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9/13/2015
CHAPTER 2
ATOMS, MOLECULES,
AND IONS
Prof. Dr. Nizam M. El-Ashgar
THE ATOMIC THEORY OF MATTER
Historically:
Greek Philosophers: matter can be subdivided into
fundamental particles.
Democritus (460–370 BC) and other early Greek
philosophers described the material world as made up
of tiny indivisible particles they called atomos,
meaning “indivisible or uncuttable.”
The “atomic” view of matter faded for many centuries.
The notion of atoms reemerged in Europe during
the seventeenth century.
That theory came from the work of John Dalton
during the period from 1803 to 1807.
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THE ATOMIC THEORY
John Dalton’s 4 Postulates:
1. Each element is composed of extremely small
particles called atoms.
2. All atoms of an element are identical to each
other, but different than atoms of other
elements.
3. Atoms of one element cannot be changed into
atoms of different elements by chemical
reactions; atoms are neither created or
destroyed in reactions.
4. Compounds are formed when atoms of different
elements combine in a definite ratio.
I) LAW OF CONSERVATION OF MASS
Mass is neither created nor destroyed in
chemical reactions.
3.25 g + 3.32 g = 6.57 g
Hg(NO3)2(aq) + 2KI(aq)
HgI2(s) + 2KNO3(aq)
4.55 g + 2.02 g = 6.57 g
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II) THE LAW OF DEFINITE PROPORTIONS
Different samples of a pure chemical substance always
contain the same elements with constant proportion
of elements by mass.
By mass, water (H2O) is:
88.8 % oxygen
11.2 % hydrogen
III) LAW OF MULTIPLE PROPORTIONS:
If 2 elements combine to form more than one compound,
and if one element presents in fixed mass the masses of
the second element exist in small whole number ratios.
Sample 1
H O
1 g 16 g
Sample 2
H O
1g 8g
H2O
H2O2
mall whole number ratios: 2/1, 1/2, 3/2, 2/3, 3/4,,,,,,,,
This follows from the postulate that individual atoms
enter into chemical combination.
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THE DISCOVERY OF ATOMIC STRUCTURE
J. J. Thomson (1898—1903)
Postulated the existence of electrons using
cathode-ray tubes.
Determined the charge-to-mass ratio of an
electron.
The atom must also contain positive particles
that balance exactly the negative charge
carried by particles that we now call
electrons.
7
Thomson’s Experiment
Voltage source
-
+
Cathode rays = radiation produced when high voltage is
applied across the tube.
The voltage causes negative particles to move from the
negative electrode (cathode) to the positive electrode
(anode).
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If no Magnetic and Electric field applied: The cathode rays
pass in straight direction>
Applying EF: The path of the cathode rays altered toward
the +ve plate.
Applying MF: The cathode rays can be deflected in opposite
direction.
By balancing both MF and EF: The rays pass in straight
direction.
Deflection of electron depends on three factors:
1) Strength of electric or magnetic field. 2) Size of negative
charge. 3)Mass of the electron.
In 1897 Thomson determined the charge-to-mass ratio of
an electron. Charge/Mass = 1.76 108 C/g.
C is a symbol for coulomb (SI of electric charge).
Either the charge or the mass of an electron would yield the
other.
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Oil Drop Experiment (Millikan, 1868–1953):
Applied a voltage to oppose the downward fall of
charged drops and suspend them.
Voltage on plates place
1.602176 x 10-19 C of
charge on each oil drop
= electron charge.
− 19
- 28
1.60× 10 C
electronmass =
=
9.10×
10
g
8
1.76× 10 C / g
RADIOACTIVITY
Radioactivity is the spontaneous emission
by an atom.
It was first observed by Henri Becquerel.
Marie and Pierre Curie also studied it.
Method:
A radioactive substance is placed in a
containing a small hole so that a beam of
emitted from the shield.
The radiation is passed between two
charged plates and detected.
