Download 1) Molecular Compounds

AP CHEMISTRY II
Chapter 2 Notes
I. HISTORY OF THE ATOM
The Ancient Greeks were the first to come up with the idea of the atom.
1) Two Greek Philosophers, Democritus and Leucippus believed that all
matter was made up of tiny particles. They observed sand on the
seashore and saw that if you broke it into smaller pieces, it was still
sand. They reasoned that if you kept breaking the sand into smaller
and smaller pieces, eventually you would get to a piece that couldn’t
be broken. This smallest piece of matter they called an atomos which
means indivisible.
2) They developed a theory of matter based on five principles.
Point #1 - All matter is made up of undividable particles called atoms
Point #2 - There is a void, which is empty space between atoms.
Point #3 - Atoms are completely solid
Point #4 - Atoms are homogeneous, with no internal structure.
a) Point #5 - Atoms vary in size, shape and weight
3) Aristotle was another Greek Philosopher. He believed that all matter
was composed of four elements: Earth, Wind, Fire, Water. All matter
was made up of different proportions of these four.
Earth = cool and heavy;
Wind = light;
Fire = hot;
Water = wet.
4) Who was right? Even though Democritus and Leucippus were closer
to being right, Aristotle won the argument. Why?
a) Greeks didn’t experiment, they argued. Aristotle was more
famous and so he won.
b) In fact, his ideas carried throughout the middle ages. The
Alchemists believed Aristotle’s theory and were trying to use
these four elements in different proportions to turn base metals
into gold.
B) Antoine Lavoisier, The Father of Modern Chemistry—demonstrated with
careful measurements that transmutation of water to earth was not
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possible, but that the sediment observed from boiling water came from
the container. He burnt phosphorus and sulfur in air, and proved that the
products weighed more than he original. Nevertheless, the weight gained
was lost from the air. Thus he established the Law of Conservation of
Mass. In a Chemical reaction Mass is neither created nor destroyed.
C) Dalton’s Atomic Theory—John Dalton—English school teacher—Unlike
the Greeks, Dalton performed experiments to test and correct his atomic
theory. He studied the ratios in which elements combine in chemical
reactions and formulated hypotheses and theories which could be tested.
1) All elements are composed of tiny indivisible particles called atoms.
2) Atoms of a given element are identical in size, mass and other
properties; atoms from different elements differ in size, mass, and
other properties.
3) Atoms of different elements can physically mix together or can
chemically combine with one another in simple whole number rations
to form compounds
4) Chemical reactions occur when atoms are separated, joined, or
rearranged. Atoms of one element, however, are never changed into
atoms of another element as a result of chemical reactions.
II. The Atom—the smallest particle of an element that retains the properties of
that element. (The element Sulfur is only made up of Sulfur atoms)
A) The Structure of an Atom—most of Dalton’s Atomic Theory is accepted
today but one major change is the fact that an atom can be divided.
1) J.J. Thomson an English physicist discovered the existence of
electrons in 1897. He performed experiments using a cathode ray
tube. A sealed glass tube containing different gases. When
connected to high voltage electricity a glowing beam called a cathode
ray is created. By bringing positively charged metal plates near the
cathode ray, the path was altered. Since unlike charges attract each
other, Thomson determined that the cathode ray was made up of
negatively charged particles. Since no difference was found by using
different gases in the tube, and different metals for the plates.
Thomson concluded that all atoms contain electrons
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Thomson could not find a positively charged particle, so he believed
that the electrons were like plums embedded in a positively charged
“pudding,” thus it became known as the “plum pudding” model.
2) Robert Millikan added to our understanding of the electron with his oil
drop apparatus. Millikan sprayed very fine drops of oil into the drum,
where they would drop through a very small hole. The drum had two
electric plates on the inside. Millikan would watch through a scope
