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Atoms, Molecules, Ions, and
Nomenclature
AP Chemistry
Dr. Daniel Schuerch
Fundamental Laws of Chemistry
•
Law of conservation of mass
–
•
Mass is neither created or destroyed
Law of Definite Proportions
–
•
In different samples of the same chemical
compound, the masses of the elements are always
present in the same proportions
Law of multiple proportions
–
When two elements form more than one compound,
the different masses of one element that combine
with the same mass of the other element are in a
ratio of small whole numbers
Law of definite proportions
•
In the compound sulfur dioxide, a food
preservative, the mass ratio of sulfur to oxygen
is 1:1. An 80-g sample of a compound made
of sulfur and oxygen contains 48 grams of
oxygen. Is the sample sulfur dioxide?
Explain?
Defining the Atom
• Early Models of the Atom
– All matter is made up of atoms
• Atoms are the smallest particle of an element that retains its identity
in chemical reactions
– Democritus’s Atomic Theory (460-370 BC)
• Atoms were indivisible and indestructible
– Agreed with later scientific theory, but did not explain chemical behavior
– Was never tested scientifically
– The modern process of discovery regarding atoms began with
John Dalton (1766-1844)
• Using experimental methods, Dalton transformed Democritus’s
ideas on atoms into a scientific theory
• Studied the ratios in which elements combine in chemical reactions
which led to hypotheses and theories to explain his observations
• The result was Daltons atomic theory
Defining the Atom
Daltons Atomic Theory
1. All elements are composed of tiny indivisible particles
called atoms
2. Atoms of the same element are identical
–
Atoms of any one element are different from those of any other
element
3. Atoms of different elements can physically mix together
or can chemically combine in simple whole-number
ratios 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
Defining the Atom
Sizing up the Atom
• Atoms are very small
– A pure copper coin contains 2.4 x 1023 copper atoms!
– The radii of most atoms fall within the range of 5x10-11
to 2x10-10 m.
– Despite their small size, individual atoms are
observed with instruments such as scanning
tunneling microscopes
http://domino.watson.ibm.com/comm/pr.nsf/pages/rsc.stm.html
http://upload.wikimedia.org
Subatomic Particles
• Three kinds of subatomic particles make
up atoms:
1. Electrons
2. Protons
3. Neutrons
Subatomic Particles
Structure of the Nuclear Atom
– Electrons are negatively charged subatomic particles
discovered by Thomson (1856-1940)
– The result of experiments with cathode rays
• A stream of electrons produced at the negative electrode
(cathode) of a tube containing a gas at low pressure
• An electron carries exactly one unit of negative charge and
has a mass of 9.11 × 10−31 𝑘𝑔
• Has a charge 1.60 × 10−19 𝐶.
Protons and Neutrons
• If cathode rays are electrons given off by atoms, what
remains?
– Atoms are neutrally charged
– Must be a positive charged particles
• Eugen Goldstein noticed rays going the opposite
direction of the cathode ray in a cathode-ray tube
– Determined those rays must be made up of positively charged
particles called protons
• Has a mass of 1.67 × 10−27 𝑘𝑔
– Because atoms are neutrally charged, the positive particles in
atom must equal the negative charges in the same atom
• James Chadwick confirmed the existence of neutrons
– Neutrons have about the same mass of an proton but no charge
• How are the electrons, protons, and neutrons arranged
in an atom?
The Atomic Nucleus
• The atomic Nucleus
– Scientist used to think subatomic
particles were evenly dispersed
throughout the atom
• Plum pudding Model
– Ernest Rutherford disproved the
plum pudding model using his GoldFoil experiment
– Based on the results of his
experiments, Rutherford developed a
new model of the atom
• Proposed the atom is mostly empty
space
• The positive charge and almost all
the mass of an atom are
concentrated in a small space, the
nucleus composed of protons and
neutrons
Distinguishing Among Atoms
Atomic Number
• All atoms are composed of protons, neutrons,
and electrons with protons and neutrons making
up the nucleus and electrons making up the
electron shell
• What makes elements different from each other?