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of radiation
lead shield
radiation is
electrically
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Three types of radiation were discovered by Ernest
Rutherford:
α particles: helium nucleus (+2 charge, large mass)
β particles: high speed electron (-ve)
γ rays: high energy light, similar to X-rays (no
charge).
The mass of an α -particle is 7300 times that of the
electron
(similar to X-rays)
THE ATOM, CIRCA 1900
Early model: the “plum
pudding” model.
Thompson: proposed a
positive
sphere
of
matter with negative
electrons imbedded in
it.
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RUTHERFORD’S EXPERIMENT
Used uranium to produce alpha particles
Aimed alpha particles at gold foil by drilling
hole in lead block
Since the mass is evenly distributed in gold
atoms alpha particles should go straight
through.
Used gold foil because it could be made atoms
thin
+
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Explanation:
Atom is mostly empty.
Small dense, positive piece at
center in the nucleus.
Alpha particles are deflected
by it if they get close enough.
Proton (p) has opposite (+)
charge of electron (-)
Mass of p is 1840 times the
mass of e- (1.67 x 10-24 g).
Protons were discovered by
Rutherford in 1919.
Neutrons were discovered by
James Chadwick in 1932.
+
SUBATOMIC PARTICLES
Protons and electrons are the only particles that
have a charge.
Protons and neutrons have essentially the same
mass.
The mass of an electron is so small we ignore it.
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Particle
Mass
(g)
Electron (e-)
9.1 x 10-28
Proton (p+)
1.67 x 10-24
Neutron (n)
1.67 x 10-24
Mass (amu)
Charge
(Coulombs)
Charge
(units)
5.486 x 10-4 -1.6 x 10-19
-1
1.0073
+1.6 x 10-19
+1
1.0087
0
0
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The angstrom is a convenient non-SI unit of length used to
denote atomic dimensions.
1 Å = 10 x 10 –10 m
Density of nucleus: 1013–1014 g/cm3
A matchbox full of material of such density would weigh over
2.5 billion tons!
SAMPLE EXERCISE 2.1
The diameter of a US dime is 17.9 mm, and the diameter of a
silver atom is 2.88 Å . How many silver atoms could be
arranged side by side across the diameter of a dime?
Conversion Factors:
1 Å ↔ 10 -10 m
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1 Ag atom ↔ 2.88 Å
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PRACTICE EXERCISE
The diameter of a carbon atom is 1.54 Å. (a) Express this diameter
in picometers. (b) How many carbon atoms could be aligned side by
side in a straight line across the width of a pencil line that is 0.20
mm wide?
Answer: (a) 154 pm, (b) 1.3 ×106C atoms
ATOMIC NUMBERS, MASS NUMBERS AND ISOTOPES
Atomic Number (Z): Number of protons in an
atom’s nucleus. Equivalent to the number of
electrons around an atom’s nucleus
Mass Number (A): The sum of the number of
protons and the number of neutrons in an atom’s
nucleus
Isotope: Atoms with identical atomic numbers but
different mass numbers
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All atoms of the same element have the same
number of protons:
The atomic number (Z)
ISOTOPES
Isotopes are atoms of the same element with different
masses.
Isotopes have different numbers of neutrons.
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11
6C
12
6C
13
6C
C−11
C − 12
C − 13
14
6C
C − 14
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ISOTOPES
carbon-12
mass number
12
C
6
6 protons
6 electrons
6 neutrons
14
6 protons
6 electrons
8 neutrons
atomic number
carbon-14
mass number
6
atomic number
C
SAMPLE EXERCISE 2.2
How many protons, neutrons, and electrons are in (a)
an atom of 197Au:
From PT: Z = 79
P= 79
e = 79
n = 197-79 =118
(b) an atom of strontium-90?
From PT: Z = 38
P= 38
e = 38
n = 90-38 = 52
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PRACTICE EXERCISE:
How many protons, neutrons, and electrons are in (a) a
138Ba atom, (b) an atom of phosphorus-31?
Answer: (a) 56 protons, 56 electrons, and 82 neutrons;
(b) 15 protons, 15 electrons, and 16 neutrons.
