PACE UNIVERSITY
Chemistry 111
Chapter 2
2.1 -The Atomic Theory
A - Summary:
Class:
John Dalton’s Modern Atomic Theory
1) Atoms are the smallest repeating units that compose an element.
Atoms that compose an element are identical in size, shape and
chemical properties.
2) Compounds are the result of combining two or more different atoms
with a fixed ratio.
3) Chemical reactions are simply the rearrangement of such atomic
entities into compounds.
Law of Definite Proportions:
A substance will have the same ratio and constituents no matter what the
source. Pure substances will always have the same chemical and
physical properties.
Law of Multiple Proportions:
Ratios of the multiple elements combine and can form more than one
compound. The ratio is an integer or a small whole number. Example:
CO
1:1
CO2
1:2
or
H2O
2:1
H2O2
2:2
Law of Conservation of Mass and Energy:
Mass and Energy can neither be created nor destroyed: They can simply
be transformed from one substance to another, or transformed from
mass into energy.
General chemistry equation
aA + bB  cC + dD
reactants = products
Both sides of the equation must be equal stoichiometrially!
B - Definitions:
1 -Law of Definite Proportions - a compound is composed of two or more
elements chemically combined in a definite ratio by weight.
2 - Law of Multiple Proportions: when any two elements (A+B) combine
to form more than one compound, the different weights of B that
combine with A have a small whole number ratio to each other.
3 - Law of Conservation of Mass: mass can neither be created or
destroyed. Weight remains constant in an ordinary chemical
reaction.
4 - John Dalton (1766-1844): an English chemist who first proposed
Atomic Theory and later worked on gas laws.
2.2 Structure of the Atom:
Class:
Radioactivity is the spontaneous emission of energy by certain natural
ores.
Radiation on the other hand is the transmission or emission of energy in
space in the form of particles or waves.
Reference to page 39 in book:
A lead container is closed on all sides except for a small hole that allows
an irradiating beam to travel out of the container. Without an electrical
field, the beam passed straight through the gold foil. With an electrical
field, the beam was split with one section going towards the anode (
alpha ( + )) while the other is deflected to the cathode (  beta ( - )). The
third part of the beam went straight ahead ( gamma (neutral)).
The particles in the beams can be ranked by energy:
 (gamma) has least energy, can be stopped by paper.
 (beta) has moderate energy; can be stopped by thin metals.
 (alpha) has the highest energy; needs a very thick metal barrier
(i.e.Pb).
This theory assumed the beams were all traveling at the same speed.
Since opposites attract the protons (+) were drawn to the cathode (-), the
electrons (-) were drawn to the anode (+) and since neutral particles in
the beam could not be deflected, the neutral neutrons went straight.
J.J.Thompson’s Atomic Theory :
He proposed that any atm must be spherical and that the positive protons
are evenly distributes throughout the atom. This is the Rahni
‘watermelon’ theory. The protons (seeds) are randomly but evenly
scattered throughout the atom.
Rutherford’s Atomic Theory:
If alpha particles () were sent at Thompson’s theory of the atom, then
most of the alpha particles should be deflected by the scattered protons.
But it didn’t happen, most of the particles traveled straight through
undisturbed. Those that were deflected were near the center of the
atom. Therefore protons must be clustered near the center in a nucleus.
Rutherford calculated that 1/1843 of the space of an atom was occupied
by the protons, the rest of the space was occupied by electrons. Based
on this Thompson’s Theory was rejected.
2
Reference Table 2.1 on page 40 in text book.
Mandelev
He found that He has 4x the weight of H but is next in line on the Periodic
Table. Therefore the electron is 1/10,000 of the weight of a proton.
Chadwick:
He bombarded Beryllium with alpha particles and got a mystery ray off of
the surface. It was like gamma rays but was much heavier. The weight
was more like the weight of a proton. He called the new particle a
neutron.
B - Definitions:
1) Atom is the smallest basic unit of an element that can enter into a chemical
reaction. It is also the smallest unit that cannot be broken down into another
chemical substance.
2) Electron is the negatively charged, subatomic particle with a very low mass.
3) Radioactivity is the spontaneous breakdown of an atom by the emission of
particles or radiation.
4) Radiation: the emission or transmission of particles through space as waves
or particles.