– Elements are different from each other
because they contain different numbers of
protons
– The number of protons an element
contains constitutes its atomic number
Distinguishing Among Atoms
Mass Number
• Most of the mass of an atom is concentrated in
the nucleus of atom
• The mass of an atom is determined by the
number of protons and neutrons in an atom
– The number of protons plus neutrons in an atom is its
mass number
– The number of neutrons in an atom is the difference
between the mass number and atomic number
Number of neutrons = mass number – atomic number
• Do all atoms of the same element have the
same amount of neutrons?
Atomic Symbols
1. What element has an atomic number of 30?
2. If the element from number one has a mass number of 65, how many
protons and neutrons does it have?
3. Uranium-235 has how many neutrons?
4. How many electrons does uranium-235 have?
Distinguishing Among Atoms
Isotopes
• All the atoms of an element have an identical
amount of protons, but they may differ in the
amount of neutrons they contain and thus have
different masses
• Atoms of the same element having different
numbers of neutrons are called isotopes
Distinguishing Among Atoms
Atomic Mass
• The actual mass of atoms are very small and hard to
work with
• It is more useful to compare the relative masses of
atoms using a reference isotope as a standard.
• The reference standard used today is the carbon-12
atom that has 6 protons and 6 neutrons
• Carbon-12 was assigned to have exactly 12 atomic
mass units
– Atomic mass unit = one twelfth the mass of a carbon-12 atom
Distinguishing Among Atoms
Atomic Mass
• Atoms of an element have different isotopes that occur at
different frequencies
• The atomic mass of an element is a weighted average mass of
the atoms in a naturally occurring sample of the element
– A weighted average mass reflects both the mass and the relative
abundance of the isotopes as they occur in nature
– To calculate the atomic mass of an element, multiply the mass of each
isotope by the natural abundance, expressed as a decimal, and add the
products.
𝒂𝒕𝒐𝒎𝒊𝒄 𝒎𝒂𝒔𝒔 = (
% 𝒐𝒄𝒄𝒖𝒓𝒓𝒆𝒏𝒄𝒆 𝒊𝒔𝒐𝒕𝒐𝒑𝒆 𝟏
𝟏𝟎𝟎
× 𝐦𝐚𝐬𝐬 𝐧𝐮𝐦𝐛𝐞𝐫 𝐢𝐨𝐬𝐨𝐭𝐨𝐩𝐞 𝟏)+
% 𝒐𝒄𝒄𝒖𝒓𝒓𝒆𝒏𝒄𝒆 𝒊𝒔𝒐𝒕𝒐𝒑𝒆 𝟐
𝟏𝟎𝟎
× 𝐦𝐚𝐬𝐬 𝐧𝐮𝐦𝐛𝐞𝐫 𝐢𝐨𝐬𝐨𝐭𝐨𝐩𝐞 𝟐 …
The Periodic Table
• A periodic table is an arrangement of elements in which
the elements are separated into groups based on a set
of repeating properties
• Each horizontal row periodic table is called a period
– Atomic number increases left to right
– There are seven periods in the periodic table
• Each vertical column of the periodic table is called a
group, or family
– Elements within a group have similar chemical and physical
properties
•
•
•
•
•
•
•
•
Group 1A is called the alkali metals
Group 2A is called the alkaline earth metals
Group 3A is called the triels
Group 4A is called the tetrels
Group 5A is called the pnictogens
Group 6A is called the chalcogens
Group 7A is called the halogens
Group 8A is called the noble gases
The Periodic Table
• All elements can be divided up into three broad
categories based on their position in the periodic
table
1. Metals (Most of the Elements and occupy the left
side of the table)
2. Metalloids aka Semimetals (Has the least number of
Elements)
3. Nonmetals (Occupy the right side of the table)
Periodic Table
• Three other categories
– Representative Elements (Group A Elements)
– Transition Metals
– Inner Transition Metals aka Rare Earth Metals
Periodic Trends: Atomic Size
Periodic Trends
– The atomic radius is one half of the distance between
the nuclei of two atoms of same element when atoms
are joined
– In general, atomic size increases from top to bottom
within a group and decreases from left to right across
a period
– Why
• Periodic Trends
– The increases of positive charge in the nucleus of the atom
draws in it’s electrons in its highest energy level as you move
left to right across a period and thus, the atomic size decreases
• Group Trends
– As you move down a group, new principle energy levels are
filled with electrons. These filled inner energy levels shield the
electrons in the outer energy levels from the positive charge of
the nucleus and thus atoms are larger down a group in the
periodic table
Periodic Trends
• Ions
– Cations are smaller than the atoms they were
formed from
– Anions are larger than the atoms they were
formed from
Periodic Trends
• Trends in Ionization Energy
