Download Chapter 2: Atoms, Molecules, and Ions

Chemistry, The Central Science, 11th edition
Theodore L. Brown; H. Eugene LeMay, Jr.;
and Bruce E. Bursten
Chapter 2:
Atoms, Molecules,
and Ions
Atomic Theory of Matter
The theory that atoms are the fundamental
building blocks of matter reemerged in the early
19th century, championed by John Dalton.
Dalton's Postulates
Each element is composed of extremely small
particles called atoms.
Dalton's Postulates
All atoms of a given element are identical to one
another in mass and other properties, but the
atoms of one element are different from the
atoms of all other elements.
Dalton’s
Book: A
New
System of
Chem.
Philosophy
Dalton's Postulates
Atoms of an element are not
changed into atoms of a different
element by chemical reactions;
atoms are neither created nor
destroyed in chemical reactions.
Dalton's Postulates
Compounds are formed when atoms of
more than one element combine; a
given compound always has the same
relative number and kind of atoms.
Law of Constant Composition
Joseph Proust (1754–1826)
• also known as the law of definite
proportions.
• It states that the elemental composition
of a pure substance never varies.
Law of Conservation of Mass
Antoine Lavoisier (1754–1826)
The total mass of substances
present at the end of a chemical
process is the same as the mass of
substances present before the
process took place.
Prob. 2.11. A chemist find that 30.82 g of
nitrogen react completely with 17.60, 35.20
70.40 or 88.00 g of oxygen to from four different
compounds.
a) Calculate the mass of oxygen per gram of
nitrogen in each compound.
b) How is the result explained by Dalton’s
atomic theory.
Exam Problem. A sample of ascorbic acid (vitamin C) is synthesized
in the laboratory. It contains 1.50 g of carbon and 2.00 g of oxygen.
Another sample of ascorbic acid isolated from citrus fruits contains
9.22 g of carbon. How many grams of oxygen does it contain?
a. 12.3
b. 6.92
c. 2.63
d. 10.5
e. 19.00
The Electron
• Streams of negatively charged particles were
found to emanate from cathode tubes.
• J. J. Thompson is credited with their
discovery (1897).
The Electron
Thompson measured the charge/mass ratio
of the electron to be 1.76 108 coulombs/g.
Millikan Oil Drop Experiment
Once the charge/mass
ratio of the electron
was known,
determination of either
the charge or the mass
of an electron would
yield the other.
Robert Millikan (U. of Chicago) was able to determing
the mass of the electron by measuring its charge in this
experiment. The mass calculated, me = 9.10 E-28 g.
Atom Structure, circa 1900
• The prevailing theory
was that of the “plum
pudding” model, put
forward by Thompson.
• It featured a positive
sphere of matter with
negative electrons
imbedded in it.
Discovery of the Nucleus, 1911
Ernest Rutherford
students (Marsden &
Geiger) shot a beam
of -particles at a thin
sheet of gold foil and
observed the pattern
caused by their
scattering
-particles are helium
nuclei, 4He2+
The Nuclear Atom
Since some particles were
deflected at large angles, thus
Thompson’s model could not
be correct.
Rutherford on YouTube
The Nuclear Atom
• Rutherford postulated a very small,
dense nucleus (1.8 fm) with the
electrons around the outside of the
atom.
• Most of the volume of the atom is empty
space.
Other Subatomic Particles
• Protons were discovered by Rutherford
in 1919.
• Neutrons were discovered by James
Chadwick in 1932.
quark structure
of the neutron
Subatomic Particles
• Protons and electrons are the only particles that
have a charge.
• Protons and neutrons have essentially the same
mass.
• The mass of an electron is so small we ignore it
when labeling particles.
Symbol
1 p
1
1
0n
0
-1e
The neutron to proton ratio is
critical in determining the
stability of a nucleus.
Elements above the belt of
stability undergo -decay and
elements below the belt of
stability undergo positron
emission or electron capture.
