Chapter 2: Atoms, Molecules, and Ions (2.

1) Atoms and the atomic theory

• (2.1) Atoms and the atomic theory Fundamental Laws of Matter
• (2.2) Components of the atom • Law of conservation of mass
• (2.3) Quantitative properties of the atom • There is no detectable change in mass in an
ordinary chemical reaction. The mass of substances
• (2.4) Introduction to the periodic table produced by a chemical reaction is always equal to
• (2.5) Molecules and ions the mass of the reacting substances. After the reaction, atoms are unchanged.
No new atoms are formed. The number of
• (2.6) Formulas of ionic compounds each atom remains the same.

• (2.7) Names of compounds • Law of constant composition

• A compound always contains the same
elements in the same proportions by mass
Building Blocks of Chemistry • Water always contains 8 g of oxygen for every 1 g of hydrogen.
• Carbon dioxide always contains 2.7 g of oxygen for every 1 g
• Help describe the structure of matter of carbon.

• Atoms – made of electrons, protons, and neutrons • Law of multiple proportions

• The masses of one element that combine with a fixed mass of the second
• Molecules – building blocks of elements and compounds element are in a ratio of small whole numbers
• CO and CO2
• Ions – species of opposite charge present on all ionic compounds • H2O and H2O2
Different combinations form in different compounds.

(2.1) Atoms and the atomic theory (2.2) Main Components of the Atom: Electrons, Protons and Neutrons
Atomic Theory – developed by John Dalton The atom contains: Proton, Neutron and Electron
1. Elements are made of tiny particles called atoms.
• Electrons – located outside the nucleus, possess
2. All atoms of a given element are identical. a unit negative charge (-1).
– accounts for the volume of the atom
3. The atoms of a given element are Atoms of different elements have
different from those of any other element. different properties (e.g. mass). • Nucleus – tiny massive center of the atom
– accounts for the mass of the atom
– consist mainly of
4. Atoms of one element can combine with atoms of
other elements to form compounds. A given • Proton, which possesses a unit
compound always has the same relative positive charge (+1) equal in
Different combinations form in different compounds.
numbers and types of atoms. magnitude to that of the electron
• Neutron, an uncharged particle which
has a mass slightly greater than that
5. Atoms are indivisible in chemical processes. of the proton
Atoms are not created or destroyed in chemical
reactions. A chemical reaction simply changes
the way the atoms are grouped together. It
involves the combination, separation or
After the reaction, atoms are unchanged.
rearrangement of atoms. No new atoms are formed. The number of
each atom remains the same.

An atom is the smallest reacting particle of an element.

(2.2) Main Components of the Atom: Electrons, Protons and Neutrons (2.3) Quantitative properties of the atom
Determination of the number of protons, number of electrons • Atomic number (Z) is the number of protons in the nucleus of an atom
and nuclear charge from the periodic table. • Neutral atoms have the same number of protons as well as electrons

number of protons = atomic number • Mass number (A) of an atom is the sum of the protons and neutrons in the
nuclear charge = + atomic number nucleus. A = number of protons + number of neutrons
Atomic number
number of electrons:
if atom is neutral or has no charge: Element symbol Isotopes
• Isotopes are atoms with the same number of protons (same element) but
number of e = atomic number Atomic mass
different numbers of neutrons.
if atom has a charge:
number of e = atomic no.  charge

Example: K Example: K+
Atomic number = 19 Atomic number = 19 • Isotopes of an element have different masses.
Number of protons = 19 Number of protons = 19
• Isotopes are identified by their mass numbers.
Nuclear charge = + 19 Nuclear charge = + 19
mass number = number of protons + number of neutrons
Number of electrons = 19 Number of electrons = 19  (+1) = 18
• In nature most elements contain mixtures of isotopes.
Atomic charge = 0 Atomic charge = + 1
• Show almost identical chemical properties; chemistry of atom is due to its electrons.