5)  Rays or particles: are helium ions with a positive charge of +2.
6)  Rays or particles are electrons or negatively charged particles.
7)  Rays or particles are high energy radiation.
8) Proton is a subatomic particle with a single positive charge. The mass of a
proton is 1840 times greater than an electron.
9) Nucleus: is the center of an atom.
10) Neutron is a subatomic particle with no net charge. It mass is slightly more
than a proton.
11) J.J. Thompson (1856 - 1940): British physicist received the Nobel Prize for
discovering the electron.
12) Wilhelm Roentgen ( 1845 - 1923): German physicist who won the Nobel
Prize for the discovery of x-rays (1901).
13) Robert Andrew Millikan (1868 - 1953): American physicist who determined
the charge of an electron. (Nobel Prize 1923).
14) Antoine Becquerel (1852 - 1908): French physicist won the Nobel Prize for
discovering radioactivity in uranium (1903)
15) Marie Curie (1867 - 1934): won the Nobel Prize twice for her work on
radioactive elements (1911) and on radioactivity (1903).
16) Ernest Rutherford (1871 - 1937): New Zealand physicist that received the
Nobel Prize for investigations into the atomic structure of the nucleus (1908).
3
17) Johannes Geiger (1882 - 1945): German physicist who worked on radiation
and atomic nucleus theory. Best as the inventor of the Geiger counter.
18) James Chadwick (1891 - 1972): British physicist who won the Nobel Prize for
the discovery of neutrons (1935).
C - Problems:
9-How many helium atoms between first and last in a cm:
cm = 1 x 10 -10 pm
1 x 10-10 pm/cm  1 x 102 pm / He atom =
1 x 108 He atoms / cm
28- Calculate charge (coulombs) and mass (grams) for 1 mole of electrons:
a- Mass / electron = 9.09 x 10-2g x 6.022 x 10 23 electrons/ mole =
5.47 x 1026 g/mole
b--1.76 x 10 8 coulombs/gram x 5.47 x 1026 g/mole =
-9.63 x 1034 coulombs/mole.
124- Calculate the molar mass of hemoglobin- C2952H4664N812O832S8Fe4:
C2952: (12.01 g/mole(1mole  6.022 x 1023 atoms/mole)) x 2952atoms =
6 x 10-20 g
H4664
(1.008g/mole(1 mole  6.022 x 1023 atoms/mole)) x 4664atoms =
8 x 10-21 g
O832
(16.00g/mole(1 mole  6.022 x 1023 atoms/mole)) x 832 atoms =
2 x10-20 g
N812
(14.01g/mole(1 mole 6.022 x 1023atoms/mole)) x 812 atoms =
2 x 10-20g
S8
(32.00g/mole(1 mole  6.022 x 1023 atoms/mole)) x 8 atoms =
4 x 10-22g
Fe4
(55.85g/mole(1 mole  6.022 x 1023 atoms/mole)) x 4 atoms =
4 x 10-22g
Molar mass of hemoglobin is 1 x 10-19 g
4
2.3 - Mass Relationships of Atoms:
Book:
The atomic number(Z) is the number of protons in the nucleus of a
given element. In a naturally neutral element, the number of protons is
equal to the electron. Each element has a unique atomic number.
Example a naturally occurring atom containing 7 protons will always be
nitrogen.
Mass number (A) is the total number of protons and neutrons in a
nucleus of an elemental atom. Therefore:
mass # (A) = # of protons + # of neutrons
= atomic # (z)+ neutrons
Therefore
A
Z
X
A - Z = # of neutrons
where A is the mass number
Z is the atomic number
Example:
1
1H
Hydrogen
2
1H
Deuterium
3
1H
Tritium
are all isotopes of hydrogen, they have the same atomic number
(1) but have different mass numbers (1,2,3).
The mole is the SI unit that is the amount of a substance that has the
same number of elemental particles, atoms or units as there are atoms in
exactly 12 grams of carbon-12 atoms. Simply remember:
Avogadro’s number = 1 mole = 6.022045 X 1023 particles
Class:
Atomic Mass is measured in amu. The amu is based on the weight of
the carbon 12 atom.