– Electrons can move to higher energy levels when atoms absorb
energy
– If there is enough energy, electrons overcome the attraction of
protons in the nucleus and escape
– The energy required to remove an electron from an atom is
called ionization energy
• Measured with when an element is in its gaseous state
• The energy required to remove the first electron from an atom is
called the first ionization energy and produces a cation with a 1+
charge
• The first ionization energy tends to decrease from top to bottom
within a group and increases from left to right across a period
– Atoms with low ionization energy levels tend to lose electrons
easily in chemical reaction becoming cations
– Atoms with high ionization energy tend to gain electrons in
chemical reaction becoming anions
Periodic Trends: Ionization Energy
Periodic Trends
• Trends in Ionization Energy
– Periodic Trends in ionization energy
• Increases from left to right due to increasing
nuclear charge, and constant shielding of the outer
most electrons
– Results in increased attraction of the outer electrons to
the nucleus and greater energy required to remove
outermost electron
– Group Trends in ionization energy
• Decreases from top to bottom in a group due to
increased shielding of outer electron by completely
filled inner energy levels and the outer electrons
further distance from the nucleus
– Results in decreased attraction of the outer electrons to
the nucleus and less energy required to remove
outermost electron
Periodic Trends: Electronegativity
Trends in Electronegativity
• Electronegativity
– When atoms react to form a compound, two kinds of bonds can
result between the atoms called ionic bonds and covalent bonds
– Whether an ionic or covalent bond forms can be predicted by the
electronegativity of the atoms involved in making the compound
• Electronegativity is the ability of an atom of an element to attract
electrons when the atom is in a compound
• Scientist use factors such as ionization energy to calculate values of
electronegativity
– In general, electronegativity values decrease from top to bottom
within a group. For representative elements, the values tend to
increase from left to right across a period
• Metals at the far left of the periodic table have low electronegativity
values
• Nonmetals at the far right of periodic table (excluding noble gases)
have high electronegativity values
• Electronegativity among transition metals are not as regular
Bond Type and Polarity
• The difference in electronegativity
between the two atoms involved in the
bond will determine the most probable
type of bond that will form
– If the difference is:
•
•
•
•
0.0—0.4 the bond is nonpolar covalent
0.4—1.0 the bond is moderately polar covalent
1.0—2.0 the bond is very polar covalent
>2.0 the bond is ionic
Ionic Compounds
• Metals have low electronegativities and lose
electrons in chemical reactions becoming
positively charged cations
• Nonmetals have high electronegativities and
become negatively charged anions when
reacting with metals
• When a metal reacts with a nonmetal, an ionic
compound results
– Ionic compounds are held together by ionic bonds
that result from the attraction of oppositely charged
ions
Molecular Compounds
• Molecular compounds result when atoms
react that do not have great differences in
electronegativities and thus electrons are
shared between the atoms
• When electrons are shared between the
atoms, the result is a covalent bond
• Covalent bonds result from the reaction of
two or more nonmetals
Structural Formulas
• Molecules can be represented in several
ways
– Chemical formulas
H
C2H6O
Cl
– Structural formulas
– Space-filling model
– Ball-and-stick model
H
O
H
C
F
H
Naming Monatomic Ions (Type I)
•
•
Monatomic ions are ions consisting of
only one atom
Monatomic ions are either cations or
anions
– Cations (Charge and Naming)
•
How is the ionic charge of a group 1A, 2A, or 3A
ion determined?
– The ionic charge is numerically equal to the group
number and positive (+)
•
The names of group 1A, 2A, or 3A metals are
the same as the name of the metal, followed
by the word ion or cation
Naming Monatomic ions (Type I)
•
Anions (Charge and Naming)
– How is the ionic charge of a group 5A, 6A,
or 7A ion determined?
•
•
The charge of an ion in groups 5A, 6A, or 7A is
determined by subtracting 8 from the group
number
Anion names start with the stem of the
element and end in ide.
