Ch 6 Section 5 Flashcards Preview

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Flashcards in Ch 6 Section 5 Deck (93):
1

Molecular geometry is the

3-dimensional arrangement of a molecule's atoms in space

2

The polarity of each bond along with the geometry of the molecule determines

Molecular polarity

3

Molecular polarity is the

Uneven distribution of molecular charge

4

Molecular polarity strongly influenced the forces that act

Between molecules in liquids and solids

5

A chemical formula reveals little information about

A molecules geometry

6

After performing many tests designed to reveal the shapes of various molecules chemists developed two different

Equally successful theories to explain certain aspects of their findings

7

One theory accounts for

Molecular bond angles

8

The other theory is used to describe the orbitals that

Contain the valence electrons of a molecules atoms

9

Diatomic molecules like h2 and HCl must be

Linear because they consist of only two atoms

10

To predict the geometries of more complicated molecules one must consider the

Locations of all electron pairs surrounding the bonded atoms

11

The abbreviation VSEPR stands for

valence-shell, electron-pair repulsion

12

The abbreviation for VSEPR refers to the repulsion between

Pairs of valence electrons of the atoms in a molecule

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VSEPR theory states that repulsion between the sets of valence level electrons surrounding an atom causes these sets to be

Oriented as far apart as possible

14

According to the VSEPR theory the shared pairs in BeF2 will be as far

Away from each other as possible

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The distance between electron pairs is maximized if the bonds to fluorine are on

Opposite sides of the beryllium afl , 180 degrees apart. Thus the molecule is linear

16

If we represent the central atom in s molecule by the letter A and we represent the atoms bonded to the central atom by the letter B then according to VSEPR theory BeF2 is an example of an

AB2 molecule which is linear

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In an AB3 molecule the three A-B bonds stay farthest apart by pointing to the corners of an

Equilateral triangle giving 120 degree angles between the bonds

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The central atoms in AB4 molecules follow the octet rule by sharing

Four electron pairs with B atoms

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In AB4 molecules the distance between electron pairs is maximized if each A-B bond points to

One of four corners of a tetrahedron

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Ammonia (NH3) and water (H2O) are examples of molecules in which the central atom has both

Shared and unshared electron pairs

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VSEPR theory postulates that the line pair of electrons occupies space around the (ammonia atom)

Nitrogen atom just as the bonding pairs do

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Thus in ammonia the electron pairs maximize their separation by assuming the

Four corners of a tetrahedron

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Lone pairs do occupy space but our description of the observed space of s molecules refers to the

Positions of atoms only

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The general VSEPR formula for molecules such as ammonia is

AB3E where E represents the unshared electron pair

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The VSEPR formula for water is

AB2E2

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For H2O VSEPR theory states that the lone pairs occupy space around the central atom but that the actual shape of the molecule is determined by the position of the

Atoms only, resulting in a bent molecule

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The bond angles in ammonia and water are smaller because the unshared electron pairs

Repeal electrons more strongly than do bonding electron pairs

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In VSEPR theory double and triple bonds are treated in the same way as

Single bonds

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In VSEPR theory poly atomic ions are treated similarly to

Molecules

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Lewis structures and VSEPR theory and molecular geometry can be used together to predict the... Of poly atomic ions as well as... With

Shapes; molecules with double or triple bonds

31

VSEPR troth does not reveal the relaid shop between a molecules

Geometry and the orbitals occupied by its bonding electrons

32

To explain the orbitals of an atom become rearranged when the atom forms covalent bonds, a different

Model, hybridization is used

33

Hybridization is the mixing of two or more atomic orbitals of similar energies on the same atom to

Produce new hybrid atomic orbitals of equal energies

34

The sp3 orbitals all have the same

Energy which is greater than that of the 2s orbitals but less than that of the 2p orbitals

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Hybrid orbitals are orbitals of

Equal energy produced by the combination of two or more orbitals on the same atom

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The number of hybrid orbitals produced equals the number of

Orbitals that have combined

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Hybridization also explains the

Bonding and geometry of many molecules formed by group 15 and 16 elements

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The linear geometry of molecules such as beryllium fluoride is made possible by hybridization involving the

S orbital and one available empty p orbital To yield sp hybrid orbitals

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The trogonal-planar geometry of molecules such as boron fluoride is made possible by hybridization involving the

S orbital one singly occupied p orbital and one empty p orbital to yield sp2 hybrid orbitals

40

Atomic orbitals: s, p
Type of hybridization:sp
Number of hybrid orbitals: 2
Geometry:

180 degrees; linear

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Atomic orbitals:s,p,p
Type of hybridization: sp2
Number of hybrid orbitals: 3
Geometry:

120 degrees: trigonal planar

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The properties of molecules depend not only on the bonding of atoms but also on

Molecular geometry

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Atomic orbitals:s,p,p,p
Type of hybridization: sp3
Number of hybrid orbitals: 4
Geometry:

