1. Chapter 5 Lecture presentation
Molecules and Compounds
Catherine E. MacGowan
Armstrong Atlantic State University

2. Elements to Molecules
When two or more elements combine, a molecule is formed.
Elements can be the same.
The molecular compound of oxygen, O2
Elements can be different.
The molecular compound of water, H2O
Molecules are compounds.
The great diversity of substances that we find in nature is a direct result of the ability of elements to form compounds.

3. Hydrogen, Oxygen, and Water
The dramatic difference between the elements hydrogen and oxygen and the compound water is typical of the differences between elements and the compounds that they form.
When two or more elements combine to form a compound, an entirely new substance results.

4. Law of Definite Proportion: Formation of Molecules
A hydrogen–oxygen mixture can have any proportions of hydrogen and oxygen gas.
Water (H2O) has a definite proportion of hydrogen to oxygen.
A water molecule always is composed of two hydrogen atoms to every one oxygen atom.
A ratio of 2 hydrogen atoms : 1 oxygen atom

5. Types of Chemical Bonds
Compounds are composed of atoms held together by chemical bonds.
Chemical bonds result from the attractions between the charged particles (the electrons and protons) that compose atoms.
Chemical bonds are classified into three types:
Ionic
Covalent
Metallic

6. Types of Chemical Bonds: Ionic versus Covalent

7. Ionic Bonds: Between a Metal Atom and Nonmetal Atom
Ionic bonds occur between metals and nonmetals.
They involve the transfer of electrons from one atom to another.
When a metal interacts with a nonmetal, it can transfer one or more of its electrons to the nonmetal.
The metal atom then becomes a cation.
The nonmetal atom becomes an anion.

8. Ionic Bonds: Ionic Compounds
These oppositely charged ions attract one another by electrostatic forces and form an ionic bond.
The result is an ionic compound, which in the solid phase is composed of a lattice (i.e., a regular three- dimensional array of alternating cations and anions).

9. Covalent Bonds: Bonds between Nonmetal Atoms
Covalent bonds occur between two or more nonmetals.
They involve the sharing of electrons between two atoms.
When a nonmetal bonds with another nonmetal, neither atom transfers its electron to the other. Instead, the bonding atoms share some of their electrons.
The covalently bound atoms compose a molecule.
Hence, they are referred to as molecular compounds.
Molecular compounds are composed of atoms covalently bonded to each other.

10. Representing Compounds: Chemical Formulas and Molecular Models
A compound is represented with its chemical formula.
Chemical formula indicates the type and number of each element present in the compound.
Water is represented as H2O.
Carbon dioxide is represented as CO2.
Sodium chloride is represented as NaCl.
Carbon tetrachloride is represented as CCl4.

11. Types of Chemical Formulas
Chemical formulas can generally be categorized into three different types:
Empirical formula
Molecular formula
Structural formula

12. Types of Chemical Formulas
An empirical formula gives the relative number of atoms of each element in a compound.
It is the simplest whole-number (ratio) representation of the type and number of elements present in a molecule.
A molecular formula gives the actual number of atoms of each element in a molecule of a compound.
(a) For C4H8, the greatest common factor is 4. The empirical formula is therefore CH2.
(b) For B2H6, the greatest common factor is 2. The empirical formula is therefore BH3.
(c) For CCl4, the only common factor is 1, so the empirical formula and the molecular formula are identical.

13. Types of Chemical Formulas: Structural
A structural formula is a sketch or diagram of how the atoms in the molecule are bonded to each other.
It uses lines to represent covalent bonds and shows how atoms in a molecule are connected or bonded to each other.
Example: The structural formula for H2O2 is shown below:

14. Problem Solving: Molecular and Empirical Formulas

15. Types of Chemical Formulas: Summary
The type of formula used depends on how much is known about the compound and how much information is to be communicated.
A structural formula communicates the most information.
It conveys the type and actual number as well as the arrangement of the atoms in the molecule.
It’s a “visual” picture of the compound.
A molecular formula conveys the actual type and number of elemental atoms in the compound.
It does not tell you how each of the atoms are bonded to each other.
An empirical formula communicates the least.
It conveys the simplest whole-number (relative) relationship of atom to atom in the molecule.

