Download Chapter 2 Law of Conservation of Mass Law of Conservation of Mass

Chapter 2
Atoms, Ions, and the Periodic Table
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Dalton’s Atomic Theory
Structure of the Atom
Ions
Atomic Mass
The Periodic Table
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Law of Conservation of Mass
• the mass of the products (overall)
always equals the mass of the
reacting substances
• Proposed by Antoine Lavoisier in
1787
• His experiments showed that no
measurable change in mass occurs
during a chemical reaction.
How could you show mass is conserved in a reaction?
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Law of Conservation of Mass
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Law of Definite Proportions
• Proposed by Joseph Proust between 1797
and 1804
• states that all samples of the same
compound always contain the same
proportions by mass of the component
elements
– For example, water is always composed of
oxygen and hydrogen in a mass ratio of 8:1
(or 16:2).
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Dalton’s Atomic Theory
•
Dalton’s Atomic Theory has 4
postulates:
1. All matter is composed of exceedingly small,
indivisible particles called atoms.
2. All atoms of a given element are identical both
in mass and in chemical properties. However,
atoms of different elements have different
masses and different chemical properties.
3. Atoms are not created or destroyed in chemical
reactions.
4. Atoms combine in simple, fixed, whole-number
ratios to form compounds.
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Structure of the Atom
• Atoms actually are divisible. They are composed of
subatomic particles.
• Subatomic particles include:
– 1 kind of particle found outside the nucleus
• Electrons
– negatively charged subatomic particles
– 2 kinds of particles found in the nucleus
(center of the atom)
• Protons
– positively charged subatomic particles
• Neutrons
– uncharged subatomic particles
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Structure of the Atom
Figure 2.10
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The Discovery of Electrons
•
The existence of the electron was demonstrated by
J.J. Thomson in 1897.
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He conducted a series of experiments with cathode ray
tubes, in which:
1. Voltage was applied by connecting each end of a
tube to a battery.
2. The electricity forms rays that flow from one end
of the tube to the other and that are visible
through specially coated glass.
3. When an electric or magnetic field was applied to
the tube (and the rays), the rays bent toward a
positively charged plate, and were deflected by a
negatively charged plate. Because like charges
repel and opposite charges attract, the particles
were negatively charged.
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The Discovery of Electrons
Figure 2.6
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The Nuclear Atom
•
From his experiments with electrons, J.J. Thomson proposed
that electrons might be embedded in a sphere of positive
charge (“plum pudding” model of the atom).
•
Ernest Rutherford designed an experiment in the early 1900’s
to test J.J. Thomson’s “plum pudding” model of the atom.
– The experiment involved bombarding a piece of
gold foil with alpha particles (positively charged
Helium atoms without the electrons).
– Alpha particles were expected to zip through the
gold foil, and most did, but some were deflected
slightly and a few bounced backwards.
– The deflected particles led to the hypothesis of the
nucleus, a concentrated, positively charged core,
while electrons occupied the volume outside of the
nucleus.
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The Nuclear Atom
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Subatomic Particles
• Protons have:
– a charge equal to +1.6022 x 10-19 C
(expressed as +1)
– a mass equal to 1.6726 x 10-24 g (approx. the
same mass as a hydrogen atom)
• Neutrons have:
– no charge
– a mass equal to 1.6749 x 10-24 g
– Neutrons were proposed by Ernest
Rutherford in 1907 (to account for a mass
discrepancy in the nucleus) and discovered
in 1932 by James Chadwick.
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Subatomic Particles Continued
• Electrons have:
– a charge equal to -1.6022 x 10-19 C
(expressed as -1)
– a mass equal to 9.1094 x 10-28 g (1836
times less than the mass of one
hydrogen atom)
– Electrons were discovered in 1897 by
J.J. Thomson.
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Subatomic Particles
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Atomic Number and Mass
Number
• Atomic Number (Z)
– the number of protons in the nucleus of an element’s
atom
– is generally found on the periodic table above the
elemental symbol
• Mass Number (A)
– the number of protons and neutrons in the nucleus of
an element’s atom
– is generally found below the elemental symbol on the
periodic table
A=Z+N
• Neutron Number (N)
– the number of neutrons in the nucleus of an element’s
atom
N=A-Z
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Atomic Number and Mass
Number
79
0
197
Au 110
Atomic Mass
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Isotopes
• An isotope of an element
– is an atom that contains a specific number of
neutrons.