(2.2) Main Components of the Atom: Electrons, Protons and Neutrons (2.3) Quantitative properties of the atom: Isotopes
Isotopic Symbol
DRILL
Isotopic symbol number of protons = Z
A
A = Mass number
• What is the atomic number of boron, B? 5 X = X-A Z = Atomic number
nuclear charge = +Z
Z number of neutrons = A – Z
• What is the atomic mass of silicon, Si? 28.09 amu mass number = number of protons + number of neutrons
• How many protons does a chlorine atom have? 17
• How many electrons does a neutral neon atom have? 10 Two Isotopes of Sodium:
• Will an atom with 6 protons, 6 neutrons, and 6 electrons be
electrically neutral? YES
• Will an atom with 27 protons, 32 neutrons, and 27 electrons
be electrically neutral? YES
• Will a Na atom with 10 electrons be electrically neutral? NO
It has a charge of +1
 A number of protons = Z (2.3) Quantitative properties of the atom: Atomic Mass
Isotopic Symbol: Z
X = X-A
nuclear charge = +Z
Atomic Mass Unit (amu or u)
Z = atomic number number of neutrons = A – Z
• An atomic mass unit (amu or u) is a unit used to express the relative masses of
A = mass number = number of protons + number of neutrons atoms. One atomic mass unit is equal to 1/12 the mass of a carbon-12 atom.
DRILL – write the isotopic symbol
• C-12 atom is the standard and is assigned a mass of 12 u.
235U 88Sr2+ 75As3-
54Fe U-235 Sr-882+ As-753-
Fe-54 26 92 38 33
• An atom with a mass equal to 1/12 the mass of a carbon-12 atom would have a
relative mass of 1 amu.
DRILL: Complete the Following Table.
No. of • An atom with a mass equal to 1/6 the mass of a carbon-12 atom would have a
Atomic Mass No. of No. of Nuclear
Neutrons Atom/Ion relative mass of 2 amu.
No. No. Protons Electrons Charge
(Z) (A) (Z) (Z  charge) (A  Z) (+Z) charge
Isotope • An atom with a mass equal to 1/3 the mass of a carbon-12 atom would have a
relative mass of 4 amu.
Kr-84 36 84 36 36 48 + 36 0
• An atom with a mass equal to twice the mass of a carbon-12 atom would have a
Al-273+ 13 27 13 10 14 +13 +3 relative mass of 24 amu.
13 – (+3) (27–13)

16 36 16 18 20 +16 -2 • An atom with a mass equal to five times the mass of a carbon-12 atom would have
S-362- a relative mass of 60 amu.
(16+20)

40 90 40 40 50 +40 0
Zr-90 (90 – 40)

(2.3) Quantitative properties of the atom: Atomic Mass (2.3) Quantitative properties of the atom: Atomic Mass
Mass Number Is NOT the Same as Atomic Mass Average Atomic Mass Determination
• The mass number refers to the number of protons + neutrons in an isotope.
• whole number Ave atomic mass = +

• The atomic mass is an experimental number determined from all naturally occurring + + …. +
isotopes. It makes use of the relative mass which is compared to a standard.
• In calculating the atomic mass, the relative amount of each isotope (% natural
abundance) is considered. Sample Problem 1. Mass spectrometer can determine isotopic masses as
well as isotopic abundances. Chlorine has two stable isotopes, Cl-35 and Cl-
37. Given the following graph obtained by mass spectrometry for Cl, find the
• Atomic masses and isotopic abundances can average atomic mass of Cl.
be determined using a mass spectrometer
From the graph:
75.53%
34.97 amu
75.53% of the atoms have a mass of 34.97 amu (Cl-35)
• Atomic mass of an element indicates how heavy, on average, an 24.47% have a mass of 36.97 amu (Cl-37)
atom of an element is when compared to an atom of another element
24.47 %
36.97 amu 75.53 24.47
• Carbon-12 scale is based on the most common isotope of carbon 34.97 amu × + 36.97 amu × = 35.46 amu
100.0 100.0
• Mass of one 12C atom = 12 atomic mass units (amu)
• An atom half as heavy as a C-12 atom weighs 6 amu
(2.3) Quantitative properties of the atom: Atomic Mass (2.3) Quantitative properties of the atom: Atomic Mass
Average Atomic Mass Determination Average Atomic Mass Determination
Sample Problem 2
Sample Problem 2 Strontium has four naturally occurring isotopes with the following
Strontium has four naturally occurring isotopes with the masses and natural abundances:
following masses and natural abundances: atomic mass % natural abundance
atomic mass % natural abundance Sr-84 83.9134 amu 0.56%
Sr-86 85.9093 amu 9.86%
Sr-84 83.9134 amu 0.56%
Sr-87 86.9089 amu 7.00% ĸ calculated from a
Sr-86 85.9093 amu 9.86% Sr-88 87.9056 amu 82.58%
Sr-87 86.9089 amu ?
Calculate b) the atomic mass of strontium.
Sr-88 87.9056 amu 82.58%
atomic mass = (83.9134 u 0.56/100) + (85.9093 u 9.86/100)

