1. Molecules and Compounds

2. Bonding
Chemical bonds are forces that cause a group of atoms to behave as a unit.
Bonds result from the tendency of a system to seek its lowest possible energy.
Bond breaking always requires energy, and bond formation always releases energy.

3. Types of Bonds
The type of bonding depends upon the nature of the atoms that are combined.
A metal and a non-metal will form ionic bonds when electrons are transferred from the metal to the non-metal. The resulting attraction between oppositely charged ions forms a stable crystal.
Ionic compounds are usually solids at room temperature, with high melting points. They conduct electricity only when dissolved in water.

4. Types of Bonds
When metals bond with each other, the valence electrons are shared by the atoms in the entire crystal. The electrons are no longer associated with a specific nucleus, and are free to move throughout the sample.
Metallic solids tend to be malleable and ductile. Their melting points can be low (mercury is a liquid at room temperature) to quite high. They conduct electricity and heat.

5. Types of Bonds
When non-metals bond with each other (water, carbon dioxide, nitrogen), electrons are shared between the atoms to form a molecule. When electrons in the bonds are shared between atoms, a covalent bond is formed.
Covalent molecules are often gaseous or liquids at room temperature. The solids tend to be soft and crumbly.

6. Lewis Structures
Lewis Structures, also known as Lewis dot diagrams, show how the valence electrons are arranged among the atoms in the molecule.
For ionic compounds, it shows the end result when the metal loses its electrons to the non-metal.

7. Lewis Structures – Ionic Compounds
A dot (. ) is used to represent each valence electron. Consider sodium chloride. The dot diagram for the sodium atom is: Na
The dot diagram for the chlorine atom is: Cl . : . : :

8. Lewis Structures – Ionic Compounds
Note that electrons (dots) representing valence electrons are written singly or in pairs on either side of the atomic symbol, and above and below the symbol.
Na Cl
It doesn’t matter which side has the unpaired electron. . : . : :

9. Lewis Structures – Ionic Compounds
Sodium donates its single valence electron to the chorine atom, and chlorine accepts the extra electron.
Na Cl
After transfer, both ions have a noble gas electron configuration.
Na+ Cl . : . : : - : : : :

10. Lattice Energy
When metals react with non-metals, a large amount of energy is usually released. These reactions are usually exothermic. The electron transfer from the metal to the non-metal usually requires energy, and is endothermic. The large release of energy comes from forming an ionic crystal.

11. Lattice Energy
The crystal will have oppositely charged ions in contact with each other and like charged ions separated from each other. The alternating charged ions form a crystal lattice. As the lattice forms, large amounts of energy are released.

12. Lattice Energy

13. Lattice Energy
Lattice energy is defined as the energy released when a mole of an ionic solid into its constituent gaseous ions. It is usually a large negative quantity.

14. Lattice Energy
Based on Coulomb’s Law, the lattice energy increases with smaller ionic size.

15. Lattice Energy
Based on Coulomb’s Law, the lattice energy increases significantly with greater ionic charge. E = constant (q1)(q2)/r

16. Melting Points
As ionic crystals melt, the lattice is disrupted, and the ions become free to move, and can conduct electricity.

17. Ionic Bonding
The strength of the attraction between ions increases significantly with increasing ionic charge, and results in high melting points.

18. Ionic vs. Covalent Compounds
Ionic Compounds
Crystalline solids (made of ions)
High melting and boiling points
Conduct electricity when melted
Many soluble in water but not in nonpolar liquids
Covalent Compounds
Gases, liquids, or solids (made of molecules)
Low melting and boiling points
Poor electrical conductors in all phases
Many soluble in nonpolar liquids but not in water

19. Properties of NaCl and CCl4

20. Naming Inorganic Compounds
1. Binary Compounds
Binary compounds contain only two elements. The elements are either a metal with a non-metal (ionic bonding), or two non-metals (covalent bonding).

21. Naming Binary Compounds
a) Metal + Non-metal:
When metals react with non-metals, the metal loses electrons and the non-metal gains electrons. The resulting attraction between oppositely charged ions creates ionic bonds.

22. Common Ionic Charges
The charges of ions of elements in groups 1A-7A (the main groups) are usually predictable.
Group 1A metals form +1 ions, group 2A metals form +2 ions, etc.
The non-metals of group 5A have a -3 charge, those of group 6A have a -2 charge, and the halogens form ions with a -1 charge.

23. Typical Ionic Charges

24. Naming Binary Compounds
For example, NaCl is called
sodium chloride
Where “chlor” is the root for the element chlorine.

25. Naming Binary Compounds
Three common transition metals also have only one ionic charge, and are also named the same way.
They are: zinc ion (always +2), silver ion (+1) and cadmium ion (+2)
ZnS is zinc sulfide, as “sulf” is the root for sulfur.

