CHAPTER 2:
ATOMS, MOLECULES,
AND IONS
Vanessa Prasad-Permaul
Valencia College
CHM 1045
1
Dalton’s Atomic Theory
 Elements are made of tiny particles called atoms.
 Each element is characterized by the mass of its atoms.
Atoms of the same element have the same mass.
 A Compound is a type of matter composed of atoms of
two or more elements chemically combined in fixed
proportions. ( Law of Definite Proportions)
 Chemical reactions only rearrange the way atoms are
combined; the atoms themselves are unchanged. (Law of
Conservation of Mass)
2

Priestly: HgO
Hg (l) + O2 (g)

Lavoisier: Law of Conservation of Mass

Dalton: Atomic Theory of Matter

Proust: Law of Definite Proportions
3
The Law of Definite Proportions: Two or more elements
form more than one compound, the masses of one
element in these compounds for a fixed mass of the
other element are in ratios of small whole numbers.
1.000g carbon react with oxygen
2 compounds
1.3321g oxygen/1.000g carbon
2.6642g oxygen/1.000g carbon
Ratio is 2:1 (2.6642
CO
CO2
1.3321)
This indicates that there is TWICE as much oxygen in
the second compound
4
Structure of an atom
 Nucleus: the atom’s central core, which is
positively charged and contains most of the
atom’s mass and 1 or more electrons.
 Electron: a very light negatively charged particle
that exists in the region around the atom’s
positively charged nucleus.
5
The Structure of Atoms
 Cathode-Ray Tube (Thomson, 1856–1940):
Cathode rays
consist of tiny
negatively
charged
particles,
now called
electrons.
6
The Structure of Atoms
DISCOVERY OF AN ELECTRON
 Atoms are NOT indivisible particles
Cathode: negative electrode
High voltage current is turned on a visible green light using zinc
sulfide) is emitted from the cathode and is attracted to the anode.
Green beam bends away from negative charged magnet
In conclusion: Cathode ray consists of a beam of negatively
charged particles (now called electrons) and that electrons are in
all matter.
7
The Structure of Atoms
 Deflection of electron depends on three factors:
 Strength of electric or magnetic field
 Size of negative charge on electron
 Mass of the electron
 Thomson calculated the electron’s charge to
mass ratio as 1.759 x 108 Coulombs per gram.
8
The Structure of Atoms
 Thomson could not obtain the mass or the charge of the
electron separately
 Millikan performed a series of experiments obtaining the
charge on the electron
 He observed how a charged drop of oil falls in the presence
and absence of an electric field
 Charge on an electron: 1.602 x 10-19 coulombs ( C, unit of
electric charge)
 Electron mass: 9.109 x 10-28g (1800x smaller than the
mass of the lightest atom, hydrogen)
9
The Structure of Atoms
• Oil Drop Experiment (Millikan, 1868–1953):
Applied a voltage to oppose the downward fall
of charged drops and suspend them.
• Voltage on plates place
1.602176 x 10-19 C of
charge on each oil drop.
• Millikan calculated the
electron’s mass as
9.109382 x 10-28 grams.
10
The Structure of Atoms
 Discovery of Nucleus: Rutherford irradiated
gold foil with a beam of alpha particles to
search for positive charged particles. Most of
the particles passed through but some were
deflected at large angles, why?
11
The Structure of Atoms
99.95% of mass is concentrated around the
positively charged center; the nucleus
12
Nuclear Structure; Isotopes
 Proton: a nuclear particle having a positive charge
equal to that of the electron and a mass more than 1800
times than that of an electron
 Atomic number (Z): the number of protons in the
nucleus of an atom
 Element: a substance whose atoms all have the same
atomic number
 Neutron: a nuclear particle having a mass almost
identical to that of a proton but no electrical charge
13
The Structure of Atoms
• Isotopes: Atoms with identical atomic
numbers, but different mass numbers.
• Average Isotopic Mass: A weighted average of
the isotopic masses of an element’s naturally
occurring isotopes.
• Atomic Mass: A weighted average of the
isotopic masses of an element’s naturally
occurring isotopes.
14

Mass number (A): the total number of protons and
neutrons in a nucleus.

