1. Chemistry 140 Fall 2002
Reserve Material
A copy of the course textbook, selected solutions manual, and study guide are available in the library in the course reserve section (behind the main desk).
They are available on loan for 2 hours at a time in order to give everyone access to them (no overnight sign out).
Often do not specify Z when writing. For example 14C, C specifies Z = 12.
Special names for some isotopes. For example hydrogen, deuterium, tritium.
Chemistry 140 Fall 2002 Dutton
Prentice-Hall © 2002

2. Things To Remember A= mass number Z = atomic number
Particle Mass Charge
kg (SI) amu Coulombs (SI) (e)
Electron x –1.602 x –1
Proton x x
Neutron x
Chemistry 140 Fall 2002 Dutton
Prentice-Hall © 2002

3. The Periodic Table Noble Gases Alkali Metals Main Group
Alkaline Earths
Halogens
Transition Metals
Lanthanides and Actinides
Chemistry 140 Fall 2002 Dutton
Prentice-Hall © 2002

4. The Periodic Table Read atomic masses.
Read the ions formed by main group elements.
Read the electron configuration.
Learn trends in physical and chemical properties.
We will discuss these in detail in Chapter 10.
Chemistry 140 Fall 2002 Dutton
Prentice-Hall © 2002

5. The Mole Physically counting atoms is impossible.
We must be able to relate measured mass to numbers of atoms. Similar to:
buying nails by the pound.
buying eggs by the dozen.
buying beer by the case.
using atoms by the gram.
Chemistry 140 Fall 2002 Dutton
Prentice-Hall © 2002

6. Avogadro’s Number
The mole is an amount of substance that contains the same number of elementary entities as there are carbon-12 atoms in exactly 12 g of carbon-12.
NA = x 1023 mol-1
NA = mol-1
Chemistry 140 Fall 2002 Dutton
Prentice-Hall © 2002

7. Molar Mass The molar mass, M, is the mass of one mole of a substance.
M (g/mol 12C) = A (g/atom 12C) x NA (atoms 12C /mol 12C)
Mass of one atom of 12C
The molar mass of an atom is the atomic weight of that element in units of gmol-1.
The molar mass of a molecule is the addition of the atomic weights of all atoms making up the molecule in units of gmol-1.
Chemistry 140 Fall 2002 Dutton
Prentice-Hall © 2002

8. Example 2-9
Combining Several Factors in a Calculation—Molar Mass, the Avogadro Constant, Percent Abundance.
Potassium-40 (40K) is one of the few naturally occurring radioactive isotopes of elements of low atomic number. Its percent natural abundance among K isotopes is 0.012%. How many 40K atoms do you ingest by drinking one cup of whole milk containing 371 mg of K?
Want atoms of 40K, need atoms of K,
Want atoms of K, need moles of K,
Want moles of K, need mass and M(K).
Chemistry 140 Fall 2002 Dutton
Prentice-Hall © 2002

9. Calculating Numbers of Particles Calculating Numbers of Moles
Useful Formula’s
Calculating Numbers of Particles
N = n x NA
N = number of particles
NA = Avogadro’s number (particles / mol)
n = number of moles (mol)
Calculating Numbers of Moles
n = m / M
n = number of moles (mol)
m = mass (g)
M = molar mass (g / mol)
Chemistry 140 Fall 2002 Dutton
Prentice-Hall © 2002

10. Strategies (Flowchart)
Moles
Number of Particles
Mass of Substance
n = N / NA
m = n x M
N = n x NA
n = m / M
N = n x NA
n = m / M
Chemistry 140 Fall 2002 Dutton
Prentice-Hall © 2002

11. Convert Strategy to Plan
and Plan Into Action
Convert mass of K(mg K) into moles of K (mol K)
n = m / M
mK(mg) x (1g/1000mg)  mK (g) x 1/MK (mol/g)  nK(mol)
nK = (371 mg K) x (10-3 g/mg)
x (1 mol K) / (39.10 g K)
= 9.49 x 10-3 mol K
Convert moles of K into atoms of 40K
N = n x NA
nK(mol) x NA  atoms K x 0.012%  atoms 40K
atoms 40K = (9.49 x 10-3 mol K) x (6.022 x 1023 atoms K/mol K)
x (1.2 x K/K)
= 6.9 x K atoms
Chemistry 140 Fall 2002 Dutton
Prentice-Hall © 2002

