Chapter 2 Notes for Chem Exam

Chapter 2: Atoms and Molecules
2.1 Symbols and Formulas:
• Element – homogenous pure substances made up of identical atoms
• 88 naturally occurring elements found in the Earth’s crust, oceans, and
atmosphere
• Each element can be characterized and identified by its unique set of
physical and chemical properties.
• Each element is therefore assigned a unique name and symbol, called an:
• Elemental Symbol – based on element’s name and consists of a single
capital letter or a capital letter followed by one lowercase letter.
Example: H – Hydrogen, He – Helium
• Compounds – pure substances made up of two or more different kinds of
atoms.
• Atoms are identical, whether in an element or compound, so the symbols
used for elements are the same used for the atoms in a compound.
• Compound Formula – Symbol for the molecule of a compound, consisting
of the symbols of the atoms in that compound.
Example: Hydrochloric Acid, (aka stomach acid), is written as HCl.
• Atoms present in numbers > 1 in a compound have that number indicated by
a subscript.
Example: Carbon Dioxide is written as CO2.
• A subscript number one is never used to indicate when only one atom is
present in a compound.
• The # 1 is implied if no other subscript is written in a compound formula.
Practice Problems: Write formulas for the following compounds.
1) Sulfuric Acid: two hydrogen atoms (H), one sulfur atom (S), and four oxygen
atoms (O)
2) Glucose: six carbon atoms (C), twelve hydrogen atoms, and six oxygen atoms.

2.2 Inside the atom:
• Atom – the limit of chemical subdivision of matter, the basic building block
of matter
• Atoms are made up of over 100 smaller (subatomic) particles
• Three major subatomic particles contribute the most influence on an atom’s
characteristics: Protons, Neutrons, and Electrons.
Table 2.3 Characteristics of important subatomic particles
Characteristics
Particle Symbol Charge Mass (g) Mass (u) Location
Electron e- -1 9.07 x 10-28 1/1836 Outside
Nucleus
Proton p, p+, H+ +1 1.67 x 10-24 1 Inside
Nucleus
Neutron n 0 1.67 x 10-24 1 Inside
Nucleus
• Nucleus – Central core of the atom. It is made up of neutrons and protons.
It contains 99.99% of the atom’s mass.
• Protons and neutrons are tightly bound together. Each nucleus has a positive
charge equal to the number of protons it contains.
• Even though the mass of a proton is 1836 times greater than the mass of an
electron, the charges of e- and p+ are of equal but opposite strength.
• So, an atom with equal #s of p+ and e- has no net charge and is considered to
be neutral.
• Electrons are negatively charged particles located outside of the nucleus.
http://www.silvershake.com/store/amethyst/images/Atomic-Structure.gif
• Electrons move very rapidly around the nucleus, throughout a relatively
large volume of space.
• Subatomic particles by themselves are relatively unstable, short-lived, and
do not display the properties of any element.
• The only way they gain long-term stability is by combining with other
particles to form an atom.
• Therefore atoms are considered the fundamental building blocks of matter.
Chem. 120/121 Chapter 2 Lecture Notes

2.3 Isotopes:
• Most atoms prefer to be neutral most of the time and are most stable when
they have no net charge (# of p+ = # of e-)
• Since neutrons have no charge, the number of neutrons can vary from the
number of protons and electrons in an atom.
• Atomic number (Z) – The number of protons in the nucleus of an atom.
• The atomic number is also the number of electrons in the neutral atom.
• ALL atoms of a specific element MUST have the same atomic number. (The
# of p+ in an atom is what give the element its identity)
• But, the number of neutrons can vary among atoms of the same element.
Example:
http://images.encarta.msn.com/xrefmedia/aencmed/targets/illus/ilt/T046738A.gif
Hydrogen (H) exists in 3 different atomic forms.
All 3 forms have the same atomic number (Z = 1 for 1 proton). They also all have
1 electron and all have zero net charge (They’re neutral).
Where they differ is these three forms contain 0, 1, and 2 neutrons respectively.
Note: all three have different names (protium, deuterium, and tritium) and different
properties (tritium is radioactive).
• Isotopes – atoms with the same atomic # but different numbers of neutrons.
• Mass number (A) – The sum total of the # of protons and # of neutrons in
the nucleus of the atom.
• Mass #s: Protium (A =1), Deuterium (A = 2), Tritium (A = 3).
• To distinguish between isotopes, the following notation is used:
• A
ZE where E is the elemental symbol, A is the mass #, and Z is atomic #
Example: 1
1H = protium, 2
1H = deuterium, 3
1H = tritium
Practice Problems:
1) What is the atomic #, mass #, and isotope symbol for an atom with 4 protons
and 5 neutrons?
2) How many neutrons are contained in an atom of chlorine-37?
2.4 Relative Masses of Atoms and Molecules:
• The masses of subatomic particles are very small and difficult to work with.
• Atomic Mass Unit (amu or u) – A unit used to express the relative masses
of atoms. One u is equal to 1/12th the mass of an atom of carbon–12.
• One atomic mass unit is ~ the weight of one proton or one neutron.
• Atomic weight – The mass of an average atom of an element expressed in
atomic mass units.
• Molecular weight (MW) – The relative mass of a molecule expressed in
atomic mass units and calculated by adding together the atomic weights of
the atoms in the molecule.
Ex: Water (H2O) has a MW of 18u: [2 x 1u (Z of H) + 16u (Z of O)] = 18u
Practice Problems: (1) Which element has atoms that are closest to twice the mass
of copper (Cu)?
(2) How many Helium (He) atoms would be required to have a mass ~ equal to the
mass of a single Neon (Ne) atom?
(3) What is the molecular weight of ethanol (C2H6O)?