and measure the speed at which the drops fell. When adding an
electrical charge he noticed that the drops fell slower and in fact he
could add enough charge to cause the drops to stop completely in
mid-air. Through his experiments, Millikan determined both the mass
and the amount of charge for the electron.
3) Conclusions from the study of Electrons:
a) Cathode rays have identical properties regardless of the element
used to produce them. All elements must contain identically
charged electrons
b) Atoms are neutral, so there must be positive particles in the atom
to balance the negative charge of the electron
c) Electrons have so little mass, that atoms must contain other
particles that account for most of the mass.
4) Ernest Rutherford and the nucleus (1910)—Rutherford believed in the
plum pudding model of the atom. He wanted to see how big atoms
were, so he used radioactivity. He shot alpha particles - positively
charged pieces given off by uranium through gold foil which can be
made a few atoms thick. He expected the alpha particles to go
straight through the gold foil. Since the positive charges in the atom
were spread out evenly and could not deflect the positively charged
alpha particles. However what he found was completely
unexpected. Many of the particles did go straight through the gold
foil, but several were deflected, some were turned at 90 degrees or
more! In his words, it was “like a howitzer shells bouncing off of
tissue paper”. A new model of the atom had to be devised.
Rutherford concluded:
a) The atom is mostly empty space
b) It has a small dense, positively charge center
c) Alpha particles are deflected when they get near this center or
nucleus
5) TRY IT YOURSELF
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6) In 1886 twelve years before Thomson, a man named E. Goldstein
observed in a cathode ray tube rays flowing in the opposite direction.
He called these rays canal rays and concluded that they were
composed of positive charges. This lead eventually to the discovery
of the proton
7) In 1932 English physicist James Chadwick confirmed the existence of
another subatomic particle, the neutron.
B) How these particles fit together
1) Electrons are negatively (-) charged subatomic particles. They are
found orbiting around the nucleus in shells.
2) In the nucleus are found:
a) Protons—have a positive (+) charge, and
b) Neutrons—have no charge, or are neutral (they are NOT neutrally
charged!)
c) The nucleus has an overall positive charge
3) In a neutral atom of any element, the number of protons in the
nucleus and the number of electrons orbiting around the nucleus is
always equal—this number is known as the Atomic Number.
4) In the Periodic Table, the elements are arranged in order of
increasing atomic number.
5) An element’s Atomic Mass number is the number of protons PLUS
the number of neutrons in the nucleus—(added together!)
6) To find the number of neutrons in the nucleus of the atom, subtract
the atomic number from the atomic mass.
Number of Neutrons = Atomic Mass – Atomic Number
The number of neutrons in the element Fluorine is 19 (mass) – 9
(number) = 10
Mercury: 200 – 80 = 120
7) Symbols: Contain the symbol of the element, the mass number and
the atomic number
Mass
Number
Atomic
Number
X
C
12
6
4
Examples:
8) Isotopes—all atoms of the same element have the same number of
protons, however the number of neutrons can vary, these varied
atoms are called isotopes. Isotope—atoms of the same element with
different number of neutrons.
Example: H has 3 isotopes
Example: C has 3 isotopes
1
1H
2
1H
3
1H
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6C
13
6C
14
6C
1 p+ , 0n, 1 e1 p+ , 1n, 1 e1 p+ , 2n, 1 e6 p+ , 6n, 6 e6 p+ , 7n, 6 e6 p+ , 8n, 6 e-
C) Size of Atoms
1) Atoms are small.
2) Measured in picometers, 10-12 meters
3) Hydrogen atom, 32 pm radius
4) Nucleus tiny compared to atom
5) If the atom was the size of a football stadium, the nucleus would be
the size of a marble.
6) Radius of the nucleus near 10-15m.
7) Density near 1014 g/cm
D) The Mass of an Atom—most of the mass of an atom is concentrated in
the nucleus
1) Proton = 1.67 x 10-24 g
2) Neutron = 1.67 x 10-24 g
3) Electron = 9.11 x 10-28 g
E) The mass of the electron is virtually negligible, since the mass of the
nucleus contains approximately 3600X more mass.
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F) Rather than using these extremely small numbers, chemists instead
