1. 06:33 PM

2. Chapter 4: Chemical
Bonds
06:33 PM

3. Overview of Chapter 4
 Electron configuration- valence electrons and stability.
 Electron-dot structures of atoms and ions; use to
describe reactions.
 Bonding in compounds: Ionic and covalent.
 Nomenclature of ionic and covalent compounds.
 Electronegativity.
 Polar versus nonpolar compounds.
 Writing electron-dot structures of molecules.
 Simple geometries of molecules.
 Intermolecular forces in states of matter and in mixtures:
Dipole forces, hydrogen bonding, dispersion forces,
forces in solutions.
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4. Bonding and Valence Electrons
Forces that hold atoms together within a
molecule, or ions together in crystals.
Involve valence electrons = Outermost
electrons.
Inner electrons (core electrons) are generally
not involved in bonding.
Electron configuration- arrangement of
electrons.
Valence
electrons +
Core
06:33 PM electrons

5. Noble Gas Configurations
Example:
Sodium (Na) losing an electron-
11p 11p + 
Na: 11 e-1 Na+1: 10 e-1 lost e-1
The sodium ion and neon are 10p
isoelectronic- They have the same
electron configuration.
Ne: 10 e-1

6. Noble Gas Configurations
Example:
Chlorine (Cl) gaining an electron-
17p +  17p
Cl: 17 e-1 gained e-1 Cl-1: 18 e-1
18p
Are the chlorine ion
and argon isoelectronic?.
Ar: 18 e-1
Ar: 18 e-1

7. Electron-Dot Structures
 Electron-dot structures (EDSs) = Lewis-dot
structures (LDSs)?
 Represent the no. of valence e-1 around an atom (dots
around symbol).
 Example 1: Give the electron-dot structure for sodium.
What is the symbol for sodium?
Na
How many valence electrons does it contain?
1
Write the symbol, then place electrons
around the symbol (in pairs if possible).
Na

8. Electron-Dot Structures of Atoms
 Give the electron-dot structure for the
following:
2.) K  K 3.) Mg 
Mg
4.) Al  Al 5.) O  O
6.) Br  Br

9. More on Electron Configurations
Noble gases are most stable group of
elements (least reactive) Why?
8 valence electrons = stable octet 
Stable electron configuration.
Octet rule- Atoms attempt to obtain 8
valence electrons.
Exceptions: Group 1A, 2A, 3A elements
and helium.

10. Electron-Dot Structures of Ions
Example: Sodium forming sodium ion.
11p 11p + 
Na: 11 e-1 Na+1: 10 e-1 lost e-1
Na Na +1

11. Electron-Dot Structures of Ions
Example: Chlorine forming chloride ion.
17p +  17p
Cl: 17 e-1 gained e-1 Cl-1: 18 e-1
Cl Cl -1

12. Ionic Bonds
 A chemical bond between two ions;
 Transfer of electrons between a metal
and a nonmetal.
 Ionic compounds contain ionic bonds;
usually crystals.
 Very strong bond indeed!

13. Sodium and Chlorine: Boom!
 Example 1: Sodium will react with chlorine
to form sodium chloride (NaCl).
Na + Cl ---- -----> NaCl
Na + Cl Na +1
Cl -1

14. Ionic Compounds
Example 2: A potassium atom reacts with a
bromine atom to form potassium bromide
(KBr).
K + Br  KBr
K + Br K +1 +
Br -1

15. Ionic Compounds
Example 3: A potassium atom reacts with an
oxygen atom to form potassium oxide
(K2O).
2 K + O  K 2O
K K +1
O + O -2
K +
K+1

16. Nomenclature of Binary
Ionic Compounds
Roman numeral system. See Table 5.2 (CFCT).
Rules:
Names of the elements in order according to chemical
formula.
Change the ending of the last element to “-ide”.
Consider ionic state of metal:
If the metal has only one ionic state, you are finished naming.
If the metal has more than one ionic state, state in
parentheses after the metal the size of the charge on the
metal ion.

17. Nomenclature of Binary Ionic
Compounds: Formula to Name
Examples:
1.) LiCl
Li+1 and Cl-1
lithium chlorine  lithium chloride
2.) CuO
Copper oxygen
Cu+2 and O-2
copper (II) oxygen  copper (II) oxide

18. Nomenclature of Binary Ionic
Compounds: Name to Formula
Examples:
1.) What is the chemical formula for sodium iodide.
Na and I 
Na+1 + I-1 
NaI
2.) What is the chemical formula for potassium
sulfide?
K and S 
K+1 + S-2  Cross-over method:
K2S

19. Name to Formula
Examples:
3.) Copper (II) oxide
Cu+2 O-2  One Cu for one O
Cu2O2  lowest-whole number ratios!
CuO
4.) Iron (III) chloride
FeCl3

20. Covalent Bonding
Electrons strongly shared, not transferred.
Covalent bonds are weaker bond than
ionic bonds; nonmetal bonded to
nonmetal.
Examples:
SO2 (sulfur dioxide)
H2S (hydrogen sulfide)
NH3 (ammonia)

21. Types of Covalent Bonds
Three major types of covalent bonds:
- Single: one pair of electrons Ex: C:H or C–H
- Double: two pairs of electrons Ex: C::O or C=O
- Triple: three pairs of electrons Ex: N:::N or N≡N
Example: How many pairs of electrons are in the
hydrogen-oxygen bond in water?
H-O-H
Answer: 1 pair

22. Rules for Naming
Binary Covalent Compounds: Ex 1.
 Start with the chemical formula: CO2
 Elements: carbon oxygen
 First element-
Use a prefix for elements
in quantity greater than one:
carbon oxygen
 Second element-
Use prefix for elements of any
quantity:
carbon dioxygen
 Add –ide carbon dioxide

23. Nomenclature for Covalent
Compounds: Name from Formula
Example 2: What is the chemical name for CO?
carbon monoxide.
Example 3: What is the chemical name for PCl3?
phosphorus trichloride.
Example 4: What is the chemical name for N2O?
dinitrogen monoxide.