Calculate a) the %natural abundance of Sr-87 + (86.9089 u 7.00/100) + (87.9056 u 82.58/100)

b) the atomic mass of strontium. = (0.46991504) + (8.47065698) + (6.083623) + (72.59244448)

= 87.62 amu

(2.3) Quantitative properties of the atom: Atomic Mass (2.3) Quantitative properties of the atom: Avogadro’s Number
Average Atomic Mass Determination Masses of Atoms in Grams
Sample Problem 2 • It is necessary to know the mass of an atom in grams so the quantity of
matter can be determined by weighing
Strontium has four naturally occurring isotopes with the • One He atom is about four times as heavy as one H atom
following masses and natural abundances: • The mass of 100 He atoms is about four times the mass of 100 H
atomic mass % natural abundance atoms
• The number of He atoms in 4.003 g helium = The number of H
Sr-84 83.9134 amu 0.56%
atoms in 1.008 g of hydrogen
Sr-86 85.9093 amu 9.86%
Sr-87 86.9089 amu ? • Avogadro’s Number (NA): Number of atoms of an element in a sample
Sr-88 87.9056 amu 82.58% whose mass in grams is numerically equal to the atomic mass of the
element
Calculate a) the %natural abundance of Sr-87 • NA = 6.022 × 1023

Solution: the total % abundance should be equal to 100% • Example: atomic mass of P = 30.97 amu
so 6.022 × 1023 atoms of P = 30.97 g
100%  [0.56% + 9.86% + 82.58%] = 7.00%
(2.3) Quantitative properties of the atom: Avogadro’s Number (2.4) Introduction to the Periodic Table
Example 2.3 (a) The Modern Periodic Table
• A periodic table is a tabular arrangement of the elements based on the periodic law.
• Consider arsenic (As), a poison usually used in crime stories. Determine
• Elements are arranged in increasing atomic numbers.
the mass (g) of an arsenic atom. (Avogadro’s number is 6.022×1023)
• Elements with similar chemical and physical properties are in the same column.
atomic mass of As = 74.92 amu so, 6.022 × 1023 atoms of As = 74.92 g • Columns are called Groups or Families - designated by a number and letter at top.

To calculate the mass of 1 As atom in grams: • Rows are called Periods.

74.92 g As
1 atom As u = 1.244 u 10 –22 g
6.022 u 1023 atoms As

DRILL. One of the isotopes of zinc, Zn-70 has an atomic mass of 69.93 amu.
What is the the mass in grams of 895 atoms of Zn-70?