16. Molecular Models: 3-D Representation of a Molecule
A molecular model is a more accurate and complete way to specify a compound.
A ball-and-stick molecular model represents atoms as balls and chemical bonds as sticks; how the two connect reflects a molecule’s shape.
The balls are typically color-coded to specific elements.

17. Molecular Models: 3-D Representation of a Molecule
In a space-filling molecular model, atoms fill the space between each other to more closely represent our best estimates for how a molecule might appear if scaled to visible size.

18. Ways of Representing a Compound

19. Lewis Structure Model: Representing a Substance’s Valence Electrons
The Lewis Model:
Valence electrons are represented as dots.
Lewis electron-dot structures (Lewis structures) depict the structural formula with its valence electrons.
Lewis structures focus on valence electrons because chemical bonding involves the transfer or sharing of valence electrons between two or more atoms.

20. Lewis Theory and Bonding
The Lewis Model:
It is used mainly to illustrate covalent-bonded molecular compounds but can be used to illustrate simple ionic compounds.
Lewis Symbols:
They can be used to represent the transfer of electrons from a metal atom (group 1A–group 3A metals) to a nonmetal atom, resulting in ions that are attracted to each other and therefore bond.

21. Problem Solving: Using Lewis Structures to Predict Compound

22. Octet Rule: A Guideline for Molecule Formation
When atoms bond, they tend to gain, lose, or share electrons to result in a noble gas–like configuration.
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Nonmetals period 2 elements must obey the octet rule (i.e., eight valence electrons around each atom in the molecule).
Exceptions to the octet rule: expanded octets
They involve the nonmetals elements located in period 3 and below.
Nonmetals (period 3 on down in the periodic table) follow the octet rule when they are not the center atom.
The center atom is the atom in the molecule that the other elements individually bond (attach) to.
When they are the center atom, they can accommodate more than eight electrons.
Using empty valence d orbitals that are predicted by quantum theory

23. Octet Rule: A Guideline for Molecule Formation
Exceptions to the octet rule:
H, Li, Be, and B attain an electron configuration like that of He.
He can have ONLY two valence electrons, a duet.
Li loses its one valence electron.
H shares or gains one electron.
Though it commonly loses its one electron to become H+
Be loses two electrons to become Be2+.
Though it commonly shares its two electrons in covalent bonds, resulting in four valence electrons
B loses three electrons to become B3+.
Though it commonly shares its three electrons in covalent bonds, resulting in six valence electrons

24. Lewis Theory Predictions for Ionic Bonding
Lewis theory predicts the number of electrons a metal atom should lose or a nonmetal atom should gain in order to attain a stable electron arrangement.
The octet rule
Octet rule (guideline) allows us to predict the
formulas of ionic compounds that result;
relative strengths of the resulting ionic bonds based on Coulomb’s law.

25. Ionic Bonding Model versus Reality
Lewis theory
implies that the positions of the ions in the crystal lattice are critical to the stability of the structure.
predicts that moving ions out of position should therefore be difficult, and ionic solids should be hard.
Ionic solids are relatively hard compared to most molecular solids.
implies that if the ions are displaced from their position in the crystal lattice, repulsive forces should occur.
predicts that the crystal will become unstable and break apart. Lewis theory predicts that ionic solids will be brittle.
Ionic solids are brittle. When struck, they shatter.

26. Ionic Bonding Model versus Reality
Lewis theory
implies that, in the ionic solid, the ions are locked in position and cannot move around.
predicts that ionic solids should not conduct electricity.
To conduct electricity, a material must have charged particles that are able to flow through the material.
Ionic solids do not conduct electricity.
implies that, in the liquid state or when dissolved in water, the ions will have the ability to move around.
predicts that both a liquid ionic compound and an ionic compound dissolved in water should conduct electricity.
Ionic compounds conduct electricity in the liquid state or when dissolved in water.