– Fluoride, chloride, bromide, iodide, oxide, sulfide
– Are the negative anions of Fluorine, Chlorine, Bromine,
Iodine, Oxygen, Sulfur
Naming Monatomic Ions (Type II)
• Mostly transition metals
– Often have more than one common
ionic charge
• Two methods are used to name ions of
the transition metals
– The stock system of naming transition metal
cations uses a Roman numeral in parentheses
to indicate the numeric value of the ionic
charge
– An older naming system uses the suffix –ous
to name the cation with the lesser charge, and
the suffix –ic to name the cation with the
greater charge
Common Type II Ions to Memorize
Fe2+
Iron (II) Cation
Fe3+
Iron (III) Cation
Cu+
Copper (I) Cation
Cu2+
Copper (II) Cation
Co2+
Cobalt (II) Cation
Co3+
Cobalt (III) Cation
Sn2+
Tin (II) Cation
Sn4+
Tin (IV) Cation
Pb2+
Lead (II) Cation
Pb4+
Lead (IV) Cation
Hg22+
Mercury (I) Cation
Hg2+
Mercury (II) Cation
Ag+
Silver Ion
Type II cation with only one charge
Zn2+
Zinc Ion
Type II cation with only one charge
Cd2+
Cadmium Ion
Type II cation with only one charge
Examples of stock and classical naming
Symbol
Stock Name
Classical Name
Cu+
Copper(I) ion
Cuprous ion
Cu2+
Copper(II) ion
Cupric ion
Fe2+
Iron(II) ion
Ferrous ion
Fe3+
Iron(III) ion
Ferric ion
Pb2+
Lead(II) ion
Plumbous ion
Pb4+
Lead(IV) ion
Plumbic ion
•
What is a major advantage of the Stock system over the old naming
system?

The stock system gives the actual charge of the ion
Use the periodic table to write the name and formula (including charge)
for each ion formed from the element in the table below.
Element
Fluorine
Calcium
Oxygen
Sodium
Nitrogen
Name
Formula
Use the periodic table to write the name and formula (including charge)
for each ion formed from the element in the table below.
Element
Name
Formula
Fluorine
Fluoride ion
F-
Calcium
Calcium ion
Ca2+
Oxygen
Oxide ion
O2-
Sodium
Sodium Ion
Na+
Nitrogen
Nitride Ion
N3-
Polyatomic Ions
•
What is a polyatomic ion?
–
A polyatomic ion is tightly bound group of
atoms that behaves as a unit and carries a
charge
•
–
–
Examples: NH4+, CO32-, OH-
The names of most polyatomic anions
(containing oxygen) usually end in -ite or ate
What is the difference between the anions
sulfite and sulfate?
–
The sulfite ion has one less oxygen atom than the
sulfate ion
Polyatomic Ions to Memorize
Table 9.3
Polyatomic ions continued
How many atoms make up the oxalate ion
and what is its charge?
– Oxalate C2O42-
Polyatomic ions continued
How many atoms make up the oxalate ion
and what is its charge?
– Oxalate C2O42•
It is made up of 6 atoms (2 carbon atoms and 4
oxygen atoms) and has a charge of 2-.
Binary Ionic compounds
•
The general name for compounds composed of
two elements is binary compounds
–
The name of a binary ionic compound is written with the name of
the cation first followed by the name of the anion
•
When writing the formula for any ionic
compound, the charges of the ions must
balance
•
What are the two methods for writing a balanced
formula?
•
•
Finding the least common multiple of the charges
The crisscross method
Least common multiple
•
What are the formulas for the compounds formed by
the following pairs of ions
•
•
•
•
Fe2+, ClCr3+, O2Na1+, S2-
2 is the least common multiple
6 is the least common multiple
2 is the least common multiple
What are the formulas for these compounds?
•
•
•
Lithium bromide
Cupric nitride
Magnesium chloride
Crisscross Method of Writing Chemical Formulas
•
•
What are the formulas for the compounds formed by the following
pairs of ions
•
Fe2+, Cl1-
•
Cr3+, O2-
•
Na1+, S2-
What are the formulas for these compounds?
•
Lithium bromide
•
Cupric nitride
•
Magnesium chloride
Chemical Names and Formulas
•
How can you tell that cobalt(II) chloride is a
binary ionic compound formed by a transition
metal with more than one ionic charge?
–
•
The name includes a Roman numeral representing
the ionic charge of the transition metal cation
Write the names of these binary compounds
–
PbS
–
MgCl2
–
Al2Se3
Chemical Names and Formulas
•
How can you tell that cobalt(II) chloride is a
binary ionic compound formed by a transition
metal with more than one ionic charge?
