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Daniel L. Reger
Scott R. Goode
David W. Ball
http://academic.cengage.com/chemistry/reger
Chapter 2
Atoms, Molecules, and Ions
History Lesson
• Democritus (460-370 BC): indivisible particles called atoms
• Plato and Aristotle challenged this view believing that matter
was continuous
• Newton (1642-1727 AD) proposed the idea of invisible
particles in the air called atoms
• Antoine Lavoisier (1743-1794 AD) conducted experiments
demonstrating mass of products = mass of reactants
• John Dalton (1766-1844 AD) proposed a model of matter
• Dalton’s Atomic Theory
Dalton’s Atomic Theory
Postulates
Assumption: Matter is discontinuous!
1: Matter is composed of atoms. An atom
is the smallest unit of an element that
has all the properties of that element.
2: An element is composed entirely of one
type of atom.
Dalton’s Atomic Theory (cont’d)
3: A compound contains atoms of two or
more different elements. The relative
number of atoms of each element in a
compound is always the same.
4: Atoms do not change identity in
chemical reactions; only the way in which
they are joined together changes.
Law of Constant Composition
• Law of constant composition: All
samples of a pure substance contain the
same elements in the same proportions
by mass.
• This observation follows from Dalton’s third
postulate (the relative numbers of atoms are
the same in the same compound).
Law if Multiple Proportions
• Law of multiple proportions: When the
same elements form more than one
compound, the masses of one element
that combines with a fixed mass of a
second element are in a ratio of small
whole numbers.
• This follows from the postulate that
individual atoms enter into chemical
combination.
Law of Conservation of Mass
• Law of Conservation of Mass: There is
no detectable change in mass when a
chemical reaction occurs.
• Dalton’s fourth postulate accounts for this
law. The atoms do not change mass or
identity when a chemical reaction takes
place.
Atomic Composition and Structure
• Experiments over many years showed
that atoms are not simple particles, but
are composed of the subatomic particles
listed below:
• Electrons
• Protons
• Neutrons
Cathode Rays
• The application of a high voltage
across a partially evacuated tube
produces cathode rays.
Electrons
• J. J. Thomson demonstrated that cathode
rays were negatively charged by applying
magnetic and electric fields to cathode rays.
• Cathode rays are electrons, negatively
charged particles that are one of the
components of an atom.
Millikan Oil Drop Experiment
• Robert A. Millikan performed experiments
that determined the charge of the electron
as 1.60 x 10-19 coulombs.
Mass of electron
• Thompson, using Milikan’s data for the charge
of an electron, determined the mass to charge
ratio of an electron.
• This allowed him to calculate the mass of an
electron.
• 9.11 x 10-31 kg (mass of electron)
Scattering of Alpha Particles by Gold
The Nuclear Model of the Atom
• Rutherford concluded that the results of
the scattering experiment required that
atoms consist of:
• a nucleus that is very small
compared to the atom, has a high
positive charge and contains most of
the mass of the atom.
• the remainder of the space in an atom
contains enough electrons to give a
neutral atom.
Atomic View of Rutherford Experiment
The Proton
• Rutherford proposed that the
hydrogen nucleus was a
fundamental particle called the
proton, which has a positive
charge equal in magnitude to the
negative charge of the electron.
• Protons account for the charge on the
nucleus of all atoms.
• The mass of the proton (1.673 x 10-27
kg) is 1836 times that of the electron.
The Neutron
• The number of protons in a nucleus,
as determined by its positive charge,
accounts for half or less of the nuclear
mass.
• Scientists inferred there must be a
massive, neutral particle also present in
the nucleus.
• This neutral particle is called the
neutron; its mass is almost the same
as that of the proton.
Particles in the Atom
Particle
Charge (C)
Mass (kg)
Relative Relative
charge mass
Electron -1.602 x 10-19 9.109 x 10-31
1-
0
Proton +1.602 x 10-19 1.673 x 10-27
1+
1
1.675 x 10-27
0
1
Neutron
0
Definitions
• Atomic number (Z) is the number of
protons in the nucleus of an atom.
• Mass number (A) is the sum of the
numbers of protons and neutrons in
the nucleus.
• The number of protons (the atomic
number) determines the identity of the
element; all H atoms contain 1 proton,
all He atoms contain 2 protons, etc.
Isotopes
• Isotopes are atoms of one element
whose nuclei contain different numbers
of neutrons (same Z, different A).
Isotopes of
Hydrogen
What about Dalton’s
Postulate #2?
Symbols of Isotopes
• A symbol to identify a specific isotope is
A
Z
X
where A = mass number, Z = atomic number,
and X is the one or two letter symbol of the
element.
