Chapter 2

Report
Chapter 2
Atoms, Ions, and Molecules: Matter
Starts Here
Chapter Outline
• 2.1 The Nuclear Model of Atomic Structure
• Electrons
• Radioactivity and the Nuclear Atom
• Protons and Neutrons
•
•
•
•
•
•
2.2
2.3
2.4
2.5
2.6
2.7
Isotopes
Average Atomic Mass
The Periodic Table of the Elements
Trends in Compound Formation
Naming Compounds and Writing Formulas
Nucleosynthesis
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Atomic Structure: Electrons
• J. J. Thomson (1897)
• Beam from cathode ray tube is deflected
toward positively charged plate.
• Atoms contain negatively charged particles
with a constant mass-to-charge ratio.
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Mass of an Electron
• Robert Millikan (1909)
• Determined the mass and charge of an
electron with his oil-drop experiment.
• e– = –1.602 x 10–19 C
• me = 9.109 x 10 –28 g
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Thomson’s Model of the Atom
• Plum-pudding model:
• e– distributed
throughout diffuse,
positively charged
sphere
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Radioactivity and the
Nuclear Atom
• Henri Becquerel (1896)
• Some materials produce invisible radiation,
consisting of charged particles.
• Radioactivity
• Spontaneous emission of high energy
radiation and particles
• Beta particles (, high energy electrons)
• Alpha particles (, +2 charge, mass = He nucleus)
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Atomic Structure: The
Nucleus
• Rutherford’s experiment:
• Bombard a thin gold foil with  particles to
test Thomson’s model of the atom.
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Rutherford’s Experiment
(b) Expected results from
plum-pudding model
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(c) Actual results
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The Nuclear Atom
• The nucleus:
• is the positively charged center of an atom,
containing nearly all of the atom’s mass.
• is about 1/10,000 the size of the atom.
• consists of two types of particles:
• Protons: positively charged subatomic particles
• Neutrons: electrically neutral subatomic particles
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The Nuclear Atom (cont.)
Nucleus has protons (+ charge) plus neutrons (neutral).
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Atomic Mass Units
• Atomic mass units (amu)
• Unit used to express the relative masses of
atoms and subatomic particles
• Equal to 1/12 of a carbon atom:
• 6 protons + 6 neutrons = 12 amu
• 1 amu = 1 Dalton (Da)
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The Nuclear Atom: Summary
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Chapter Outline
• 2.1 The Nuclear Model of Atomic Structure
• 2.2 Isotopes
• Positive Ray Analyzer Experiments
• Isotopes: Atoms of the Same Element, but Different Masses
•
•
•
•
•
2.3 Average Atomic Mass
2.4 The Periodic Table of the Elements
2.5 Trends in Compound Formation
2.6 Naming Compounds and Writing Formulas
2.7 Nucleosynthesis
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Aston’s Positive-Ray Analyzer
Ne gas ions of different
masses strike the
detector in different
locations.
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Aston’s Experiment
• Positive-ray analyzer results:
• Two different kinds of neon gas atoms exist.
• 90% = 20 amu
• 10% = 22 amu
• Aston proposed theory of isotopes.
• Isotopes
• Atoms of an element containing the same #
of protons but different numbers of neutrons.
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Symbols of Isotopes
Atomic Mass (A) – total number of nucleons
(protons + neutrons) in the nucleus.
A
Z
X
Elemental Symbol – a one- or
two-letter symbol to identify
the type of atom.
Atomic Number (Z) – the number of protons in the
nucleus; determines the identity of the element.
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Practice: Isotopic Symbols
• Use the format AX to write the symbol for
the nuclides having 26 protons and 30
neutrons.
- Collect and Organize: We have a nuclide with 26
protons and 30 neutrons.
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Practice: Isotopic Symbols
• Use the format AX to write the symbol for
the nuclides having 26 protons and 30
neutrons.
- Analyze: The number of protons represents the
atomic number and identity of the element, and
the total number of nucleons (protons +
neutrons) represents the atomic mass.
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Practice: Isotopic Symbols
• Use the format AX to write the symbol for
the nuclides having 26 protons and 30
neutrons.
- Solve: The number of protons (26) identifies this
nuclide as Fe (iron). The total number of
nucleons is 26 + 30 = 56 and is the mass of the
nuclide. So, the symbol for this nuclide would
be 56Fe.
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Practice: Isotopic Symbols
• Use the format AX to write the symbol for
the nuclides having 26 protons and 30
neutrons.
- Think About It: The mass of Fe in the periodic
table is 55.85, which is close to the mass of this
nuclide, so this represents a reasonable result
for an isotope of iron.