Fusion processes in stars have been shown to
form nuclei up to 26 protons and 30 neutrons
(5626Fe).
Supernova
SN 1054
produced the
Crab Nebula
Supernova produce temperatures (energies)
capable of synthesizing elements through Z=94
Radioactivity
• Radioactivity is the spontaneous emission
of radiation by an atom.
• It was first observed by Henri Becquerel in
1896.
• Marie and Pierre Curie studied it
extensively and isolated the first
radioactive elements (polonium and
radium).
Radioactivity
• It is not uncommon for some nuclides of
an element to be unstable, or
radioactive.
• We refer to these as radionuclides.
• There are several ways radionuclides
can decay into a different nuclide.
Radioactivity
• Three types of radiation were discovered by
in the early 1900s in Ernest Rutherford’s lab:
 particles
 particles
 rays
Writing the Symbol of an Isotope
A
ZX
• A is the mass number (p + n)
• Z is the atomic number (# protons)
• X is the atomic symbol
Alpha Decay
Alpha decay is the loss of an -particle (a
helium nucleus).
4
2
He
238
92
U
234
90
U
+ 4He
2
Beta Decay
Beta decay is the loss of a -particle (a high
energy electron).
0
−1
131
53
I
or
131
54
0
−1
e
Xe
+ 0e
−1
Positron Emission
Some nuclei decay by emitting a
positron, a particle that has the same
mass as but an opposite charge to that
of an electron.
0
1
11
6
C
e
11
5
B
+ 0e
1
Electron Capture (K-Capture)
Addition of an electron to a proton in the
nucleus is known as electron capture or Kcapture.
– The result of this process is that a proton is
transformed into a neutron.
1
1
p
+ 0e
−1
1
0
n
Gamma Emission
This is the loss of a -ray, which is highenergy radiation that almost always
accompanies the loss of a nuclear
particle.
0
0
Writing the Symbol of an Isotope
A
ZX
• A is the mass number (p + n)
• Z is the atomic number
• X is the atomic symbol
Symbols of Elements
Elements are symbolized by one or two
letters. For this class you should know the
names and symbols for the first 36 elements.
Atomic Number
All atoms of the same element have the same
number of protons:
The atomic number (Z)
Atomic Mass
The mass of an atom in atomic mass units
(amu) is the total number of protons and
neutrons in the atom.
Atomic Mass
Atomic and
molecular masses
can be measured
with great accuracy
with a Mass
Spectrometer.
Isotopes
• Isotopes are atoms of the same element with
different masses.
• Isotopes have different numbers of neutrons.
11
C
6
12
C
6
13
C
6
14
C
6
Oxygen isotopes and decay modes
Name
Symbol
Mass (amu)
Oxygen-13
13
13.02810
O
8
Oxygen-14
14
15
O
14.008595
16
% Natural
Abundance
e
8.9 ms
----
0
e
70.6 s
---
e
122 s
---
1
O
15.003065
8
Oxygen-16
0
Half-Life
1
8
Oxygen-15
Decay
Mode
0
1
O
15.994915
---
---
99.76
O
16.999131
---
---
0.04
O
17.999160
---
---
0.20
O
19.003577
0
e
26.9 s
---
e
13.5 s
---
e
≈3 s
---
8
Oxygen-17
17
8
Oxygen-18
18
8
Oxygen-19
19
8
Oxygen-20
20
-1
O
20.004075
8
Oxygen-21
21
8
0
-1
O
21.008730
0
-1
Average Mass
• Because in the real world we use large
amounts of atoms and molecules, we
use average masses in calculations.
• Average mass is calculated from the
isotopes of an element weighted by
their relative abundances.
Example
Neon is found in three isotopes in nature.