ATOMIC WEIGHTS:
The Atomic Mass Scale : Early: It was related to H
mass.
Consider 100 g of water:
• Upon decomposition 11.1 g of hydrogen and 88.9 g of
oxygen are produced.
• The mass ratio of O to H in water is 88.9/11.1 = 8.
• Therefore, the mass of O is 2 x 8 = 16 times the mass of
H.
• If H has a mass of 1, then O has a relative mass of 16.
We can measure atomic masses using a mass
spectrometer.
1H mass = 1.6735 x 10–24 g
16O mass = 2.6560 x10–23 g.
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Atomic mass units (amu): are convenient units to use
when dealing with extremely small masses of individual
atoms.
The amu is 1/12 the mass of one 12C atom.
1 amu = 1.66054 x 10–24 g
1g
= 6.02214 x1023 amu
By definition, the mass of 12C is exactly 12 amu.
Now all the present atoms are assigned according to C12 isotopes
A 24Mg atom has a mass approximately twice that of the
12C atom, so its mass is 24 u.
A 4He atom has a mass approximately 1/3 that of the
12C atom, so its mass is 4 u.
1H atom has a mass of 1.0078 amu.
16O atom has a mass of 15.9949 amu.
MASS SPECTROPHOTOMETER
Atomic
and
molecular
masses
can be measured
with great accuracy
with
a
mass
spectrometer.
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ATOMIC WEIGHT
Most elements occur in nature as mixtures of isotopes.
We can determine the average atomic mass of an
element, usually called the element’s atomic weight.
We average the masses of isotopes to give average
atomic masses.
Example: Naturally occurring C consists of:
12
6
C
13
6
C
Atomic mass:
12.0 amu
13.00335 amu
Abundance:
98.93 %
1.07 %
• The average mass of C is:
(0.9893)(12 amu) + (0.0107)(13.00335 amu) = 12.01 amu.
Atomic weights are listed on the periodic table
Because in the real world we use large amounts of
atoms and molecules, we use average masses in
calculations.
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HW: Chlorine has two naturally occurring isotopes: Cl35 with an isotopic mass of 34.969 amu, and Cl-37 with
an isotopic mass of 36.966 amu. The atomic weight of
chlorine is 35.5 . What is the relative abundances of the
two isotopes?
37
17 Cl
35
17 Cl
Atomic mass
Abundance:
34.969 amu
X
36.966 amu
100-X
ELEMENTS IN PERIODIC TABLES
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SAMPLE EXERCISE 2.3
Magnesium has three isotopes, with mass numbers 24, 25, and
26.
Write the complete chemical symbol (superscript and
subscript) for each of them.
24
Mg
12
25
Mg
12
26
12
Mg
(b) How many neutrons are in an atom of each isotope?
The numbers of neutrons in an atom of each isotope are
therefore 12, 13, and 14, respectively.
PRACTICE EXERCISE
Give the complete chemical symbol for the atom that
contains 82 protons, 82 electrons, and 126 neutrons.
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PRACTICE EXERCISE
Three isotopes of silicon occur in nature: 28Si (92.23%), which has an
atomic mass of 27.97693 amu; 29Si (4.68%), which has an atomic
mass of 28.97649 amu; and 30Si (3.09%), which has an atomic mass
of 29.97377 amu. Calculate the atomic weight of silicon.
Answer: 28.09 amu
PERIODIC TABLE
It is a systematic catalog of the elements.
Elements are arranged in order of atomic number.
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Periods: The rows on the periodic chart.
Columns: The groups on the periodic chart.
Elements in the same group have similar chemical
properties.
Periods (Raws)
Groups (columns)
GENERAL ARRANGEMENT
Metallic elements, or metals:
Are located on the left-hand side of the periodic table (most
of the elements are metals).
Metals tend to be: malleable, ductile, and lustrous and are
good thermal and electrical conductors.
Nonmetallic elements, or nonmetals:
Are located in the top right-hand side of the periodic table.