Example: Atomic mass of H2O =
2 (1.008 g) + 1 (16.01 g) = 18.02 g
Atomic mass of C6H12O6 =
6(12.01 g) + 12(1.008 g) + 6(16.0 g) = 180.18g glucose
Mole is an SI unit having a constant number of an entity. The constant is
6.022x 10 23 of that entity. It could be atoms, molecules particles, etc.
Example: 1 mole of hydrogen = 6.022 x 1023 atoms of hydrogen present.
5
The number of moles is the mass (g) times the atomic number or the
molar mass:
moles = mass(g) X 1 mole
atomic ( or molar ) mass
Example: 64 g of O2 Atomic mass = 16.01 g
moles = 64 g X 1 mole
16 g
moles = 4 moles
Therefore the number of O2 atoms in 4 moles =
6.022 X 1023 (4 mole O )
1 mole
Number of oxygen atoms = 2.40 x 1024 in 4 mole O
B - Definitions:
1)
Atomic Number is the number of protons in the nucleus.
2) Atomic Mass is the mass of an atom expressed in amu.
3) Isotope is an atom with the same atomic number but different mass
numbers.
4) Atomic Mass Units a mass equal to the exact mass of the carbon 12 atom.
5) Mole(mol) is a unit of quantity that consists of 6.022 X 1023 particles
6) Avogadro’s Number is 6.022 x 1023 particles per mole. This can be atoms,
molecules, etc.
7) Molar Mass is the mass in grams or kilograms of one mole of particles,
atoms or molecules.
C - Problems:
17-Determine 1)the number of protons and neutrons and 2) the neutron to
proton ratio and explain the resulting trend:
Neutrons:
a- 24He
b-.2010Ne
c- 3684 Kr
d.-54112Xe
Protons:
2
10
48
58
Ratio:
4
10
36
54
1:2
1:1
4:3
19:17
The ratio of neutrons to protons increases as atomic mass increases.
6
19- Give chemical symbols for the indicated isotope:
a-. 1123 Na
b-.2864N
c- 74186W
d- 80201 Hg
29- Explain clearly what is meant by the atomic mass off Au is 197.0amu:
Au has a mass 197.0 times greater then the SI standard for amu,the mass
of a single carbon-12 atom
31- How many years to count the particles?
5.00 x 10 9 people x 6.31 x 10 7 particles / person/year =
3.15 x 10 17 particles / year
6.0 x1023 particles  3.15 x 1017 particles / year = 1.9 x 106 years
39- Mass of a single atom of:
a- Hg:
200.6 g/mole  6.022 x 1023 atoms/ mole = 3.3 x 10 -22 g/atom
b- Ne
20.18 g/mole  6.022 x 1023 atoms / mole = 3.4 x 10-23 g/atom
c-As:
74.92 g/mole  6.022 x 1023 atoms/ mole = 1.24 x 10-22 g / atom
d- Pb:
207.2 g/mole  6.022 x 10 23 atoms/mole = 3.4 x 10-22 g/atom
42-Calculate which has more atoms:1.10 g of H or 14.7 g of Cr:
1.10 g  1.008 g / mole = 1.09moles x 6.022 x 1023 atoms/mole=
6.57 x 1023atoms of H
14.7g  52.0 g/mole = 2.8 x 10-1 moles x 6.022 x 10 23 atoms/mole=
1.7 x 10 23 atoms of Cr
Hydrogen has more atoms than chromium in this problem.