• The three isotopes of hydrogen are:
1
1
H
2
1
H
3
1
H
Symbols of Isotopes
• Oxygen also has three isotopes, containing 8,
9, and 10 neutrons respectively. The symbols
are:
16
17
18
8
8
8
O
O
O
• Since the value of Z, and the symbol, both
identify the element, Z is often omitted from
the symbol:
16
O
17
O
18
O
Example: Symbols of Atoms
• Write the symbol for the isotopes with:
(a) 15 protons and 16 neutrons.
(b) 21 protons and 24 neutrons.
Ions
• In many chemical reactions, atoms gain
or lose electrons, producing charged
particles called ions.
• A cation has a positive charge and forms
when an atom loses one or more electrons.
• An anion has a negative charge and forms
when an atom gains one or more electrons.
Symbols for Ions
• The number of protons in the
nucleus determines the symbol used
for an ion.
• The element’s symbol is followed by
a superscript number and a sign that
shows the charge on the ion in
electron charge units.
• If the ionic charge is one unit, the
number is omitted, e.g. Na+ is the
symbol for a sodium cation.
Example: Symbols of Ions
• Write the symbol for the ions that
contain:
(a) 9 protons, 10 neutrons, 10
electrons.
(b) 19 protons, 20 neutrons, 18
electrons.
Practice
• Write the symbols for the particles
containing:
(a) 8 protons, 9 neutrons, 10 electrons
(b) 13 protons, 14 neutrons, 13
electrons
Test Your Skill
• Write the symbols for the particles containing:
(a) 8 protons, 9 neutrons, 10 electrons
(b) 13 protons, 14 neutrons, 13 electrons
Answer: (a)
17
8
O
2
(b)
27
13
Al
Example: Components of Ions
• Fill in the blanks.
Symbol 23 Na 
11
Atomic number
Mass number
Charge
no. of protons
no. of neutrons
no. of electrons
____
____
____
____
____
____
The Atomic Mass Unit (u)
• A relative mass scale has been
established to express the masses
of atoms.
• The atomic mass unit (u) is 1/12
the mass of one 12C atom.
Experimentally to three significant
digits:
1 u = 1.66 x 10-27 kg
Masses of Atoms in u
• The masses of both the proton and the
neutron are approximately 1 u.
• A 24Mg atom has a mass approximately
twice that of the 12C atom, so its mass is
24 u.
• A 4He atom has a mass approximately
1/3 that of the 12C atom, so its mass is 4
u.
Atomic Mass and Mass Number
• Factors other than the mass of the
protons and neutrons affect the mass of
atoms, so the actual mass of atoms are
not whole numbers. (24Mg = 23.98504 u;
4He = 4.002603 u)
• When the accurate atomic mass of an
atom is rounded to a whole number, it
equals the mass number.
Natural Distribution of Isotopes
• About 75% of the elements occur in
nature as mixtures of isotopes.
• Usually, the relative abundance of
isotopes of an element is the same
throughout nature.
• In all natural samples of Li, 7.42% of the
atoms are 6Li and the remaining 92.58%
are 7Li.
Atomic Masses of the Elements
• Isotopic mass is the mass in u, of a
particular isotope of an element.
• Different isotopes of an element all react
essentially the same, so a weighted
average of isotopic masses can be used
in calculations.
• The atomic mass is the weighted
average mass, of the naturally occurring
element.
atomic mass = fractionA x isotopic massA +
fractionB x isotopic massB + . .
Example: Calculating Atomic Mass
• A mass spectrometer was used to
determine that gallium is 60.11%
69Ga (isotopic mass = 68.9256 u)
and 39.89% 71Ga (isotopic mass =
70.9247 u). Calculate the atomic
mass of Ga.
The Periodic Table
• Proposed independently by Dimitri
Mendeleev and Lothar Meyer.
• Periodic table: arranges the
elements in rows that place elements
with similar properties in the same
column.
• Period: a horizontal row
• Group: a column - contains chemically
similar elements
Atomic Number and Atomic Mass
• The atomic number and atomic
mass for each element is given on
the periodic table.
38
Atomic number
Sr
87.62
Atomic mass
Important Groups of Elements
• Metal: a material that is shiny and is a
good electrical conductor; metallic
elements are on the center and left
side of the periodic table.
• Nonmetal: an element that is typically
a nonconductor; nonmetals are in the
top right part of the periodic table.
• Metalloid: an element that has
properties of both metals and
nonmetals.
