Unit 2:
Atoms, Molecules and Ions
The Language of Chemistry
Atoms
◦Composed of electrons, protons and
neutrons
Molecules
◦Combinations of atoms
Ions
◦Charged particles
2
Laws of Chemical Composition
(1790) Antoine Lavoisier, The Father of Modern Chemistry
Law of Conservation of Matter
Total mass remains constant during a chemical
reaction; or
Total mass of reactants = total mass of products.
3
Law of Conservation of Mass:
A Conceptual Example
Jan Baptista van Helmont (1579–1644) first
measured the mass of a young willow tree and,
separately, the mass of a bucket of soil and then
planted the tree in the bucket. After five years, he
found that the tree had gained 75 kg in mass even
though the soil had lost only 0.057 kg. He had added
only water to the bucket, and so he concluded that
all the mass gained by the tree had come from the
water. Explain and criticize his conclusion.
4
Laws of Chemical Composition
Joseph Proust, Law of Constant Composition
(a.k.a. Law of Definite Composition or Definite
Proportions)
All samples of a compound have the same
composition, or all samples have the same
proportion by mass of the elements present.
5
Law of Constant Composition: Example
Example: CuHCO3 is ALWAYS
57.48% Cu, 5.43% C, 0.91% H and 36.18% O by mass
6
John Dalton and the Atomic
Theory of Matter
Importance:
Dalton explained the Laws of Conservation
of Mass and Constant Composition and
extended them to cover another law…
The Law of Multiple Proportions
7
Main ideas of Dalton’s model
1. All matter consists of small, indivisible particles called
atoms.
2. All atoms of a given element are alike but atoms of any one
element are different from the atoms of every other
element.
3. Compounds are formed when atoms of different elements
unite in small, whole-number ratios.
4. Chemical reactions involve rearrangement of atoms; no
atoms are created, destroyed or broken apart in a
chemical reaction.
According to Dalton, atoms are indivisible and
indestructible.
8
Dalton’s Atomic Theory:
Conservation of Mass and Definite Proportions
Six fluorine atoms and four
hydrogen atoms before reaction …
… six fluorine atoms and four
hydrogen atoms after reaction.
Mass is conserved.
HF always has one H atom
and one F atom; always
has the same proportions
(1:19) by mass.
9
Another Important Law
Law of Multiple Proportions
A given set of elements may combine to produce
two or more different compounds, each with a unique
composition.
Example: H2O (water) &
H2O2 (hydrogen peroxide)
10
Law of Multiple Proportions (cont’d)
Four different oxides of nitrogen can be formed by
combining 28 g of nitrogen with:
16 g oxygen, forming Compound I
48 g oxygen, forming Compound II
64 g oxygen, forming Compound III
Compounds I–IV
are N2O, N2O3,
N2O4, N2O5
80 g oxygen, forming Compound IV
What is the ratio 16:48:64:80
expressed as small whole numbers?
11
Dalton’s Model of the Atom
NO subatomic particles!
In modern atomic theory, the atom is
divided into protons, neutrons and
electrons.
12
1897 JJ Thomson
Thomson
experimented with
CATHODE RAY
TUBES
Cathode Ray Tube Demonstration
-Caused stream of
negative particles
that were always
the same, no
matter what gas
was used
13
1897 JJ Thomson
Thomson known as discoverer of the
ELECTRON—led to the “plum pudding model”
of the atom
14
15
Robert Millikan:
Oil Drop Experiment
Obtained the charge of an electron, which coupled
with Thomson’s work, allowed the calculation of the
mass of an electron.
Oil Drop
Experiment
Demo
16
Millikan’s Conclusions
Measured the charge of an electron:
1.602 x 10-19 coulomb (C)
Calculated the mass of an electron:
9.109 x 10-31 kg
17
The modern view of the atom was
developed by Ernest Rutherford of
New Zealand (1871-1937).
18
Ernest Rutherford
Canterbury
University in
Christchurch, NZ
Rutherford laboratory
19
Gold Foil Experiment
Gold Foil
Experiment
Demonstration
20
Rutherford’s Main Conclusions
1. The atom is mostly empty space.
2. All of the positive charge, and most of the mass,
is concentrated in a very small volume:
THE NUCLEUS!
3. Electrons are outside the nucleus.
21
Protons
The mass of a proton is about the same as
the mass of an H atom (1 atomic mass unit)
In a neutral atom…
Positive charge from protons =
negative charge from electrons
22
Neutrons
(James Chadwick, 1932)
The nucleus also contains neutrons:
particles with masses almost identical to
protons but with no charge
***Neutrons also help disperse the strong
repulsion of positive charges within the
nucleus due to the protons in the nucleus.