56- What is the mass in amu’s of:
CH4:
12.01amu + 4(1.008)amu = 16.04 amu
H2O:
16.00amu + 2(1.008)amu = 18.02 amu
H2O2:
7
2(16.00)amu + 2(1.008)amu = 34.02 amu
C6H6:
6(12.01)amu + 6(1.008)amu = 78.11 amu
Pcl5:
30.97amu + 5(35.45)amu = 208.22 amu
58- Calculate molar mass of C55H72MgN4O5:
55(12.01)g + 72(1.008)g + 24.31g + 4(14.01)g + 5(16.00)g =
893.48g
62- How many molecules present in 2.56 g of H2O at 4C:
1.0 g/mL x 2.56 mL = 2.56 g H2O
2(1.008) + 16.00 = 18.016 amu
2.56 g  18.02 g/mole = 0.142 moles x 6.022 x 1023 molecules/mole =
8.56 x 1022 molecules H2O
120- Arrange the following in increasing mass:
a- 16 H2O molecules:
(18.02g/mole 6.022 x 1023molecules/mole) x 16 molecules =
4.79 x 10-23 grams of H2O
b- 2 atoms of Pb:
(207.2 g/mole  6.022 x 1023 atoms/mole) x 2 atoms =
6.9 x 10-22 g of Pb
-23
c- 5.1 x 10 moles He:
5.1 x 10-23 moles x 4.003 g/mole = 2.0 x 10 -22 g He
16 H2O molecules : 5.1 x 10-23moles He : 2 atoms Pb
2.4 - Molecules: Atoms in Combinations:
Class:
Empirical formula is the lowest denominator of the combination of
molecules. To determine the empirical formula calculate the moles for
each element in the molecule. Reduce these numbers to the lowest
common denominator and use them as the subscripts for each element.
Example: Empirical formula for CHSO. (Percents are given for each )
Assume that the percents are grams of each element.
B - Definitions:
1) Molecule: a combination of two or more atoms in a definite arrangement that
are held together by special forces (i.e. bonds).
2) Chemical Formulas are expressions that show the chemical composition of a
compound using the chemical abbreviations for the elements found in the
periodic table.
3) Molecular Formula is an expression used to show the exact number of atoms
of each element in a given molecule.
4) Diatomic Molecules are molecules made up of just two atoms.
8
5) Polyatomic Molecules are molecules made up of two or more atoms.
6) Allotropes are two or more forms of the same element that have significant
differences in chemical and physical properties.
7) Empirical Formula is an expression showing the different elements present in
a compound as well as the ratios of the different atoms present.
8) Molecular Mass is the sum of all the atomic masses present in a particular
molecule expressed in amu’s.
C - Problems:
a) Find formula mass for Ca(OH)2 :
1 Ca (atomic mass = 40)
= 40
2 O (atomic mass =16)= 2(16)
= 32
3 H (atomic mass = 1) = 3(1)
= 3
Formula mass for Ca(OH)2
= 74
b) Find the formula mass for CuSO4  5H2O:
1 Cu (atomic mass = 64)
= 64
1 S (atomic mass = 32)
= 32
4 O (atomic mass = 16)
= 64
(4 x 16)
5 H2O (molecular mass = 18) = 90
(5 x 18)
Formula mass for CuSO4  5H2O
=250
2.5 - Ions and Ionic Compounds:
Compounds that contain negatively charged ions (anions) and positively charged
ions (cations) are considered ionic. In chemical reactions the lose or gain of
electrons by an atom or molecule results in a charged particle called an ion. Ion
that contain more than one atom is a polyatomic ion. Example: NH4+ or NO3-.
An ion with only on atom is a monatomic ion. Example: Fe3+ or S2- .
B - Definitions:
1) Monatomic Ions are ions that contain just one atom.
2) Ionic Compounds are any neutral compounds that contain a balanced
number of anions and cations.
3) Polyatomic Ions :a group of chemically combined atoms that react as a unite
and have an electrical charge. Examples: SO42NH4+ PO4 34) Ion is a charged particle formed when an atom or group of atoms either gain
or give up electrons.
5) Anion is an ion with a net negative charge.
6) Cation is an ion with a net positive charge.
C - Problems:
9
2.70 Give the number of protons and electrons in each of the following
ions:
Ion
Na+
Ca+
Al3+
Fe2+
I-
F-
S2-
O2-
N3-
# Protons
11
20
13
26
53
9
16
8
7
# Electrons
10
18
10
24
54
10
18
10
10
2.72 Which of the following compounds are likely to be ionic? Which are
likely to be molecular?
Compounds of metals with nonmetals are usually ionic therefore :
LiF
BaCl2
KF
are ionic.
Non-metal and non-metal compounds are usually molecular, therefore:
SiCl4
B2H6
C2H4
are molecular.
2.6 - Experimental Determination of Atomic and Molecular Masses
The most accurate way to determine the atomic and molecular mass of an atom
or compound is using a mass spectrometer. This is based on the idea that
differences in masses cause differences in the amount of bending that occur
when a beam of ions passes through a magnetic field. F.W.Aston received a
Nobel Prize for this work in 1922. See Fig. 2.13 on page 56 in book.
