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Charles E. Ophardt, Professor of Chemistry, Elmhurst College, Elmhurst, IL 60126, charleso at elmhurst.edu, Copyright 2004
 

TOPIC 3: CHEMICAL COMPOUNDS AND BONDING

ON-LINE Lecture Discussion Requirement: Answer 3 Questions.

First Question: Do one question from the QUES 1 - 9. Student ID number assignments are as follows:

 Questions  Ques. 1  Ques. 2  Ques. 3  Ques. 4  Ques. 5
 Student ID  9, 17, 27  10, 18, 28  1, 19, 29  2, 20, 30  3, 11, 21, 8
 Questions  Ques. 6  Ques. 7  Ques. 8  Ques. 9  
 Student ID  4, 12, 22  5, 13, 23, 16  6, 14, 24, 26  7, 15, 25  


Second Question: Pick one more question from 1-9 - your choice.

A third question may be to do another question of your choice or respond or comment to someone else, or use General Questions to ask a general question the Prof or others might answer.

Check answers already completed in
Blackboard - Discussion Pages

Write out answers to questions in a WORD PROCESSOR
and then copy and paste into
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Requirements for the Lecture On-Line Discussion
Method to list references and citations.

Text Readings: Chap 4

1. Atom and Ion Electron Arrangements

 A. Bohr Diagram Review  Text p. 85-87
ProfONotes: Positive Ions

 Notes: The Bohr model provides the initial insights about the arrangement of electrons in energy levels. The maximum number of electrons in an energy level is calculated as follows:
energy level maximum number of electrons
1 2
2 8
3 18
4 32

A Bohr Diagram is used to show the number and arrangement of electrons in various energy levels. The arrangement of all electrons in a particular atom can be determined by using the periodic table and the principle that electrons fill the lowest energy levels first.
A simple BOHR DIAGRAM can be used to represent the atoms:

For example: Li atom with atomic number 3 has 3 total electrons is: 2e- 1e- in energy levels 1 and 2
For example: Na atom with atomic number 11 has 11 total electrons is: 2e- 8e- 1e- in energy levels 1, 2, and 3
See text for more details.

 B. Electron-Dot Structures or Lewis Symbols  Text p. 93-95
ProfONotes: Lewis Diagrams

Notes: If the previous examples are examined carefully for relationships with the periodic table, the following facts should become evident: lithium, and sodium all have one electron in the outer energy level and these elements are all in Group I.

Electron and Group Number Principle: The number of electrons in the outer energy level can be determined by finding the group number for that element.
The number of electrons in the outer energy level is the fundamental basis for the fact that elements in the same group have similar properties. The formation of ions and compounds are related to how many electrons are in the outer energy level of the atoms.

Lewis Symbols for Elements:
Elemental properties and reactions are determined only by electrons in the outer energy levels. Electrons in completely filled energy levels are ignored when considering properties. Simplified Bohr diagrams which only
consider electrons in outer energy levels are called LEWIS SYMBOLS. A Lewis symbol consists of the element symbol surrounded by "dots" to represent the number of electrons in the outer energy level. The
number of electrons is determined by looking at the Group number.
See text for more details.

QUES. 1: Give the Bohr diagrams (just list the number of electrons for each energy level) and the Lewis Symbols for the following elements: H, C, N, O, F, Ne, Mg, Al, S, Cl. Each student should do 4 different ones and you may pick other elements with atomic numbers from 1 to 18.  