Isotope
Mass (amu)
Natural abundance (%)
Neon-20
19.9924
90.4838
Neon-21
20.99395
0.2696
Neon-22
21.9914
9.2465
Average atomic mass of neon = (19.9924 x
0.904838) + (20.99395 x 0.002696) +
(21.9914 x 0.092465) = 20.1797 amu
Another Isotope Problem…
Using the following table of abundances and masses of
the three naturally occurring argon isotopes, calculate
the exact mass of 40Ar. Symbol Exact Mass (amu) Natural
Abundance (%)
36Ar
35.96755
0.337
38Ar
37.96272
0.063
40Ar
?
99.60
Ave. 39.948
Ans: 39.96 amu
Mendeleev’s Periodic Table
• Dmitrii I. Mendeleev
(1869) arranged
elements in the periodic
table by their chemical
and physical properties.
• He left open spaces in
his periodic table to
account for elements
not yet discovered.
Periodic Table
• It is a systematic
catalog of the
elements.
• Elements are
arranged in order
of atomic number
, Z, (rather than
atomic weights).
• For this class, you
should know the
names and
symbols for
elements 1-36
Periodic Table
• The rows on the
periodic chart are
called Periods.
• Columns are Groups.
• Elements in the same
group have similar
chemical properties.
Groups
These five groups are known by their historic
group names.
Periodic Table
Nonmetals are
on the right
side of the
periodic table
(with the
exception of
H).
Periodic Table
Metalloids
(simiconductors)
border the stairstep line (with
the exception of
Al, Po, and At).
Periodic Table
Metals are
on the left
side of the
chart.
Exam Question:
Potassium is a __________ and chlorine is a
__________.
A) metal, nonmetal
B) metal, metal
C) metal, metalloid
D) metalloid, nonmetal
E) nonmetal, metal
Chemical Formulas
The subscript to the right
of the symbol of an
element tells the number
of atoms of that element
in one molecule of the
compound.
Chemical Formulas
Molecular compounds
are composed of
molecules and almost
always contain only
nonmetals, i.e. elements
on the right side of the
periodic table.
The bonding between
non-metals is more
covalent (sharing of
electrons)
Diatomic Molecules
These seven elements occur naturally as
molecules containing two atoms.
Types of Formulas
• Empirical (experimental) formulas give
the lowest whole-number ratio of atoms
of each element in a compound.
• Molecular formulas give the exact
number of atoms of each element in a
compound.
Types of Formulas
• Structural formulas show the
order in which atoms are
bonded.
• Perspective drawings also
show the three-dimensional
array of atoms in a
compound.
• Space-filling model
represent the boundaries of
electron density.
Ions
• When atoms lose or gain electrons, they
become ions.
– Cations are positive and are formed by elements
on the left side of the periodic chart.
– Anions are negative and are formed by elements
on the right side of the periodic chart.
Ionic Bonds
Ionic compounds (such as NaCl) are
generally formed between metals and
nonmetals by transfer of charge (electrons).
Exam Question
When a metal and a nonmetal react, the __________
tends to lose electrons and the __________ tends to
gain electrons.
A) metal, metal
B) nonmetal, nonmetal
C) metal, nonmetal
D) nonmetal, metal
E) None of the above, these elements share electrons.
Writing Formulas
• When compounds are electrically neutral, one
can determine the formula of an ionic
compound this way:
– The charge on the cation becomes the subscript
on the anion.
– The charge on the anion becomes the subscript
on the cation.
– If these subscripts are not in the lowest wholenumber ratio, divide them by the greatest common
factor.
Exam Question
What is the formula of the compound formed
between strontium ions and nitrogen ions?
A) SrN
B) Sr3N2
C) Sr2N3
D) SrN2
E) SrN3
Common Cations (Table 2.4)*
*Commit to Memory
Common Anions*
*Memorize…the anions as well. Note that many
common anions are polyatomic.
Inorganic Nomenclature (i.e.
naming inorganic compounds)
• Write the name of the cation.
• If the anion is an element, change its
ending to -ide; if the anion is a
polyatomic ion (see previous, simply
write the name of the polyatomic ion.
• If the cation can have more than one
possible charge (Fe2+ vs. Fe3+) write the
charge as a Roman numeral in
parentheses.