Atomic number: number of protons

A nuclide: an atom characterized by a definite atomic
number and mass number
15
Exercise 2.1
A nucleus consists of 17 protons and 18 neutrons.
What is its nuclide symbol?
17+18 = 35 = atomic mass = Cl
35
Cl
17
16
 Dalton’s hydrogen-based atomic mass scale replaced by
oxygen-based scale.
 Later replaced with present day carbon-12 mass scale
(1961).
 One amu (atomic mass unit) is a mass unit equal to
exactly 1/12 the mass of a carbon-12 atom.
 Atomic mass is the average atomic mass for the
naturally occurring element expressed in atomic mass
units.
17
The Structure of Atoms
Exercise 2.2
Chlorine has two naturally occurring isotopes:
35
17
Cl
with an abundance of 75.771% and an isotopic mass
of 34.96885 amu, and
37
17
Cl with an abundance of
24.229% and an isotopic mass of 36.96590 amu. What is
the atomic mass of chlorine?
18
Exercise 2.2 cont…
34.96885 amu x 0.75771 = 26.496247
36.96590 amu x 0.24229 = 8.956467
26.496247 + 8.956467 = 35.452714 = 35.453 amu
Therefore the atomic mass unit for chlorine is 35.453amu
19
The Structure of Atoms
 Dmitri Mendeleev: founder of the Periodic Table of
elements.
 Elements arranged in order of atomic masses; horizontal
rows one row under the other (PERIOD).
 Elements in each vertical column have similar properties
(GROUP).
20
 Group I: Alkali metals: soft metals that react easily with
water.
 Group II: Alkali earth metals: metals that form oxides and
have an earthy texture to yield alkaline solutions.
 Inner group: Transition metals: characteristic luster or shine
and is generally a good conductor of electricity.
 Group 7: Halogens are highly reactive “salt” formers.
 Group 8: Inert gases are non-reactive gases.
21
Periods: Seven horizontal rows
Groups: 18 vertical columns, based on similar chemical properties
22
Metals: element with characteristic luster or shine and is generally
a good conductor of electricity.
 Solids at room temperature, fairly malleable and ductile.
 Mercury is the only liquid metal @ room temperature.
Metalloid: element showing both metal and non-metal properties.
Si and Ge are good semi-conductors (poor conductors at
room temperature but good conductors at higher
temperatures.
Nonmetal: element that does not exhibit the metallic properties.
 Most nonmetals are gases ( chlorine, oxygen) or solids
(phosphorus, sulfur).
 Solid nonmetals are usually hard, brittle substances.
 Bromine is the only liquid non-metal.
23 23
The Periodic Table
24
Exercise 2.3
By referring to the Periodic table, identify the group and the
period to which each of the following elements belong. Then
decide whether the element is metal, nonmetal or a metalloid.
A) Se: Group VIA, Period 4, nonmetal
B) Cs:
C) Fe: Group VIIIB; Period 4, metal
D) Cu:
E) Br: Group VIIA, Period 4, nonmetal (Halogen)
25
The Periodic Table
What are the atomic numbers for the following
elements and how many protons in each?
A) Copper = 29 = 29 protons
B) Sodium
C) Sulfur = 16 = 16 protons
D) Oxygen
E) Hydrogen = 1 = 1 proton
26
The Periodic Table
What are the atomic masses for the following
Elements and how many protons and neutrons
in each?
A) Iron = 55.845amu = 26 protons (56-26=30) 30 neutrons
B) Magnesium =