2.5 Isotopes and Atomic Weights:
• Atomic Weight –the average mass of all of the atoms of a particular
element.
• Protons and Neutrons both have masses of 1u and the mass of e- are ~ 0.
• So, the 3 isotopes of H have different masses that are the sum of the p+ and
n in the nucleus of each atom (1 u, 2 u, 3 u).
• According to the Periodic Table, Hydrogen has an atomic weight of 1.008 u.
• Where does that # come from?
• It is the average weight of all H atoms.
Example: I have 3 different types of poker chips. I have 60 chips that weigh 11g,
30 chips that weigh 8g, and 10 chips that weigh 2g.
What is the average weight of my poker chips?
60% of my chips weigh 11g, 30% weigh 8g, and 10% weigh 2g, so multiply the
weight of each by its percentage and add the weights together.
11g (0.60) + 8g (0.30) + 2g (0.10) = 9.2g
Practice Problem:
1) Chlorine has two isotopes, 35Cl and 37Cl. 75.53% of all Chlorine is 35Cl (mass =
34.97 u) and 24.47% is 37Cl (mass = 36.97 u). Calculate the atomic weight.
2) Mg has 3 isotopes, 24Mg (23.99 u, 78.70%), 25Mg (24.99 u, 10.13%), and 26Mg
(25.98 u, 11.97%) Calculate the atomic weight.
Chem. 120/121 Chapter 2 Lecture Notes

2.6 – Avogadro’s Number - The Mole:
• Mole (mol) – The number of particles (atoms or molecules) contained in a
sample of element or compound with a mass in grams equal to the atomic or
molecular weight.
• 1 Mole = 6.022 x 1023 particles
• What the heck does that mean?
1 mol C atoms = 6.022 x 1023 C atoms = 12.01 g C
and
1 mol O atoms = 6.022 x 1023 O atoms = 16.00 g O
and
1 mol CO2=6.022 x 1023 CO2 molecules=44.01 g CO2
• How is that possible?
• Using modern scientific equipment, masses of indiviual atoms have been
determined.
• One atom of C has a mass of 1.99 x 1023 g
• One atom of O has a mass of 2.66 x 1023g
• One molecule of CO2’s mass is 7.31 x 1023g
Math check: 1.99g + 2.66g + 2.66g = 7.31g
How many atoms of C, O, and elements of CO2 would it take to equal the
Atomic or Molecular Masses?
• Using the Periodic Table, The atomic mass of C is 12.01u, the atomic mass
of O is 16.00u, and the molecular mass of CO2 is 44.01u.
12.01 g C 1 atom C = 6.02 x 1023 atoms C
1.99 x 10-23 g of C
16.00 g O 1 atom O = 6.02 x 1023 atoms O
2.66 x 10-23 g O
44.0g CO2 1 molecule CO2 = 6.02 x 1023 atoms CO2
7.31 x 10-23g CO2
• The number 6.022 x 1023 is a constant for all atoms or molecules.
• 6.022 x 1023 is ALWAYS the number of particles (atoms or molecules)
contained in a sample of element or compound with a mass in grams equal
to the atomic or molecular weight.
• A mole is a defined number, just like a dozen or a score. A dozen atoms =
12 atoms, a mole of atoms = 6.022 x 1023 atoms.
• Don’t get freaked out because it is a really big number and has a strange
name!

Practice Problems:
1) What is the mass in grams of 1.42 mol of Na?
2) How many moles of P atoms are in 67 g of P?
3) What is the mass in grams of one atom of N?
4) How many S atoms are in 98.6 g of S?

2.7 – The Mole and Chemical Formulas:
• The Chemical Formula of water is H2O.
• All water molecules contain H and O in a 2:1 ratio.
Using this ratio:
- 2 H2O molecules have 4 H atoms & 2 O atoms.
- 200 H2O molecules have 400 H atoms & 200 O atoms.
- 6.022 x 1023 H2O molecules have 12.04 x 1023 H atoms & 6.022 x 1023 O atoms.
- 1 mol of H2O molecules has 2 mol of H atoms & 1 mol of O atoms.
• The ratios of atoms in compounds will always equal the ratio of moles of
that atom in moles of that compound.
Example: 1 mol of H2SO4 has 2 mol H, 1 mol S, and 4 mol of O atoms.
However, 5 mol H2SO4 has 10 mol H, 5 mol S, and 20 mol O atoms.
• What kind of fun can we have with this?
Chem. 120/121 Chapter 2 Lecture Notes
Practice Problems:
1) How many mols of horns, hooves, and tails are there in 1 mol of cows? How
many in 1.5 mols?
2) How many mols of cave trolls could 3 mols of Legolas kill if it takes 5 arrows to
kill each troll and each Legolas has 100 arrows?
3) How many mol of each atom are found in 1 mol of CHCl3?
4) How many mol of each atom are found in 0.5 mol of glucose (C6H12O6)?
5) How many mol of each atom are found in 3 mol of baking soda (NaHCO3)?