Lecture Notes

The listed bold chapters from Brown & Lemay will be covered for sure in Chem 1A. Other topics may be covered as time permits. Go to the bottom of this page for lecture notes.

Chapter 1 Introduction: Matter and Measurement

Chapter 2 Atoms, Molecules, and Ions

Chapter 3 Stoichiometry: Calculations with Chemical Formulas and Equations

Chapter 4 Aqueous Reactions and Solution Stoichiometry

Chapter 10 Gases

Chapter 6 Electronic Structure of Atoms

Chapter 7 Periodic Properties of the Elements

Chapter 8 Basic Concepts of Chemical Bonding

Chapter 9 Molecular Geometry and Bonding Theories

Chapter 5 Thermochemistry

Chapter 11 Intermolecular Forces, Liquids, and Solids

Chapter 13 Properties of Solutions

Chapter 15 Chemical Equilibrium

Chapter 16 Acid-Base Equilibria

Chapter 17 Additional Aspects of Aqueous Equilbria

Chapter 12 Modern Materials


Here are some general chemistry review topics from Perdue.


Lecture notes with hyperlinks below:

Substances, Properties, and Measurements

Origin of Atomic Theory

Elements and Compounds: Use of Symbols

Law of Combining Volumes and Avogadro’s hypothesis

Discoveries of the Electron and The Nucleus

Relative Atomic Masses, the Atomic Mass Unit, and Atomic Molar Mass




Substances, Properties, and Measurements


Chemistry:  A Science of Substances



Matter:  What is matter?  What are substances?  How do we classify matter and substances?


Demo:  examples of substances (elements, compounds), and mixtures (homogeneous, heterogeneous).  What is the difference between a homogeneous mixture and a compound?


Chemistry is fundamentally about substances and how substances may change into other substances by way of chemical reactions or into other forms of the same substance by way of physical transformations.


A substance is either an element (simplest kind of substance) or a compound (chemical combination of two or more elements).  At the submicroscopic level the book says that an element, the simplest kind of substance, is made of only one kind of atom.  (This is not strictly true as we will see when we learn about isotopes.)  It is more accurate to say that an element consists only of atoms with the same nuclear charge.


In the early days of chemistry the existence of the nucleus, of nuclear charge, and or even of atoms was not known about!  So at that time an element was a substance that could not be broken down into or synthesized from simpler substances.


Example:  100.00 g of the substance called cinnabar (a red mineral) can be decomposed by heating to 86.22 g of mercury and 13.78 g of sulfur.  Mercury and sulfur cannot be broken down into simpler substances via chemical reactions so these substances were deemed elements by the early chemists.


Substances are classified and described by observing and measuring their properties.   This is what descriptive chemistry is about.  A particular substance can be identified by its chemical properties and by its intensive physical properties.


Chemical Property:  any chemical reaction involving the substance of interest

It is a chemical property of cinnabar that it can be chemically decomposed (broken down) into mercury and sulfur.  Another chemical property of cinnabar is that it can be synthesized from mercury and sulfur.


It is a chemical property of hydrogen that it can be reacted with oxygen to synthesize water.  Another chemical property of hydrogen is that it can be produced by reacting magnesium with hydrochloric acid.


In summary, the chemical properties of a substance are all possible chemical reactions involving the substance.



Physical Property:  any property which pertains only to the substance itself without reference to other substances.   (i.e. any property not having to do with chemical reactions.)


Melting point temperature, boiling point temperature, density, color, electrical conductivity, specific heat, thermal conductivity, etc. are all physical properties which can be used to identify a substance.  These also happen to be intensive physical properties which do not depend on the amount of substance present.  Temperature and pressure are also intensive properties.


Volume, mass, number of moles, energy, enthalpy, entropy, etc. are extensive physical properties which always depend on the amount of substance present.




The ratio of two extensive properties is always an intensive property which can be used for substance identification!  Examples:


mass/volume = density


mass/(number of moles) = molar mass


volume/(number of moles) = molar volume


These kind of intensive properties (involving ratios of two extensive properties) also serve as conversion factors between the extensive properties involved.  Calculate the volume of a piece of cinnabar weighing 1.00 kg (density of cinnabar is 8.19 g/cc).