Nonmetals tend to be brittle as solids, dull in appearance,
and do not conduct heat or electricity well.
Metalloids: Are located at the interface between the
metals and nonmetals.
Elements with properties similar to both metals and
nonmetals
These include the elements B, Si, Ge, As, Sb and Te
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NAMES OF SOME GROUPS IN THE PERIODIC TABLE
These five groups are known by their names.
IMPORTANT FAMILIES OF ELEMENTS
Representative Elements:
The elements in the A groups (1,2, 13-18).
Transition Metals:
The elements in B groups (3-12).
Inner Transition Metals:
The two rows of metals (lanthanides and actinides)
set at the bottom of the periodic table.
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SAMPLE EXERCISE 2.5
Which two of the following elements would you expect to
show the greatest similarity in chemical and physical
properties: B, Ca, F, He, Mg, P?
Solution:
Ca and Mg should be most alike because they are in the
same group (2A, the alkaline earth metals).
PRACTICE EXERCISE
Locate Na (sodium) and Br (bromine) on the periodic table. Give the
atomic number of each, and label each a metal, metalloid, or
nonmetal.
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2.6 MOLECULES AND MOLECULAR COMPOUNDS
Only Nobel gases are found in nature as isolated atoms.
Most matter is composed of molecules or ions.
A molecule: consists of two or more atoms bound tightly
together.
Types:
1) Molecular Elements: Same two or more atoms combined
with each others.
Diatomic molecules: made up of two same kind of atoms.
Examples: N2, O2, Cl2
Polyatomic molecules: made up of more than two atoms of
same atoms.
Examples: S8, P4, O3
Allotropes: Different forms of an element, which have
different chemical formulas
Allotropes differ in their chemical and physical properties.
Examples: Ozone (O3) and “normal” Oxygen (O2)
C (diamond) and C (Graphite)
2) Molecular compounds: Composed of molecules which
contain more than one type of nonmetallic atoms.
Examples:
- Diatomic: HCl, CO.
- Polyatomic: H2SO4, HCl, CO2, H2O2, NH3
Molecules
H2
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H2O
NH3
CH4
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CHEMICAL FORMULA
Each molecule has a chemical formula.
The chemical formula indicates :
1. which atoms are found in the molecule.
2. what proportion they are found.
Types:
1) Molecular formulas (MF):
These formulas give the actual numbers and types
of atoms in a molecule.
Examples: H2O, CO2, CO, CH4, H2O2, O2, O3, and C2H4.
2) Empirical formulas (EF):
These formulas give the relative (simplest whole
numbers ratio) numbers and types of atoms in a
molecule
Examples: H2O, CO2, CO, CH4, HO, CH2.
MF = n (EF)
Example:
C2H6 = 2 (CH3)
3) Structural Formulas (SF):
Structural formulas show the order in which atoms are
bonded.
Examples: Methane
Ethane
H
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H
H
C
C
H
H
H
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Different types of formulae of some compounds
Compound
Empirical
formula
Molecular
formula
Structural
formula
Carbon
dioxide
Water
CO2
CO2
O = C =O
H2O
H2O
Methane
CH4
CH4
Glucose
CH2O
C6H12O6
O
H
H
H
H
C
H
H
OH
O H
H
H
OH
H
H
OH
HO
Sodium
fluoride
NaF
Not
applicable
OH
Na+F-
PICTURING MOLECULES
Molecules occupy three-dimensional space.
However, we often represent them in two dimensions by
SF which usually does not depict the actual geometry
of the molecule, that is, the actual angles at which
atoms are joined together.
Perspective drawings: use dashed lines and wedges to
represent bonds receding and emerging from the plane
of the paper.
Molecular Modeling:
Ball-and-stick models : show atoms as contracted
spheres and the bonds as sticks. • The angles in the
ball-and-stick model are accurate.
Space-filling
models:
give
an
accurate
representation of the 3-D shape of the molecule.
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SAMPLE EXERCISE 2.6
Write the empirical formulas for the following molecules:
(a) glucose, whose molecular formula is C6H12O6.
