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Chapter #2 – Atoms, Molecules and Ions
2.1 The Early History of Chemistry
2.2 Fundamental Chemical laws
2.3 Dalton’s Atomic Theory
2.4 Cannizzaro’s Interpretation
2.5 Early experiments to Characterize the Atom
2.6 The Modern View of Atomic Structure: An
Introduction
2.7 Molecules and Ions
2.8 An Introduction to the Periodic Table
2.9 Naming Simple Compounds
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Chapter 2 Topics
1: The Observations That Led to an Atomic View
of Matter
2: The Observations That Led to the Nuclear Atom
Model
3: Dalton’s Atomic Theory and Today’s Version
4: Molecules and Ions
5. Elements: A First Look at the Periodic Table
6: Compounds: Introduction to Bonding
7: Compounds: Formulas and Names
2
1
Law of Conservation of Mass
The total mass of substances does not change during a
chemical reaction.
3
Law of Definite Proportions
No matter what its source, a particular
chemical compound is composed of
the same elements in the same parts
(fractions) by mass.
4
2
Law of Definite Proportions
WATER  H2O
No matter what the source
water is ALWAYS
2 parts hydrogen to 1 part oxygen
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Law of Definite Proportions
Chemical analysis of a 9.07 g sample of calcium
phosphate shows that it contains 3.52 g of Ca.
How much Ca could be obtained from a 1.000 kg
sample?
1.
2.
3.
4.
0.388 kg
3.52 kg
38.8 kg
0.38 kg
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3
Law of Definite Proportions
Chemical analysis of a 9.07 g sample of calcium
phosphate shows that it contains 3.52 g of Ca.
How much Ca could be obtained from a 1.000 kg
sample?
Mass fraction of Ca =
3.52 g Ca = 0.388 * 100% = 38.8%
9.07 g sample
(i.e., ANY sample of Ca3 (PO4)2 is 38.8% Ca by mass )
Mass of Ca in 1.000 kg of sample =
1.000 kg sample 38.8 kg Ca
100 kg sample
= 0.388 kg Ca
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Law of Multiple Proportions
In a nutshell, two (or more) compounds can
contain different relative amounts of the same
elements:
If elements A and B react to form two compounds,
the different masses of B that combine with a fixed
mass of A can be expressed as a ratio of small
whole numbers.
(Evidence of the existence tiny individual particles.)
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4
Law of Multiple Proportions
Mass of Oxygen that Combines
with 1.00 g of Carbon
Compound #1
1.33 g O per g C
Compound #2
2.66 g O per g C
mass of O in compound #2 = 2.66 g = 2
mass of O in compound #1
1.33 g
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EXACT 2:1 RATIO
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10
5
Atomic Theory
Protons, Neutrons and Electrons
Theory and Discovery
11
Daltons Atomic Theory
1.
All matter consists of tiny particles called
atoms.
2.
Atoms of an element are identical in mass
and other properties and are different from
atoms of any other element.
3.
Compounds result from the chemical
combination of a specific ratio of atoms of
different elements.
4.
Chemical reactions involve reorganization
of the atoms – changes in the way they are
bonded. Atoms of one element do not
change and cannot be converted into
atoms of another element during chemical
reactions.
John Dalton
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6
Thomson and Cathode Rays
Aim:
• To study the structure of the
atom.
• Investigate the electrical
discharges of atoms.
Procedure:
• Use a partially evacuated
tube to apply high voltage to
a screen coated with a
chemical compound.
• Measure the deflection.
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Thomson and Cathode Rays
• Thomson used partially
evacuated glass tubes to
discover the existence of
negatively charged particles
called electrons.
• Developed the plum
pudding model.
Plum Pudding Model
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7
Millikan and Oil Drops
Aim:
• To study the electronic
properties of the atom.
• Investigate the mass of an
electron.
Procedure:
• Use x-rays to produce
charges on oil drops.
• Measure the magnitude of
the electron charge.
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Millikan and Oil Drops
• With Thomson’s
cathode ray
experiment,
determined the
mass of an
electron,
• 9.11x10-31 kg.
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8
Rutherford Experiment
Aim:
• To study the internal
structure of the atom.
• Investigate the mass
distribution in the atom.
Procedure:
• Use a radioactive source to
bombard a thin piece of
gold foil.