Measurements:  units, uncertainty, and significant figures


Since we are studying substances, and want to characterize or identify them according to intensive properties as described above, we need to make measurements on extensive properties!  Then we can take the ratio of two extensive properties to get an intensive property useful for identification.  (Some intensive properties we can measure directly like temperature or pressure.)  Of course any measurement always involves some kind of measuring device.  We might measure length with a ruler, volume with a graduated cylinder, mass with an analytical balance, etc.  To be useful, a value obtained from a measuring device needs to be communicated with units telling us what the measurement means and with some indication of the uncertainty to tell us how good the measurement is.



(Mars Climate Orbiter crashed in 1999) Lockheed delivered thrust data in English units instead of metric units.  (Thrust data was sent in pounds instead of newtons.  1 pound = 4.45 newtons)


Metric Base Units vs English Base Units

              distance                mass                     time                   temperature         amount

SI           meter  “m”            kilogram “kg”       second “s”         Kelvin “K”              mole “mol”


English    foot                     pound                     second               Farenheit degree    mole


SI Derived Units              

             volume                         speed                                            density

SI          cubic meter “m3”         meters per second “m/s”             kilograms per cubic meter “kg/m3


Metric Units for Laboratory Scale work

SI base units and units derived from base units are too large quantities for use in the chemistry lab.  Instead will commonly use the following metric units:


distance                              mass         volume                                                                        

centimeter “cm”  gram “g”   cubic centimeter “cc”, “cm3”, or “mL”



gram per milliliter “g/cc”, “g/cm3”, “g/mL”               (mL = cc = cm3 is the same unit of volume)


Of course you should know how to do unit conversions between metric units, i.e. between laboratory scale units and SI base units.  Study the prefix table on p. 14 of the text.  Other useful unit conversion factors (including English-metric conversions) can be found under the back flap of your book.






Uncertainty in Measurements

Measured Quantity (p.22)                1.03 x 104 g                        1.030 x 104 g                      1.0300 x 104 g


Absolute Uncertainty                        0.01 x 104 g                        0.001 x 104 g                      0.0001 x 104 g


Relative Uncertainty                         0.01                                     0.001                                   0.0001


Number of Sig. Figs.                         3                                           4                                           5



1 – log(relative uncertainty)             3                                           4                                           5



When the uncertainty increases by a factor of ten, the number of sig. figs. decreases by one.  When the uncertainty decreases by a factor of ten, the number of sig. figs. increases by one.  Get it?


Absolute Uncertainty is the actual uncertainty of the measured quantity and depends on the device used.  The absolute uncertainty has the same units as the measured quantity!


Relative Uncertainty = Absolute Uncertainty / Measured Quantity and can also be called the “fractional uncertainty”.  The Relative Uncertainty has no units.


The Number of Significant Figures is closely related to the log of the Relative Uncertainty (see above)!  For each unit increase of the Number of Sig. Figs. there is a 10 fold decrease in the Relative Uncertainty!  More Sig. Figs. always implies smaller uncertainties!



Significant Figure Propagation Rules are really about uncertainty propagation in calculations.


Multiplication or Division:  Answer has least number of sig. figs. of the input quantities.


Addition or Subtraction:  Answer has least number of decimal places of the input quantities.




Origin of Atomic Theory

Origin of the Atomic Theory


When early chemists evolved in sophistication to the point where they were making careful mass measurements before and after chemical reactions, the following patterns were observed in the experimental data:


Law of Conservation of Mass: total mass remains constant in the course of a chemical reaction. (Antoine Lavosier)


Law of Definite Proportions: in a given chemical compound, the mass ratios of elements making up the compound are always fixed, independent of how the compound was prepared. (Joseph Proust)


(Remember that a Law is a very general experimental result!)