2. Ionic Compounds

 A. Types of Compounds  ProfONotes: Types of compounds
Notes: The molecules may be divided into two groups. Those molecules that consist of charged ions with opposite charges are called IONIC. These ionic compounds are generally solids with high melting points and conduct electrical current. The other type of molecules are called COVALENT and do not consist of ions. Covalent compounds have low melting points and do not conduct electric current.
 B. Formation of Ions  Text p. 95-97
ProfONotes: Positive Ions
ProfONotes: Negative Ions
Notes: The IONS have properties that are completely different from those of their individual element atoms. Whereas elements are neutral in charge, IONS have either a positive or negative charge depending upon whether there is an excess of protons (positive ion) or excess of electrons (negative ion).
 C. Octet Rule  Text p. 95
Notes: Octet Rule: Elemental atoms generally lose, gain, or share electrons with other atoms in order to achieve the same electron structure as the nearest rare gas with eight electrons in the outer level.
 D. Ionic Compounds  Text p. 95-99
ProfONotes: Ionic Compounds
 Notes: The Octet Rule is the basis for the predictions about the charges on ions. Ionic compounds are formed as the result of the formation of positive and negative ions. Electrons are actually transferred from one atom to another to form rare gas electron structures for each ion. The atom which forms a positive ion loses electrons to the atom which gains electrons to form a negative ion. A compound is not stable unless the number of electrons which are lost and gained are equal.

 

QUES. 2: Text p. 113 - Do Exercise 14 or 26 or 27 or 28. Each student pick one. Just describe the process and give the number of electron dots for the starting elements and then the ions.
QUES. 3: Text p. 114 - Do Exercise 37 or 39. Each student pick one. Just describe the process and give the number of electron dots for the starting elements and give the final formula.  


3. Covalent Compounds
 A. Simple Diatomic Molecules Text p. 100
ProfONotes: Types of compounds
  Notes: Covalent compounds share two electrons in forming a bond between atoms. Covalent compounds are formed only by the interactions of non-metal atoms. The number of atoms which make up covalent molecules is determined by the number of electrons in outer levels and the Octet Rule.

The simplest covalent molecule is the diatomic hydrogen molecule. The Lewis Symbols are: H . and H . ; The "octet" for hydrogen is only 2 electrons since the nearest rare gas is He. The diatomic molecule is formed because individual hydrogen atoms containing only a single electron are unstable. Since both atoms are identical a complete transfer of electrons as in ionic bonding is impossible -- How would you determine which atom should lose an electron and which atom should gain an electron? Instead the two hydrogen atoms share both electrons. Other examples are fluorine, chlorine, bromine.
B. Polyatomic Covalent Molecules

Text p. 101-102

ProfONotes: Lewis Diagrams

 Notes: Compounds which need more than one electron to complete the octet will share as many electrons as necessary in order to complete the octet. This translates into forming two or more bonds to atoms. The water molecule is a good example of this. The Lewis Symbols are:

H (1 electron) and O (6 electrons) = H2O

Hydrogen needs only l electron while oxygen needs 2 electrons to complete the octet. Since hydrogen can only share one electron, two hydrogen atoms with one electron each will be needed by oxygen to complete the octet.
C. Multiple Bonds  Text p. 103
Notes: Atoms may share more than 2 electrons. Diatomic oxygen shares 4 electrons and diatomic nitrogen shares 6 electrons.
Example: Determine the formula for an oxygen molecule.
Solution:
Lewis Symbols: O (6 electrons) + O (6 electrons)

Two electrons in each oxygen are needed for an octet, therefore, they need to share a total of 4 electrons between them. This makes two sets of two electrons or two bonds or a double bond.  

QUES. 4: Describe the electrons in bonding for any TWO of the following molecules: sodium chloride, magnesium chloride, potassium oxide chlorine, ammonia, nitrogen, carbon tetrachloride, methane, carbon dioxide, ethylene, and hydrogen chloride. Answers are found somewhere in text chap 4. Text p . 96-105


General Questions and Comments

4. Shapes of Molecules
Molecular Shapes ProfONotes: Molecular Geometry
 Notes: Molecules have three dimensional shapes which may greatly influence the physical and chemical properties. The water molecule is a prime example and will be discussed in detail in a future topic.