27. Ionic Bonding and the Crystal Lattice
The extra energy that is released comes from the formation of a structure in which every cation is surrounded by anions and vice versa.
This structure is called a crystal lattice.
The crystal lattice is held together by the electrostatic attraction of the cations for all the surrounding anions.
Electrostatic attraction is nondirectional force.
Therefore, there is no ionic molecule.
The chemical formula is an empirical formula, simply giving the ratio of ions based on charge balance.
The crystal lattice maximizes the attractions between cations and anions, leading to the most stable arrangement.

28. Energetics of Ionic Bond Formation: Using NaCl as an Example
The ionization energy of the metal is endothermic.
Na(s) → Na+(g) + 1 e─ ΔH° = +496 kJ/mol
The electron affinity of the nonmetal is exothermic.
½Cl2(g) + 1 e─ → Cl─(g) ΔH° = −244 kJ/mol
Generally, the ionization energy of the metal is larger than the electron affinity of the nonmetal; therefore, the formation of the ionic compound should be endothermic.
But the heat of formation of most ionic compounds is exothermic and generally large. Why is this?
Na(s) + ½Cl2(g) → NaCl(s) ΔH°f = −411 kJ/mol

29. Crystal Lattice and Lattice Energy of NaCl
The extra stability that accompanies the formation of the crystal lattice is measured as the lattice energy.
It is the energy released when the solid crystal forms from separate ions in the gas state.
Always exothermic
Lattice energy is measured directly but is calculated from knowledge of other processes.
It depends directly on the size of charges and inversely on distance between ions.

30. Conductivity of NaCl
In NaCl(s), the ions are stuck in position and not allowed to move to the charged rods.
In NaCl(aq), the ions are separated and allowed to move to the charged rods.

31. Ionic Compounds
Ionic compounds are composed of cations (metals) and anions (nonmetals) bound together by ionic bonds.
Examples of ionic compounds:
NaBr, Al2(CO3)3, CaHPO4, and MgSO4
The basic unit of an ionic compound is the formula unit, the smallest, electrically neutral collection of ions.
Example:
The ionic compound table salt, with the formula unit NaCl, is composed of Na+ and Cl– ions in a one-to-one ratio.
Summarizing Ionic Compound Formulas
Ionic compounds always contain positive and negative ions.
In a chemical formula, the sum of the charges of the positive ions (cations) must equal the sum of the charges of the negative ions (anions).
The formula of an ionic compound reflects the smallest whole-number ratio of ions.

32. Naming Ionic Compounds
Ionic compounds can be categorized into two types, depending on the metal in the compound.
The first type contains a metal whose charge is invariant from one compound to another.
Whenever the metal in this first type of compound forms an ion, the ion always has the same charge.

33. Metals with Invariant
Cation Charges
Common Nonmetal Anions

34. Naming Binary Ionic Compounds of Type I Cations
Binary compounds contain only two different elements. The names of binary ionic compounds take the following form:
For example, the name for KCl consists of the name of the cation, potassium, followed by the base name of the anion, chlor, with the ending -ide.
KCl is potassium chloride.
The name for CaO consists of the name of the cation, calcium, followed by the base name of the anion, ox, with the ending -ide.
CaO is calcium oxide.

35. Problem Solving: Naming Type I Ionic Compounds

36. Problem Solving: Naming Type II Ionic Compounds

37. Multivalent Metals: Naming Type II Ionic Compounds
The metals in this category tend to have multiple charges (i.e., multivalent cations):
Their charge cannot be predicted as in the case of most representative elements and must be noted in their name.
Multivalent metals
Transition and inner transition metals
Iron (Fe) forms a 2+ cation in some of its compounds and a 3+ cation in others.
FeSO4: Here iron is a +2 cation (Fe2+).
Fe2(SO4)3: Here iron is a +3 cation (Fe3+).
Many of the p-block metals
Not all p-block metals are multivalent.
Some main-group metals, such as Pb, Tl, and Sn, form more than one type of cation.