B - Definitions:
1) Mass Spectrometer is a device used for determining masses of electrically
charged particles by separating then into distinct beams by means of
magnetic deflection.
C - Problems:
99-How many peaks on mass spec for H2S:
1
32
2
32
1H +
16S = 1 + 32 = 33
1H +
6S = 2 + 32 = 34
1
33
2
33
1H +
16S = 1 + 33 = 34
1H +
16S = 2 + 33 = 35
1
34
2
34
1H +
16S = 1 + 34 = 35
1H +
16S = 2 + 34 = 36
1
36
2
36
1H +
16H = 1 + 36 = 37
1H +
16H = 2 + 36 = 38
Since the mass spec only sees mass the number of peaks would be six
(eight minus the duplicate masses).
2.7 Percent Composition of Compounds:
Chemical formulas represent the makeup of a compound and state the type of
atoms present and their relative numbers. By taking the atomic masses of the
elements present and adding up the weights of each atom present, the resulting
mass is the mass of the compound. The formula mass is determined by
multiplying the atomic mass of a given atom by the subscript for that atom and
summing up all the masses. The total mass of the compound or formula divided
into the total individual mass of the individual element gives that elements
percent of the composition of total mass.
total mass of element in compound X 100 % = percent composition
10
total formula mass
of that element
Example: % composition of Ca in Ca(OH)2 :
mass of element
X 100 % = percent of total composition by mass
formula mass
Ca
=
40 (mass of element)
X 100 % = 54% calcium
Ca(OH)2 = 74 (formula mass)
B- Definitions:
1) Percent Composition is the percent by mass of each of the elements present
in a given compound.
Problems:
78- Calculate percent composition by mass of the following:
a-NaF :
22.99g + 19.00g= 41.99g
(22.99g 41.99g) x 100=
54.75% Na
45.25% F
b- NaCl :
22.99g + 35.45g = 58.44g
(22.99g  58.44g) x 100 =
39.34% Na
60.66% Cl
c- NaBr :
22.99g + 79.90g = 102.89g
(22.99g  102.89g) x 100=
22.34%Na
77.66%Br
d- NaI :
23.0g + 126.9g = 149.9g
(23.0g 149.9) x 100 =
15.3%Na
84.7%I
80- a-Calculate empirical formula and b- show molecular formula of Allicin:
a-44.44% C
44.44g  12.01 g/mole = 3.70 moles C  0.62 moles =5.97 = 6
6.21% H
6.21g  1.008g/mole = 6.16 moles H  0.62 moles
= 9.94 moles = 10
39.5% S
39.5g  32.1g/mole = 1.2 moles S  0.6 moles = 2
9.86% O
9.86g  16.00 g/mole=0.62 moles O  0.62 moles
=1
C6H10S2O
b- 162.28g  162.00 = 1
Therefore empirical formula equals the molecular formula.
11
83- Which substance has greater mass of Cl:
a- 5g Cl2:
b- 60g NaClO3:
22.99g + 35.45g + 32.00g = 90.44g/ mole NaClO3
35.45g/mole  90.44g/mole =0.392 x 60.00 g =23.53 g Cl  2 =
11.76 Cl2
c- 0.1 mole KCl:
39.10g + 35.45g = 74.55g / mole KCl
35.45 g  74.55g = 0.48 x 7.46 g = 3.55g Cl  2 = 1.77g Cl2
d- 30g MgCl2
24.31g + 70.90g = 95.21g /mole
70.90g/mole  95.21 g/mole = 0.75 x 30.00 g = 22.50 gCl2
e- 0.5 mole Cl2:
70.90g/mole x 0.5 mole = 34.35 g Cl2+
e has the most grams of Cl2
2.8 Naming Inorganic Compounds:
Book:
Rules:
Examples
Acids and Salts::
-ic acids form -ate salts
Sulfuric acid forms sulfate salts
-ous acids form -ite salts
Sulfurous acid forms sulfite salts
hydro- (stem) -ic acids form -ide salts
hydrochloric acid forms
chloride salts
Molecular Compounds:
Binary molecules: name first element and end the second element with ide :
Examples : HBr
NaCl
Hydrogen bromide
Sodium chloride
One pair of elements but multiple atoms: Use Greek numbers to name parts:
Examples:
CO
Carbon monoxide
CO2
Carbon dioxide
PCl3
Phosphorus trichloride
N2O4
dinitrogen tetroxide
mon-
one
bi-
two
Know:
12
note a on tetra- dropped
tri
three
tetra-
four
penta-
five
hexa-
six
hepta-
seven
octa-
eight
nona-
nine
deca-
ten
Exceptions to Greek number rules:
B2H6
diborane
CH4
methane
SiH4
silane
NH3
ammonia
PH3
phosphine
H2S
hydrogen sulfide
H2O
water
Naming Bases:
Addition of (OH)- ion adds hydroxide to simple bases. Examples:
NaOH
sodium hydroxide
KOH
potassium hydroxide
Ba(OH)2
barium hydroxide
13
Naming Hydrates :
Addition of water to an anhydrous molecule yields a hydrate. Examples:
LiCl  H2O
lithium chloride monohydrate
BaCl2  2H2O
barium chloride dihydrate
MgSO4  7H2O
magnesium sulfate heptahydrate
NOTE the use of Greek numbers for the number of waters present.