QUES. 5:  Read carefully the rules for determining the shapes of the molecules. Do a or b or c. Second and third student does the ones not already done.

a. Explain the difference between H2O and CO2 (both molecules have one central atom with two others attached). Use the ideas of bonding and non-bonding pairs of electrons.

b. Explain the difference between H2O and O3 (both molecules have one central atom with two others attached). Use the ideas of bonding and non-bonding pairs of electrons.

c. Explain the difference between BH3 and NH3 (both molecules have one central atom with three others attached. Use the ideas of bonding and non-bonding pairs of electrons.

 ProfONotes: Molecular Geometry

5. Covalent Compounds - Polar vs. Non-polar

Non-polar bond = equal electron pair sharing; recognize by two identical atoms and carbon-hydrogen is a very common exception. p. 100-101
ProfONotes: Non-Polar Bonds
 Polar bond = unequal electron pair sharing; recognize by two different non-metal atoms  p. 103-104, 106-107
ProfONotes: Polar Bonds

QUES. 6: Classify the following covalent bonds as polar or non-polar: H-H, H-Cl, H-O, H-C, H-N, O-O, N-N, Cl-Cl.

If the molecule is polar, describe which atom has a partial positive and negative charge.

Each student should pick 2 or 3 that have not been done.

In addition, explain how carbon dioxide can have polar bonds, but be a non-polar molecule.

 p. 100-107
ProfONotes: Non-Polar Bonds
ProfONotes: Polar Bonds

6. Compounds - Changes of State

Properties of Solids, Liquids, and Gases  Text p. 4-7
ProfONotes: Physicial Properties
 Changes of State: Solid to liquid, liquid to gas  

Do a or b.
QUES. 7a:
At the molecular level, contrast and explain how the molecules behave or move in the solid, liquid, and gas state.

or

QUES. 7b: At the molecular level, contrast and explain what happens to the molecules when a solid melts OR when a liquid boils or vaporizes. Include statements relating to temperature and energy.

 Text p. 4-7
ProfONotes: Physicial Properties

7. Changes of State - Intermolecular Forces

Ionic Forces - full positive and negative charges in ionic crystals - strongest of all forces. Positive and Negative charges attract. Example: Salt - melts at 800 C. p. 107, 110
ProfONotes: Intermolecular Forces
Dipole Forces - caused by partial positive and negative charges which attract. Example: Hydrogen Chloride - melts at -112 C  p. 107, 110
ProfONotes: Intermolecular Forces
Hydrogen Bonds - a special case of dipole forces but generally stronger. The hydrogen attached to O, N, or F in one molecule is attracted to the O, N, or F or a different molecule. Example: Water - melts at 0 C.  p. 107, 110
ProfONotes: Intermolecular Forces
ProfONotes: Hydrogen Bonding
Dispersion Forces - non-polar molecules do not have any positive or negative charges to hold them together. Very weak attractive forces can be induced by the movement of the electrons in the bonds to set up very very weak positive and negative charges.  p. 107, 110
ProfONotes: Intermolecular Forces

QUES. 8: Compare, describe, and contrast any of the pairs below:

a. ionic vs. dipole forces
b. dipole vs. hydrogen bonds
c. dipole vs. dispersion forces

Each student pick one above.

 Text p. 107, 110

8. Solutions - Intermolecular Forces

Solutions: A solute is dissolved into a solvent to make a homegeneous solution.  p. 177- 178
ProfONotes: Solubility

Solubility Rule: "Likes dissolve Likes".

Ionic and polar substance dissolve in polar water. Examples: salt in water, alcohol in water.

Non-polar substances do not dissolve in water. Example: oil does not dissolve in water.

Non-polar substances dissolve in non-polar solvents. Example: oil dissolves in gasoline

 p. 177- 178
ProfONotes: Solubility

QUES. 9: Explain at the molecular level how salt dissolves in water.  Text p. p. 177- 178
ProfONotes: Solubility
ProfONotes: Solubility of Salts