– Many elements have multiple isotopes.
– Specific isotopes have many applications, particularly
in medical testing, imaging, and treatment.
• An isotope symbol (Nuclide Symbol)
– is a common notation that represents the mass
number, atomic number, and elemental symbol.
• The subscript in the isotope symbol is the atomic
number.
• The superscript in the isotope symbol is the mass
number.
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Isotopes
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Practice – Isotope Symbols
•
Practice writing the isotope symbols
for the following isotope pairs.
1.
2.
3.
4.
Carbon-13 and carbon-14
Chlorine-35 and chlorine-37
Uranium-235 and uranium-238
Lithium-6 and lithium-7
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Practice Solutions – Isotope
Symbols
•
Practice writing the isotope symbols for the
following isotope pairs.
1. Carbon-13 and carbon-14
13
6
C and
14
6
C
2. Chlorine-35 and chlorine-37
35
17
Cl and
37
17
Cl
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Practice Solutions – Isotope
Symbols
•
Practice writing the isotope symbols for the
following isotope pairs.
3. Uranium-235 and uranium-238
235
92
U and
238
92
U
4. Lithium-6 and lithium-7
6
3
Li and 73 Li
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Ions
• An ion
– is a charged atom that contains more or less electrons
than protons.
– The overall charge is represented as a superscript to
the right of the elemental symbol.
• Ions can be classified as cations or anions.
– Cations
• are ions with a positive charge
• have less electrons than protons
– Anions
• are ions with a negative charge
• have more electrons than protons
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Ions
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Practice – Counting Protons and
Electrons
•
Write the number of protons and
electrons for the following ions:
1. Na+
2. Cl3. O24. Al3+
5. P32 - 24
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Practice Solutions – Counting
Protons and Electrons
•
Write the number of protons and electrons for
the following ions:
1. Na+
Sodium has an atomic number equal to 11.
Thus, it has 11 protons. It also has a +1
charge, and therefore has 1 less electron
than proton. Thus, it has 10 electrons.
2. ClChlorine has an atomic number equal to 17.
It has 17 protons. It also has a -1 charge,
and therefore has 1 more electron than
proton. Thus, it has 18 electrons.
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Practice Solutions – Counting
Protons and Electrons
•
Write the number of protons and electrons for
the following ions:
3. O2Oxygen has an atomic number equal to 8. It
has 8 protons. It also has a -2 charge, and
therefore has 2 more electrons than protons.
Thus, it has 10 electrons.
4. Al3+
Aluminum has an atomic number equal to
13. It has 13 protons. It also has a +3
charge, and therefore has 3 less electrons
than protons. Thus, it has 10 electrons.
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Practice Solutions – Counting
Protons and Electrons
•
Write the number of protons and
electrons for the following ions:
5. P3Phosphorus has an atomic number
equal to 15. It has 15 protons. It also
has a -3 charge, and therefore has 3
more electrons than protons. Thus, it
has 18 electrons.
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Atomic Mass
• Individual atomic masses are determined
by mass spectrometry.
• Instead of expressing atomic masses in
grams (a very small number), chemists
express atomic masses in atomic mass
units.
– An atomic mass unit (amu) is equal to
1/ the mass of a carbon-12 atom.
12
1 amu = 1/12 x mass of 1 12C atom
1 amu = 1.6606 x 10-24 g
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Atomic Mass
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Relative Atomic Mass
• Because most elements have multiple
isomers, the masses on the periodic table
cannot describe only 1 isotope’s individual
atomic mass.
• Therefore, the mass numbers on the periodic
table are relative atomic masses:
– Relative atomic mass is the average mass of
the individual isotopes of an element, taking
into account the naturally occurring relative
abundance of each.
– To find the relative atomic mass for an element,
sum the mass contributions from each isotope
of the element.
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Relative Atomic Mass
Mass contribution from isotope = Isotope mass x relative
abundance
Relative Atomic Mass = Mass contribution from 1st
isotope + Mass contribution from 2nd isotope + …
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Example
69. An unknown element (X) discovered on a planet in another
galaxy was found to exist as two isotopes. Their atomic
masses and percent abundances are listed in the following
table. What is the relative atomic mass of the element?