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Chapter Outline
• 2.1 The Nuclear Model of Atomic Structure
• 2.2 Isotopes
• 2.3 Average Atomic Mass
• Weighted Average of Isotopic Masses
• Natural Isotopic Abundances
•
•
•
•
2.4 The Periodic Table of the Elements
2.5 Trends in Compound Formation
2.6 Naming Compounds and Writing Formulas
2.7 Nucleosynthesis
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Average Atomic Masses
• Average atomic mass:
• Weighted average of masses of all isotopes
of an element
• Calculated by multiplying the natural
abundance of each isotope by its mass in
amu and then summing these products
• Natural abundance:
• proportion of a particular isotope; usually
expressed as a percentage relative to all the
isotopes for that element in a natural sample
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Weighted Average Example
Neon has three naturally occurring isotopes.
• Average atomic mass of neon:
(19.9924 x 0.904838) + (20.99395 x 0.002696) +
(21.9914 x 0.092465) = 20.1797 amu
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Chapter Outline
•
•
•
•
2.1 The Nuclear Model of Atomic Structure
2.2 Isotopes
2.3 Average Atomic Mass
2.4 The Periodic Table of the Elements
• Periodicity and Mendeleev’s Table of the Elements
• Navigating the Modern Periodic Table
• 2.5 Trends in Compound Formation
• 2.6 Naming Compounds and Writing Formulas
• 2.7 Nucleosynthesis
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Mendeleev’s Periodic
Table
Dmitri Mendeleev (1872)
• Ordered elements by atomic
mass
• Arranged elements in
columns based on similar
chemical and physical
properties
• Left open spaces in the table
for elements not yet
discovered
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The Modern Periodic Table
• Also based on a classification of elements
in terms of their physical and chemical
properties
• Horizontal rows – called periods (1 → 7)
• Columns – contain elements of the same
family or group (1 →18)
• Several groups have names as well as
numbers
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Groups of Elements
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Nonmetals
Metalloids
Metals
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Broad Categories of Elements
• Metals (left side and bottom of the table)
• Shiny solids; conduct heat and electricity;
are malleable and ductile
• Nonmetals (right side and top of the table)
• Solids, liquids and gases; nonconductors;
solids are brittle.
• Metalloids (between metals/nonmetals)
• Shiny solids (like metals); brittle (like
nonmetals); semiconductors
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Groups of Elements (cont.)
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Chapter Outline
•
•
•
•
•
2.1 The Nuclear Model of Atomic Structure
2.2 Isotopes
2.3 Average Atomic Mass
2.4 The Periodic Table of the Elements
2.5 Trends in Compound Formation
• Law of Multiple Proportions
• Molecular Compounds
• Ionic Compounds
• 2.6 Naming Compounds and Writing Formulas
• 2.7 Nucleosynthesis
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The Composition of Compounds
• Law of Multiple Proportions
• The ratio of the two masses of one element
that react with a given mass of another
element to form two different compounds is
the ratio of two small whole numbers.
• Examples:
• SO2, SO3
• NO, NO2
• NO2 = 22.8 g O/10 g N
• NO = 11.4 g O/10 g N
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Molecular Compounds
• Molecular compounds
• Composed of atoms held together in
molecules by covalent bonds
• Covalent bonds
• Bond between two atoms created by sharing
one or more pairs of electrons
• Molecular compounds are composed of
nonmetals.
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Molecular Compounds (cont.)
• Molecular formula:
• Shows the number and type of atoms present
in one molecule of a compound
• Empirical formula:
• Shows the smallest whole-number ratio of
elements in a compound
• Example: Glucose
• Molecular formula = C6H12O6
• Empirical Formula = CH2O
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Ionic Compounds
• Ionic compounds
• Consist of charged particles (ions) formed by
transfer of electrons between atoms
• Ions held together by electrostatic forces
Cations = ions with
positive charge
Anions = ions with
negative charge
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Ionic Compounds (cont.)
• Ionic compounds are made of a metal and
a nonmetal.
• Metals form cations; nonmetals form anions.
• Charges on ions depend on location in the
periodic table.
• e.g., Group 1 Metals = +1; Halogens = –1
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Ionic Compounds (cont.)
