1. Chapter 4 Atoms and Chemical Bonding

2. Continuous Spectra Roy G. Biv Red Orange Yellow Green Blue Indigo
Violet

3. Line Spectrum

4. Model of an Atom
Must explain many things, one of which is the line spectra
First to come up with a decent model was Bohr
Visualized the electrons in an atom orbiting around the nucleus like a planet around a sun
Each orbit was associated with a definite energy (no in-between levels)
Took a “quantum” of energy (fixed amount) to move from one level to another

5. Electron Configuration of the Atom Bohr (1885-1962)

6. Electron Energy Levels

7. Bohr’s Model
Energy levels each can hold a different (maximum) number of electrons
Energy level (n) = 1, 2, 3, 4, …
# electrons (2n2)= 2, 8, 18, 32, …
Example: Na has 11 electrons
– 2e – 8e – 1e
(1st) (2nd) (3rd) -> has 17 empty spots in 3rd level

8. Bohr’s Model Worked well for hydrogen
Worked okay for about the next 20 atoms
Didn’t work well at all for anything past the transition metals
Need new model…

9. Wave Mechanical Model of the Atom Schrodinger (1881-1961)
Retains Bohr’s energy level concept
Distinguishes orbitals at each energy level
Orbitals identified as: s, p, d, f

10. Orbitals in the First Four Energy Levels in the Wave-Mechanical Model of the Atom

11. Building Atoms with Wave-Mechanical Model
Electrons are added starting at the first level
Superscripts are used to indicate the number of electrons in each orbital

12. Electron Arrangements of the First 20 Elements in the Periodic Table

14. Energy Levels of Oribitals

15. An Aufbau Diagram

16. Using the Periodic Table to Determine Electron Configuration

17. The Shape of an s Orbital

18. The Shape of a p Orbital

19. Chemical Properties and the Periodic Table
Chemical properties and electron configuration correlate
Alkali metals all have one s electron in their highest energy level
Li 1s2 2s1
Na 1s22s22p63s1
K 1s22s22p63s23p64s1
Rb 1s22s22p63s23p64s23d104p65s1

20. Valence Electrons Lewis (1916)
Valence electrons are the electrons in the atom’s highest numbered energy level.

21. Octet Rule
Atoms tend to gain or lose electrons to have eight valence electrons
Same as noble gases
He and H are exceptions, get only two valence electrons

22. Stable Electron Configurations
Valence electrons – outermost level with electrons
Core electrons – all other electrons in an atom
Isoelectronic – same number of valence electrons
Example: O2-, F-, and Ne all have 18 e-’s

23. Electron-Dot Structures
Valence electrons represented by dots
Electron-dot symbols
– Examples: Na•, •Mg•, …

24. Lewis Dot Structures
The Lewis dot representation (or Lewis dot formulas)
convenient bookkeeping method for valence electrons
electrons that are transferred or involved in chemical bonding

25. Chemical Bonds
Forces responsible for holding together atoms in molecules and ions in crystals
Determine shape of molecules
Predict chemical and physical properties of materials
Related to arrangement of electrons in compounds
The electrons involved in bonding are usually those in the outermost (valence) shell.

26. Chemical Bonding Chemical bonds are classified into two types:
Ionic bonding results from electrostatic attractions among ions, which are formed by the transfer of one or more electrons from one atom to another.
Covalent bonding results from sharing one or more electron pairs between two atoms.

27. Ionic Bonding Remember cations or positive (+) ions
atoms have lost 1 or more electrons
anions or negative (-) ions
atoms have gained 1 or more electrons

28. Ionic Bonding Atoms lose or gain electrons to form ions
Cations are positive ions
Anions are negative ions
Ionic compounds are held together by electrostatics- the positive charge of the cation attracting the negative charge of the anion.