Important Groups of Elements
• Representative Elements: the elements
in the A groups (1,2, 13-18).
• Transition Metals: the elements in B
groups (3-12).
• Inner Transition Metals: the two rows of
metals (lanthanides and actinides) set
at the bottom of the periodic table.
Important Groups of Elements
• Alkali Metals: soft, reactive metals
in group 1A.
• Alkaline Earth Metals: elements in
group 2A.
• Halogens (salt formers): reactive
nonmetals in group 7A.
• Noble Gases: the stable, largely
inert, gases in group 8A.
Elements and Biology
Molecules
• A molecule is a combination of
atoms joined so strongly that they
behave as a single particle.
• The simplest molecules are
diatomic - they contain two atoms.
Elements
• If all the atoms in a molecule are the
same, the substance is an element.
Molecules
• If two or more elements form a
molecule, it is a molecular compound.
Molecular Formulas
• A molecular formula gives the number of
every type of atom in the molecule.
• The elements present in the molecule are
identified by their symbols.
• A subscript number follows each symbol,
giving the number of atoms of that element
present in the molecule; the subscript is
omitted if only one atom of the element is
present.
• A structural formula shows how the
atoms are connected in the molecule.
Molecular Formulas
Molecular Mass
• The molecular mass is the sum of the
atomic masses of all atoms present in
the molecular formula, expressed in
atomic mass units (u).
• The diagram shows the strategy for
calculating molecular mass.
Example: Calculate Molecular Mass
• One substance present in smog is
dinitrogen tetroxide (N2O4). Calculate its
molecular mass.
Practice
• What is the molecular mass of the fuel
propane (C3H8 )?
Ionic Compounds
• An ionic compound is composed of
cations and anions joined to form a
neutral species.
• Ionic compounds generally form from the
combination of metals with nonmetals.
• In ionic compounds each cation is
surrounded by several anions and vice
versa.
Structure of Sodium Chloride
Formulas of Ionic Compounds
• The formula of an ionic compound is an
empirical formula that uses the smallest
whole number subscripts to express the
relative numbers of ions.
• The relative numbers of ions in the empirical
formula balances the charges to zero.
• The formula of sodium chloride is NaCl,
because the 1+ ions have to be present in a 1:1
ratio.
• The formula of sodium oxide is Na2O, because
the charge of the Na+ and O2- ions balance to
zero in a 2:1 ratio.
Formulas of Ionic Compounds
• The position of an element in the periodic
table can be used to determine the
charges of some ions.
• The metallic elements in Groups 1A, 2A,
3B, and Al (Group 3A) all form cations with
a charge equal to the Group number.
• The nonmetals in Groups 6A, 7A, and N in
group 5A form anions with a charge of 2-, 1and 3-, respectively.
Charges on Common Ions
1A
2A
Li+
Be2+
Na
+
+
K
Rb
Cs
+
+
Mg
Ca
Sr
2+
2+
2+
Ba
3B
2+
3A
Al
Sc
3+
3+
Y
La
3+
3+
5A
6A
7A
N3-
O2-
F-
2-
-
S
Se
2-
Cl
Br
I
-
-
Example: Ionic Compounds Formulas
• Write the empirical formulas of the
compound formed by
(a) the cation of Ca and the anion of Br.
(b) the cation of Al and the anion of O.
Polyatomic Ions
• Polyatomic ion: a group of atoms with a
net charge that behaves as a single
particle.
• The ammonium ion (NH4+) is the most
common polyatomic cation.
• There are many important polyatomic
anions.
Some Polyatomic Anions
Name
Formula
Name
Formula
Acetate
CH3CO2-
Nitrate
NO3-
Carbonate
CO32-
Nitrite
NO2-
Bicarbonate
HCO3-
Permanganate
MnO4-
Chlorate
ClO3-
Phosphate
PO43-
Perchlorate
ClO4-
Hydrogen
phosphate
HPO42-
Chromate
CrO42-
Dihydrogen
phosphate
H2PO4-
Cyanide
CN-
Sulfate
SO42-
Dichromate
Cr2O72-
Bisulfate
HSO4-
Hydroxide
OH-
Sulfite
SO32-
Example: Polyatomic Ions Formulas
• Write the formulas of the compounds that
contain:
(a) the calcium ion and nitrate ion.
(b) the ammonium ion and the
dichromate ion.
Formula Mass of Ionic Compounds
• Formula mass is the sum of the atomic
masses of all atoms in the empirical
formula of an ionic compound.