atomic mass of 70Zn = 69.93 amu so, 6.022 × 1023 atoms of 70Zn = 69.93 g

belongs to the 6th period

belongs to the 7th period

(2.4) Introduction to the Periodic Table (2.4) Introduction to the Periodic Table
Mendeleev's Periodic Law
• ordered elements by atomic mass; observed a repeating pattern of properties.
• Mendeleev’s Periodic law – when the elements are arranged in order of
library.thinkquest.org
increasing relative atomic mass, certain sets of properties recur periodically
1 2 3 4 5 6 7 8 9 10 11 12 13 14 15 16 17 18 19 20
H He Li Be B C N O F Ne Na Mg Al Si P S Cl Ar K Ca
H He Li Be B C N O F Modern Periodic law – when the elements
are arranged in order of increasing atomic
Ne Na Mg Al Si P S Cl number, certain sets of properties recur
periodically
Ar K Ca
Dmitri Mendeleev began development of the periodic table. He left empty spaces in his table to
place new elements yet to be discovered. • Separated by the diagonal line (stairway)
Property Silicon’s Tin’s Predicted Measured
properties properties value value
• Line begins from the bottom of the cell of B and ends at the right of Po
Mendeleev's Predictions for atomic mass 28 118 72 72.6
Ekasilicon (Germanium) • Metals: Elements present to the left of the diagonal line in the periodic table
color gray white gray gray-white
• Possess the ability to conduct electricity
density 2.32 7.28 5.5 5.4
• Nonmetals: Above and to the right of the stairway
reaction resists acid, reacts acid, resists resists
with acid reacts with resists with both both • Metalloids: Elements along the dividing line in the periodic table that possess
and base base base
properties between those of metals and non-metals
oxide SiO2 SnO2 Eks1O2 GeO2
• Commonly classified as metalloids: B, Si, Ge, As, Sb, and Te
(2.4) Introduction to the Periodic Table (2.4) Introduction to the Periodic Table
Metals, Nonmetals and Metalloids Drill 1 – Classify Each Element as Metal, Nonmetal, or Metalloid. If it is a metal
mercury, Hg copper, Cu classify it further as alkali, alkaline earth, transition or inner transition metal
METALS
• Solids at room temperature, except Hg • Conduct heat and electricity 1. Xenon, Xe – Nonmetal
• Reflective surface - shiny • Lose electrons and form cations in reactions
2. Tungsten, W – Metal (transition metal)
• Malleable – can be shaped • About 75% of the elements are metals
• Ductile - drawn or pulled into wires • Lower left on the table 3. Bromine, Br – Nonmetal
4. Arsenic, As – Metalloid
NONMETALS
5. Cerium, Ce – Metal (inner transition metal)
• Found in all 3 states.
neon, Ne 6. Potassium, K – Metal (alkali metal)
• Poor conductors of heat and electricity. helium, He
• Solids are brittle. 7. Strontium, Sr – Metal (alkaline earth metal)
• Gain electrons in reactions to become anions.
chlorine bromine iodine
• Upper right on the table. Drill 2. Identify the element that belongs to
• Except H.
A. Group IIA (2) and period 4 Ÿ Ca, calcium

B. Group IB (11) and period 5 Ÿ Ag, silver

METALLOIDS Properties of Silicon:
• Show some properties of metals
9Shiny C. Group VIA (16) and period 3 Ÿ S, sulfur
9Conducts electricity
and some of nonmetals.
9Does not conduct heat well D. period 5 and is a noble gas Ÿ Xe, xenon
• Also known as semiconductors. 9Brittle
E. alkali metal family and period 6 Ÿ Cs, cesium
www.webelements.com/silicon/

(2.4) Introduction to the Periodic Table (2.5) Molecules and Ions

Group or Family Chemical Formulas Describe Compounds
• A group or family is a • Compound – distinct substance that is composed of the atoms of
vertical column of
elements that have similar two or more elements and always contains exactly the same
chemical properties. relative masses of those elements.
• Traditional designation
uses a Roman numeral
and a letter (either A or B)
• Chemical Formulas – expresses the types of atoms and the
at the top of the column. number of each type in each unit (molecule) of a given compound.
• Modern (but not
universally-used) Inner Transition
designation uses only a
number from 1 to 18.
Metals
Molecules

• Main-group elements (representative elements = “A” groups): 1, 2, 13,14,15,16,17, and 18 • Two or more atoms may combine to form a molecule
• Transition metals (“B” groups) : 3, 4, 5, 6, 7, 8, 9, 10, 11, and 12 • Atoms involved are often nonmetals
• Post-transition metals: Elements in Groups 13, 14, and 15 to the right of the transition • Covalent bonds are strong forces that hold the atoms together
metals e.g. Ga, In, Tl, Sn, Pb, Bi, Po
• Alkali metals: Elements in Group 1, at the far left of the periodic table • In a molecular formula, the number of each atom is indicated by a
• Alkaline earth metals: Elements in Group 2 subscript
• Halogens: Elements in Group 17 • A molecular substance such as water is represented by the formula H2O
• Noble gases: Constitute Group 18
• Inner transition elements = rare earth elements = metals; really belong in periods 6 and 7
(2.5) Molecules and Ions (2.5) Molecules and Ions
Rules for Writing Formulas
CATIONS
1. Each atom present is represented by its element symbol.
– positive charge ions Ÿ more protons than electrons
2. The number of each type of atom is indicated by a subscript written to
– formed by losing electrons
the right of the element symbol.
– metals tend to lose one or more electrons to form cations
3. When only one atom of a given type is present, the subscript 1 is not
written.