These very general experimental results led John Dalton to propose his Atomic Theory as an explanation, namely that


1. Matter consists of indivisible atoms.


2. All atoms of a given element have the same mass (and other properties).


3. Different elements have different kinds of atoms, and in particular atoms with different masses.


4. Atoms cannot be created or destroyed, and maintain their identities in chemical reactions.


5. A compound forms from different elements by the combination of atoms (of those elements) in simple whole number ratios.


6. Chemical reactions involve combination, separation, or rearrangement of atoms (into new substances).


So, thanks to Dalton we to represent atoms and compounds with symbols! H, O, C, N, CO, CO2, NO, NO2, N2O etc.


Assuming his atomic theory to be correct, Dalton suggested that another law would also be true, namely the


Law of Multiple Proportions: When more than one compound is formed from the same two elements, then the masses of one element that combine with a fixed mass of the other element are in the ratios of small whole numbers to each other.


For example it is observed that two compounds can be formed by combining the elements carbon and oxygen. For a fixed mass of carbon, the mass of oxygen in the first compound is twice the mass of oxygen in the second compound. How is this explained by Dalton’s Atomic Theory?




Elements and Compounds: Use of Symbols


Symbols for Elements & Compounds – Relative Atomic Mass Scale

Symbols for elements are typically the first or the first and second letters of the common name for the element.  Examples:  H (hydrogen), He (helium), B (boron), Be (beryllium).  Usually the English name but sometimes the Latin or German name will be used.  (Paul, King, & Farinholt 1967.)


Non-obvious chemical symbols:

English                   Latin                       Symbol

Antimony              Stibium                   Sb

Copper                   Cuprum                  Cu

Gold                        Aurum                    Au

Iron                         Ferrum                    Fe

Lead                       Plumbum                Pb

Mercury                 Hydragyrum          Hg

Silver                      Argentum              Ag

Tin                          Stannum                                Sn


English                   German                   Symbol

Potassium              Kalium                    K

Sodium                   Natrium                  Na

Tungsten               Wolfram W


Chemical formulas:  What do they mean?  Use molecular formula for molecules, empirical formula for salts.


Chemical formulas in molecules

Molecular formulae:  H­2O, CO2, CO, NH3, NO, NO2    Symbols refer to atoms in a molecule.

Show models here.


Chemical formulas in salts

Empirical formulae:  NaCl,  MgCl2, NaBr, KCl, NaNO3    Symbols refer to atoms in a formula unit.

Show models here.


The whole idea behind chemical formulae is to symbolically represent Dalton’s atomic theory, that

  1. 1.        Matter consists of indivisible atoms.
  2. 2.        These atoms are unique in properties (like mass) peculiar to a given element.
  3. 3.        Atoms combine in various combinations of simple whole number ratios to form compounds.


Dalton’s mistake

So, for example, early chemists including Dalton noticed that when combining hydrogen and oxygen, 1g of hydrogen always combines with 8 g of oxygen.  Thinking that the simplest compounds formed from two elements consisted of on atom from each element, he erroneously took the formula for water to be HO.


Avogadro proposed that equal volumes of a gas contained equal numbers of molecules and noticed that when hydrogen combined with oxygen to form water, two volumes of hydrogen always combined with one volume of oxygen.   Hence Avogadro correctly inferred the formula for water to be H2O.  Now we know that both hydrogen and oxygen are diatomic:  i.e. in their elemental gaseous state they consist of the molecules H2 and O2 respectively.  Symbolically, hydrogen and oxygen recombine chemically to form water according to the reaction equation:


2H2 + O2 à 2H2O


The subscripted numbers refer to numbers of atoms in a molecule (or in a formula unit in the case of ionic compounds).  The numbers in front of the molecules give the ratios between molecules combined and/or formed in the reaction and are called stoichiometric coefficients.   Stoichiometric coefficients arise from the balancing of chemical equations.  We balance reaction equations because in chemical reactions atoms are neither created nor destroyed according to Dalton’s Atomic Theory (which explains the Law of Mass Conservation).