Solution:
Divide by 6 so EF is CH2O.
(b) nitrous oxide, a substance used as an anesthetic and
commonly called laughing gas, whose molecular formula is
N2O.
Solution: EF is N2O
Practice Exercise:
Give the empirical formula for the substance called
diborane, whose molecular formula is B2H6.
Answer: BH3
IONS AND IONIC COMPOUNDS
Ions: Species pear charge (-ve or +ve) and formed by
adding or removing electrons. •
Cations: Formed when an atom or molecule
loses electrons and becomes positively charged.
Metals tend to form Cations by losing electrons.
Anions: Formed when an atom or molecule gains
electrons and becomes negatively charged.
Nonmetals tend to form Anions by gaining
electrons.
Generally atoms gain or lose electrons to attain the
noble gas elements (stable uncreative elements).
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MONOATOMIC IONS
Anions or Cations of only one atom.
Cations of Representative elements have a (+ve)
charge = group number. (GI A =1+, G2A=2+, Al=3+)
Cations formed from metal atoms have the same name as the metal:
Na+ sodium ion
Zn2+ zinc ion
Al3+ aluminum ion
If a metal can form cations with different charges, the positive
charge is indicated by a Roman numeral in parentheses
following the name of the metal:
Cation
Cu+
Cu2+
Stock Name Old Name
copper(I) ion Cuprous ion
copper(II) ion Cupric ion
Cation
Fe2+
Fe3+
Stock Name
Iron(II) ion
Iron(III) ion
Old Name
Ferrous ion
Ferric ion
Cations formed from nonmetal atoms have names that end in -ium:
NH4+ ammonium ion
H3O+
hydronium ion
Anions of Representative elements have a (–ve)
charge = 8- group number.
The names of monatomic anions are formed by replacing
the ending of the name of the element with -ide:
H- hydride ion
O2- oxide ion
N3- nitride ion
A few polyatomic anions also have names ending in -ide:
OH- hydroxide ion CN- cyanide ion
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O2- peroxide ion
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PREDICTING THE CHARGE OF IONS
Cations: are positive and are formed by elements on the
left side of the periodic chart.
Anions: are negative and are formed by elements on the
right side of the periodic chart.
POLYATOMIC IONS
When molecules lose electrons, polyatomic ions are
formed:
٢٩
NH4+
ammonium
SO42-
sulfate
CO32-
carbonate
SO32-
sulfite
HCO3-
bicarbonate
NO3-
nitrate
ClO3-
chlorate
NO2-
nitrite
Cr2O72-
dichromate
SCN-
thiocyanate
CrO42-
chromate
OH-
hydroxide
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SAMPLE EXERCISE 2.7
Give the chemical symbol, including mass number, for each of
the following ions:
(a) The ion with 22 protons, 26 neutrons, and 19
electrons.
48
3+
22 Ti
(b) The ion of sulfur that has 16 neutrons and 18
electrons.
32
16
S2-
Practice Exercise:
How many protons, neutrons, and electrons does the 79Se2– ion
possess?
Answer: 34 protons, 45 neutrons, and 36 electrons
SAMPLE EXERCISE 2.8
Predict the charge expected for the most stable ion of
barium and for the most stable ion of oxygen.
Solution:
Ba → 2+
Symbol: Ba2+cation.
O → 2Symbol: O2-cation.
Practice Exercise:
Predict the charge expected for the most stable ion of :
(a) aluminum and (b) fluorine.
Answer: (a) 3+; (b) 1–
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IONIC COMPOUNDS
An ionic compound is composed of cations and anions joined
to form a neutral species.
Ionic compounds generally form from the combination of
metals with nonmetals.
In ionic compounds each cation is surrounded by several
anions and vice versa.
Crystal
WRITING FORMULAS
OF IONIC
COMPOUNDS
The formula of an ionic compound is an empirical
formula that uses the smallest whole number subscripts
to express the relative numbers of ions.
The relative numbers of ions in the empirical formula
balances the charges to zero.