compare the masses of atoms using simple whole numbers called amu
or Atomic Mass Units
1) Mass of 1 proton = 1 amu
2) Mass of 1 neutron = 1 amu
3) 1.67 x 10-24 g = 1 amu
G) The Atomic Mass (mass number or molar mass) of an element is the
weighted mathematical average of the masses of all the isotopes of that
element.
Example 1
Chlorine (Cl)
35
17Cl
Cl-35 has a relative abundance of 77.5% and Cl-37 has a relative abundance of
22.5%. To calculate the atomic mass of Cl:
35 x (.775) + 35 x (.225) = 27.125 + 8.325 = 35.45
Check your Periodic Table to verify the atomic mass of Cl is 35.45.
Example 2
An unknown element, X, has 2 isotopes—one with a mass number of 10 that
has a relative abundance in nature of 20% and the other with a mass
number of 11 that has a relative abundance in nature of 80%. What is the
atomic mass of this element? What is its atomic number and identity?
Atomic mass = 10 x (.20) + 11 x (.80) = 2.0 + 8.8 = 10.8
Atomic number = 5 = Boron
III. THE PERIODIC TABLE
A) Dmitri Mendeleev—developed the first periodic table based on Atomic
Mass. He observed that the different groups of elements had similar
properties, and that these properties periodically repeated when
arranged by atomic mass. So he arranged the elements in a table by
their increasing mass in horizontal rows, and put elements with similar
properties in vertical columns. Thus creating the first Periodic Table
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B) This periodic repetition of physical and chemical properties is called the
Periodic Law
C) The Periodic Table used to day is based on increasing Atomic number
instead of the Mass Number
D) Horizontal row is called a period—numbered from 1 to 7 tell the outer
most electron shell
E) Vertical column is called a group or family—numbered I-VIIIA (0) and
I-VIIIB or 1-18—same number of electrons in the outermost shell—
similar properties.
F) Representative Elements—are the group A elements, they illustrate the
entire range of chemical properties
G) Metals
1) Found on the left side of the periodic table—Groups IA and IIA and
some in IIIA as well as all Group B or Transition Metals and Inner
transitions metals which are the rare earth elements or Lanthanide
and Actinide series elements.
a) Alkali Metals—IA
b) Alkaline Earth Metals—IIA
c) Transition Metals—IIIB-IIB
d) Inner Transition Metals—Rare Earth Elements—Lanthanide and
Actinide series
2) Highly lustrous when clean
3) Good electrical conductors
4) Ductile—can be drawn into wire
5) Malleable—can be beaten into thin sheets
6) Metals tend to lose electrons in a chemical reaction
H) Non-Metals
1) Found on the right side of the Periodic Table –from Group IVA and
over
a) Halogens—VIIA
b) Noble Gases—0 or VIIIA
2) Non lustrous (dull)
3) Poor electrical conductors
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4) Can be gases (N, O, F, Cl, I, and all of Group VIIIA [the noble gases])
or liquids (Br) or brittle solids (C, S, P, Se, etc)
5) They tend to gain electrons in a chemical reaction
I) Metalloids—elements with properties of both metals and non-metals—
found along the zig-zag line groups IIIA-VIIA—(Semi-conductors)
IV. Atoms and Ions
A) According to atomic theory, elements are composed of atoms of the
same kind—example the element Carbon is composed of only Carbon
atoms.
B) The charge on an atom is neutral (zero) because the number of protons
(+) is equal to the number of electrons (-)
C) When atoms of different elements combine to form chemical compounds,
they may either gain, lose or share electrons.
D) When electrons are gained or lost by and atom, it is no longer neutral, it
now has differing numbers of protons and electrons—it is called an Ion
E) An ion is an atom or group of atoms that has a positive or negative
charge.
F) Whether or not an atom of a particular elements gains or loses electrons
is based on the number of valence electrons it has.
V. Electron Configuration
A) Electrons are found in shells that surround the nucleus
B) There are 7 shells each with sub shells. The number of the shell
indicates the theoretical number of sub shells that shell contains.
(Example: shell 1 has only 1 sub-shell, shell 6 has 6 sub-shells)
C) Each sub shell can contain a different number of electrons and the
number increases with each of the sub-shells
D) The sub-shells are named s, p, d, f… we only use these four even
though theoretically there could be more g, h, I, j,…
1) An s sub-shell can contain only 2 electrons—they are described as