 Law of Combining Volumes and Avogadro’s hypothesis

Law of Combining Volumes and Avogadro’s hypothesis



Dalton’s Atomic Theory, introduced in 1808, was readily accepted by chemists since it neatly explained the experimental results summarized in the Law of Conservation of Mass, the Law of Definite Proportions, and the Law of Multiple Proportions.


The problem was now to determine the relative masses of the elements and the formulae of simple molecules and compounds. If you already know the formula of a binary compound, then it is easy to determine the relative masses of the elements that make up the compound. Conversely if you know the relative masses, then it is easy to determine the formula (to an integer multiple) of the compound. The problem the early chemists had was to determine both the relative masses and the formulae simultaneously! To this end, Dalton (wrongly) proposed his “rule of greatest simplicity”, namely that the simplest compound consisting of elements A & B consisted of one atom of A plus one atom of B. So Dalton thought the formula for water was HO and therefore that an oxygen atom’s mass was 8 times greater than a hydrogen atom’s mass.


At about this time, a French chemist named Joseph Gay-Lussac was carrying out experiments reacting simple gases together to form new gases. By carefully measuring the volumes of both the reactant and product gases he discovered that the volumes of the two reacting gases (at fixed P and T) are always in the ratio of simple whole numbers (integers). Additionally it was always the case that the ratio of the volume of the product gas to any reactant gas was also a ratio of simple whole numbers. This general experimental result was called the Law of Combining Volumes.




2 volumes of hydrogen combines with 1 volume of oxygen to give 2 volumes of water vapor.


3 volumes of hydrogen combine with 1 volume of nitrogen to give 2 volumes of ammonia.


Avogadro, thinking about this data, proposed that equal volumes of different gases (at fixed T and P) contain equal numbers of particles. This statement is now known as Avogadro’s Hypothesis (or Avogadro’s Law). The question was then whether the particles were the same as Dalton’s atoms. Avogadro’s opinion was that they were not, i.e. some elements could exist as diatomic molecules. In this way he could explain Gay-Lussac’s experimental results. In particular, for the above two reactions Avogadro proposed the equations:


Avogadro where Dalton would have written


2 H2 + O2 à 2H2O H + O à HO


3 H2 + N2 à 2NH3 H + N à HN



It is easy to see that, given Avogadro’s hypothesis, the equations on the left agree with Gay-Lussac’s results where the equations on the right do not.



Armed with Avogadro’s hypothesis, chemists were now able to correctly determine formulas for simple compounds and the relative atomic masses for many elements.




Discoveries of the Electron and The Nucleus

History of discovery of the electron and the nucleus:J.J. Thompson is credited with discovering the electron in 1897: http://www.aip.org/history/electron/

Dalton thought the atom was an indivisible particle. He had no idea about the existence of electrons, protons, neutrons, or atomic nuclei. In the very late 1800s, Thompson used a cathode ray tube fitted with variably charged plates and a magnet to infer the existence of the electron and then to measure its charge/mass ratio. He inferred the existence of particles with a particular charge and mass because cathode rays behaved analogously to large charged bodies moving in electric and magnetic fields: that is these rays acted like streams of charged objects whose trajectory obeyed Newton’s laws of motion when subjected to forces given by Maxwell’s equations of electromagnetism.

For those of you with some physics background, Thompson’s calculations are reproduced here: http://dbhs.wvusd.k12.ca.us/webdocs/AtomicStructure/Disc-of-Electron-Results.html

Then in 1909, Robert Millikan measured the charge on an electron in his famous oil drop experiment: http://online.cctt.org/physicslab/content/phyapb/lessonnotes/dualnature/Millikan.asp

More Homework: Calculate the mass of the electron from Thompson’s e/m ratio of 1.759 x 10**11 Coulomb/Kg and Millikan’s charge on the electron of 1.602 x 10**-19 Coulombs. (This is just a simple conversion factor – dimensional analysis problem.)

J.J. Thompson’s “plum pudding” model of the atom – Before protons were known, Thompson thought that in the atom, electrons moved around in evenly distributed “positively charged pudding”:


Rutherford’s experiment bombarding gold foil with alpha particles proved the “plum pudding” model to be wrong!


