Download Chapter 2 Atoms, Molecules and Ions

Chapter 2
Atoms, Molecules and Ions
• 2.1 The Atomic Theory of Matter
• 2.2 The Discovery of Atomic Structure
• 2.3 The Modern View of Atomic Structure
• 2.4 The Periodic Table
• 2.5 Molecules and Molecular Compounds
• 2.6 Ions and Ionic Compounds
• 2.7 Naming Inorganic Compounds
2.1 The Atomic Theory of Matter
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Greeks: 4 “elements”
Democritus and Leucippus - atoms
Aristotle- elements
Alchemy
1660 - Robert Boyle- experimental definition
of element.
• Lavoisier- Father of modern chemistry.
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.
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Dalton's Postulates
1.
Each element is composed of extremely small
particles called atoms.
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Dalton's Postulates
2.
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.
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Dalton's Postulates
3.
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.
Law of Conservation of Mass
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.
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Dalton’s Postulates
4. 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 (also known as Proust’s Law or the law
of definite proportions)
It states that the elemental composition of a pure substance never
varies.
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• Modern Version of the
Experiment: Using
electricity, water is
electrolyzed into its
components, hydrogen
and oxygen.
• Results: Water is always
found to contain hydrogen
and oxygen in a mass ratio
of 1:8 (volume ratio of
2:1) no matter how it is
formed or where it is
found.
2.2 The Discovery of
Atomic Structure
The Electron
Voltage source
In 1897 J.J. Thomson discovered that passing an electric
current through an evacuated sealed glass tube, makes a
stream of particles move from the negative to the
positive end.
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Thomson’s cathode ray tube experiment
Voltage source
+
-
By adding an electric field, he found that the moving
pieces were negative.
Thompson measured the charge/mass ratio of the
electron to be 1.76  108 coulombs/g.
Plum Pudding Model
The atom was thought to be
a positive sphere of matter
with negative electrons
imbedded in it.
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Robert Milikan’s
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.
Milikan sprayed oil drops above above a
charged plate containing a small hole.
As the oil drops fall through the hole,
they are given a negative charge. When
a drop is perfectly balanced, the weight
of the drop is equal to the electrostatic
force of attraction between the drop
and the positive plate.
Using this experiment, Millikan
determined the charge on the electron
to be 1.60  10-19 C.
Henri Becquerel discovers radioactivity
Radioactivity is the spontaneous emission of radiation by an atom.
Also studied by Marie and Pierre Curie.
Ernest Rutherford noted three spots on the detector:
a spot in the direction of the positive plate ( particles)
a spot which is not affected by the electric field ( rays)
a spot in the direction of the negative plate ( particles)
Rutherford’s gold foil
experiment
A source of α-particles was placed
at the mouth of a circular detector.
The α -particles were shot through
a piece of gold foil.
Rutherford’s nuclear atom
Since some particles were deflected at large
angles, but most went straight through,
Thompson’s model could not be correct.
Evidence  Theory
1) Some alpha particles deflected greatly 
presence of very small, dense positive
nucleus (with the electrons around the
outside)
2) Most alpha particles went straight through 
most of the volume of the atom is empty
space.
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Niels Bohr
• Electrons travel in
orbit around nucleus
• Idea of electron
transitions between
energy levels,
emitting/absorbing
photons
Bohr’s energy levels
• The energy levels in an atom are
sort of like steps of a ladder.
• The more energy an electron has,
the farther away from the nucleus it
usually will be.
• The energy levels are not evenly
spaced. They get closer together as
you travel farther away.
• To move from one “rung” to
another requires a “quantum” of
energy.
Bohr Atomic Model
The difference between continuous and quantized energy levels.
continuous energy levels
quantized energy levels
James Chadwick discovers the neutron
A thin sheet of beryllium was
bombarded with alpha particles, and a
very high energy radiation was
emitted.
The radiation was electrically neutral,
had no charge but a slightly larger
mass than a proton.
Heisenberg’s Uncertainty Principle
The Modern View
Quantum Mechanical Model
Quantum
Mechanical
Model
2.3 The Modern View of Atomic
Structure
Subatomic particles
We define: mass of 12C = exactly 12 amu.
Using atomic mass units:
1 amu = 1.66054 x 10-24 g
1 g = 6.02214 x 1023 amu
Symbols of Elements
Elements are symbolized by one or two letters.
First letter is always capital.
Second letter (if there is one) is lower case.
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Atomic Number
All atoms of the same element have the same
number of protons:
The atomic number (Z)
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Mass Number
The mass of an atom in atomic mass units (amu) is
the total number of nucleons (protons + neutrons)
The mass number (A)
By convention, for element X, we write ZAX.
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Isotopes
• Isotopes are atoms of the same element with different
masses.
• In terms of subatomic particles, isotopes have different
numbers of neutrons.
11
C
6
12
C
6
13
C
6
14
C
6
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Average Atomic 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.