38. Type II Cation

39. Naming Type II Binary Ionic Compounds
For these types of metals, the name of the cation is followed by a roman numeral (in parentheses) that indicates the charge of the metal in that particular compound.
For example, we distinguish between Cu2+ and Cu+ as follows:
Cu2+ Copper(II)
Cu+ Copper(I)
The full names for compounds containing metals that form more than one kind of cation have the following form:
Cu2O Copper(I) oxide
CuO Copper(II) oxide

40. Naming Type II Binary Ionic Compounds—Example: CrBr3
To name CrBr3,determine the charge on the chromium.
Total charge on cation + total anion charge = 0
Cr charge + 3(Br– charge) = 0
Since each Br has a –1 charge, then
Cr charge + 3(–1) = 0
Cr charge + (–3) = 0
Cr = +3
Hence, the cation Cr3+ is called chromium(III), and Br– is called bromide.
The name for CrBr3 is chromium(III) bromide.

41. Naming Type II Binary Ionic Compounds—Example: SnCl2
To name SnCl2, determine the charge on the tin.
Total charge on cation + total anion charge = 0
Sn charge + 2(Cl– charge) = 0
Since each Cl has a –1 charge, then
Sn charge + 2(–1) = 0
Sn charge + (–2) = 0
Sn = +2
Hence, the cation Sn2+ is called tin(II), and Cl– is called chloride.
The name for SnCl2 is tin(II) chloride.

42. Problem Solving: Naming Type II Ionic Compounds

43. Oxyanions
Most polyatomic ions are oxyanions, anions containing oxygen and another element.
Notice that when a series of oxyanions contains different numbers of oxygen atoms, the oxyanions are named according to the number of oxygen atoms in the ion.
If there are two ions in the series,
the one with more oxygen atoms has the ending -ate; and
the one with fewer has the ending -ite.
For example,
NO3– is nitrate SO42– is sulfate
NO2– is nitrite SO32– is sulfite

44. Oxyanions
If there are more than two ions in the series, then the prefixes hypo-, meaning less than, and per-, meaning more than, are used.
ClO– hypochlorite BrO– hypobromite
ClO2– chlorite BrO2– bromite
ClO3– chlorate BrO3– bromate
ClO4– perchlorate BrO4– perbromate

45. Naming Ionic Compounds Containing Polyatomic Ions
Ionic compounds that contain a polyatomic ion rather than a simple anion (e.g., Cl–) are named in the same manner as binary ionic compounds, except that the name of the polyatomic ion used.
For example, NaNO2 is named according to
its cation, Na+, sodium; and
its polyatomic anion, NO2–, nitrite.
Hence, NaNO2 is sodium nitrite.

46. Common Polyatomic Ions

47. Problem Solving: Naming Ionic Compounds Containing Polyatomic Anions

48. Problem Solving: Naming Ionic Compounds Containing Polyatomic Anions

49. Hydrated Ionic Compounds
Hydrates are ionic compounds containing a specific number of water molecules associated with each formula unit.
For example, the formula for epsom salts is MgSO4 • 7H2O.
Its systematic name is magnesium sulfate heptahydrate.
Another example: CoCl2 • 6H2O is cobalt(II) chloride hexahydrate.

50. Hydrates Common hydrate prefixes
hemi = ½
mono = 1
di = 2
tri = 3
tetra = 4
penta = 5
hexa = 6
hepta = 7
octa = 8
Other common hydrated ionic compounds and their names are as follows:
CaSO4 • ½H2O is called calcium sulfate hemihydrate.
BaCl2 • 6H2O is called barium chloride hexahydrate.
CuSO4 • 6H2O is called copper sulfate hexahydrate.