Rutherford concluded that:

1. An atom’s positive charge and mass must be concentrated in a very small positively charged nucleus as only a very small number of alpha particles either deflected or rebounded off the foil.

2. Most of the atom must be empty space. This space must contain the negatively charged electrons.

Rutherford Atom, today’s view:http://www.chemistry.ohio-state.edu/~woodward/ch121/ch2_atoms.htm

Some student questions answered……The following important questions were emailed several semesters ago:Question: What is the difference between atomic mass and mass number?Atomic mass is the actual mass of an atom (in amu or g/mol). Mass number, A, is the sum of number of protons + number of neutrons for an atom. Atomic mass is a measured quantity (has uncertainty) while the mass number is an exact or counted quantity (has no uncertainty).The mass number turns out to be a decent estimate of the measured atomic mass of a given isotope to within 2 or 3 significant figures. Why? The proton and neutron both have masses nearly but not precisely the same as 1 amu.Question: What is the difference between atomic number and mass number?Atomic number, Z, is the number of protons in an atom’s nucleus. All atoms of a given element have the same atomic number but not necessarily the same mass number. Question you should be able to answer: How do isotopes of the same element differ?



Relative Atomic Masses, the Atomic Mass Unit, and Atomic Molar Mass


Atomic mass scale:


The relative atomic mass scale in amu’s was worked out (in the 1800s before chemists new how to weigh or count atoms directly) from experimentally observed mass ratios and volume ratios in the formation of simple binary compounds from pairs of elements as we showed in class. (Avogadro’s Law was used in this this endeavor to compare masses of equal volumes of gases.) Thus, setting a hydrogen atom equal to 1 amu (atomic mass unit), the relative weights of other elements were determined relative to hydrogen even though hydrogen’s absolute mass was not known.


Nowadays we can measure masses of individual isotopes of elements very precisely with mass spectrometers. Now the amu is based on the carbon 12 isotope rather than hydrogen: 1 amu is defined as exactly 1/12 the mass of a single carbon 12 atom. Avogadro’s number, NA, or a mole, is the number of amus in a gram. Alternatively, Avogadro’s number, NA, or a mole, is the number of carbon 12 atoms in exactly 12 grams of carbon 12.


For any element there are Avogadro’s number, i.e. one mole, of atoms in the atomic mass of that element expressed in grams. Example: one oxygen atom has a mass of 16.00 amu while a mole of oxygen atoms has a mass of 16.00 grams. Can you prove this? So the atomic molar mass is a conversion factor between grams and moles. Thinking macroscopically, we say oxygen has an atomic molar mass of 16.00 g/mol (grams per mole). The same goes for any other element. This is why the atomic mass is also known as a molar mass.


Early chemists knew Avogadro’s number was important even though they were unable to measure it at the time. Avogadro’s number (also known as a mole) is 6.022 x 1022. You don’t need to know the number to convert from grams to moles or vice versa. You do need to know the number to calculate the number of atoms (or molecules) in a given mass of the pure substance in question.



Molar mass of a compound: We can calculate the mass of a molecule ( in amu) by adding up the masses of all the atoms in that molecule. Thinking macroscopically, the molar mass of a molecule is the sum of the molar masses of the constituent atoms. So we can easily determine the mass to mole conversion factor for any compound.


Example: the molecular mass of an H2O molecule is 2 x 1.008 amu + 1 x 16.00 amu = 18.01 amu


So the molar mass of H2O is 2 x 1.008 g/mol + 1 x 16.00 g/mol = 18.01 g/mol


For ionic compounds for which the word “molecule” does not apply because we are dealing with a lattice, simply substitute the term “formula unit” for the word “molecule” in the above discussion. For a molecular substance the molar mass is also known as the molecular mass. For an ionic compound, the molar mass is also known as the formula mass.


Molecular Formula: Gives the numbers of atoms in the molecule


Empirical Formula or Formula Unit: Gives the simplest whole number ratios between atoms in any compound.


Determining the Empirical Formula from mass data: know how to do this.