26. Writing Formulas of Binary Compounds
Compounds have no net charges, so the formulas of ionic compounds must contain equal numbers of positive and negative charges.
Magnesium bromide, made from magnesium ion (Mg2+) and bromide ion (Br1-) has the formula
MgBr2

27. Binary Compounds with Variable Charge Metals
Most transition metals and the metals on the lower right side of the periodic table can have several ionic charges.
The properties of the ion vary greatly with charge, so the charge must be specified in naming the ion or its compounds.

28. Typical Ionic Charges

29. Binary Compounds with Variable Charge Metals

30. Binary Compounds with Variable Charge Metals
If an ion has variable charges, you must specify the charge in naming the metal.
If an ion has only one charge, it is incorrect to specify its charge.

31. Naming Fe2O3
Fe2O3 is an iron oxide, but we must specify the charge of the iron ion.
We know each oxide has a -2 charge, so three oxide ions have a total charge of -6.
The two iron ions therefore have a charge of +6, with each iron having a charge of +3.
The name of the compound is iron(III) oxide.

32. Unusual Ions
Mercury forms two ions, mercury(I) and mercury(II). The mercury(I) ion is polyatomic, and exists as two mercury(I) ions bonded together. Its formula is Hg22+.
Oxygen in compounds usually exists as the oxide ion, O2-. Oxygen also exists as the peroxide ion, O22-, with each oxygen having a -1 charge.

33. Naming Polyatomic Ions
There are many ions, such as sulfate or nitrate, that contain more than one element.
Many of these ions contain oxygen and a non-metal.
These ions can be found in a group of acids called the oxy acids (such as sulfuric acid, nitric acid, etc.).

34. Polyatomic Ions
The bonding within these polyatomic ions (such as nitrate, sulfate and phosphate) is covalent. The ionic charge results from the loss of one or more H ions to water, resulting in a negative charge on the anion formed.
In water, the covalently bonded hydrogen is donated to water, forming hydronium ions and the corresponding anion.

35. Naming the Oxy Acids
The easiest way to learn the names of the ions is to memorize a short list of oxy acid names and their formulas.
The names of the ions are derived from the names of the acids.
Keep in mind that the acids must be aqueous solutions.

36. Common Oxy Acids Acid Name HNO3 Nitric acid H2SO4 Sulfuric acid
HClO3 Chloric acid (or iodic or bromic acid)
H3PO4 Phosphoric acid
H2CO3 Carbonic acid

37. Naming Complex Ions
Once the list of acids is learned, the names of other acids and ions can be derived.
Removal of the hydrogens in the acid as H+ ions results in ions that end in ate.
HNO3 minus one H+ ion gives NO31-, the nitrate ion.
The oxy acids that end in ic, produce ions that end ate.

38. Naming Complex Ions
Sulfuric acid is H2SO4. Removing two H+ ions produces SO42-, the sulfate ion.
Keep in mind that the formula of the ions must include the charge.
If only one of the H+ ions is removed from sulfuric acid, HSO41- is produced. This is called the hydrogen sulfate ion, also commonly known as the bisulfate ion.

39. Naming Complex Ions Carbonic acid, H2CO3, produces two ions:
HCO31-, the hydrogen carbonate or bicarbonate ion
and
CO32-, the carbonate ion

40. Naming Complex Ions
Some of the oxy acids previously listed also exist with one more oxygen in the formula.
HClO3, HBrO3 and HIO3 , in aqueous solution are chloric, bromic and iodic acid respectively.
Adding an oxygen to the formulas provides the formulas for the per root ic acid.
HClO4 is perchloric acid. The ion, ClO41- is the perchlorate ion.

41. Naming Complex Ions
Several of the oxy acids listed previously can have one less oxygen atom in the formula. These acids have names that end in ous, and ions that end in ite.
HNO3 is nitric acid. HNO2(aq) is nitrous acid. The ion NO21- is the nitrite ion.

42. Naming Complex Ions
Sulfuric acid, phosphoric acid, chloric, bromic and iodic acids all can have one less oxygen atom. The acids are sulfurous acid, phosphorous acid, chlorous acid, bromous acid and iodous acid.
The ions are called sulfite, phosphite, chlorite, bromite and iodite ion.

43. Naming Complex Ions
The halogen oxy acids HClO3, HBrO3, and HIO3 also exist with two less oxygen atoms in the formula. The name of the resulting acid has the name
hypo root ous acid.
HClO(aq) is hypochlorous acid, and ClO1- is the hypochlorite ion.