23
Summary
24
Atomic Symbols
Atomic symbol – the letter or letters that
represent an element.
13
Al
26.981
Atomic number
Atom symbol
Atomic mass or weight
25
Mass Number, A
The Mass Number (A) = # protons + # neutrons
A Boron atom can have:
A = 5 p + 5 n = 10 amu
A
10
Z
5
B
Named as boron-10
26
Atomic Number, Z
Atomic number, Z =
the number of protons in the nucleus.
(same for every atom of that element)
13
Al
26.981
Atomic number
Atom symbol
Atomic mass or weight
27
Isotopes
Atoms of the same element (same Z)
with different mass numbers (A).
Boron-10 has 5 p and 5 n: 105B
Boron-11 has 5 p and 6 n: 115B
11B
10B
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Hydrogen Isotopes
Hydrogen has 3 isotopes:
1 H
1
1 proton and 0
neutrons, protium
2 H
1
1 proton and 1
neutron, deuterium
3 H
1
1 proton and 2
neutrons, tritium
***radioactive
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Isotopes &
Their Uses
Heart scans with
radioactive
technetium-99.
99
43Tc
Emits gamma rays
30
Sample Problem
Example 2.1
Write the atomic symbols for the following species:
a. the isotope of carbon with a mass of 13
b. the nuclear symbol when Z = 92 and the number of
neutrons = 146.
31
Solution to Problem
13
C
6
238
U
92
32
Ions: particles w/ a charge
Atoms GAIN electrons to become negative ions, or anions.
Atoms LOSE electrons to become positive ions, or cations.
How are ions represented?
Charges are always shown to upper right of symbol.
33
Sample Problem
Example 2.2 Write the atomic symbols for the following:
a. a species with 16 protons, 16 neutrons and 18 electrons
b. the phosphide ion (P) with an overall charge of -3
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Solution
32
S 2-
16
31
P 3-
15
35
Atomic Mass
An atomic mass unit (amu or u) is defined as
exactly one-twelfth the mass of a carbon-12 atom
1 u = 1.66054 × 10–24 g
The atomic mass of an element is the relative mass
of an atom compared to a standard (carbon-12
atom). It is NOT equal to the mass number!
36
Atomic Mass Is Not Equal
to Mass Number!!
The atomic mass is a weighted average of the
masses of the naturally occurring isotopes.
(Atomic mass is also called atomic weight)
13
Al
26.981
Atomic number
Atom symbol
Atomic mass or weight
37
Atomic Mass
The weighted average is the addition of the
contributions from each isotope.
Isotopic Abundance is the percent or
fraction of each isotope found in nature.
38
Most Abundant Isotope
Usually can round atomic mass on the
periodic table to nearest whole number
(but not always!!)
13
Al
26.981
Atomic number
Atom symbol
Atomic mass or weight
39
Example 2.3 Determine the average atomic mass of
magnesium which has three isotopes with the following
masses: 23.98 (78.6%), 24.98 (10.1%), 25.98 (11.3%).
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Radioactivity
Radioactive isotopes are unstable
◦ These isotopes decay over time
◦ Emit other particles and are transformed into
other elements
◦ Radioactive decay is not a chemical process!
Particles emitted
◦ Beta (β) particles: High speed electrons
◦ Alpha (α) particles: helium nuclei
◦ Gamma (γ) rays: high energy light
41
Nuclear Stability
Nuclear stability depends on the
neutron/proton ratio.
◦ For light elements, n/p is approximately 1
◦ For heavier elements, n/p is approximately 1.4/1
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Figure 2.5 – The Nuclear Belt of Stability
43
The Periodic Table:
Elements Organized
Know location and description of:
◦ groups or families
◦ periods or series
◦ metals, metalloids, nonmetals and their
properties
◦ main group elements
◦ transition metals
◦ lanthanides and actinides
44
Groups or Families
Vertical columns are groups
◦ Numbered as 1-18 (new)
◦ Old system uses Roman numerals and A,B
45
Group Names to Memorize
Group 1 (IA): alkali metals
Group 2 (IIA) : alkaline earth metals
Group 17(VIIA): halogens
Group 18 (VIIIA): noble gases
46
Group 1: Alkali Metals
Li, Na, K, Rb, Cs
Reaction of
potassium + H2O
Cutting sodium metal
Alkali Metal Video
Potassium in Water Demo
47
Group 2: Alkaline Earth Metals
Be, Mg, Ca, Sr, Ba, Ra
Magnesium
Magnesium
oxide
Burning Magnesium Ribbon Demo
48
Group 17: Halogens
F, Cl, Br, I, At
49
Group 18: Noble Gases
He, Ne, Ar, Kr, Xe, Rn
50
Periods or Series
Horizontal rows are periods
◦7 periods total
◦First period is H and He
◦Second period is Li - Ne
◦ Etc.