• Collect the radiation.
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9
Original and Revised Theories
Original Theory:
– Plum pudding model
Revised Theory:
– Nucleus (dense positive
charge) at the center of the
atom.
– Large amount of space
between nucleus and
electron.
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•Moving electron
cloud surrounding
nucleus.
__________________
•Almost all the
mass is in the
nucleus!
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10
1+
+
1-
-
Notes:
• mass of e- tiny relative to p+, n.
• p+, n have same mass (almost).
• e-, p+ have same charge, opposite sign.
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Atomic Definitions I: Symbols, Isotopes, Numbers
A
Z
X
The Nuclear Symbol of the Atom
X = Atomic symbol of the element
Z = The atomic number: the number of protons in the nucleus
(All atoms of the same element have the same # of protons.)
A = The mass number (protons plus neutrons: A = Z + N)
N = The number of neutrons in the nucleus N = A - Z
Isotopes = atoms of an element with the same number of protons,
but different number of neutrons in the nucleus
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11
Figure 2.14 Isotopes of sodium
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What is the nuclear particle
with no charge?
1.
2.
3.
4.
5.
Electron
Proton
Neutron
Positron
Alpha
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12
Atomic Nucleus
• contains protons with positive charge
• contains neutrons with no charge
• chemistry of atom is determined by its electronic
structure
• atoms with same number of protons have
identical chemical properties
• isotopes are atoms with same number of protons
but different number of neutrons
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Which of the following statements are true?
I. The number of protons is the same for all neutral
atoms of an element.
II. The number of electrons is the same for all neutral
atoms of an element.
III. The number of neutrons is the same for all neutral
atoms of an element.
1.
2.
3.
4.
5.
I, II, and III
I and II only
II and III only
I and III only
None
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Available on the course website under “Exam Info” and “Lecture Notes”
27
A neutral atom of
rhenium(Re)-185 contains:
1. 110 p+, 75 n,
110 e2. 75 p+, 111 n,
75 e3. 75 p+, 110 n,
75 e4. 75 p+, 110 n,
74 e28
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Neutral ATOMS
If neutral, then # e- = # p+ = atomic number.
Remember: # n = A - # p+
Numbers of each particle:
51 Cr
= p+ (24), e- (24 ), n ( 27)
239 Pu
15 N
= p+ (94), e- (94), n (145)
= p+ (7), e- (7), n (8)
56 Fe
= p+ (26), e- (26), n (30)
235 U
= p+ (92), e- (92), n (143)
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Modern Reassessment of the Atomic Theory
1. All matter is composed of atoms. Although atoms are composed
of smaller particles (electrons, protons, and neutrons), the atom
is the smallest body that retains the unique identity of the element.
2. Atoms of one element cannot be converted into atoms of another
element in a chemical reaction. Elements can only be converted into
other elements in nuclear reactions in which protons are changed.
3. All atoms of an element have the same number of protons and
electrons, which determine the chemical behavior of the element.
Isotopes of an element differ in the number of neutrons, and thus
in mass number, but not in chemical behavior (much). A sample of
the element is treated as though its atoms have an average mass.
4. Compounds are formed by the chemical combination of two or more
elements in specific ratios, as originally stated by Dalton.
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15
What separates isotopes of the
same element?
1. Different number of
protons
2. Different number of
electrons
3. Different number of
neutrons
4. Different overall
charge
5. None of the above
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Definitions
Element - The simplest type of substance with unique physical and
chemical properties. An element consists of only one type of
atom. It cannot be broken down into any simpler substances by
physical or chemical means.
Compound - A substance composed of two or more elements that are
chemically combined.
Pure Substances - Their compositions are fixed! Elements and
compounds are examples of pure substances.
Mixture - Is a group of two or more elements and/or compounds that
are physically intermingled.
Molecule - A structure that consists of two or more atoms that are
chemically bound together and thus behave as an independent
unit.
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16
Definitions
Chemical Formula –
The symbols of for the elements are used to indicate the
types of atoms present, and the subscripts are used to indicate the
relative numbers of atoms present.
Structural Formula –
A formula in which the bonds are shown along with the
elemental symbols and order of atom arrangement.
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See figures
2.15-2.17 in
the textbook.
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17
More on amu
in Ch 3.