C) Bromine = 79.904amu = 35 protons = (80-35=45) 45 neutrons
D) Xenon =
E) Carbon = 12.0107amu = 6 protons = (12-6=6) 6 neutrons
27
The Periodic Table
What are the mass numbers for the following
elements?
1)
Chlorine = 35
2)
Nitrogen =
3)
Fluorine = 19
4)
Zinc =
5)
Silicon = 28
28
The Structure of Atoms
 The isotope
79
34 Se
is used medically for
diagnosis of pancreatic disorders. How many
protons, neutrons, and electrons does an
atom of
79
34 Se
have?
Protons = 34 = electrons
Neutrons = 45
79 - 34 = 45
29
The Structure of Atoms
 An atom of element X contains 47 protons and
62 neutrons. Identify the element, and write the
symbol for the isotope in the standard format.
47
X
109
47
47
Ag
109
62 neutrons
Ag
108
61 neutrons
30
Periodic Table
Element
Electrons
Protons
Neutrons
P
15
15
16
Na+
12
S2-
18
Li
3
Ca2+
18
Cl-
16
16
20
20
17
31
Molecular and Ionic Substances
 Molecule: definite group of atoms that are
chemically bonded together (tightly
connected by attractive forces).
 Molecular formula: gives the exact number of
different atoms of an element in a molecule.
 H2 O 2
 NH3
 C2 H 6 O
32
Atoms, Molecules, and Ions
 Covalent Bonding (Molecules): The most
common type of chemical bond is formed
when two atoms share some of their
electrons.
(non-metal -- non-metal)
33
Atoms, Molecules, and Ions
Naming Binary Molecular Compounds:
 The more cationlike element uses its
elemental name.
 The more anionlike element substitutes the
second half of its elemental name with –ide.
 Use the Greek prefixes to express the number
of each element present.
34
Greek Prefixes
35
Atoms, Molecules, and Ions
Examples:
CO carbon monoxide
CO2 carbon dioxide
SF4 sulfur tetrafluoride
Name:
NCl3 =
Write formulas:
Disulfur dichloride =
P4O6 = tetraphosphorus hexoxide
Iodine monochloride = ICl
S2F2 =
Nitrogen trioxide =
36
Atoms, Molecules, and Ions
37
Atoms, Molecules, and Ions
 Ionic Bonding (Ionic Solids): These are formed
by a transfer of one or more electrons from one
atom to another. (metal -- non-metal)
38
Example 9
Which of the following drawings represents an
ionic compound? Molecular compound?
39
Atoms, Molecules, and Ions
Naming Binary Ionic Compounds:
 Identify the positive ion and then the negative
ion.
 The positive ion uses its elemental name.
 The negative ion substitutes the second half of
its elemental name with –ide.
 Do not use Greek prefixes such as mono–, di–,
or tri–.
 Use roman numerals for transition metals
40
Atoms, Molecules, and Ions
41
Naming Ionic Compounds
1.
NaCl = sodium chloride
2.
MgS =
3.
Ba3N2 = barium nitride
4.
CaO =
5.
K2S = potassium sulfide
6. FeCl2 = iron(II) chloride
7.
FeCl3 =
8. CrO2 = chromium(II) oxide
9. ZnCl2 =
10. V2O3 = vanadium oxide
42
Ionic Compounds
1. Calcium chloride
2. Copper (II) sulfide =
CuS
3. Sodium nitride
4. Silver bromide =
AgBr
5. Nickel (II) phosphide =
Ni3P2
6. Cesium oxide
7.
Strontium iodide =
SrI2
8.
Cobalt (II) sulfide
43
Atoms, Molecules, and Ions
44
Atoms, Molecules, and Ions
 Naming Ionic Compounds Containing
Polyatomic Ions :
 Same as binary ionic compounds
 But use the name provided for the polyatomic
ion
45
Atoms, Molecules, and Ions
 Examples
 CaCO3
 FeCrO4
 KOH
Calcium carbonate
Potassium hydroxide