24. Nomenclature for Covalent
Compounds: Formula from Name
Example 1: What is the chemical formula for
dihydrogen monoxide?
H2O
Example 2: What is the chemical formula for sulfur
trioxide.
SO3
Example 3: What is the chemical formula for
tetraphosphorus trisulfide?
P4S3

25. Electronegativity
The ability for a nucleus to attract
electrons.
 Atoms of different elements have
different abilities of attracting electrons.
 See table in text, and lecture guide.

26. Applying Electronegativities:
Overview
 Take differences between the two
elements in a bond to determine
predominant character of bond.
 Nonpolar covalent bond: < 0.5.
 Polar covalent bond: between 0.5 and 2.0.
 Ionic: > 2.0.

27. Applying Electronegativities
Examples:
Hydrogen (H2): H—H
Difference = 0; nonpolar covalent
Hydrogen chloride (HCl): H—Cl
Difference = 0.96; polar covalent
Sodium chloride (NaCl):
Difference = 2.23; ionic  Na+1 Cl-1

28. Applying Electronegativities
Examples: How is the electron density
around the molecule distributed for a
hydrogen molecule?
Hydrogen atoms  hydrogen
molecule
H + H  H–H

29. Applying Electronegativities
Example: Give the partial charges on the
atoms in the hydrogen-chlorine bond in
HCl. Show the dipole moment in the
bond.
How is the electron
δ+ δ-
density distributed?
H – Cl
2.2 3.16 H – Cl

30. Applying Electronegativities
 Example: Show ions given their
electronegativities.
Na Cl
0.93 3.16
Na+1 Cl-1

31. Polyatomic molecules
 See LG (p. 169) for rules on preparing
molecules.
 Go through examples in the Lewis-Dot
Structures Worksheet for Covalent
Molecules in the LG (p. 171-175)
 Consult rules for writing Lewis-dot
structures as reference.

32. Polyatomic Ions
 Polyatomic ions: ions that contain more than one
atom.
 CFCT: Table 5.4 on p. 235.
 Be familiar with these ions.
 Examples include:
 Carbonate (CO3-2)
 Bicarbonate (HCO3-1)
 Phosphate (PO4-3)
 Sulfate (SO4-2)
 Hydroxide (OH-1)
 Nitrate (NO3-1)
 Ammonium (NH +1)

33. Writing Lewis-dot Structures
 SeeRules for “Writing Lewis-Dot
Structures” (p. 203 in Lecture Guide).

34. Determining the Central Atom in
Lewis-Dot Structures of Molecules
 Consider how many free pairs of electrons an atom has
before bonding. The higher this number, the more potential
for bonding. This means it is more likely to be a central atom.
Abridged Periodic Table of Elements – Lewis Dot Structures
1A 2A 3A 4A 5A 6A 7A 8A
H He
Li Be B C N O F Ne
Na Mg Al Si P S Cl Ar

35. Free Radicals
 Molecules/atoms with an unpaired
electron.
Cl N O
 Importance in terms of health: Can cause
damage to tissue in the body.
 Antioxidants are used to counteract free
radicals: Sources include blueberries and
green tea.

36. Types of Molecular Geometries

37. Polar vs. Nonpolar Molecules
Dipole: a molecule that has
unequally distributed charges.
Polar- having unequally distributed charge.
Nonpolar- having equally distributed charge.
Polar Example: Hydrogen
Chloride
δ+ δ- δ+ δ-
H Cl H Cl

38. Water: Notice how the
electron pairs spread
out and cause a bent
geometry. This also
causes the formation
of a dipole, making
water polar.

39. Water as a Polar Molecule
O
H H
Molecular Geometry: Bent

40. Water as a Polar Molecule
δ-
Bond Dipole
moment
O
H H
δ+ δ+
Molecular Geometry: Bent

41. Water as a Polar Molecule
δ-
Overall Dipole
O Moment = Polar
H H
δ+ δ+
Molecular Geometry: Bent

42. Ammonia (NH3): Notice
how the electron
pairs spread out and
cause a pyramidal
geometry. This also
causes the formation
of a dipole, making
ammonia polar.

43. Ammonia as a Polar Molecule
N
H H
H
Molecular Geometry: Pyramidal

44. Ammonia as a Polar Molecule
δ-
N
H H
δ+ H δ+
δ+
Molecular Geometry: Pyramidal

45. Ammonia as a Polar Molecule
δ-
Overall Dipole
Moment = Polar
N
H H
H δ+
δ+
δ+
Molecular Geometry: Pyramidal

46. Methane (CH4): Notice
how the electron pairs
spread out and cause a
tetrahedral geometry.
This causes the
molecule to be nonpolar
since the overall dipole
moments (see +) to
counteract each other.

47. Methane as a Nonpolar
Molecule
δ+
H
δ-
C
H H δ+
δ+ H
δ+
Molecular Geometry: Tetrahedral

48. Methane as a Nonpolar
Molecule
H
C
H H
H
No Dipole Moment = Nonpolar!
No positive or negative side overall
Molecular Geometry: Tetrahedral