35
Review: Based on the Law of
Definite Proportions, what H:O
ratio do you expect to find in
water (H2O)?
1.
2.
3.
4.
1:1
2:1
1:2
2:2
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18
Forming Bonds
Sharing and Transfer of Electrons
37
Chemical Compounds and Bonds
Chemical Bonds – The forces that hold the atoms
of elements together in the compound.
Covalent Bonds – Electrons are shared between
atoms of different elements to form covalent
compounds.
Ionic Bonds- Attractive forces between two
oppositely charged ions to form ionic compounds.
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19
Chemical Compounds and Bonds
Cation
• An atom that has lost electron(s) to form a “+” ion.
• Common for metal elements.
Anion
• An atom that has gained electron(s), to form a “-” ion.
• Common for nonmetal elements.
Monatomic (single atom) ions form binary ionic compounds.
Polyatomic ions have many atoms per ion and an overall charge.
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A: No interaction, atoms too far
apart.
B: Atoms move closer, electron
clouds distort.
C: Covalent bond forms.
D: The protons from each
nucleus share the overlapping
electron cloud.
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20
Forming an Ionic Bond
Step 1: Electron transfer from sodium atom
(neutral sodium atom becomes sodium “+” ion)
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Copyright © Houghton Mifflin Company. All rights reserved.
Forming an Ionic Bond
Step 2: Electron added to chlorine atom
(neutral chlorine atom becomes chloride “-” ion)
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Copyright © Houghton Mifflin Company. All rights reserved.
21
Arrangement of sodium ions and chloride
ions
43
Figure 2.18: Sodium metal
reacts with chlorine gas…
2 Na(s) + Cl2(g)  2NaCl(s)
44
22
How is an ion formed?
1.
2.
3.
4.
5.
By either adding or
subtracting protons from
the atom.
By either adding or
subtracting electrons
from the atom.
By either adding or
subtracting neutrons
from the atom.
All of the above are
true.
Two of the above are
true.
45
PERIODIC TABLE
“MENDELEEV TABLE”
A tabular arrangement of the elements based upon their
chemical properties
• most elements are metals (left side of line) and tend to lose
electrons
• nonmetals (right side) tend to gain electrons
• vertical columns are groups - similar properties such as:
alkali metals, halogens, noble gases, etc.
• horizontal rows are periods
46
23
Figure 2.20: The periodic table
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Groups in the Periodic Table
Main Group Elements (Vertical Groups)
Group 1A - Alkali Metals
Group 2A - Alkaline Earth Metals
Group 3A - Boron Family
Group 4A - Carbon Family
Group 5A - Nitrogen Family
Group 6A - Oxygen Family
Group 7A - Halogens
Group 8A - Noble Gases
Other Groups (Vertical and Horizontal Groups)
Groups 3-12 - Transition Metals
Period 6 - Lanthanides (Rare Earth Elements)
Period 7 - Actinides (Radiocative & Artificial)
48
24
POSITION of ELEMENTS on
the PERIODIC TABLE
Tells you about:
• Properties
• Reactivity with other elements
• Most probable oxidation state(s)
• Composition of compounds
49
WRAP TABLE AROUND!
nonmetals
tend to gain
electrons
metals tend
to lose
electrons
50
25
What monatomic ion would
chlorine (Cl) form?
1.
2.
3.
4.
Cl2ClCl+
Cl2+
51
Predicting the Ion that an Element will Form
Problem: What monatomic ions will each of the elements form?
(a) barium(z=56) (b) sulfur(z=16) (c) titanium(z =22) (d) fluorine(z=9)
Plan: We use the “z” value to find the element in the periodic table and
which is the nearest noble gas. Elements that lie after a noble gas will
lose electrons and those before a noble gas will gain electrons.
Solution:
(a) Ba2+, barium - an alkaline earth element (Group 2A)
- will lose two electrons
- attains the same number of electrons as the xenon (Xe)
(b) S2-, sulfur - in the oxygen family (Group 6A)
- expected to gain two electrons
- attains the same number of electrons as argon (Ar)
(c) Ti4+, titanium - in Group 4B
- expected to lose 4 electrons
- attains the same number of electrons as argon (Ar).
(d) F-, fluorine - a halogen (Group 7A)
- expected to gain one electron
52
- attains the same number of electrons as neon (Ne).