109.5 degrees; tetrahedral

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Atoms bonded to central atom: 2
Lone pairs of electrons: 0
Type of molecule: AB2
Molecular shape:

Linear

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Atoms bonded to central atom: 3
Lone pairs of electrons: 0
Type of molecule: AB3
Molecular shape:

Trigonal-planar

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Atoms bonded to central atom: 2
Lone pairs of electrons: 1
Type of molecule: AB2E
Molecular shape:

Bent or angular

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Atoms bonded to central atom: 4
Lone pairs of electrons: 0
Type of molecule: AB4
Molecular shape:

Tetrahedral

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Atoms bonded to central atom: 3
Lone pairs of electrons: 1
Type of molecule: AB3E
Molecular shape:

Trigonal-pyramidal

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Atoms bonded to central atom: 2
Lone pairs of electrons: 2
Type of molecule: AB2E2
Molecular shape:

Bent or angular

50

Atoms bonded to central atom: 5
Lone pairs of electrons: 0
Type of molecule: AB5
Molecular shape:

Trigonal-bipyramidal

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Atoms bonded to central atom: 6
Lone pairs of electrons: 0
Type of molecule: AB6
Molecular shape:

Octahedral

52

As a liquid is heated the kinetic energy of its particles

Increases

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At the boiling point the needy is sufficient to overcome the force of

Attraction between the liquids particles

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The particles pull away from each other and enter the

Gas phase

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Boiling point therefore is a good measure of the force of

Attraction between particles of a liquid

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The higher the boiling point the stronger the

Forces between particles

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The forces of attraction between molecules are known as

Intermolecular forces

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Intermolecular forces. Art in strength but are generally weaker than bonds that join

Atoms in molecules, ions in ionic compounds, or metal atoms in solid metals

59

The values for ionic compounds and metals are much higher than those for

Molecular substances

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The strongest intermolecular forces exist between

Polar molecules

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Polar molecules act as tiny dipoles because of their

Uneven charge distribution

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A dipole is created by equal but opposite charges that are

Separated by a short distance

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The direction of a dipole is from the dipoles

Positive pole to its negative pole

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A dipole is represented by an arrow with a head pointing toward the

Negative pole and a crossed tail situated at the positive pole

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The negative region in one polar molecule attracts the

Positive region in adjacent molecules

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The forces of attraction between polar molecules are known as

Dipole-dipole forces

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The forces of attraction between polar molecules are known as

Dipole-dipole forces

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Dipole-dipole forces are short range forces, acting only between

Nearby molecules

69

The polarity of diatonic molecules such as ICl is determined by just

One bond

70

For molecules containing more than two atoms molecular polarity depends on both the

Polarity and the orientation of each bond

71

Because the molecule is bent the polarities of these two bonds combine to make the molecule

Highly polar

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In some molecules individual bond dipoles

Cancel one another causing the molecular polarity to be zero

73

A polar molecule can induce a dipole in an Nonpolar molecule by

Temporarily attracting its electrons

74

The result is a short range intermolecular force that is somewhat

Weaker than the dipole-dipole force

75

Some hydrogen containing compounds have unusually high

Boiling points

76

These high boiling points is explained by the presence of a particularly strong type of

Dipole-dipole force

77

In compounds containing H-F, H-O, or H-N bonds, the large electronegativity differences between hydrogen atoms and fluorine, oxygen, or nitrogen atoms make the bonds connecting them

Highly polar

78

The high polarity gives the hydrogen atom a positive charge that is almost half as large as that of a

Proton

79

The small size of the hydrogen atom allows the atom to come very close to an

Unshared pair of electrons on an adjacent molecule

80

The intermolecular force in which a hydrogen atom that is bonded to a highly electronegative atom is attracted to an electronegative atom in a nearby molecule is known as

Hydrogen bonding

81

Hydrogen bonds are usually represented by dotted lines connecting the hydrogen bonded hydrogen to the unshared electron pair of the

Electronegative atom to which it is attracted

82

Even noble gas atoms and molecules that are Nonpolar experience a

Weak intermolecular attraction

83

In any atom or molecule the electrons are in

Continuous motion

84

Thus at any instant the electron distribution may be slightly

Uneven

85

The momentary uneven charge creates a positive pole in one part of the atom or molecule and a

Negative pole in another

86

This temporary dipole can then induce a dipole in an

Adjacent atom or molecule

87

The two are held together for an instant by the weak attraction. Between the

Temporary dipoles

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The intermolecular attract drinks resulting from the constant motion of electrons and the creation of instantaneous dipoles are called

London dispersion forces

89

London dispersion forces are named after

Fritz London who first proposed their existence in 1930

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London forces act between all

Atoms and molecules

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London forces are the only intermolecular forces acting among

Noble gas atoms and Nonpolar molecules

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Because London forces are dependent on he motion of electrons their strength increases with the number of

Electrons in the interacting atoms or molecules

93

London forces increase with increasing

Atomic or molar mass