Example:
If these subscripts are not in the lowest whole-number
ratio, divide them by the greatest common factor.
Al 3+ + PO43- = Al3(PO4)3 simplified to AlPO4
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SAMPLE EXERCISE 2.9
Which of the following compounds would you expect to be
ionic:
N2O:
Molecular
Na2O:
Ionic.
CaCl2:
Ionic
SF4:
Molecular
Practice Exercise:
Which of the following compounds are molecular: CBr4,
FeS, P4O6, PbF2?
Answer: CBr4 and P4O6
SAMPLE EXERCISE 2.10
What are the empirical formulas of the compounds
formed by:
(a) Al3+ and Cl– ions:
AlCl3
(b) Al3+and O2– ions:
Al2O3
(c) Mg2+and NO3–ions:
Mg(NO3)2
Practice Exercise:
Write the empirical formulas for the compounds formed by the
following ions:
(a) Na+and PO43–, (b) Zn2+and SO42–, (c) Fe3+and CO32–.
Answer: a) Na3PO4 , (b) ZnSO4, (c) Fe2(CO3)3
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COMMON CATIONS
COMMON ANIONS
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CHEMISTRY AND LIFE:
• Of the known elements, only about 29 are required for life.
• Water accounts for at least 70% of the mass of most cells.
• More than 97% of the mass of most organisms comprises just six elements
(O, C, H, N, P and S). These are the most important elements for life
• Carbon is the most common element in the solid components of cells.
• The next most important ions are Na+, Mg2+, K+, Ca2+, and Cl– .
• The other required 18 elements are only needed in trace amounts (green);
they are trace elements.
NAMING OF INORGANIC
(NOMENCLATURE)
Inorganic Compounds:
1- Ionic
2- Molecular
COMPOUNDS
3- Acids and bases
Naming Guide of Ionic Compounds:
Write the name of the cation (same English name of
neutral element) .
If the anion is an element: change its ending to -ide;
If the anion is a polyatomic ion: simply write the name
of the polyatomic ion.
If the cation can have more than one possible charge:
write the charge as a Roman numeral in parentheses.
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NAMING IONIC COMPOUNDS
Binary Ionic Compounds:
1) With Main Groups Metals: G 1A, 2A, and Al
NaCl:
Sodium Chloride
MgO:
Magnesium Oxide
Al2O3:
Aluminum Oxide
2) With Transition and Post Transition Metals: Use Roman
numerals in parentheses to indicate the charge on metals
that have more than one cation.
Fe2O3 :
Iron(III) Oxide
SnCl2 :
Tin(II) Chloride
PbF2 :
Lead(II) Fluoride
Some transition metals form more than one cation
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PATTERNS IN OXYANION NOMENCLATURE
Oxyanions: Polyatomic anions containing oxygen have names
ending in either -ate or -ite .
If two oxyanions involving the same element:
The one with fewer oxygens ends in -ite.
NO2− : nitrite
SO32− : sulfite
The one with more oxygens ends in -ate.
SO42− : sulfate
NO3− : nitrate
For CO32-only one anion: carbonate
For halogens 4 anions present: (Cl as an example)
ClO4ClO3ClO2ClO-
perchlorate ion
chlorate ion
chlorite ion
hypochlorite ion
(one more O atom than chlorate)
(one O atom fewer than chlorate)
(one O atom fewer than chlorite)
Polyatomic anions containing oxygen with additional
hydrogens:
Anions from diprotic acids (H2CO3, H2SO3, H2SO4)
CO32–
carbonate anion.
–
HCO3
hydrogen carbonate (or bicarbonate) anion.
HSO3hydrogen sulfite (or bisulfite) anion
2SO3
sulfite anion.
HSO4
hydrogen sulfate (or bisulfate) anion
SO42sulfate anion.
Anions from polyprotic acids (H3PO4 )
PO43–
phosphate ion.
H2PO4–
dihydrogen phosphate anion.
2–
HPO4
hydrogen phosphate anion.