having different spins  and .
2) The p sub-shell can contain 6 electrons—they are described as x, y,
and z and  and .
3) The d sub-shell can contain 10 and the f sub-shell 14.
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4) Groups IA and IIA are filling up the s sub-shell of that shell, Groups
IIIA-VIIIA are filling up the p sub shell of that shell.
5) Octet Rule—atoms want 8 electrons in the outer shell—(except for 2
in shell #1)
6) Electrons are located in the shell that requires the least amount of
energy, so the shells fill in the following order.
7) Diagonal Rule
1s
2s 2p
3s 3p 3d
4s 4p 4d 4f
5s 5p 5d 5f
6s 6p 6d
7s
VI. Chemical Compounds and Formulas
A) Compounds are formed when atoms of 2 or more different elements
combine chemically.
B) Follow the Law of Definite Proportion—the elements always combine in
the same proportion by mass
C) Two types of compounds—Molecular compounds and ionic compounds
1) Molecular Compounds
a) Made of molecules—a neutral group of atoms that act as a single
unit
b) Most molecular compounds are composed of 2 or more
nonmetals
c) They have low melting and boiling points
d) Most exist as either gases or liquids at room temperature (some
solids)
2) Ionic Compounds
a) Composed of positive and negative ions (Cations and Anions)
b) Most are combinations of Metals and nonmetals
c) They have extremely high melting points
d) Most exist as crystalline solids at room temperature
3) Ionic Compounds Con’t
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a) Positive ions are called Cations
(1) Formed by losing electrons
(2) More protons than electrons
(3) Metals form cations
b) Negative ions are called Anions
(1) Formed by gaining electrons
(2) Non metals form anions
c) The charge on the ion is written as a superscript on the right (K+1
or O-2)
d) The electrons lost by the cation are gained by the anion
e) The cation and anions surround each other—there is no
distinction between molecules so the smallest piece of an ionic
compound is called a FORMULA UNIT
f) A formula unit represents the lowest whole-number ration of ions
in an ionic compound
Two Types of Compounds
Smallest
piece
Types of
elements
State
Melting
Point
Ionic
Molecular
Formula Unit
Molecule
Metal and
Nonmetal
Nonmetals
solid
Solid, liquid
or gas
High >300ºC
Low <300ºC
4) Chemical Formulas
a) Chemical Formula shows the kinds and numbers of atoms in the
smallest representative unit of a substance.
b) Two Types
(1) Molecular formula- number and kinds of atoms in a molecule.
Examples: CO2 , C6H12O6
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(2) Formula Unit--the smallest whole number ratio of atoms in an
ionic compound. Examples: NaCl , K2O , MgBr2
c) Law of Multiple Proportions—whenever 2 elements can form more
than 1 compound, the different masses of 1 element that can
combine with the same mass of the other element are in the ratio
of small whole numbers
5) Charges on ions
a) For most of the Group A elements, the Periodic Table can tell
what kind of ion they will form from their location.
b) Elements in the same group have similar properties including the
charge when they are ions
c) Groups 1A, 2A, 3A lose electrons becoming positively charged.
d) Their ionic charge is the same as their group number +1, +2, or
+3.
e) Groups 5A, 6A, and 7A gain electrons becoming negatively
charged.
f) Their ionic charge is equal to 8 minus their group number -3, -2,
-1
g) Group 4A elements either lose 4 electrons (+4) or gain 4 electrons
(-4) depending on the other element they combine with
h) Group 8A or 0 elements are called the Noble Gases. Their
electron structure is stable, so they rarely form ions, and so rarely
react to anything
6) What about the others?
a) Transition elements are metals—they lose electrons when
combining chemically with other atoms
b) Because of their electrons structure, most transition metals have
more than one common ionic charge
c) They are named the same as the neutral atom, plus a Roman
numeral to indicate their ionic charge
(1) Example: Fe+2 = Iron (II), Fe+3 = Iron (III)
VII. Naming Ions
A) Cation- if the charge is always the same (Group A) just write the name of
the metal (and ion)
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B) Transition metals need the name plus the Roman numeral (and ion)
Write Formulas for these
Name these
 Na+1
 Ca+2
Potassium ion
 Magnesium ion
 Copper (II) ion
 Chromium (VI) ion
 Barium ion
 Mercury (II) ion