51
Regions of the Periodic Table
Metals are on
the left of stair
step line
NON-METALS
are on the right
of stair step line
52
Exception:
Group 1A: Hydrogen is a Non-metal!
Shuttle main engines
use H2 and O2
53
Properties of Metals/
Non-metals/Metalloids
Metals – shiny, smooth, solid at room temperature,
good conductors of heat and electricity, malleable and
ductile.
Metalloids (along stair step line) – physical and
chemical properties of both metals and nonmetalsB, Si, Ge, As, Sb, Te
Nonmetals – low melting and boiling points, brittle,
dull-looking solids, poor conductors of heat and
electricity.
54
The Periodic Table:
Elements Organized
Main group elements – tall columns
(Groups 1, 2, 13, 14, 15, 16, 17, 18)
Transition metals – short columns (10 grps
total, Groups 3-12)
Lanthanides and actinides – long rows
below main part of table
55
Transition Elements
Lanthanides and actinides
Iron in air gives
iron(III) oxide
56
Periodic Table
Dmitri Mendeleev
developed the modern
periodic table. He argued
that element properties are
periodic functions of their
atomic weights.
57
Germanium:
Prediction vs. Observation
58
Henry Moseley
A student of Rutherford’s
Arranged the periodic table in
order of increasing atomic number
59
Periodic Table
Periodic Law:
We now know that element
properties are periodic
functions of their ATOMIC
NUMBERS.
60
Molecules
A molecule is a group of two or more atoms
held together in a definite shape by
covalent bonds.
61
Empirical and Molecular Formulas
Empirical formula: the simplest whole
number ratio of elements in a compound
Molecular formula: gives the ACTUAL
number of each kind of atom in a molecule.
Example: Molecular formula of glucose – C6H12O6
Can divide all subscripts by 6, so empirical formula is CH2O
62
Structural Formulas
Structural formulas show how atoms are attached
to one another.
63
Ions: Atoms with a Charge
Definition:
Cations: positive ions
Anions: negative ions
Polyatomic ion: A group of atoms with a charge
◦ You must memorize all the polyatomic ions (structure,
name, and charge) found on your flashcard sheet!
64
Ionic Compounds are cations
and anions held together by
electrostatic attraction.
Their formulas are the
simplest ratio of numbers of
atoms (called an empirical
formula) and represent one
formula unit.
There is NO net charge
in an ionic compound!
65
Solutions of Ionic Compounds
are strong electrolytes: their
solutions conduct electricity.
Non-electrolytes do not
conduct electricity in water
solution. (sugar, molecular
compounds)
There is NO net charge
in an ionic compound!
66
•Atoms that are close to a noble gas (group 18) form
ions that contain the same number of electrons as the
neighboring noble gas atom
•Applies to Groups 1, 2, 16 and 17, plus Group 13
metals (e.g., Al 3+) and Group 15 nonmetals/metalloids
(e.g., N 3-)
67
Monatomic Ions
• Group IA metals form ions of +1 charge.
• Group IIA metals form ions of +2 charge.
• Aluminum, a group IIIA metal, forms ions with a
+3 charge.
• Nonmetal ions of groups V, VI, and VII usually
have charges of:
• VA: -3
• VIA: -2
• VIIA: -1
68
Writing Formulas for Ionic Compounds
• We can use the charges of the ions in a
compound to easily determine the formula of
the compound.
• All you do is crisscross the charges to make
them subscripts
– Ex: magnesium nitride
+2
-3
Mg3 N2
Don’t forget to drop the positive/negative signs
once they’re subscripts!!
PRACTICE, PRACTICE, PRACTICE!!