29. Ionic Bonding Continued

30. Ionic Bonds Na+ and Cl– Opposite charges attract
Ions organize themselves in orderly manner
Crystal of NaCl

31. Structures of Ionic Compounds
extended three dimensional arrays of oppositely charged ions
high melting points because coulomb force is strong

32. Ionic Bonding
We can also use Lewis formulas to represent the neutral atoms and the ions they form.

33. Naming Ions For cations, simple positive ions Add the word ion
Examples: Na+ – sodium ion
Al3+ – aluminum ion
For anions, simple negative ions
Change the usual ending to -ide
Examples: Cl– – chloride
S2– – sulfide

35. Ionic Compound
Ionic compounds are formed primarily when metals on the left side of the periodic table react with nonmetals on the right side of the periodic table.
Transition metals also form ionic compound
Their behavior is less predictable
Iron forms Fe2+ or Fe3+
Copper forms Cu+ or Cu2+

36. Naming Binary Ionic Compounds
Two components in compound

37. Common Ions and Their Position in the Periodic Table

38. Polyatomic Ions
Polyatomic means “many-atom” ion
Memorize These

39. Polyatomic Ions “Many-atom” Ions Example: NH4+, OH-, CH-
Covalently bonded groups of atoms, that tend to stay together
Na O H O H Na+

40. Equations + - + - - +

41. Polyatomic Ions
Charged groups of atoms that remain together through most chemical reactions

42. Covalent Bonds Bond formed by a shared pair of electrons
Gives atom an octet of electrons
Shared pair of electrons – bonding pair
Other electrons not involved in bonding – nonbonding pairs

43. Covalent Bonding covalent bonds formed when atoms share electrons
share 2 electrons - single covalent bond
share 4 electrons - double covalent bond
share 6 electrons - triple covalent bond
attraction is electrostatic in nature
lower potential energy when bound

44. Covalent Bonding extremes in bonding: pure covalent bonds
electrons equally shared by the atoms
pure ionic bonds
electrons are completely lost or gained by one of the atoms
most compounds fall somewhere between these two extremes

45. Covalent Bonds
How do two identical atoms bond to form molecules such as H2, N2, or O2?
Neither atom is more likely than the other to transfer an electron
The two atoms have to share electrons

46. Covalent Bonding

47. Number of Covalent Bonds/Element
Follow the electron-dot rules for the following elements

48. Covalent Bonding
Lewis dot representation
H molecule formation

49. Covalent Bonding
Lewis dot representation
H molecule formation
HCl molecule formation

50. Lewis Dot Structures
homonuclear diatomic molecules
hydrogen, H2
fluorine, F2
nitrogen, N2

51. Lewis Dot Structures
heteronuclear diatomic molecules
hydrogen halides
hydrogen fluoride, HF
hydrogen chloride, HCl
hydrogen bromide, HBr

52. Lewis Dot Structures
water, H2O

53. Lewis Dot Structures
ammonia molecule , NH3

54. Lewis Dot Structures
ammonium ion , NH4+

55. N - A = S rule
N = # of electrons needed to be noble gas
8 for everything except H or He
Only 2 for H or He
A = # of electrons available in outer shells
equal to group #
8 for noble gases

56. N - A = S rule
for ions
add one e- for each negative charge
subtract one e- for each positive charge
central atom in a molecule or polyatomic ion is determined by:
atom that requires largest number of electrons
for two atoms in same group - less electronegative element is central

57. Drawing Lewis Dot Formulas
Example: Write Lewis dot and dash formulas for hydrogen cyanide, HCN.
N = = 18
A = = 10
S =

58. Drawing Lewis Dot Formulas
Example: Write Lewis dot and dash formulas for hydrogen cyanide, HCN.
N = = 18
A = = 10
S =

59. Drawing Lewis Dot Formulas
Example: Write Lewis dot and dash formulas for the sulfite ion, SO32-.
N = 8 + 3(8) = 32
A = 6 + 3(6) + 2 = 26
S = 6

60. Drawing Lewis Dot Formulas
Example: Write Lewis dot and dash formulas for the sulfite ion, SO32-.
N = 8 + 3(8) = 32
A = 6 + 3(6) + 2 = 26
S = 6

