1. Chapter 4 Compounds and Their Bonds
4.6
Electronegativity and Bond Polarity

2. Electronegativity The electronegativity value
indicates the attraction of an atom for shared electrons.
increases from left to right going across a period on the periodic table.
is high for the nonmetals, with fluorine as the highest.
is low for the metals.

3. Some Electronegativity Values for Group A Elements
Electronegativity decreases `
Electronegativity increases
High values `
Low values
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4. Nonpolar Covalent Bonds
A nonpolar covalent bond
occurs between nonmetals.
is an equal or almost equal sharing of electrons.
has almost no electronegativity difference (0.0 to 0.4).
Examples:
Electronegativity
Atoms Difference Type of Bond
N-N = Nonpolar covalent
Cl-Br = Nonpolar covalent
H-Si = Nonpolar covalent

5. Polar Covalent Bonds A polar covalent bond
occurs between nonmetal atoms.
is an unequal sharing of electrons.
has a moderate electronegativity difference (0.5 to 1.7).
Examples:
Electronegativity
Atoms Difference Type of Bond
O-Cl = 0.5 Polar covalent
Cl-C = 0.5 Polar covalent
O-S = 1.0 Polar covalent

6. Comparing Nonpolar and Polar Covalent Bonds
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7. Ionic Bonds An ionic bond occurs between metal and nonmetal ions.
is a result of electron transfer.
has a large electronegativity difference (1.8 or more).
Examples:
Electronegativity
Atoms Difference Type of Bond
Cl-K – 0.8 = Ionic
N-Na 3.0 – = Ionic
S-Cs 2.5 – 0.7 = Ionic

8. Electronegativity and Bond Types

9. Predicting Bond Types

10. Examples
Use the electronegativity difference to identify the type of bond [nonpolar covalent (NP), polar covalent (P), or ionic (I)] between the following:
A. K-N
B. N-O
C. Cl-Cl
D. H-Cl

11. Chapter 4 Compounds and Their Bonds
4.7
Shapes and Polarity of Molecules
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12. VSEPR In the valence-shell electron-pair repulsion theory
(VSEPR), the electron groups around a central atom
are arranged as far apart from each other as possible.
have the least amount of repulsion of the negatively charged electrons.
have a geometry around the central atom that determines molecular shape.

13. Guide to Predicting Molecular Shape

14. Four Electron Groups In a molecule of CH4,
there are 4 electron groups around C.
repulsion is minimized by placing 4 electron groups at angles of 109°, which is a tetrahedral arrangement.
the shape with four bonded atoms is tetrahedral.
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15. Three Bonding Atoms and One Lone Pair
In a molecule of NH3,
3 electron groups bond to H atoms, and the fourth one is a lone (nonbonding) pair.
repulsion is minimized with 4 electron groups in a tetrahedral arrangement.
with 3 bonded atoms, the shape is pyramidal.

16. Two Bonding Atoms and Two Lone Pairs
In a molecule of H2O,
2 electron groups are bonded to H atoms and 2 are lone pairs (4 electron groups).
4 electron groups minimize repulsion in a tetrahedral arrangement.
the shape with 2 bonded atoms is bent.

17. Shapes with 4 Electron Groups
Electron Pairs
Bonded Atoms
Lone Pairs
Molecular Shape
Example 4
Tetrahedral
CH4 3 1
Pyramidal
NH3 2
Bent
H2O 17

18. Examples
State the number of electron groups, lone pairs, and use VSEPR theory to determine the shape of the following molecules or ions. 1) tetrahedral 2) pyramidal 3) bent A. PF3 B. H2S C. CCl4

19. Study Tip
To determine shape, 1. draw the electron-dot structure. 2. count the electron pairs around the central atom. 3. count the bonded atoms to determine shape. 4 electron pairs and 4 bonded atoms = tetrahedral 4 electrons pairs and 3 bonded atoms = pyramidal 4 electron pairs and 2 bonded atoms = bent

20. Polar Molecules A polar molecule contains polar bonds.
has a separation of positive and negative charge called a dipole, indicated with + and -.
has dipoles that do not cancel.
+  • •
H–Cl H—N—H
dipole H
dipoles do not cancel

21. Nonpolar Molecules A nonpolar molecule contains nonpolar bonds.
Cl–Cl H–H
or has a symmetrical arrangement of polar bonds.
O=C=O Cl
Cl–C–Cl Cl
dipoles cancel

22. Determining Molecular Polarity
STEP 1: Write the electron-dot formula.
STEP 2: Determine the shape.
STEP 3: Determine if dipoles cancel or not.
Example: H2O
. .
H─O: H2O is polar │ H
dipoles do not cancel

23. Examples
Identify each of the following molecules as 1) polar or 2) nonpolar. Explain. A. PBr3 B. HBr C. Br2 D. SiBr4

24. Chapter 4 Compounds and Their Bonds
4.8
Attractive Forces in Compounds

25. Ionic Bonds In ionic compounds, ionic bonds
are strong attractive forces.
hold positive and negative ions together.

26. Dipole-Dipole Attractions
In covalent compounds, polar molecules
exert attractive forces called dipole-dipole attractions.
form strong dipole attractions called hydrogen bonds between hydrogen atoms bonded to F, O, or N, and other very electronegative atoms.

27. Dispersion Forces Dispersion forces are
weak attractions between nonpolar molecules.
caused by temporary dipoles that develop when electrons are not distributed equally.

28. Attractive Forces

29. Melting Points and Attractive Forces
Ionic compounds require large amounts of energy to break apart ionic bonds. Thus, they have high melting points.
Hydrogen bonds are the strongest type of dipole-dipole attractions. They require more energy to break than other dipole-dipole attractions.
Dispersion forces are weak interactions and very little energy is needed to change state.

30. Melting Points of Some Substances30

31. Examples
Identify the main type of attractive forces for each: 1) ionic 2) dipole-dipole 3) hydrogen bonds 4) dispersion A. NCl3 B. H2O C. Br-Br D. KCl E. NH3