Patterns in Oxyanion Nomenclature
• When there are two oxyanions involving
the same element:
– The one with fewer oxygens ends in -ite.
• NO2− : nitrite; SO32− : sulfite
– The one with more oxygens ends in -ate.
• NO3− : nitrate; SO42− : sulfate
Exam Question:
The correct formula of iron(III) bromide is
__________.
A) FeBr2
B) FeBr3
C) FeBr
D) Fe2Br2
E) Fe3Br
What is the systematic name for Ca(NO3)2?
Exam Question:
The correct name for CaH2 is __________.
A) hydrocalcium
B) calcium dihydride
C) calcium hydroxide
D) calcium dihydroxide
E) calcium hydride
What is the systematic name for (NH4)2SO4?
Exam Question:
50) The correct name for CCl4 is __________.
A) carbon chloride
B) carbon tetrachlorate
C) carbon perchlorate
D) carbon tetrachloride
E) carbon chlorate
What is the systematic name for Co(OH)2?
NaH2PO4?
Write a chemical formula for cobalt(II) hydrogen
carbonate.
Patterns in Oxyanion
Nomenclature
• The one with the second fewest oxygens ends in -ite.
– ClO2− : chlorite
• The one with the second most oxygens ends in -ate.
– ClO3− : chlorate
Patterns in Oxyanion Nomenclature
• The one with the fewest oxygens has the prefix hypoand ends in -ite.
– ClO− : hypochlorite
• The one with the most oxygens has the prefix per- and
ends in -ate.
– ClO4− : perchlorate
Acid Nomenclature
• If the anion in the acid
ends in -ide, change
the ending to -ic acid
and add the prefix
hydro- .
– HCl: hydrochloric acid
– HBr: hydrobromic acid
– HI: hydroiodic acid
Acid Nomenclature
• If the anion in the acid
ends in -ite, change
the ending to -ous
acid.
– HClO: hypochlorous
acid
– HClO2: chlorous acid
Acid Nomenclature
• If the anion in the acid
ends in -ate, change
the ending to -ic acid.
– HClO3: chloric acid
– HClO4: perchloric acid
Exam Question:
The correct name for H2SO3 is __________.
A) sulfuric acid
B) sulfurous acid
C) hydrosulfuric acid
D) hydrosulfic acid
E) sulfur hydroxide
Nomenclature of Binary
Molecular Compounds
• The less electronegative
atom is usually listed first.
• A prefix is used to denote
the number of atoms of
each element in the
compound (mono- is not
used on the first element
listed, however) .
Nomenclature of Binary
Compounds
• The ending on the more
electronegative element
is changed to -ide.
– CO2: carbon dioxide
– CCl4: carbon tetrachloride
Nomenclature of Binary
Compounds
• If the prefix ends with a or o
and the name of the
element begins with a
vowel, the two successive
vowels are often elided into
one.
N2O5: dinitrogen pentoxide
not…dinitrogen pentaoxide
Exam Question:
The name of PCl3 is __________.
A) potassium chloride
B) phosphorus trichloride
C) phosphorous(III) chloride
D) monophosphorous trichloride
E) trichloro potassium
Nomenclature of Organic
Compounds
• Organic chemistry is the study of carbon
compounds.
• Organic chemistry has its own system of
nomenclature (i.e. naming).
Nomenclature of Organic
Compounds
The simplest hydrocarbons (compounds
containing only carbon and hydrogen) are
alkanes. Alkanes contain only single bonds
and are often referred to as “saturated”
Nomenclature of Organic
Compounds
The first part of the names above correspond
to the number of carbons (meth- = 1, eth- = 2,
prop- = 3, etc.).
Nomenclature of Organic
Compounds
• When a hydrogen in an alkane is replaced with
something else (a functional group, like -OH in
the compounds above), the name is derived from
the name of the alkane.
• The ending denotes the type of compound.
– An alcohol ends in -ol.