Name:
Ba3(PO4)2 =
Na2SO4 = sodium sulfate
Sn(ClO4)4 =
Write the Formula:
 Iron(II) permanganate =
 Cesium nitrate = CsNO3
 Zinc acetate =
46
Exercise 2.4
Potassium chromate is an important compound of
chromium. It is composed of K+ and CrO42- ions.
Write the formula of the compound.
K+
CrO42K2CrO4
47
Exercise 2.5
Write the names of the following compounds
a) MgO =
b) PbCrO4 =
c) Mg3N2 =
d) CrSO4 =
48
Exercise 2.6
A compound has the name thallium (III) nitrate.
What is its formula?
Tl3+
NO3-
Tl(NO3)3
49
Ionic and Covalent Bonding
Which of the following is a covalent compound?
1) NaCl
2) NaOH
3) H2O
4) AlCl3
50
Atoms, Molecules, and Ions
Acid: A substance that provides H+ ions in H2O
Base: A substance that provides OH- in H2O
Oxoacid: Contain oxygen and hydrogen and
another element
51
Atoms, Molecules, and Ions
Naming acids: When acid is dissolved in water
gives one or more H+ and a polyatomic oxoanion,
(has to have (aq))
Name of acid is based on the oxoanion
52
Atoms, Molecules, and Ions
53
Atoms, Molecules, and Ions
Name the following acids:
(a) HBrO3(aq) = bromic acid
(b) HCN(aq) =
(c) HIO3(aq) = iodic acid
(d) HMnO4(aq) =
(e) H2CrO4(aq) = chromic acid
54
Balancing Chemical Equations
 A balanced chemical equation represents the
conversion of the reactants to products such
that the number of atoms of each element is
conserved.
Calcium carbonate
CaCO3(s)
calcium oxide + carbon dioxide
CaO(s) + CO2(g)
55
Balancing Chemical Equations
 Balancing Equations: write unbalanced
equation
A2 + B2
A2B
 Use coefficients to indicate how many
formula units are required to balance the
equation:
2 A2 + B2
2 A2B
56
Balancing Chemical Equations
 Method 1 (suggested)
 Balance those atoms which occur in only one
compound on each side
 Balance remaining atoms
 Reduce coefficients to smallest whole integers
 Check your answer
57
Balancing Chemical Equations
• Method 2
– Identify most complex compound
– Balance this compound by placing 1 before it
– Balance remaining compounds using
fractions
– Multiply fractions to obtain integers
58
Balancing Chemical Equations
 Balance the following equations
C6H12O6
Fe + O2
NH3 + Cl2
C2H6O + CO2
Fe2O3
N2H4 + NH4Cl
KClO3 + C12H22O11
KCl + CO2 +
H2 O
59
Exercise 2.7
Name the following compounds:
a) N2O4 =
b) Cl2O6 =
c) PCl3 =
d) PCl5 =
60
Exercise 2.8
Give formulas for the following compounds:
a) Disulfur dichloride =
b) Tetraphosphorus trisulfide =
c) Carbon disulfide =
d) Sulfur trioxide =
61
Exercise 2.10
Selenium has an oxoacid H2SeO4, called selenic
acid. What is the formula and name of the
corresponding anion?
SeO42- selenate ion
What are the name and formula of the anion
corresponding to perbromic acid HBrO4?
BrO4- perbromate ion
62
Exercise 2.11
Epsom salts has a formula MgSO4.7H20. What is the
chemical name of this compound?
Washing soda has the formula Na2CO3.10H2O. What is
the chemical name of this compound?
What is the chemical formula for sodium thiosulfate
pentahydrate?
63
Exercise 2.12
Balance the following equations:
a) H3PO3
H3PO4
b) Ca + 2H2O
+ PH3
Ca(OH)2 + H2
c) Fe2(SO4)3 + NH3 + H2O
d) O2 + PCl3
POCl3
e) P4 +
P4O6 + N2
f)
N2O
As2S3 + O2
g) Ca3(PO4)2 + H3PO4
Fe(OH)3 +
(NH4)2SO4
As2O3 + SO2
Ca(H2PO4)2
64