Drill: Write Formulas for Each of the Following Compounds Mg o Mg2+ + 2e

• Acetone  Each molecule contains three carbon atoms, six hydrogen

atoms, and one oxygen atom. C H O3 6

– generally named by using the name of the parent atom

• Table sugar  Each molecule contains twelve carbon atoms, twenty two
Na = sodium atom Ca = calcium atom
hydrogen atoms, and eleven oxygen atoms.
C12H22O11 Na+ = sodium ion Ca2+ = calcium ion

(2.5) Molecules and Ions (2.5) Molecules and Ions

Ions ANIONS
• When atoms or molecules lose or gain electrons, they form charged – negative charge ions Ÿ more electrons than protons
particles called ions – formed by gaining electrons
• There is no change in the number of protons in the nucleus when an ion – nonmetals tend to gain one or more electrons to form anions
forms

• Positively charged ions (formed by losing electrons) are called CATIONS

• Na ĺ Na+ + e-
• Ba ĺ Ba2+ + 2e-
• Al ĺ Al3+ + 3e-

Cl + e o Cl
• Negatively charged ions (formed by gaining electrons) are called ANIONS
• Br + e– ĺ Br –
• S + 2e– ĺ S2– – named by using the root of the atom name followed by the suffix –ide
• N + 3e– ĺ N3– Br = bromine atom S = sulfur atom

Br = bromide ion S2 = sulfide ion
(2.5) Molecules and Ions (2.5) Molecules and Ions: Monoatomic Ions
Monatomic and Polyatomic Ions Type I Cations: Metal atoms that form only one type of charge
• Monatomic ions
• monoatomic anion – derived from a single atom by the gain of electron(s)
Type I cations
example: S2–, N3–, Cl–
• monoatomic cation – derived from a single atom by the loss of electron(s)
example: Na+, Ca2+, Pb4+, Zn2+, Ga3+, Fe3+

• Polyatomic ions are those that possess more than one atom
• OH, NH4+, Hg22+

Two Types of Monatomic Cations:

• Type I Cations – metal atoms that form only one type of charge
example: Na+, Ba2+, Ag+

• Type II Cations – metal atoms that can form more than one type of
Type II cations – form more
positive charge example: Cu = Cu+ , Cu2+ Fe = Fe2+ , Fe3+
than one type of + charge