• Average atomic mass =
Σ (isotope mass)( relative abundance)
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Atomic Mass Calculation 1
Isotope
Atomic Mass
Relative Abundance
28Si
27.976
92.2297%
29Si
28.976
4.6832%
30Si
29.973
3.0872%
• Using the isotope
information for Silicon.
Find the average atomic
mass.
Atomic Mass Calculation 2
Silver consists of two isotopes 107Ag and 109Ag. Its
average atomic mass is 107.87. Calculate the
percentage of each isotope in naturally occurring
silver. (Assume that the masses are 107.00 and
109.00 respectively.)
Atomic Mass
Atomic and molecular
masses can be
measured with great
accuracy with a mass
spectrometer.
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2.4 The Periodic Table
Periodicity
When one looks at the chemical properties of
elements, one notices a repeating pattern of
reactivities.
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Periodic Table
Metals are on
the left side of
the chart.
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Good conductors
Lose e- to form cation
Malleable
Ductlie
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Periodic Table
Metalloids
border the
stair-step line
• Intermediate
properties
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Periodic Table
Nonmetals are
on the right
side of the
periodic table
(except H).
• Poor conductors
• Gain e- to form anion
• Brittle
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Alkali Metals
Very reactive
1+
Alkaline Earth Metals
Reactive
2+
Transition metals
Form colored compounds
Variable charges
Inner Transition Metals
Lanthanides or rare earth
Actinides are all radioactive
Z ≥ 92 are manmade
Chalcogens
Oxygen family
6 valence electrons
Halogens
Very reactive
Only group containing elements in s, l, g @STP
Form acids when bonded with H
Noble Gases
Low reactivity
Full octet
Colorless, odorless gas
2.5 Molecules and Molecular Compounds
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.
Molecular compounds are
composed of molecules and
almost always contain only
nonmetals.
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Allotropes
Different forms of the same element.
Tin: White (metallic tin), Gray Tin
Oxygen: Oxygen gas, Ozone
Carbon:
Graphite, Diamond, “Bucky Balls”
Diatomic Molecules
These seven elements occur naturally as
molecules containing two atoms.
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Molecular vs. Empirical formula
• Empirical formulas give the lowest wholenumber ratio of atoms of each element in a
compound. NO2
• Molecular formulas give the exact number of
atoms of each element in a compound. N2O4
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Picturing molecules
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2.6 Ions and Ionic Compounds
• Ions – charged particle
monoatomic or polyatomic
• Cations – positively
charged particle
• Anions – negatively
charged particle
Background: Predicting Charges
Common Cation Charges
Common Anions
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Ionic Compounds
Ionic “bond” is the attraction between oppositely
charged ions (generally metals and nonmetals)
Called salts
Examples: Table salt (NaCl), baking soda (NaHCO3),
Epsom salts (MgSO4)
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Writing Formulas
Because compounds are electrically neutral, one can
determine the formula of a 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.
– IONIC ONLY: If these subscripts are not in the lowest
whole-number ratio, divide them by the greatest
common factor.
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2.7 Naming Inorganic Compounds
IONIC
Naming Cations
• Monoatomic cations have same name as metal
Na+
Al3+
• If metals can form > 1 charge use roman numerals
Fe2+
Fe3+
• Nonmetal polyatomic cations –ium
NH4 +
H3O +
• Naming Anions
• Monoatomic anions drop ending add -ide
HO2• Oxyanions –ate or -ite
NO3-
NO2-
• Adding H+ to oxyanions prefix = hydrogen or dihydrogen
CO3HCO3H2CO3-
Ionic Compounds & Formulas
Name the compound
Write the formula
KCl
Mg3N2
Na2SO4
(NH4)2CO3
CuO
Cu2O
FePO4
Sodium fluoride
Calcium nitrate
Barium nitrite
Lead (II) hydroxide
Managanese (IV) sulfide
Tin (II) sulfite
Sodium nitride
Names and Formulas of Acids
Name these acids
HCN
HNO3
HNO2
HClO4
HClO3
H2SO3
HCl
HBr
HI
Writing formulas for acids
• If it has hydro- in the name it has no oxygen (Anion
ends in –ide)
• No hydro in name, anion ends in -ate or -ite
• Write anion and add enough H to balance the
charges.
hydrofluoric acid
carbonic acid
phosphorous acid
dichromic acid
hydrophosphoric acid
perchloric acid
Names and Formulas of Binary Molecular
Compounds (covalent)
• 2 non metals
• Cation followed by anion (-ide)
• No roman numerals
• Prefixes (no mono on cation)
mono, di, tri, tetra, penta, hexa, septa, octa, nona, deca
Sulfur dioxide
diflourine monoxide
nitrogen trichloride
diphosphorus pentoxide
CO2
CO
CCl4
N2O4
XeF6
P2O10
N4O4
Hydrates
• Hydrates are compounds that contain discrete water
molecules as part of the crystal lattice structure.
• CuSO4•5H2O is called copper(II)sulfate pentahydrate.
• In the formula you put a dot and then write the
number of molecules of water, using the Greek
prefixes.
• Calcium chloride dihydrate
• Cr(NO3)3 6H2O