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ACID NOMENCLATURE
Acids containing anions whose names end in -ide :
Are named by changing the –ide ending to -ic, adding the
prefix hydro- to this anion name, and then following with
the word acid:
FClBrIS2Se2CN-
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(fluoride)
(chloride)
(bromide)
(iodide)
(sulfide)
(chloride)
(cyanide)
HCl(aq) (hydrochloric acid)
HF (aq) (hydrofluric acid)
HBr (aq) (hydrobromic acid)
HI (aq) (hydroiodic acid)
H2S( aq) (hydrosulfuric acid)
HSe(aq ) (hydroselenic acid)
HCN(aq ) (hydrocynic acid)
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Acids containing anions whose names end in -ate or –ite:
are named by changing -ate to -ic and -ite to -ous and then
adding the word acid.
ClO4- (perchlorate)
ClO3- (chlorate)
ClO2- (chlorite)
ClO- (hypochlorite)
SO32- (sulfite)
SO42- (sulfate)
PO43- (phosphate)
NO32- (nitrate)
NO22- (nitrite)
HClO4
HClO3
HClO2
HClO
H2SO3
H2SO4
H3PO4
HNO3
HNO2
(perchloric acid)
(chloric acid)
(chlorous acid)
(hypochlorous acid)
(sulforous acid)
(sulfuric acid)
(phosphoric acid)
(nitric acid)
(nitrous acid)
NOMENCLATURE OF BINARY MOLECULAR COMPOUNDS
The less electronegative
atom is usually listed first.
A prefix is used to denote
the number of atoms of each
element in the compound
(mono- is not used on the
first
element
listed,
however) .
The ending on the more
electronegative element is
changed to -ide.
Successive vowels are often
elided into one.
e.g. a and o→ o
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9/13/2015
EXERCISE
Cl2O
NF3
N2O4
P4S10
CO
CO2
P2O5
S2Cl4
NO2
N2O5
HCl(g)
HI(l)
N2Cl4
dichlorine monoxide
nitrogen trifluoride
dinitrogen tetroxide
tetraphosphorus decasulfide
carbon monoxide
carbon dioxide
diphosphorus pentoxide
disulfur tetrachloride
nitrogen dioxide
dinitrogen penoxide
hydrogen chloride gas
hydrogen iodide liquid
dinitrogen tetrachloride
A base can be defined as a substance that yields hydroxide
ions (OH-) when dissolved in water.
٣٩
NaOH
sodium hydroxide
KOH
potassium hydroxide
Ba(OH)2
barium hydroxide
9/13/2015
NAMING EXERCISE
Al2(S2O3)3
P4O10
Cu(NO2)2
NaMnO4
CS2
Fe2(CrO4)3
KCl
MgBr2
CaO
CuBr
FeS
PbO2
SF6
N2 O 4
Aluminum thiosulfate
Tetraphosphorous decaoxide
Copper(II) nitrite
Sodium permanganate
Carbon disulfide
Iron(III) chromate
Potassium chloride
Magnesium bromide
Calcium oxide
Copper(I) bromide
Iron(II) sulfide
Lead(IV) oxide
Sulfur hexafluoride
Dinitrogen tetroxide
NOMENCLATURE OF ORGANIC COMPOUNDS
Organic chemistry is the study of carbon.
Organic chemistry has its own system of nomenclature.
The simplest hydrocarbons (compounds containing only
carbon and hydrogen) are alkanes.
The first part of the names above correspond to the
number of carbons (meth- = 1, eth- = 2, prop- = 3, etc.).
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9/13/2015
When a hydrogen in an alkane is replaced with
something else (a functional group, like -OH in the
compounds above), the name is derived from the name
of the alkane.
The ending denotes the type of compound.
An alcohol ends in -ol.
HOME WORK
1, 3, 4,6, 10, 14, 19, 21, 24, 28, 32, 34, 37, 40, 44, 48, 53, 55,
59, 62, 67, 72, 75, 78, 84, 88, 94, 98, 104
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‫‪9/13/2015‬‬
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