The formula mass of Ca(NO2)2 is:
1(Ca) x 40.08
= 40.08
2(N) x 14.01
= 28.02
4(O) x 16.00
= 64.00
Formula mass
=132.10 u
Chemical Nomenclature
• Chemical nomenclature is the
organized system for naming
compounds.
• Some of the basic rules of nomenclature
are given here for:
• Ionic compounds
• Acids
• Molecular compounds
• Organic compounds
Naming Ionic Compounds
• The name of the cation is given first,
followed by the name of the anion.
• For monatomic ions:
• the name of the cation is the same as
the name of the element.
• the name of an anion is formed from
the name of the element by changing
the ending to “ide”.
• The names given in the table are
used to name polyatomic ions.
Common Monatomic Anions
Anion
-
H
3-
Name
Hydride
Anion
Name
-
Fluoride
-
Chloride
Bromide
F
N
Nitride
Cl
O2-
Oxide
Br-
S2-
Sulfide
I-
Iodide
Naming Cations
• The elements in groups 1A, 2A, and 3B
form only one stable ion, but most other
metals form more than one cation.
• For metals that form more than one cation,
a Roman numeral equal to the charge of
the ion is shown in parentheses following
the name of the element.
• The name of the compound FeCl3 is
iron(III) chloride; that of FeCl2 is iron(II)
chloride.
Example: Ionic Compounds Names
• Name the following ionic compounds:
(a) NH4Br
(b) Ca(NO3)2 (c) MnSO4
• Give the formula of the following ionic
compounds:
(a) chromium(III) nitrate
(b) potassium sulfate
(c) ammonium dichromate
Acids
• An acid is a compound that produces
hydrogen ions when dissolved in water,
and for the present can be considered as
hydrogen cations combined with one of
the anions already discussed.
• For example HCl, HNO3 and H2SO4 are all
acids in water solution.
Naming Acids
• If the anion name ends in “ide”,
change the ending to “ic” and add
the prefix “hydro”. This is followed
by the word acid.
Acid
Name
Anion
Name
HBr
hydrobromic acid
Br-
Bromide
H2S
hydrosulfuric acid
S2-
Sulfide
HCN
hydrocyanic acid
CN-
Cyanide
Naming Acids (cont’d)
• If the polyatomic anion name ends in
“ate”, change the ending to “ic”; if it
ends in “ite” change the ending to
“ous”. This is followed by the word acid.
Acid
Name
Anion
Name
H3PO4 Phosphoric acid
PO43-
Phosphate
HClO4
Perchloric acid
ClO4
Perchlorate
HNO2
Nitrous acid
NO2-
Nitrite
Names of Molecular Compounds
• Many molecular compounds have
nonsystematic common names; e.g.
water (H2O), ammonia (NH3), and
methane (CH4).
Order of Element Names
• The order of the elements in the names
and formulas of molecular compounds is:
• The element farther to the left in the periodic
table appears first.
• The element closer to the bottom within any
group is first.
• Hydrogen is first when combined with 6A and 7A
elements; it is named second when combined
with groups 1A through 5A elements.
• Oxygen is second, except when combined with
fluorine.
Numerical Prefixes in Names
• Often the same elements form more
than one compound. Numerical
prefixes are used to give the number of
atoms present in the molecule.
Number
one
two
three
four
five
six
Prefix
mono- (often omitted when first)
ditritetrapentahexa-
Example: Naming Compounds
• What is the name of the following
compounds?
(a) H2SO4
(b) SF6
(c) C3O2
(d) TiO2
Ionic and Molecular Compounds
• Ionic compounds are usually
combinations of metals and nonmetals,
while molecular compounds usually
contain only nonmetals.
• Ionic compounds are usually hard,
brittle solids with high melting points;
molecular compounds have lower
melting points, and may be liquids or
gases at room temperature.
Dissociation of Ionic Compounds
• Most ionic compounds dissociate into
individual cations and anions when dissolved
in water.
• NaCl dissociates into Na+ and Cl- in water.
Electrolytes
• An electrolyte is a substance that
produces ions in water solution.
• Ionic compounds are electrolytes - they
conduct electricity when dissolved in
water.
• Ionic compounds heated until they melt to
form a liquid also conduct electricity.
Nonelectrolytes
• Water and compounds that dissolve in
water as neutral molecules are
nonelectrolytes, they do not conduct
electrical current.
• Most molecular compounds are also
nonconducting.
Electrical Conductivity
Measured conductivity of
(a) ionic solids, (b)
melted or (c) dissolved
ionic compounds and (d)
molecular compounds.
Melted or dissolved ionic
compounds conduct.
Three Phases of Bromine
Molecular compounds are frequently gasses
or low melting solids.