44. Naming Complex Ions
If you memorize the list of acids ending in ic, you can derive the names and formulas for many other acids and ions.
Acid Name
HNO3 Nitric acid
H2SO4 Sulfuric acid
HClO3 Chloric acid (or iodic or bromic acid)
H3PO4 Phosphoric acid
H2CO3 Carbonic acid

45. Naming Complex Ions
In naming the ions from the acids on the list, remember that ic  ate.
If there is one additional oxygen atom, the acid has the name per root ic, and the ion has the name per root ate.
If there is one less oxygen atom, the acid has a name ending in ous. The ions will have names ending in ite. (ous ite)

46. Naming Complex Ions
If an acid has two less oxygen atoms than the “ic” list, its name has the form hypo root ous. The ion will have the name hypo root ite.

47. Other Common Formulas CH3COOH Acetic acid CH3COO1- Acetate ion
NH3 Ammonia
NH4+ Ammonium ion
OH1- Hydroxide ion
H3O+ Hydronium ion
MnO41- Permanganate ion
CrO42- Chromate ion
Cr2O72- Dichromate ion

48. Hydrated Ionic Compounds
Hydrates are compounds that contain a specific number of water molecules associated with each formula unit. The water molecules are called waters of hydration, and can usually be removed by heating the compound. CuSO4.5H2O is copper(II) sulfate pentahydrate

49. Covalent Bonding
Covalent compounds exist as discrete molecules, whereas ionic compounds consist of an aggregate of cations and anions.

50. Covalent Bonds
When two (or more) non-metals form bonds, electrons are shared. The result is a covalent bond.
Covalent bonds form because the attraction of electrons for the nuclei in the atoms is greater than the electron-electron repulsion or the nucleus-nucleus repulsion.

51. Types of Bonds
There is usually an optimum bond length or internuclear distance where attractions between electrons and the nuclei are optimized and repulsions are minimized.

52. Covalent Bonding

53. Bond Energy
Bond formation releases energy, and bond breaking requires energy.

54. Molecules
Molecules are neutral combinations of two or more atoms that are bonded together and act as a unit. Molecules may be elements (H2, O2, O3, F2), or compounds containing atoms of different elements bonded together. Molecules typically contain elements that are non-metals.

55. The Diatomic Elements
Several common elementsprimarily exist in nature as molecules containing two atoms. The diatomic elements are: H2, N2, O2, F2, Cl2, Br2 and I2

56. Lewis Structures – Covalent Molecules
In many covalent molecules, the non-metals share valence electrons. Electrons are shared so that each atom in the molecule has a full valence shell.
H. .H
H:H or H..H or H-H
Each hydrogen has access to two electrons, as does the noble gas helium.

57. Lewis Structures – Covalent Molecules
For elements in period 2, the non-metals generally share enough valence electrons so that each atom obtains the same number of valence electrons as neon (a total of eight electrons).
This may involve making multiple (double or triple) bonds between atoms.

58. Lewis Structures – Covalent Molecules
Provide Lewis structures for elemental nitrogen, oxygen and fluorine.

59. Molecular Shapes
Lewis structures are drawn on a flat piece of paper, many molecules are three dimensional. In addition, the Lewis structure doesn’t necessarily give an accurate idea of the bond angles.

60. Molecular Shapes
Although the Lewis structure suggests a linear molecule, the bond angle in water is approximately 105o.

61. Naming Covalent Binary Compounds
When two non-metals form a compound, they share electrons, rather than transfer them. The resulting bond is called a covalent bond.
The naming of these compounds is fairly simple. The first element is named first, and the second element is named as the root + ide.
Prefixes are used to indicate the number of each atom present.

62. Naming Covalent Binary Compounds
These prefixes are used only for compounds containing two non-metals.
The prefix mono is never used for the first element in the compound.

63. Naming Covalent Binary Compounds
The prefix mono is never used for the first element. CO2 is carbon dioxide.
If the prefix ends in an a or o, and the element that follows begins with a vowel, the last letter of the prefix is usually dropped. N2O5 is called dintrogen pentoxide (and not pentaoxide).

64. Naming Covalent Binary Compounds
Note that these prefixes are only used for binary covalent compounds. It is incorrect to use them for compounds containing a metal and a non-metal.

65. Naming Covalent Binary Compounds
There are some compounds of metalloids or metals in very high oxidation states that are sometimes named using this system.

66. Naming Covalent Binary Compounds
For example, compounds such as SnCl4 or PbCl4 are covalent in nature, and not ionic solids. They may be called tin(IV) chloride or tin tetrachloride [or lead(IV) chloride or lead tetrachloride.

67. Binary Compounds with Hydrogen
With metals, hydrogen can form ionic compounds in which the hydrogen has a -1 ionic charge. These compounds are named like any binary ionic compound. NaH is sodium hydride CaH2 is calcium hydride

68. Binary Compounds with Hydrogen
With non-metals, the bonding is covalent. Hydrogen never forms a positive ion in nature. Many of the compounds containing hydrogen have common names that do not follow the usual nomenclature rules.