–
•
The name includes a Roman numeral representing
the ionic charge of the transition metal cation
Write the names of these binary compounds
–
PbS
Lead(II) Sulfide
–
MgCl2
Magnesium Chloride
–
Al2Se3
Aluminum Selenide
Compounds with Polyatomic ions
•
What is an polyatomic ion?
–
•
How do you write the formula for a compound
containing a polyatomic ion
–
•
An ion containing more than one element
Write the symbol for the cation followed by the
formula for the polyatomic ion and balance the
charges
Why are parentheses used to write the formula
Al(OH)3
–
The parentheses indicate how many polyatomic
ions are needed in the formula
•
Complete the table for these ionic compounds
containing polyatomic ions
Cation
Anion
NH4+
S2-
Fe3+
CO32-
Ag+
NO3-
K+
CN-
Name
Formula
•
Complete the table for these ionic compounds
containing polyatomic ions
Cation
Anion
Name
Formula
NH4+
S2-
(NH4)2S
Fe3+
CO32-
Ag+
NO3-
K+
CN-
Ammonium
Sulfide
Iron III
Carbonate
Silver I
Nitrate
Potassium
Cyanide
Fe2(CO3)3
AgNO3
KCN
Naming Binary Molecular Compounds
•
•
•
•
Two nonmetallic elements form binary
molecular compounds
Two nonmetallic elements often can bind
in more than one way
Prefixes are used to distinguish between
different molecular compounds that
contain the same element
mono-, 1; di-, 2; tri-, 3; tetra-, 4; penta-,5;
hexa-, 6; hepta-, 7; octa-, 8; nona-, 9
Naming Binary Covalent Compounds (Type III)
• Confirm the compound is a binary molecular compound
(two nonmetals)
• The name must identify the elements in the molecule
and indicate the number of each atom
• Name the element in the order listed in the formula
• Use prefixes to indicate the number of each kind of atom
• Omit the prefix mono- when the formula contains only
one atom of the first element in the name
• The suffix of the name of the second element is -ide
Naming Binary Molecular Compounds
•
What are the names of the following compounds?
•
•
•
•
•
BF3
N2O4
P4S7
What are the formulas for the following compounds?
•
Use the prefixes in the name to tell you the subscript of each element
in the formula. Then write the correct symbols for the two elements
with the appropriate subscripts
•
•
•
•
Carbon tetrabromide
Nitrogen triiodide
Iodine monochloride
tetraiodine nonaoxide
Notice, the less electronegative atom is always first in the
chemical formula and name
Naming and Writing Formulas for
Acids and Bases
•
Acids produce hydrogen ions when
dissolved in water
•
When naming acids, you can consider
them to be combinations of anions
connected to as many hydrogen ions as
are necessary to create an electrically
neutral compound
Rules for Naming Acids
• When the name of anion (X) ends in –ide, the
acid name begins with the prefix hydro-. The
stem of the anion has the suffix –ic and followed
by the word acid
– HCl = hydrochloric acid is formed from the anion
chloride
• When the anion name ends in –ite, the acid
name is the stem of the anion with the suffix –
ous, followed by the word acid.
– H2SO3 = sulfurous acid is formed from the anion
sulfite
• When the anion name ends in –ate, the acid
name is the stem of the anion with the suffix –ic
followed by the word acid
– HNO3 = nitric acid is formed the anion nitrate
5.
Writing Formulas Acids
Acid Name
Acetic acid
Carbonic acid
Hydrochloric acid
Nitric acid
Phosphoric acid
Sulfuric acid
Formula
Anion Name
Writing Formulas Acids
Acid Name
Formula
Anion Name
Acetic acid
HC2H3O2
acetate
Carbonic acid
H2CO3
carbonate
Hydrochloric acid
HCl
Chloride
Nitric acid
HNO3
Nitrate
Phosphoric acid
H3PO4
Phosphate
Sulfuric acid
H2SO4
Sulfate
Names and formulas for Bases
•
Another group of compounds is the
bases
– A base is a compound that produces
hydroxide ions when dissolved in water
– Bases are named in the same way as other
ionic compounds—the name of the cation is
followed by the name of the anion
•
Example (Sodium hydroxide)