26
The Periodic Table of the Elements
+1
Most Probable Oxidation States
(more on these in Ch 4)
0
H +2
+3 + 4 - 3 - 2 - 1 He
Li Be
B
Na Mg +3 +4 +5
C N O
+1 + 2 Al Si P
F Ne
S
Cl Ar
K Ca Sc Ti V Cr Mn Fe Co Ni Cu Zn Ga Ge As Se Br Kr
Rb Sr Y Zr Nb Mo Tc Ru Rh Pd Ag Cd In Sn Sb Te I Xe
Cs Ba La Hf Ta W Re Os Ir Pt Au Hg Tl Pb Bi Po At Rn
Fr Ra Ac Rf Db Sg Bh Hs Mt Ds
+3 Ce Pr Nd Pm Sm Eu Gd Tb Dy Ho Er Tm Yb Lu
+3 Th Pa U Np Pu Am Cm Bk Cf Es FmMd No Lr
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Naming Simple Compounds
• BINARY: comprised of two elements
• IONIC: cation + anion (metal + nonmetal; two types of
metals: Type I and Type II)
• COVALENT: between two nonmetals
• POLYATOMIC: several elements
54
27
55
(Type II: metals that have >1 common oxidation state)
56
28
Rules for Naming Binary Ionic
Compounds
Always list cation (+) name first, anion (-) name second.
Monatomic anion’s name = first part of element name + “-ide”
If the ion is not listed in
the Type II table, assume
it’s Type I.
57
Which of the following compounds is not
properly named?
1.
2.
3.
4.
5.
6.
FeCl3 : iron(III) chloride
NaBr : sodium bromide
Ca3N2 : calcium nitride
PbO : lead(II) oxide
MgH2 : magnesium
hydride
None of the above
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29
59
Which of the following compounds is not
properly named?
1.
Fe2(SO4)3 : iron(III) sulfate
2.
Ba(OH)2 : barium hydroxide
3.
CaCrO4 : calcium chromate
4.
Cr(NO2)3 : chromium(III)
nitrite
5.
KClO3 : potassium chloride
6.
RbHCO3 : rubidium
hydrogen carbonate
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30
Naming Binary Covalent Compounds
• The element further to the left in
the periodic table goes first,
named as if it were the cation.
• Second element is named as if it
were the anion.
• Use prefixes to say how many
atoms of each element are
present, EXCEPT never use
“mono-” on first element.
EXAMPLES:
N2O dinitrogen monoxide
NO
nitrogen monoxide
N2O4 dinitrogen tetroxide
61
Which of the following compounds is not
properly named?
1.
CCl4 : carbon
tetrachloride
2.
CO2 : carbon dioxide
3.
H2S : dihydrogen
sulfide
4.
NF3 : nitrogen
trifluoride
5.
SF6 : sulfur
hexafluoride
6.
None of the above
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HYDRATES
Compounds containing WATER molecules
MgSO4
· 7H2O
magnesium sulfate heptahydrate
CaSO4 · 2H2O
calcium sulfate dihydrate
Ba(OH)2 · 8H2O
barium hydroxide octahydrate
CuSO4 · 5H2O
copper(II) sulfate pentahydrate
Na2CO3
· 10H2O
sodium carbonate decahydrate
63
Acids
Many common anions, when combined with H+ as the
cation(s), make a very reactive compound called an acid,
which dissociates in water to give the two separate ions
both dissolved in water.
Examples:
H3PO4(aq)  H+(aq) + H2PO4-(aq)
phosphoric acid
HCl (aq)  H+(aq) + Cl-(aq)
hydrochloric acid
H2SO4 (aq)  2 H+(aq) + SO42-(aq)
sulfuric acid
Litmus paper – indicates if it’s an acidic solution
pH = quantitative measure of H+ concentration in water 64
32
Rules for naming acids
Acids = Molecules which dissociate when
dissolved in water to give H+(aq)
Binary acids
Oxoacids
Keep any
prefixes (hypoor per-)
65
Which of the following compounds is not
properly named?
1.
HNO3 : nitric acid
2.
HCl : hydrochloric
acid
3.
H2SO4 : sulfuric acid
4.
HClO4 : chloric acid
5.
H3PO4 : phosphoric
acid
6.
None of the above
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