Al+3
 Fe+3
 Fe+2
 Pb+2
 Li+1

C) Anions—change the element name by dropping the last few letters and
adding “-ide”
Name these
Cl-1
 N-3
 Br-1
 O-2
 Ga+3

Write these
Sulfide ion
iodide ion
 phosphide ion
 Strontium ion


D) Polyatomic ions
1) Groups of atoms that stay together and act as a single ion
2) You must memorize these:
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Polyatomic ions

Acetate C2H3O2-1
 Nitrate NO3-1
 Nitrite NO2-1
 Hydroxide OH-1
 Permanganate MnO4-1
 Cyanide CN-1

VIII.
Polyatomic ions
Sulfate SO4-2
 Sulfite SO3-2

Phosphate PO4-3
 Phosphite PO3-3

Carbonate CO3-2
Chromate CrO4-2
 Dichromate


Ammonium NH4+1
Cr2O7-2
Ionic Compounds
A) Naming Binary Ionic Compounds
1) Binary Compounds contain only 2 elements
2) They contain one cation and one anion
3) To write the names just name the two ions-the positive ion ALWAYS
comes first, the negative ion second.
4) With transition metals, you must also include Roman numeral
representing the charge.
Naming Binary Ionic Compounds

Write the names of the following
KCl
Na3N
N

Sc3P2

PbO

PbO2

Na2Se



B) Naming Ternary Ionic Compounds
1) They contain at least 3 elements--polyatomic ions
2) Name like binary compounds but use polyatomic ion name
3) The positive and negative charges must balance each other. If while
balancing charges you need more than one polyatomic ion, put a
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parentheses around the polyatomic, then a subscript outside the
parentheses
Example Ca+2 + NO3-1
Ca(NO3)2
4) Naming polyatomic ions
a) If oxygen and only one other element are present the ion will end
in “-ate”
Examples: Bromate (Br and O) Br)3-1
Carbonate (C and O) CO3-2
b) The “-ite” ending, means there is one less Oxygen atom than the
corresponding “-ate” ion, but the ionic charge is the same.
Example: Chlorate ClO3-1
Chlorite Cl02-1
c) The prefix “hypo-” in front of the ion means there is even one less
Oxygen than “-ite”
Example: Chlorite Cl02-1
Hypochlorite Cl0-1
d) The prefix “per-“ in front of the ion means there is one more
Oxygen atom than the ion without the “per-” prefix.
Example: Chlorate Cl03-1
Perchlorate Cl04-1
e) If a Hydrogen atom is added to the front of the ion, add
“hydrogen” to the front of the name. The ion’s charge is reduced
by “1”.
Example: Sulfate SO4-2
Hydrogen Sulfate HSO4-1
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Naming Ternary Ionic
Compounds
 NaNO3
 LiCN
 CaSO4
 Fe(OH)3
 CuSO3
 (NH4)2CO3
 (NH4)2O
 NiPO4
IX. Molecular Compounds
A) These compounds are composed of 2 non-metals—from groups 4A, 5A,
6A, 7A—occasionally, hydrogen can also act as a non-metal (when it
bonds with Oxygen)
B) These compounds are molecules, and do not contain ions—so the ionic
charges of the individual elements are not used in writing formulas
C) Non-metals often combine in more than one way, and each form has
different physical and chemical properties
Example: CO2 and CO
SO2 and SO3
PCl3 and PCl5
D) Naming Molecular Compounds
1) The second element always ends in “-ide”
2) Prefixes are used in both element’s names to signify how many
atoms are in each molecule
# of Atoms
Prefix
1
Mono2
di3
tri4
tetra5
penta6
hexa7
hepta8
octa9
nona10
deca15
3) Exception: when there is only 1 atom of the first element, drop the
“mono-“
Examples: CO2 Carbon Dioxide
4) Also don’t use double vowels when writing names (oa, oo)
Not monooxide but monoxide.
Name These
N2O
 NO2
 Cl2O7
 CBr4
 CO2
 BaCl2
Write formulas for these
diphosphorus pentoxide
 tetraiodide nonoxide
 sulfur hexaflouride
 nitrogen trioxide
 Carbon tetrahydride
 phosphorus trifluoride
 aluminum chloride


X. Acids
A) Acids give off H+ ions when dissolved in water—bases give off OH- ions
B) The general formula for an acid is:
HX (X=a monatomic or polyatomic negative ion)
C) Naming Acids
1) When the negative ion (anion) ends in “-ide”, the acid begins with the
prefix “hydro-“ + “stem of anion” + “-ic” + the word “acid”
Examples: Cl-: Chloride HCl: hydrochloric acid
2) When the anion ends in “-ate”, the acid is named with the “stem of the
anion” + “ic” + the word “acid”
Example: NO3-: Nitrate HNO3: Nitric acid
3) When the anion ends in “-ite”, the acid is named with the “stem of the
anion” + “-ous” + the word “acid”
a) Example: SO3-2: Sulfite H2SO3: Sulfurous acid
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Name these
HF
H3P
 H2SO4
 H2SO3
 HCN
 H2CrO4


D) Writing formulas
1) Hydrogen will always be first
2) The name will identify the anion
3) Make the charges cancel out
4) If it starts with “hydro-“ there is no Oxygen only Hydrogen and the
element—HCl
5) If it does not start with “hydro-“ it does contain Oxygen
a) If it ends in “-ic” then the ion would end with “-ate” H2SO4
b) If it ends in “-ous” then the ion would end with “-ite” H2SO3
Write formulas for these
hydroiodic acid
acetic acid
 carbonic acid
 phosphorous acid
 hydrobromic acid


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Writing Formulas

Write the formula for calcium chloride.
Calcium is Ca+2

Chloride is Cl-1

Ca+2 Cl-1 would have a +1 charge.
 Need another Cl-1
 Ca+2 Cl2-1

Write the formulas for these
Lithium sulfide
tin (II) oxide
 tin (IV) oxide
 Magnesium fluoride
 Copper (II) sulfate
 Iron (III) phosphide
 gallium nitrate
 Iron (III) sulfide
Write the formulas for these
Ammonium chloride
ammonium sulfide
 barium nitrate




E) Things to Look For:
1) If cations have (), the number is their charge.
2) If anions end in -ide they are probably off the periodic table
(Monoatomic)
3) If anion ends in -ate or -ite it is polyatomic
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