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Chapter 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 (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
• Explained Laws of Conservation of
Mass and Constant Composition and
extended them to cover another law.
7
Main ideas of Dalton’s model
1. All matter consists of 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) and
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
• 80 g oxygen, forming Compound IV
What is the ratio 16:48:64:80
expressed as small whole numbers?
• Compounds I–IV are N2O, N2O3, N2O4, N2O5
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
Cause
stream of
negative
particles that
are always
the same, no
matter what
gas is used
Thomson experimented with
CATHODE RAY TUBES
13
1897 JJ Thomson
Mass to charge ratio for an Known as
electron:
discoverer of
m/c = 5.69 x 10-9g/coulomb the
ELECTRON—
led to the
“plum pudding
model” of the
atom
14
Millikan
• Obtained the charge of an electron, which
coupled with Thomson’s work, allowed the
calculation of the mass of an electron.
15
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
16
The modern view of the atom was
developed by Ernest
Rutherford of New Zealand
(1871-1937).
17
Ernest Rutherford
Canterbury
University in
Christchurch, NZ
Rutherford laboratory
18
Gold Foil Experiment
Screen 2.9
19
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
Mass of proton about same as an H
atom (1 atomic mass unit)
Positive charge = negative charge
from electrons in a neutral atom.
22
Neutrons (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
23
Summary
24
Atomic Symbols
An atomic symbol represents the
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, is 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) but different mass
number (A).
• Boron-10 has 5 p and 5 n: 105B
• Boron-11 has 5 p and 6 n: 115B
11B
10B
28
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
29
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
Definition:
•
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.
19 1
F
9
33
Sample Problem
• Example 2.2 Write the atomic symbols for
the following:
• a. a species having 16 protons, 16
neutrons and 18 electrons
• b. the phosphide ion (P) with an overall
charge of -3
34
Solution
32
S 2-
16
31
P 3-
15
35
Atomic Mass
• F. 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). 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.
(also called atomic weight)
13
Al
26.981
Atomic number
Atom symbol
Atomic mass or weight
37
Atomic Mass
• 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
p.t. to nearest whole number
13
Al
26.981
Atomic number
Atom symbol
Atomic mass or weight
39
Atomic Mass
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%).
40
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
– High speed electrons: β (beta) particles
– Alpha (α) particles: helium nuclei
– Gamma (γ) rays: high energy light
41
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
42
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
Periods or Series
• Horizontal rows are periods
• 7 periods
total
– First period is H and He
– Second period is Li to Ne
– Etc.
46
Group Names to Memorize
Group 1 (IA): alkali metals.
- Group 2 (IIA) : alkaline earth
metals.
- Group 17(VIIA): halogens.
- Group 18 (VIIIA): noble
gases
-
47
Group 1A: Alkali Metals
Li, Na, K, Rb, Cs
Reaction of
potassium + H2O
https://www.youtube.com/watch?v=oqMN3y8k9So
https://www.youtube.com/watch?v=Jy1DC6Euqj4
Cutting sodium metal
48
Group 2A: Alkaline Earth Metals
Be, Mg, Ca, Sr, Ba, Ra
Magnesium
Magnesium
oxide
https://www.youtube.com/watch?v=qSr39UwpELo
49
Group 7A:
Halogens
F, Cl, Br, I, At
50
Group 8A: Noble Gases
He, Ne, Ar, Kr, Xe, Rn
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/Nonmetals/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 nonmetals- B, 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)
• 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. Argued that
element properties are
periodic functions of their
atomic weights.
57
Periodic Table
Periodic Law:
• We now know that element
properties are periodic
functions of their ATOMIC
NUMBERS.
58
Germanium:
Prediction vs. Observation
59
Henry Moseley
• A student of Rutherford’s
• Arranged the periodic table in
order of increasing atomic
number
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 the 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 purple flashcard sheet!
64
Ionic Compounds
• 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
There is NO net charge
unit.
in an ionic compound!
65
Solutions of Ionic Compounds
• 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
Charge Balance of Ionic
Compounds
• See Handout and practice worksheets.