51. Covalent Bonding: Bonding and Lone Pair Electrons
Electrons that are shared by atoms are called bonding pairs.
Electrons that are not shared by atoms but belong to a particular atom are called lone pairs.
Also known as nonbonding pairs

52. Single Covalent Bonds
When two atoms share one pair of electrons, the result is called a single covalent bond.
Two electrons
One atom may use more than one single bond to fulfill its octet.
To different atoms
H only duet

53. Double Covalent Bond
When two atoms share two pairs of electrons, the result is called a double covalent bond.
Four electrons between the two atoms
Example: O2
Elements that can double-bond with each other and themselves are C, N, O, S, and P.

54. Triple Covalent Bond
When two atoms share three pairs of electrons, the result is called a triple covalent bond.
Six electrons between the two atoms
Example: N2
Elements that can triple-bond with each other and themselves are C, N, O, and S.

55. Covalent Bonding: Model versus Reality
Lewis theory
implies that some combinations should be stable, whereas others should not.
Stable combinations result in “octets.”
allows us to predict the formulas of molecules of covalently bonded substances.
Hydrogen and the halogens are all diatomic molecular elements, as predicted by Lewis theory.
Oxygen generally forms either two single bonds or a double bond in its molecular compounds.
There are some stable compounds in which oxygen has one single bond and another in which it has a triple bond, but it still has an octet.

56. Covalent Bonding: Model versus Reality
Lewis theory of covalent bonding
implies that the attractions between atoms are directional.
The shared electrons are most stable between the bonding atoms.
predicts that covalently bonded compounds will be found as individual molecules.
Rather than an array like ionic compounds
Compounds of nonmetals are made of individual molecule units.

57. Problem Solving: Ionic versus Molecular Compounds

58. Molecular Compounds: Formulas and Names
The formula for a molecular compound cannot readily be determined from its constituent elements because the same combination of elements may form many different molecular compounds, each with a different formula.
Nitrogen and oxygen form all of the following unique molecular compounds: NO, NO2, N2O, N2O3, N2O4, and N2O5.

59. Molecular compounds are composed of two or more nonmetals.
Names of Molecular Compounds:
Write the name of the element with the smallest group number first.
If the two elements lie in the same group, then write the element with the greatest row number first.
The prefixes given to each element indicate the number of atoms present.

60. Binary Molecular Compounds
These prefixes are the same as those used in naming hydrates:
mono = 1 hexa = 6 di = 2 hepta = 7 tri = 3 octa = 8 tetra = 4 nona = 9 penta = 5 deca = 10
If there is only one atom of the first element in the formula, the prefix mono- is normally omitted.

61. Problem Solving: Naming Molecular Compounds

62. Formula Mass/Molecular Mass of a Compound
The mass of an individual molecule or formula unit is known as molecular mass or molecular weight of the compound.
It is the mass of ONE MOLE of that compound.
Determining a Compound’s Molecular Mass:
Sum of the masses of the atoms in a single molecule or formula unit
Example: What is the molecular mass of water (H2O)?
2(1.01 amu H) amu O = amu
One mole of water has a molecular mass of grams.

63. Problem Solving: Calculating Formula Mass

64. Using Molar Mass to Count Molecules by Weighing
Molar mass in combination with Avogadro’s number can be used to determine the number of atoms in a given mass of the element.
Use molar mass to convert to the amount in moles. Then use Avogadro’s number to convert to number of molecules.

65. Problem Solving: Using the Mole Concept to Convert between Mass and Number of Particles

66. Composition of Compounds
A chemical formula, in combination with the molar masses of its constituent elements, indicates the relative quantities of each element in a compound.
The percentage of each element in a compound can be determined from
the formula of the compound; and
the experimental mass analysis of the compound.
molecular mass of element Z
% mass of element Z = × 100%
mass of 1 mol of compound
The percentages may not always total to 100% due to rounding.