Write formulas for the following ionic compounds:
Magnesium chloride
Calcium bromide
Sodium oxide
Sodium chloride
Potassium iodide
Polyatomic Ions
• There are some ions that are made of more
than one element
– These ions are called POLYATOMIC IONS
• “poly” meaning more than 1, “atomic” meaning atoms 
more than 1 type of atom in the ion
• Examples: (see chart in your notes for more)
– Sulfate
SO42– Hydroxide
OH– Nitrate
NO3-
Oxyanions: anions are composed of oxygen and one other element
Ex: SO42- (sulfate), NO2- (nitrite) , MnO4- (permanganate)
If there are two oxyanions of the same element:
(a) The anion with the smaller number of oxygens uses the
roots of the element plus “ite”
(b) The higher number use the root plus “ate”
Ex: SO32- sulfite, NO2- nitrite, PO3-3 phosphite
SO42- sulfate, NO3- nitrate, PO4-3 phosphate
72
There are four oxyanions containing Cl
The middle two are named as two oxyanions
The one with one less oxygen than the chlorite has a prefix of hypo-
The one with one more oxygen than the chlorate has a prefix of per
Ex:
ClO-
: hypochlorite
ClO2- : chlorite
ClO3- : chlorate
ClO4- : perchlorate
73
Writing Formulas for Compounds with
Polyatomic Ions
• You still crisscross the charges like you did for basic
ionic compounds but you use the charges of the metal
ion and the polyatomic ion
• You may need to use parenthesis to prevent
confusion!!
Example:
Barium phosphate
PRACTICE WITH POLYATOMICS!
Write formulas:
calcium phosphate
lithium carbonate
beryllium bicarbonate
 Some metal ions have more than one
possible charge. A Roman numeral is
used for the charge.
 If a metal only has ONE charge, a Roman
numeral is NOT used.
76
Symbols and Periodic Table Locations of
Some Monatomic Ions
Copper forms either
Titanium forms both
titanium(II) and
titanium(IV) ions.
copper(I) or copper(II) ions.
77
Writing Formulas w/ Transition Metals
Writing Formulas of Compounds w/ Transition Metals:
1. Write the metal’s atomic symbol with the value of the
Roman numeral as a positive charge
2. Write the value of the anion (negatively charged ion)
with it’s charge
3. Crisscross the charges to get your formula
Example:
Lead (IV) hydroxide
6.4 Naming Ionic Compounds
• When naming ionic compounds, we take the
name of the metal followed by the name of
the non-metal, only we drop the last syllable
and add –ide as an ending.
– Examples: NaCl: Sodium Chloride
Al2O3: Aluminum Oxide
PRACTICE, PRACTICE, PRACTICE!
Write names for the given ionic compounds:
CaCl2
LiBr
BeO
AlCl3
Ra3N2
Naming Ionic Compounds with
Polyatomic Ions
• The rules for naming polyatomic ions are the
same rules for naming normal ionic
compounds except you use the unchanged
polyatomic ion name for the second word
Example:
MgSO4 – magnesium sulfate
LiNO3 – lithium nitrate
Ba(OH)2 – barium hydroxide
PRACTICE WITH POLYATOMICS!
Name the compounds:
RbNO3
NaHCO3
Mg(OH)2
Naming & Writing Formulas w/
Transition Metals
• Transition metals – all elements in groups 3-12
and elements in groups 13, 14, and 15 below the
stair step line
– Can have more than one charge in ion form!
• Except Ag+ and Zn2+
• When naming compounds involving transition
metals, you need to include a roman numeral in
parenthesis to indicate the charge of the ion
Naming Compounds w/ Transition Metals
1. Name the metal from its symbol
2. Determine the charge on the metal by
“uncrossing” the charges on the positive and
negative ions
3. Write the charge as a roman numeral in
parenthesis
4. Write the name of the anion either as an element
with an “–ide” ending or as the unchanged name
of the polyatomic ion
Charge of metal = |charge of anion × anion subscript|
Naming Compounds w/ Transition Metals
Examples:
CrCl6
Fe(OH)2
NiO
Writing Formulas for Covalent Compounds
(2 non-metals bonded together)
1. Write the symbol for the first element listed
• If there is a prefix on the first element, write the
prefix value as a subscript attached to the first
element’s symbol
2. Write the symbol for the second element listed
• If the prefix on the second element is “di-” or
greater, write the prefix value as a subscript attached
to the second element’s symbol
Writing Formulas for Covalent Compounds
(2 non-metals bonded together)
Examples:
Selenium dioxide
Carbon tetrachloride
Nitrogen trihydride
Naming Covalent Compounds
(2 non-metals bonded together)
Given the formula:
1. Write the names of the 2 elements
2. Attach a Greek prefix (see chart) to the
beginning of the first element only if it has a
subscript greater than or equal to 2
3. Attach a Greek prefix to the second element
no matter what!
4. Change the ending of the second element’s
name to an –ide ending
Naming Covalent Compounds
(2 non-metals bonded together)
Examples:
P2O6
CO2
NO