12. Chapter 2 Questions
3, 4, 5, 9, 11, 21, 22, 24, 25, 33, 51, 55, 63, 65, 83.
Chemistry 140 Fall 2002 Dutton
Prentice-Hall © 2002

13. Chapter 3: Chemical Compounds
General Chemistry
Principles and Modern Applications
Petrucci • Harwood • Herring
8th Edition
Chapter 3: Chemical Compounds
Philip Dutton
University of Windsor, Canada
N9B 3P4
Prentice-Hall © 2002
(modified 2003 by Dr. Paul Root and 2005 by Dr. David Tramontozzi)
Chemistry 140 Fall 2002 Dutton
Prentice-Hall © 2002

14. Contents 3-1 Molecular and Ionic Compounds 3-2 Molecular Mass
Chemistry 140 Fall 2002
Contents
3-1 Molecular and Ionic Compounds
3-2 Molecular Mass
3-3 Composition
3-4 Oxidation States
3-5 Names and formulas
Focus on Mass Spectrometry
Chemistry 140 Fall 2002 Dutton
Prentice-Hall © 2002

15. Molecular Compounds Molecular Compound, discrete units of molecules.
Chemistry 140 Fall 2002
Molecular Compounds
Molecular Compound, discrete units of molecules.
Chemical formula, symbolic representation of compound showing numbers and types of atoms.
Simplest formula of compound
Actual formula of compound
Chemical formula – relative numbers of atoms of each element present
Empirical formula – the simplest whole number formula
Structural formula – the order and type of attachements
– shows multiple bonds
- may show lone pairs
- hard to show 3-d
Actual attachment of atoms
Chemistry 140 Fall 2002 Dutton
Prentice-Hall © 2002

16. Standard Color Scheme Chemistry 140 Fall 2002 Dutton
Prentice-Hall © 2002

17. Condensed Structural Formulas
Some Molecules
H2O2
CH3CH2Cl
P4O10
CH3CH(OH)CH3
HCO2H
Condensed Structural Formulas
Chemistry 140 Fall 2002 Dutton
Prentice-Hall © 2002

18. Chemistry 140 Fall 2002
Ionic Compounds
Atoms of almost all elements can gain or lose electrons to form charged species called ions.
Compounds composed of ions are known as ionic compounds.
Metals tend to lose electrons (lose negative charge) to form positively charged ions called cations.
Non-metals tend to gain electrons (gain negative charge) to form negatively charged ions called anions.
Positive and negaive ions joined together by electrostatic forces
Metals tend to lose electrons to form cations
Non-metals tend to gain electrons to form anions
Ionic solids formulae are reported as the formula unit – inappropriate to call it a molecular formula
Chemistry 140 Fall 2002 Dutton
Prentice-Hall © 2002

19. Sodium chloride (Table Salt)
Chemistry 140 Fall 2002
Sodium chloride (Table Salt)
Formula Unit, smallest electrically neutral collection of ions. Smallest whole number ratio giving a neutral ionic compound.
Extended array of Na+ and Cl- ions.
FORMULA
UNIT
Na loses one electron to form the sodium ion
Cl gains one electron to form the chloride ion
Centers of ions are shown in the ball and stick model for clarity
Space filling model shows how the ions are actually in contact with one another.
We will discuss face centered cubic and other types of packing in chapter 13
Ball and stick crystal arrangement
Space filling arrangement
Chemistry 140 Fall 2002 Dutton
Prentice-Hall © 2002

20. Inorganic Molecules P4 S8
Chemistry 140 Fall 2002
Inorganic Molecules S8 P4
Some inorganic compounds for molecules
Sulfur and phosporous for example.
They come in various forms called allotropes – these are one allotrope of each
Some inorganic molecules are made of clusters of identical atoms.
Need to be familiar with the difference between atomic mass and molecular mass.
Chemistry 140 Fall 2002 Dutton
Prentice-Hall © 2002