Atoms, Molecules,
and Ions
Chapter 2 BLB 12th
Expectations
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Recognize important steps in the discovery
of the atom and its structure.
Work with isotopes.
Learn about the periodic table.
Differentiate between molecular and ionic
compounds.
Name compounds (molecular and ionic).
2.1 The Atomic Theory of Matter
Of what is matter comprised?
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Democritus (400 BC) – tiny, indivisible particles, atomos
Plato, Aristotle – NOT!
Newton (17th century) – favored atoms as invisible particles
Boyle (1660) – gas experiments with pressure & volume
Priestly (1774) – isolated oxygen
Lavoisier (1789) – Law of Conservation of Mass: Mass is
neither created or destroyed. (p.78)
Proust (1800) – Law of Definite Proportions (or constant
composition): A compound always contains the same
proportion of elements. (p. 10)
Dalton’s Atomic Theory (1808)
1. Elements are composed of small particles
called atoms.
2. All atoms of a given element are identical.
3. Atoms of an element are not changed in a
chemical reaction.
4. Compounds are formed when different
atoms combine.
>> Atoms are the building blocks of matter.<<
2.1 The Atomic Theory of Matter
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Dalton – Law of Multiple Proportions: element
mass proportions in a compound are in a ratio
of small whole numbers. (p. 40)
Avogadro (1811) – equal volumes of gases
contain the same number of particles (p. 401)
2.2 The Discovery of Atomic Structure
Subatomic particles
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J.J. Thomson (1897) – cathode ray tube experiments;
electrons; charge-to-mass ratio of the electron
(1) plum-pudding model of atoms (Fig. 2.9, p. 43)
Robert Millikan (1909) – oil-drop experiment; charge and
mass of electron (9.10939 x 10-28 g)
Henri Becquerel, Marie Curie (1896, 1899) – radioactivity
Ernest Rutherford (1911) – gold foil experiment; nucleus &
protons (1919); (2) nuclear model of atom
3 types of radioactivity: α (heaviest, 2+ charge), β (highspeed electrons, 1− charge), g (lightest, high E, 0 charge)
James Chadwick (1932) - neutrons
Separation of Radioactive Particles
Rutherford’s Gold Foil Experiment
pp. 42-43
2.3 The Modern View of Atomic Structure
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Subatomic particles
Particle
Function
Position
(to nucleus)
Charge
Mass (kg)
electron chemistry
outside
−1
9.11 x 10-31
proton
attract
electrons
inside
+1
1.67 x 10-27
neutron
nuclear glue inside
0
1.67 x 10-27
(neutral)
Atoms
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Atomic masses ~10-23 g
Atomic diameters (e- cloud) ~10-10 m = 1 Å
Atomic nuclei ~10-4 Å (very small and dense)
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Atoms are neutral: # protons = # electrons
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Practice Exercise 2.1
How many carbon atoms can be placed side
by side across the width of a pencil line that
is 0.20 mm wide? C atom diameter = 1.54 Å
Isotopes, atomic and mass numbers
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Isotopes – atoms with same number of protons
but different numbers of neutrons
mass#  12
atomic#  6
C  element symbol
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Nuclide – a single atom of a particular isotope
2.4 Atomic Weights
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Based on 12C (assigned a mass of exactly 12 amu)
1 amu = 1.66054 x 10-24 g
(1/12 mass of a 12C atom)
Weighted average atomic mass =
Σ(% abundance)(mass of isotope)
Atomic mass determined using a mass
spectrometer (p. 49)
Mass Spectrometer
Mass spectrum
of Cl
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Calculate the (weighted) average mass of
magnesium (in amu).