61. Drawing Lewis Dot Formulas
Example: Write Lewis dot and dash formulas for sulfur trioxide, SO3.
N = 8 + 3(8) = 32
A = 6 + 3(6) = 24
S = 8

62. Naming Covalent Compounds
Use prefixes to indicate the number of each kind of atom
Examples:
Carbon Monoxide
Carbon Dioxide
Trinitrogen Pentoxide

63. Polar and Nonpolar Covalent Bonds
electrons are shared equally
symmetrical charge distribution
must be the same electronegativity to share exactly equally (typically by being the same element) H2 N2

64. Polar and Nonpolar Covalent Bonds
unequally shared electrons
assymmetrical charge distribution
different electronegativities

65. Electronegativity Pauling (1901-1994)
Electronegativity is the relative tendency of an atom in a molecule to attract a shared pair of electrons in a bond to itself.
The most electronegative element is fluorine and it is given a value of 4.0.
The higher the electronegativity value of an atom, the greater is the ability of an atom of that element to attract electrons to itself.

66. Electronegativity Values for the Representative Elements

67. Polar Molecules
When hydrogen and chlorine react to form HCl, a polar molecule is formed.

68. Continuous Range of Bonding Types
all bonds have some ionic and some covalent character
HI is about 17% ionic
greater the electronegativity differences the more polar the bond

69. Polar Covalent Bonds
If elements do not have the same electronegativity, they get unequal sharing of electrons

70. Polar and Nonpolar Covalent Bonds

71. Polar and Nonpolar Covalent Bonds
Electron density map
of HF
blue areas – low
electron density
red areas – high
Polar molecules have separation of centers of negative and positive charge

72. Polar and Nonpolar Covalent Bonds

73. Polar and Nonpolar Covalent Bonds
Electron density map
of HI
blue areas – low
electron density
red areas – high
Notice that the charge separation is not as big as for HF
HI is only slightly polar

74. Molecule with an overall partial charge
Polar Molecules
Molecule with an overall partial charge
Can have polar bonds and be non-polar molecule
O = C = O O
H H
Overall charge - Asymmetrical
No overall charge - Symmetrical

75. Polar Molecules Continued
Carbon Dioxide, CO2 has polar bonds but because of its symmetrical shape it is a nonpolar molecule.
O=C=O

76. Intermolecular Forces
Glue that holds matter together
Melting and boiling points measure the relative strength
Ionic forces – strongest
Found in salts
NaCl melts at 800°C

77. Hydrogen Bonds

78. Hydrogen Bonding Hydrogen must be attached to electronegative atom
N, O, F
Plays important role in biological systems

79. Hydrogen Bonding Particularly strong dipole-dipole interaction
Occurs between Hydrogen and:
Oxygen, Nitrogen and Flourine
Why? High electronegativity of O, N, F

80. Dipole Moments in molecules some nonpolar molecules have polar bonds
2 conditions must hold for a molecule to be polar

81. Dipole Moments
There must be at least one polar bond present or one lone pair of electrons.
The polar bonds, if there are more than one, and lone pairs must be arranged so that their dipole moments do not cancel one another.

82. Dipole Forces Not as strong as ionic forces Must have polar molecule
HCl melts at –112°C

83. Dipole-Dipole Interaction
Caused by permanent dipoles (partial charges) on the molecule + - + - Cl H Cl H
Cl H
Cl H
 +
 +

84. London Forces

85. Dispersion Forces Present in all molecules
Weak momentary attractive forces
Arise for electrons moving about in molecules and atoms
Strong in larger molecules
Important in nonpolar compounds

86. HW Suggestion
Ch 4:
3, 4, 6, 9, 11, 13, 18, 19, 20, 23, 24, 27, 29, 33, 34, 35, 37, 39, 40, 43, 45, 46, 49, 53, 55
Wednesday: Lab! Bring finished prelab writeup, wear closed-toe shoes, hair up, bring goggles, etc…!