add electrons if it an anion, and
    subtract electrons if it is a cation.
  • The first element in the compound will be the
    central atom. Exception hydrogen will never be
    the central atom.

37
The Nature of Covalent Bonding
  • Steps for Drawing Lewis-dot Structures
  • 3. Use one pair of electrons to bond each outer
    or terminal atom to the central atom.
  • 4. Make all outer or terminal atoms stable using
    the valence electrons.
  • 5. Put any remaining electrons around the central
    atom as lone pairs.

38
The Nature of Covalent Bonding
  • Draw the Lewis structure for
  • NH3
  • PO43-
  • CHFClBr
  • PF5-2

39
The Nature of Covalent Bonding
  • Single covalent bond a bond in which two atoms
    share a pair of electrons.
  • Double covalent bond a bond in which two atoms
    share two pairs of electrons.
  • Triple covalent bond a bond in which two atoms
    share three pairs of electrons.

40
The Nature of Covalent Bonding
  • If you have used up all of the valence electrons
    and you still need two more electrons to make the
    central atom stable, you must have one double
    bond.
  • If you still need four more electrons to make the
    central atom stable, you must have either one
    triple bond or two double bonds.
  • Double and triple bonds exist most commonly
    between C, N, O, and S atoms.

41
The Nature of Covalent Bonding
  • Draw Lewis structures for
  • NOCl
  • CO2
  • N2
  • SiO3-2

42
The Nature of Covalent Bonding
  • Resonance structures molecules or ions that can
    have two or more different Lewis structures.
    They must contain a double bond to have any
    resonance structures.
  • Resonance structures dont truly have a single
    bonds or a double bond, but a hybrid mixture of
    bonds where the extra bond is spread equally
    among the other single bonds.

43
The Nature of Covalent Bonding
  • Draw Lewis structures for
  • NOCl
  • CO2
  • N2
  • SiO3-2

44
The Nature of Covalent Bonding
  • Single bonds are longer (length between the
    atoms) than double and triple bonds.
  • Double bonds are longer than triple bonds.
  • Single bonds are not as strong as double bonds,
    and can be broken much easier than double bonds.
  • Triple bonds are stronger than double bonds.

45
Bonding Theory
  • The valence-shell electron pair repulsion (VSEPR)
    model predicts the shapes of molecules and ions
    by assuming that the valence shell electron pairs
    are arranged as far from one another as possible.
  • To predict the relative positions of atoms around
    a given atom using the VSEPR model, you first
    note the arrangement of the electron pairs around
    that central atom.

46
Predicting Molecular Geometry
  • The following rules and figures will help discern
    electron pair arrangements.
  • Draw the Lewis structure
  • Determine how many bonding pairs are around the
    central atom. Count a multiple bond as one pair.
  • Determine how many lone pairs, if any, are around
    the central atom.
  • All diatomic molecules have a linear shape.

47
Arrangement of Electron Pairs About an Atom
2 pairs Linear
48
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49
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50
Molecular Geometry Examples
  • NH3
  • PO43-
  • CHFClBr
  • PF5
  • SeF6
  • NOCl
  • CO2
  • SF2
  • N2
  • SiO3-2

51
Polar Bonds and Molecules
  • Nonpolar covalent bond equal sharing of
    electrons between two atoms.
  • Polar covalent bond unequal sharing of
    electrons between two atoms.
  • In polar covalent bonds the electrons are pulled
    closer to the atom with the larger
    electronegativity value.
  • This creates a partial positive and a partial
    negative pole within the bond.

52
Polar Bonds and Molecules
  • Polar bonds can create polar or nonpolar
    molecules and ions.
  • If the centers of partial positive and partial
    negative charge are in the same location, the
    molecule or ion is nonpolar.
  • If the centers of partial positive and partial
    negative charge are in different locations, the
    molecule or ion is polar.

53
Polar Bonds and Molecules
  • If the central atom has lone pairs of electrons,
    the molecule or ion is polar.
  • If the central atom does not have any lone pairs
    of electrons, the molecule or ion is nonpolar (as
    long as the atoms around the central atom are all
    the same).
  • If there is more than one type of atom around the
    central atom, the molecule or ion will be polar
    regardless of whether or not it has lone pair(s)
    on the central atom.

54
Examples Polar or Nonpolar?
  • Determine whether each of the following molecules
    or ions are polar or nonpolar
  • NO2-1
  • CFCl3
  • N2
  • CN-1
  • CH4
  • SO3-2
  • NOCl

55
Polar Bonds and Molecules
  • Attractions Between Molecules
  • Molecules are attracted to one another by a
    variety of forces.
  • These intermolecular forces are weaker than ionic
    or covalent bonds.
  • These forces are responsible for whether or not a
    molecular compound is a solid, liquid, or a gas.

56
Polar Bonds and Molecules
  • van der Waals forces consist of dispersion
    forces and dipole interactions (dipole-dipole
    moments).
  • Dispersion forces weakest of all intermolecular
    forces. They are caused by the motion of
    electrons. The strength of dispersion forces
    increases with the increasing number of electrons
    in a molecule.

57
Polar Bonds and Molecules
  • All molecules contain dispersion forces.
  • As molar mass and the number of electrons
    increase, dispersion forces increase.
  • Halogens are the most common molecules to have
    dispersion forces. Fluorine is a gas, Bromine is
    a liquid and Iodine is a solid.

58
Polar Bonds and Molecules
  • Dipole interactions occur when polar molecules
    or ions are attracted to one another. This
    occurs when a partial positive charge and a
    partial negative charge come close to each other.
  • Dipole interactions are very similar to, but much
    weaker than ionic bonds.

59
Polar Bonds and Molecules
  • Hydrogen bonds force exerted between a hydrogen
    atom bonded to an F, O, or N atom in one molecule
    and an unshared pair on another F, O, or N atom
    in a nearby molecule.
  • Hydrogen bonds can have a great effect on the
    boiling point of a substance.

60
Intermolecular Forces Examples
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