B- Definitions:
1) Organic Compounds are compounds made up of strictly H, O and C
molecules.
2) Inorganic Compounds are compounds that are not based on just C, H and O
atoms.
3) Binary Compounds are compounds formed of just two elements.
4) Ternary Compounds compounds containing three elements such as H2SO4 .
5) Alfred Stock
6) Acids are substance that give up H+ ions when dissolved in water.
7) Oxoacids is an acid containing oxygen and hydrogen around a third central
element
8) Base is a substance that when dissolved in water gives up hydroxide ions
(OH-).
9) Hydrate are compounds that have a specific number of water molecules
attached to them.
C - Problems:
2.108 Name the following Compounds:
a) KH2PO4
Potassium dihydrogen phosphate
b) K2HPO4
Potassium hydrogen phosphate
c) CHBr(g)
Hydrogen bromide (molecular compound)
d) HBr(aq)
Hydrobromic acid
e) Li2CO3
Lithium carbonate
f) K2Cr2O7
Potassium dichromate
g) NH4NO2
Ammonium nitrite
h) PF3
Phosphorus tri-fluoride
i) PF5
Phosphorus penta-fluoride
j) P4O6
Tetra-phosphorus hexoxide
k) CdI2
Cadmium iodide
l) SrSO4
Strontium sulfate
m) Al(OH)3
Aluminum hydroxide
n) KClO
Potassium hypochlorite
o) Ag2CO3
Silver carbonate
14
p)
q)
r)
s)
t)
u)
v)
w)
x)
y)
z)
FeCl2
KMnO4
CsClO3
KNH4SO4
FeO
Fe2O3
TiCl4
NaH
Li3N
Na2O
Na2O2
Iron(II) chloride
Potassium permanganate
Cesium chlorate
Potassium ammonium sulfate
Iron(II) oxide
Iron(III) oxide
Titanium(IV) chloride
Sodium hydride
Lithium nitride
Sodium oxide
Sodium peroxide
2.109 Write the formulas for the following compounds:
a) Rubidium nitrite
RbNO2
b) Potassium sulfide
K2SO
c) Sodium hydrogen sulfide
NaHS
d) Magnesium phosphate
Mg(PO4 )2
e) Calcium hydrogen phosphate
CaHPO4
f) Potassium dihydrogen phosphate
KH2PO4
g) Iodine heptaflouride
IF7
h) Ammonium sulfate
(NH4)2S
i) Silver perchlorate
AgClO4
j) Iron (III) chromate
Fe2(CrO4)3
k) Copper (I) cyanide
CuCN
l) Strontium chlorite
Sr(ClO2)2
m) Perbromic acid
HBrO4
n) Hydroiodic acid
HI(aq)
o) Disodium ammonium phosphate
Na2(NH4)PO4
p) Lead (II) carbonate
PbCO
q) Tin (II) fluoride
SnF2
r) Tetraphosphorus decasulfide
P4S10
s) Mercury (II) oxide
HgO
t) Mercury iodide
Hg2I2
u) Copper(II) pentahydrate
CuSO4 5H2O
15