• Formula unit
• Smallest electrically
neutral unit of an
ionic compound
e.g., NaCl
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Charges on Ions
Figure 2.17
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Chapter Outline
•
•
•
•
•
•
2.1 The Nuclear Model of Atomic Structure
2.2 Isotopes
2.3 Average Atomic Mass
2.4 The Periodic Table of the Elements
2.5 Trends in Compound Formation
2.6 Naming Compounds and Writing Formulas
•
•
•
•
•
Binary Molecular Compounds
Binary Ionic Compounds
Binary Compounds of Transition Metals
Polyatomic Ions
Acids
• 2.7 Nucleosynthesis
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Naming Compounds
• Binary molecular compounds (e.g., SO3)
• Compounds consisting of two nonmetals
• First element in the formula is named first.
• S = sulfur
• Second element name is changed by adding suffix
-ide.
• O = oxygen → oxide
• Add prefixes to identify quantity of atoms (see
Table 2.2).
• SO3 = sulfur trioxide
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Rules for Using Prefixes
1. Do not use the prefix
mono- when naming first
element:
SO3 monosulfur trioxide
2. Prefixes ending with o- or
a- are modified when used
with elements beginning
with vowels:
P4O10 tetraphosphorus
decaoxide
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Binary Ionic Compounds
• Binary ionic compounds consist of
cations (usually metals) and anions
(usually nonmetals), e.g.,MgCl2.
• Cation named first using name of element.
• Mg = magnesium
• Anion named by adding the -ide suffix to the
name of the element.
• Cl = chlorine → chloride
• Formulas for ionic compounds must always
be neutral: Mg2+ + (Cl–) x 2
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Binary Ionic Compounds
(cont.)
• For metals that form cations with different
charges, a Roman numeral is added to
indicate the charge of the cation.
FeCl2:
Iron(II) chloride
FeCl3:
Iron(III) chloride
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Practice: Ionic Compounds
Write the name of the following
compounds.
a) NaCl
b) CrCl3
Write the chemical formula of the following
compounds.
a) Zinc nitride
b) Copper(I) oxide
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Common Polyatomic Ions
• Polyatomic ions
• Charged group of two or more atoms joined
together by covalent bonds
• Oxoanions
• Polyatomic anions containing oxygen in
combination with one or more other elements
• Examples: acetate (C2H3O2–), nitrate (NO3–),
carbonate (CO32–), perchlorate (ClO4–), sulfate
(SO42–) from Table 2.3, page 60
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Practice: Polyatomic Ions
•
Write the names of the following
compounds.
a) Cr(ClO4)3
b) NH4NO3
• Write the chemical formulas for the following
compounds.
a) Lithium bicarbonate
b) Calcium hypobromite
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Naming Binary Acids
• Binary acids
• Contain hydrogen and a monoatomic anion
(e.g., Cl–, S2–)
• Most common binary acids contain halogens.
(e.g., HCl, HBr)
• Acid names:
• The prefix “hydro-” + the halogen base name +
the suffix “-ic” + the word acid.
• Example: HBr is hydrobromic acid.
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Oxoanions & Related
Acids
• If oxoanion name ends in:
•
•
-ate
-ite
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• the corresponding acid
ends in:
•
-ic
•
-ite
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Practice: Naming Compounds
and Acids
• Identify each of the following as a
molecular compound, an ionic compound,
or an acid. Name or give formulas for the
compounds.
•K2Cr2O7
•Na3N
•NO2
•H2CrO4
•Sodium carbonate
•Sulfurous acid
•Iron(II) phosphate
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Chapter Outline
•
•
•
•
•
•
•
2.1 The Nuclear Model of Atomic Structure
2.2 Isotopes
2.3 Average Atomic Mass
2.4 The Periodic Table of the Elements
2.5 Trends in Compound Formation
2.6 Naming Compounds and Writing Formulas
2.7 Nucleosynthesis
• Primordial Nucleosynthesis
• Stellar Nucleosynthesis
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The Big Bang Revisited
H and He atoms in stars
fuse to form heavier
elements.
Subatomic particles fuse
to form H and He nuclei.
Existence of subatomic
particles
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Nucleosynthesis
• Nucleosynthesis
• Energy from Big Bang transformed into
matter (more details of this matter/energy
relationship in Chapter 21).
• Fusing of fundamental/subatomic particles
(protons/neutrons) created atomic nuclei.
1
1
p 
1
0
n 
2 d 
2
1
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2
2
1
d

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Nucleosynthesis (cont.)
• Stellar nucleosynthesis:
• High density and temperature in stars caused
additional fusion reactions to create elements
heavier than H, He:
3  
4
2
12
6
12
6
C  
4
2
C
16
8
O
• Stellar core forms shells of heavier elements
produced from fusion of lighter elements.
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Development of
Planets
• Fusion reactions in stars release energy,
increase density of stars.
Gravitational
compression reheats
core, resulting in
supernova; exploding
star disperses matter,
which coalesces into
planets.
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