(2.5) Molecules and Ions: Monoatomic Ions (2.5) Molecules and Ions: Monoatomic Ions
MEMORIZE THESE ION Type I Cations: Metal atoms that form only one type of charge
NAMES AND SYMBOLS! (Memorize these ion names and symbols.)
Group 1A (1) – forms +1 Group 2A (2) – forms +2
Metal Ion Symbol Name of Ion Metal Ion Symbol Name of Ion
lithium Li+ lithium beryllium Be2+ beryllium
sodium Na+ sodium magnesium Mg2+ magnesium
Group VIA (16): forms 2 potassium K+ potassium calcium Ca2+ calcium
MONOATOMIC ANIONS: Common Nonmetal Ion Symbol Name of Ion rubidium Rb+ rubidium strontium Sr2+ strontium
anions formed by nonmetals. oxygen O2 oxide cesium Cs+ cesium barium Ba2+ barium
Name of non-metal ion: sulfur S2 sulfide
stem of nonmetal name + -ide selenium Se2 selenide Group 3A (13) – forms +3 Group 1B (+1) Transition Metal
tellurium Te2 telluride Metal Ion Symbol Name of Ion Metal Ion Symbol Name of Ion
Group VIIA (17): forms 1
aluminum Al3+ aluminum silver Ag+ silver
Nonmetal Ion Symbol Name of Ion Group VA (15): forms 3
gallium Ga3+ gallium Group 2B (+2) Transition Metal
fluorine F fluoride Nonmetal Ion Symbol Name of Ion
indium In3+ indium Metal Ion Symbol Name of Ion
chlorine Cl chloride nitrogen N3 nitride
bromine Br bromide phosphorus P3 phosphide zinc Zn2+ zinc
iodine I iodide arsenic As3 arsenide cadmium Cd2+ cadmium
(2.5) Molecules and Ions: Monoatomic Ions (2.5) Molecules and Ions: Polyatomic Ions
Some Common Polyatomic Ions (Cont’n)
Type II Cations Common Type II Cations (incomplete list)
Metal Ion Symbol Name of Ion (MEMORIZE these ion names and symbols.)
• Metals in these Older Name
Systematic Name
compounds can form C. OXYANION (anion containing oxygen) ends in –ate
copper Cu+ copper(I) cuprous
more than one type
Cu2+ copper(II) cupric
of positive charge. Polyatomic Anions
iron Fe2+ iron(II) ferrous
• Charge on the metal Fe3+ iron(III) ferric Name Ion Symbol
ion must be chromium Cr2+ Chromium(II) chromous acetate C2H3O2
specified.
Cr3+ chromium(III) chromic permanganate MnO4
• Roman numeral cobalt Co2+ cobalt(II) cobaltous chromate CrO42
indicates the charge Co3+ cobalt(III) cobaltic
dichromate Cr2O72
of the metal cation. mercury *Hg2 2+ *mercury(I) *mercurous
carbonate CO32
Hg2+ mercury(II) mercuric
• Transition metal hydrogen carbonate (or bicarbonate) HCO3
cations usually lead Pb2+ lead(II) plumbous
require a Roman Pb4+ lead(IV) plumbic oxalate C2O42
numeral. tin Sn2+ tin(II) stannous silicate SiO32
Sn4+ tin(IV) stannic arsenate AsO43
*Mercury(I) ions always occur bound together in pairs to form Hg22+

(2.5) Molecules and Ions: Polyatomic Ions (2.5) Molecules and Ions: Polyatomic Ions
• Polyatomic ions are charged entities composed of several atoms Some Common Polyatomic Ions (Cont’n)
bound together. C. OXYANION ends in –ite or ate (MEMORIZE these ions!)
• They have special names and must be MEMORIZED! For series of 2 oxyanions:
ate – higher O number ite – lower O number
Some Common Polyatomic Ions.
A. POLYATOMIC CATIONS Name of Polyatomic Anion Symbol of Ion
Polyatomic Cation (+1) nitrate NO3
Name Ion Symbol nitrite NO2
ammonium NH4+ sulfate SO42
mercury(I) Hg22+ sulfite SO32
hydrogen sulfate (or bisulfate) HSO4
B. POLYTOMIC ANIONS ending in -ide
Polyatomic Anions ending in -ide
hydrogen sulfite (or bisulfite) HSO3
Name Ion Symbol dihydrogen phosphate H2PO4
hydroxide OH hydrogen phosphate HPO42
peroxide O22 phosphate PO43
cyanide CN
phosphite PO33
(2.5) Molecules and Ions: Polyatomic Ions (2.5) Molecules and Ions: Ionic Compounds
Some Common Polyatomic Ions (Cont’n)
(MEMORIZE these ion names and symbols.)

C. OXYANION ends in –ite or ate

For series of 4 oxyanions:
ate – higher O number ite – lower O number
prefixes used: per – “more than” ; hypo – “less than”

Polyatomic Anions Polyatomic Anions Polyatomic Anions

strong weak
Name Ion Name Ion Name Ion electrolyte electrolyte
nonelectrolyte

Symbol Symbol Symbol

hypochlorite ClO hypobromite BrO hypoiodite IO
The conductivity of a solution depends on the number of ions present
chlorite ClO2 bromite BrO2 iodite IO2
• Strong electrolytes readily produce ions in aqueous solution
chlorate ClO3 bromate BrO3 iodate IO3
• Weak electrolytes produce a relatively lesser number of ions in
perchlorate ClO4 perbromate BrO4 periodate IO4
aqueous solution
• Nonelectrolytes are those that do not produce ions in aqueous solution

