1. Bettelheim, Brown, Campbell and Farrell Chapter 3
Chemical Bonds
Bettelheim, Brown, Campbell and Farrell
Chapter 3

2. Ionization Energy
Energy required to remove outermost electron (most loosely held)

3. Noble Gas Configurations
Noble gas configuration s2p6 very stable

4. The Octet Rule
Octet rule: Group 1A-7A elements to achieve an outer shell of eight valence electrons
Anion: Negative ion formed when an atom gains electrons
Cation: Positive ion formed when an atom loses electrons

5. The Octet Rule—Cations
Cation: Sodium atom loses an electron to form a sodium ion, which has the same electron configuration as neon
Na (11 electrons): 1s2 2s2 2p6 3s1
Na+ (10 electrons): 1s2 2s2 2p6

6. The Octet Rule—Anions
Anion: Chlorine atom gains an electron to form a chloride ion, which has the same electron configuration as argon
chlorine atom (17 electrons): 1s2 2s2 2p6 3s2 3p5
chloride ion (18 electrons): 1s2 2s2 2p6 3s2 3p6

7. Cation Names Groups 1A, 2A, and 3A
The name of the element followed by the word “ion”

8. Transition Metal Cations
Cations derived from transition and inner transition elements more than one type of cation
Stock System (IUPAC):
Use Roman numerals to show charge:
Fe2+ is Iron (II) Fe3+ is Iron (III)
Cu+ is Copper (I) Cu2+ is Copper (II)
Old System:
Use the suffix -ous to show the lower positive charge and the suffix -ic to show the higher positive charge
Fe2+ is Ferrous Fe3+ is Ferric
Cu+ is Cuprous Cu2+ is Cupric

9. Transition Metal Ion Names

10. Anion Names
Add “ide” to the root name of the element

11. Polyatomic Ions Contain two or more atoms
Common names often used (in parentheses)

12. Naming Ionic Compounds
Name the positive ion first, then the negative ion
Number of each ion not included
NaBr Al2O3
MgSO4 K2S
(NH4)3PO4

13. Naming Ionic Compounds
NaBr Sodium bromide
Al2O3 Aluminum oxide
MgSO4 Magnesium sulfate
K2S Potassium sulfide
(NH4)3PO4 Ammonium phosphate

14. Formulas of Ionic Compounds
The total number of positive charges must equal the total number of negative charges
Li+ and Br- form LiBr (+1) + (-1)
Ba2+ and I- form BaI2 (+2) + 2(-1)
Al3+ and S2- form Al2S3 2(+3) + 3(-2)

15. Forming Chemical Bonds
Ionic bond: the force of electrostatic attraction between a cation and an anion
Atom loses or gains electrons to make a filled valence shell (octet) and become an ion.
Covalent bond: a pair of electrons that are shared by two atoms
Atom shares electrons to make a filled valence shell (octet)

16. Forming an Ionic Bond--NaCl
Formation of sodium chloride, NaCl
Single-headed curved arrow used to show the transfer of the electron

17. Ionic Bonds Force of attraction between a cation and an anion.
Depends on electronegativity
measure of an atom’s attraction for shared pair of electrons in chemical bond with another atom)

18. Covalent Bonds
Result of one or more pairs of electrons that are shared by two atoms
Each atom has full valence shell (octet)
In H2, each hydrogen contributes one electron to the single bond

19. Molecular Compounds Molecular compound: only covalent bonds
Naming molecular compounds
the less electronegative element is named first (it is generally written first in the formula)
prefixes “di-”, tri-”, etc. are used to show the number of atoms of each element; the prefix “mono-” is generally omitted
Exception: carbon monoxide
NO is nitrogen oxide (nitric oxide)
SF2 is sulfur difluoride
N2O is dinitrogen oxide (laughing gas)

20. Electronegativity F has highest value Noble gases have 0 value
Ionic bonds form when electronegativity difference ≥ 1.9

21. Polarity of Bonds Nonpolar: Electrons are shared equally
Polar: Electrons are NOT shared equally

22. Polarity of Covalent Bonds

23. Polarity of Covalent Bonds
More electron density shown by δ- or the head of a crossed arrow
Less electron density shown by δ+ or the tail of a crossed arrow

24. Polarity of Molecules Polar molecule
has polar bonds, and
Has partial positive and partial negative charges in different parts of molecule, i.e., is a dipole (has two poles)
Carbon dioxide, CO2, has two polar bonds but, because of its geometry, the pulls balance out so it is a nonpolar molecule

25. Polarity of Molecules
Water, H2O, has two polar bonds and, because of its geometry, is a polar molecule

26. Polarity of Molecules
Both dichloromethane, CH2Cl2, and formaldehyde, CH2O, have polar bonds and are polar molecules

27. Lewis Structures
Used to decide on the arrangement of atoms in the molecule
Bonding (shared) electrons are shown as bonds (lines)
Nonbonding electrons are represented as a pair of Lewis dots

28. Drawing Lewis Structures
1. Determine the number of valence electrons in the molecule
2. Decide on the arrangement of atoms in the molecule
3. Connect the atoms by single bonds
4. Show bonding electrons as a single line; show nonbonding electrons as a pair of Lewis dots
5. In a single bond, atoms share one pair of electrons; in a double bond, they share two pairs, and in a triple bond they share three pairs.

29. Exceptions to the Octet Rule
H and He have a maximum of 2 electrons (duet)
Period 2 elements have a maximum of 8 electrons (use 2s and 2p orbitals)
Atoms of period 3 elements may have more than 8 electrons

30. Lewis Structures
Examples: (the number of valence electrons is given in parentheses after the molecular formula

31. Lewis Structures
Examples
NH3
CH3OH
CH3COOH

33. .

34. Valence-Shell Electron-Pair (VSEPR) Model
Because like charges repel each other, the various regions of electron density around an atom spread so that they are as far away from each other as possible
CH4: measured H-C-H bond angles are 109.5°

35. Summary of Molecular Geometry
# Electron Lone Bond Angle Shape
regions pairs
o Linear
o Planar
Angular (bent)
o Tetrahedral
Trigonal pyramidal
Angular (bent)

37. VSEPR Model the measured H-N-H bond angles are 107.3°
the unshared pair is not shown on this model
the measured H-O-H bond angle is 104.5°