67. Problem Solving: Mass Percent Composition

68. Conversion Factors from Chemical Formula
Chemical formulas contain within them inherent relationships between numbers of atoms and molecules.
Or moles of atoms and molecules
These relationships can be used to determine the amounts of constituent elements and molecules.
Such as percent composition

69. Problem Solving: Using Chemical Formulas as Conversion Factors

70. Problem Solving: Using Chemical Formulas as Conversion Factors

71. Determining a Chemical Formula from Experimental Data
Empirical Formula:
Simplest whole-number ratio of the atoms of elements in a compound
Can be determined from elemental analysis
Masses of elements formed when a compound is decomposed, or that react together to form a compound
Combustion analysis
Percent composition
Note: An empirical formula represents a ratio of atoms or a ratio of moles of atoms, not a ratio of masses.

72. Finding an Empirical Formula
1. Convert the percentages to grams.
If not given, assume you start with 100 g of the compound.
Example: 24.5% C means 24.5 g C.
2. Convert mass (grams) to moles.
Use molar mass of each element.
Example: 24.5 g C × (1 mol C/12.01 grams) = 2.00 mol C
3. Divide all by the smallest number of moles to obtain the atom-to-atom ratio for each of the elements in the compound.
If the result is within 0.1 of a whole number, round to the whole number.
4. Multiply all mole ratios by a number to make all whole numbers.
If ratio is .5, multiply all by 2; if the ratio is .33 or .67, multiply all by 3; and so on.
b) Skip if already whole numbers.

73. Problem Solving: Obtaining Empirical Formula from Experimental Data

74. Problem Solving: Obtaining Empirical Formula from Experimental Data

75. From Empirical to Molecular Formulas for Compounds
The molecular formula is a multiple of the empirical formula.
It is the actual formula of the compound.
Knowing the molecular formula, you can determine the molecular mass of the compound.
To determine the molecular formula, you need to know the empirical formula and the molar mass of the compound.
Molecular formula = (empirical formula)n,
where n is a positive integer.

76. Molecular Formulas for Compounds
The molar mass is a whole-number multiple of the empirical formula molar mass, the sum of the masses of all the atoms in the empirical formula:
molar mass
empirical formula molar mass
n =

77. Problem Solving: Obtaining Empirical Formula from Experimental Data

78. Combustion Analysis
A common technique for analyzing compounds is to burn a known mass of compound and weigh the amounts of product made.
This is generally used for organic compounds containing C, H, or O.
By knowing the mass of the product and composition of constituent element in the product, the original amount of constituent element can be determined.
All the original C forms CO2, the original H forms H2O, and the original mass of O is found by subtraction.
Once the masses of all the constituent elements in the original compound have been determined, the empirical formula can be found.

79. Problem Solving: Obtaining Empirical Formula from Combustion Analysis Experimental Data

80. Problem Solving continued

81. Organic Compounds
Early chemists divided compounds into two types: organic and inorganic.
Compounds from living things were called organic; compounds from the nonliving environment were called inorganic.
Organic compounds are easily decomposed and could not be made in the lab.
Inorganic compounds are very difficult to decompose but are able to be synthesized.

82. Modern Organic Compounds
Today organic compounds are commonly made in the lab and we find them all around us.
Organic compounds are mainly made of C and H, sometimes with O, N, P, S, and trace amounts of other elements
The main element that is the focus of organic chemistry is carbon.

83. Carbon Bonding Carbon atoms bond almost exclusively covalently.
Compounds with ionic bonding C are generally inorganic.
When C bonds, it forms four covalent bonds:
four single bonds, two double bonds, one triple bond and one single bond, etc.
Carbon is unique in that it can form limitless chains of C atoms, both straight and branched, and rings of C atoms.

84. Carbon Bonding of Organic Molecules

85. Common Hydrocarbons