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
•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
Grp 13 metals (e.g., Al 3+) and Grp 15
non-metals/metalloids (e.g., N 3-)
69
• Some metal ions have > one
possible charge. A Roman
•
numeral are used for the charge.
• If a metal only has ONE charge, a
Roman numeral is NOT used.
70
Symbols and Periodic Table
Locations of Some Monatomic Ions
Titanium forms both
titanium(II) and
titanium(IV) ions.
Copper forms either
copper(I) or copper(II) ions.
71
Polyatomic Ions
• See handouts: MUST MEMORIZE!!!
72
Polyatomic Ions
• Oxyanions: the anions are composed of
oxygen and one other element
73
Nitrate and Sulfate
74
Polyatomic Ions
• Oxyanions: the anions are composed of oxygen
and one other element
Ex:SO42- (sulfate), NO2- (nitrite) , MnO4(permanganate)
• 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: SO32sulfite,
NO2- nitrite, PO3-3
phosphite
SO42sulfate, NO3- nitrate, PO4-3
phosphate
75
Four oxyanions
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
76
Naming binary compounds
• Use name of metal with no changes
• Change the name of the anion by taking
the “stem” and add the suffix –ide
ex: CI2: chlorine
Cl-: chlorine ( ine ) + ide = chloride
• Examples:
– NaCl - sodium chloride
– MgCl2 - magnesium chloride
77
Naming Binary Ionic Compounds
• Name the following binary ionic compounds
•
Metal nonmetal compound
name
• KI
potassium iodine potassium iodide
• Li2S lithium
sulfur
lithium sulfide
• Mg3N2 magnesium nitrogen magnesium nitride
78
Metals with multiple oxidation
states
• Two methods: Stock and “classical” system
Stock system (used at CHS)
 metal name and the oxidation state in Roman numbers in
parenthesis
Ex: Fe2+ = iron (II)
• Form compound by balancing charge of metal with
correct number of nonmetals
Ex: CoCl3 = cobalt(III) chloride
Charge of metal = |charge of anion x subscript|
subscript of cation
79
Sample Problem
Example 2.4 Write names or formulas for:
•
•
•
•
rubidium bromide
barium nitride
cobalt (II) bromide
Strontium oxide
AlCl3
Ca3P2
NaI
PbS2
80
Solution
•
•
•
•
RbBr
Ba3N2
CoBr2
SrO
Aluminum Chloride
Calcium Phosphide
Sodium Iodide
Lead (IV) sulfide
81
Sample Problem
Example 2.5 Write names or
formulas for:
•
•
•
•
(NH4)2S
K2Cr2O7
Al(NO2)3
Fe(CN)2
strontium hydroxide
cobalt (II) sulfate
calcium phosphate
tin (IV) carbonate
82
Solution
•
•
•
•
ammonium sulfide
potassium dichromate
aluminum nitrite
Iron (II) cyanide
Sr(OH)2
CoSO4
Ca3(PO4)2
Sn(CO3)2
83
Binary Molecular Compounds
• (Two nonmetals bonded together; may
also include a metalloid in formula)
e.g., CO, NO, HF, SiO2
• a. First symbol is usually element to
furthest left in p.t.
• b. Numbers of atoms indicated by
subscripts are written as prefixes.
84
Binary Molecular Compounds
• The name consists of two words.
• Directions:
•
1. Write name of first element
preceded by prefix EXCEPT do not write
mono- if only have ONE of first element.
•
2. Name of second element ends with
-ide; is also preceded by prefix.
85
Names of Binary Compounds
86
Example
• Consider the compounds CO and CO2
Name the element that appears first in the formula:
CARBON
The second element has an altered name: retain the stem
of the element name and replace the ending by -ide
OXYGEN
 OXIDE
87
Sample Problem
Example 2.4 Write names or formulas
for the following:
B2O3
tetraphosphorus pentachloride
AsO5
dihydrogen monoxide
As2O7
88
Solution
• Diboron trioxide
• Arsenic pentoxide
• Diarsenic heptoxide
P4Cl5
H 2O
89
Acknowledgements
• Thomson/Brooks Cole (Textbook
Publishers)
• Mark P. Heitz, State University of New
York at Brockport (Prentice Hall, Book
Publishers)
90