69. Binary Compounds with Hydrogen
Examples include: water H2O ammonia NH3 phosphine PH3 hydrogen sulfide H2S Note that the order in which the elements are written is also irregular.

70. Binary Compounds with Hydrogen- Acids
Hydrogen also forms binary compounds that act as acids in water. These compounds dissociate in water to donate a proton to water. HCl(g) + H2O(l)  H3O+(aq) + Cl–(aq) hydrogen chloride hydronium

71. Naming Binary Acids
The naming of the pure compound and its aqueous acid solution differ.
HCl is a gas called hydrogen chloride. HCl(aq) is an acid called hydrochoric acid.

72. Naming Binary Acids
Name the following acids:
H2S(aq) , HBr(aq)

73. Molecules
Scientists studying the nature of matter focused much of their research in 1800s on the composition of compounds. Since molecules are much too small to observe, they typically observed the reactions of larger amounts of matter and used mass measurements to develop their theories of matter.

74. Empirical Formulas
The studies of water and hydrogen peroxide lead to empirical formulas. These are based on experiment, and represent to simplest way of expressing the ratio of atoms in a compound. The early scientists analyzed new chemical compounds to determine their composition and chemical formulas. Modern analytical laboratories still provide this service.

75. Molecular Formulas
Molecular formulas show the exact number of each type of atom in a molecule. For example, hydrogen peroxide has a molecular formula of H2O2. Its empirical formula shows that there is one hydrogen for every oxygen, so it is OH or HO. Neither of these formulas provides the structure or arrangement of the bonds in the molecule.

76. Percent Composition
A chemical formula can be used to calculate the percent composition of a compound. Likewise, the percent composition can be used to determine the empirical formula of a compound. This is extremely useful when trying to determine the formula of a new, or unknown compound.

77. Chemical Composition
The formula obtained for the compound is the simplest whole number ratio of the elements in the compound, or the empirical formula. It may differ from the actual formula. For example, hydrogen peroxide is H2O2, but chemical analysis will provide an empirical formula of HO.

78. Percent Composition
To calculate the % composition of a known compound, you determine the total mass of the molecule, and the mass due to each of the elements in the compound. % by mass of element A = total mass of A (100%) molecular mass

79. Percent Composition
To calculate the % composition of a known compound, you determine the total mass of the molecule, and the mass due to each of the elements in the compound. % by mass of element A = total mass of A (100%) molecular mass

80. % Composition Problem
Problem: Calculate the percent composition of ammonia.

81. Determining Formulas
It is generally more useful to obtain percent composition data (usually from a laboratory), and determine the empirical formula of a compound. This will be the simplest whole number ratio of the elements, and provides no information about the structure of the compound.

82. Determining Empirical Formulas
If given % composition:
1. Assume a quantity of 100 grams of the compound.
2. Determine the number of moles of each element in the compound by dividing the grams of each element by the appropriate atomic mass.
3. To simplify the formula into small whole numbers, divide the moles of each element by the smallest number of moles.

83. Determining Empirical Formulas
4. If necessary, multiply each number of moles by a factor that produces whole number subscripts.
5. If you know the approximate molar mass of the compound, determine the molecular formula.

84. Chemical Composition
Usually, the compound is combusted in the presence of oxygen. Any carbon in the compound is collected as carbon dioxide (CO2), and any hydrogen is collected as water (H2O).

85. % Composition Problem
Caffeine contains 49.48% carbon, 5.15% hydrogen, 28.87% nitrogen, and 16.49% oxygen. The compound has a molar mass of Determine the empirical and molecular formula of caffeine.

86. Determining Empirical Formulas
If given combustion data:
The ultimate goal is to get the simplest whole number ratio of the elements in the compound. Usually the compound contains carbon, hydrogen and perhaps oxygen or nitrogen.
1. Use the information about CO2 to determine the moles and mass of carbon in the compound.

87. Determining Empirical Formulas
If given combustion data:
2. Use the information about H2O to determine the moles and mass of hydrogen in the compound.
3. The mass and moles of oxygen (or a third element) can be obtained by difference.
4. Once moles of each element is obtained, find the relative number of moles and empirical formula (as with % composition).

88. Formulas from Combustion Data
This method assumes the compound contains only C and H.

89. Empirical Formula using Combustion Data- Problem
A compound, which contains C, H and O, is analyzed by combustion. If mg of the compound produces mg of carbon dioxide and 4.37 mg of water, determine the empirical formula of the compound.
If the compound has a molar mass of g/mol, determine the molecular formula of the compound.