Isotope
Mass (amu)
Natural
Abundance (%)
22X
21.995
75.00
20X
19.996
25.00
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Mendeleev’s Table
• Russian chemist Dmitri Mendeleev
developed and published the basic
arrangement of the periodic table between
1869 and 1871.
• Mendeleev arranged the elements in order of
increasing relative atomic mass (protons
had not been discovered yet). The elements
on the modern periodic table are arranged in
order of increasing atomic number.
• He also grouped elements with similar
properties into columns and rows so that the
properties of the elements varied in a regular
pattern (periodically).
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Mendeleev’s Table
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The Modern Periodic Table
• The elements in the modern periodic table are
arranged by increasing atomic number (Z) and
in columns and rows to emphasize periodic
properties.
• The columns are collectively called families or
groups and are designated in two ways:
1. A Roman numeral (I through VIII) and a letter (A
or B)
2. An Arabic number (1-18)
• The rows are collectively called periods and
are designated by an Arabic number (1-7).
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The Modern Periodic Table
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Metals, Nonmetals, and
Metalloids
• The periodic table has many classifications.
Groups and Periods are one classification.
Another classification denotes metals,
nonmetals, and metalloids.
– A stair-step line starting at boron (B)
separates metals (to the left of the line) from
nonmetals (to the right of the line).
– The metalloids exist along the line.
• Metalloids are elements that have physical
properties resembling a metal, but the
chemical reactivity of a nonmetal.
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Metals, Nonmetals, and
Metalloids
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Main-Group Elements and
Transition Metals
• Main-group elements (also called
representative elements) contain any
element in the eight groups designated with
the letter A. (In the Arabic numbering,
groups 1, 2, and 13-18)
• Transition metals contain any element in the
10 groups designated with the letter B. (In
the Arabic numbering, groups 3-12)
• Inner-transition metals contain the
lanthanides and actinides listed separately
at the bottom of the table. (14 Groups)
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Main-Group Elements and
Transition Metals
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Common Group Names
• Some groups have descriptive names that are
commonly used instead of their group numbers.
– Alkali metals
• Group 1 (IA) metals (hydrogen is a nonmetal)
• are considered reactive because the react
readily with other elements and compounds
– Alkaline earth metals
• Group 2 (IIA) metals
• are more reactive than the transition metals
but less reactive than alkali metals
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Common Group NamesContinued
• Some groups have descriptive names that are
commonly used instead of their group numbers.
– Halogens (Halides)
• Group 17 (VIIA) nonmetals
• exist naturally as diatomic molecules
– Noble gases
• Group 18 (VIIIA) nonmetals
• are also called inert gases
• are so named because they do not
chemically react with other elements (with
the exception of krypton and xenon)
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Common Group Names
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Ions and the Periodic Table
• The noble gases are the most stable (least
reactive) elements on the periodic table.
– Their stability is associated with the number
of electrons they contain (8 electrons in their
outermost layer (or shell)).
– Many atoms in the main-group elements gain
or lose electrons to achieve similar stability.
• Metals tend to lose electrons, and
therefore become cations.
• Nonmetals tend to gain electrons, thereby
becoming anions.
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Ions and the Periodic Table
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Practice – Predicting Charges
for Ions
•
Write the symbol for the ion that each of
the following elements is predicted to
form:
1.
2.
3.
4.
5.
Beryllium
Aluminum
Phosphorus
Chlorine
Oxygen
Appendix B in Lab Book - Ions to Learn
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Practice Solutions – Predicting
Charges for Ions
•
Write the symbol for the ion that each of the
following elements is predicted to form:
1. Beryllium
Beryllium is in group IIA (2), so it will lose two
electrons to form Be2+.
2. Aluminum
Aluminum is in group IIIA (13), so it will lose
three electrons to form Al3+.
3. Phosphorus
Phosphorus is in group VA (15), so it will gain
three electrons to form P3-.
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Practice Solutions – Predicting
Charges for Ions
•
Write the symbol for the ion that each of the
following elements is predicted to form:
4. Chlorine
Chlorine is in group VIIA (17), so it will gain one
electron to form Cl-.
5. Oxygen
Oxygen is in group VIA (16), so it will gain two
electrons to form O2-.
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