isotope % abund. Mass (amu)
24Mg
78.99
23.98504
25Mg
10.00
24.98584
26Mg
11.01
25.98259
2.5 The Periodic Table
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1st table developed by Mendeleev and
Meyer in 1869
Group, period, regions, group names
Physical properties of metals and nonmetals
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Seaborg (p. 52)
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Physical Properties
Metals
 High electrical conductivity
 High thermal conductivity
 Metallic luster
 Most are solids
 Malleable, ductile
 Metallic bonding
Nonmetals
 Poor electrical conductivity
 Good heat insulator
 No metallic luster
 Solids, liquids, and gases
 Brittle in solid state
 Covalently bonded
molecules; noble gases
monoatomic
2.6 Molecules and Molecular Compounds
Chemical bonds – forces that hold atoms together in
molecules and compounds
 covalent bonds – sharing of electrons
Molecules – discrete units of covalently bonded atoms;
typically nonmetals, e.g. H2O, CO2, NH3, C2H6
Diatomic elements: H2, N2, O2, F2, Cl2, Br2, I2 (p. 53)
Polyatomic elements: O3, S8, P4
(allotropes – different forms of the same element in the
same state, p. 273)
Molecules, cont.
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Representation of
molecules, CH4
Empirical & Molecular Formulas
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Molecular formula – actual number of atoms in a
compound
Empirical formula – smallest whole number ratio
of atoms
Molecular Empirical
C2H4
CH2
P4O10
P2O5
H2O2
HO
2.7 Ions and Ionic Compounds
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Ionic bond – attraction between oppositely
charged ions; results from a transfer of
electrons
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cation – positively charged ion (metals)
anion – negatively charged ion (nonmetals)
Common ions (Fig. 2.20, p. 56)
Ionic Bonds
Ionic compounds (such as NaCl) are generally
formed between metals and nonmetals.
© 2009, Prentice-Hall, Inc.
Predicting ionic charges
Atoms will lose or gain electrons to attain a noble gas configuration.
P3–
Ionic Compounds
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Ionic compounds – consist of ions; form
crystal lattices
+ and − charges balance
Formula unit – ratio of
cation to anion
2.8 Naming Inorganic Compounds
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1957 IUPAC (Int’l Union of Pure and
Applied Chemistry) – devised systematic
rules for naming compounds
Binary compounds – consist of two different
elements
Don’t capitalize compound or element
names.
Ionic compounds - cations
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Cations (Table 2.4, p. 60)
1.
Single metal, single charge
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2.
Single metal, multiple charges
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3.
Na+, sodium ion
Al3+, aluminum ion
Cr2+, chromium(II) ion
Cr4+, chromium(IV) ion
Polyatomic ions
Ionic compounds - anions
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Anions (Table 2.5, p. 63)
1.
Monoatomic, -ide ending
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2.
Oxyanions
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3.
Cl-, chloride ion
O2-, oxide ion
NO3-, nitrate ion
NO2-, nitrite ion
H+ + oxyanion
P3–
Phosphide ion
NO2–
Nitrite ion
PO33-
Phosphite ion
Polyatomic Ions to Memorize
(formula & charge)
Name
Formula
Name
Formula
acetate ion
C 2 H 3 O2 -
hypochlorite ion
ClO-
ammonium ion
NH4+
nitrate ion
NO3-
carbonate ion
CO32-
nitrite ion
NO2-
chlorate ion
ClO3-
perchlorate ion
ClO4-
chlorite ion
ClO2-
permanganate ion
MnO4-
chromate ion
CrO42-
phosphate ion
PO43-
cyanide ion
CN-
phosphite ion
PO33-
hydroxide ion
OH-
sulfate ion
SO42-
sulfite ion
SO32-
Ionic compounds
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Cation first, anion second
Charges (+ and -) must balance
Acids (p. 64)
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Acid – substance which produces a H+
when dissolved in water
If anion ends in ____, acid ends with ____.
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-ide
-ate
-ite
-ic
-ic
-ous
Molecular Compounds
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Name as written in
formula.
Prefixes denote
number of each atom.
Exceptions:
H2O water
NH3 ammonia
CH4 methane
2.9 Some Simple Organic Compounds
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Hydrocarbon – contain only C and H
Alkanes – saturated hydrocarbons with only
C−C single bonds
Alkane derivatives:
−OH
alcohol
−COOH
carboxylic acid
−COOC−
ester
−COC−
ketone
Organic compounds, cont.
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Unsaturated hydrocarbons:
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Alkenes – contain at least one C=C double bond
Alkynes – contain at least one C≡C triple bond