(2.5) Molecules and Ions: Ionic Compounds (2.6) Formulas of Ionic Compounds

Ionic Compounds – compounds containing both anions and cations like Formulas of Ionic Compounds
table salt or NaCl consists of Na+ and Cl– ions in equal proportion • Write the cation element symbol followed by the anion element symbol.
• The number of cations and anions must be correct for their charges to
Ionic compounds – held together sum to zero.
by strong forces called ionic bonds
• ionic bonds – attraction
between cations (+) and
anions (-)
• basic unit – formula unit (they
do NOT exist as molecules)
• electrically neutral (the sum of the charges
of the anions and cations is equal to zero) • Balancing charge
• hard but brittle • Each positive charge must have a negative charge to balance it
• have high melting points • For magnesium chloride, MgCl2 composed of 1 Mg2+ and 2 Cl
• the +2 charge of Mg2+ requires 2 Cl ions for charge balance
• conduct electricity when molten or when
dissolved in water (due to moving ions) • According to the principle of electrical neutrality, the total positive charge
of cations in the formula must equal the total negative charge of the anions
(2.6) Formulas of Ionic Compounds (2.6) Formulas of Ionic Compounds

Drill Drill
Predict the chemical formula for the compound formed by each of the
following pairs.
a) magnesium and nitrogen 1. A compound contains an unknown atom X and has the formula XCl2. The
X ion contains 20 electrons. What is the identity of the X ion?
b) zinc and sulfur
c) silver and phosphorus a) Ti2+ b) Sc+ c) Ca2+ d) Cr2+
d) copper(II) and phosphate
e) ammonium and oxalate

Identify the ions composing the following ionic compounds.

a) SnO2
b) Cr3P2 2. A compound contains an unknown atom Y and has the formula Al2Y3.
The Y ion contains 36 electrons. What is the identity of ion Y?
c) CuO
d) Cu2O a) P3 b) Se2 c) As3 d) Sr2
e) FeS
f) Fe2S3
g) BaSO4

(2.6) Formulas of Ionic Compounds (2.7) Names of Compounds

Drill Naming of Compounds:

Predict the chemical formula for the compound formed by each of the
following pairs. A. Binary Compounds – composed of two kinds of elements
a) magnesium and nitrogen Mg3N2
1. Binary Ionic Compounds – composed of metal – nonmetal
b) zinc and sulfur ZnS
c) silver and phosphorus Ag3P a. Binary ionic compound with Type I metal (Type I)
d) copper(II) and phosphate Cu3(PO4)2
b. Binary ionic compound with Type II metal (Type II)
e) ammonium and oxalate (NH4)2C2O4
2. Binary Covalent Compounds – composed of two nonmetals (Type III)
Identify the ions composing the following ionic compounds.
a) SnO2 Sn4+ and O2
b) Cr3P2 Cr 2+ and P3 B. Ionic Compounds containing polyatomic ions
c) CuO Cu2+ and O2
d) Cu2O Cu+ and O2 C. Binary Acids – acids composed of H and a nonmetal
e) FeS Fe 2+ and S2 D. Oxyacids – acids formed by oxyanions
f) Fe2S3 Fe3+ and S2
g) BaSO4 Ba2+ and SO42
(2.7) Names of Compounds (2.7) Names of Compounds

Naming of Binary ionic compound with Type I metal (Type I) Naming of Binary ionic compound with Type II metal (Type II)
1. The cation is always named first and the anion second. 1. The cation is always named first and the anion second.

2. A simple cation takes its name from the name of the element. 2. Because the cation can assume more than one charge, the charge
is specified by a Roman numeral in parentheses.
3. A simple anion is named by taking the first part of the element name
(the root) and adding –ide. 3. A simple anion is named by taking the first part of the element name
(the root) and adding –ide.
Examples:
KCl potassium chloride
MgBr2 magnesium bromide Formula: Fe2O3
CaO calcium oxide Name: iron(III) oxide

Examples: CuBr copper(I) bromide

FeS iron(II) sulfide
PbO2 lead(IV) oxide

(2.7) Names of Compounds (2.7) Names of Compounds

Naming of Binary ionic compound with Type I metal (Type I) Naming of Binary ionic compound with Type II metal (Type II)

Drill Drill
1. What is the name of the compound SrBr2? 1. What is the name of the compound CrO2?

a) strontium bromine c) sulfur bromide a) chromium oxide c) chromium(IV) oxide

b) strontium dibromide d) strontium bromide b) chromium(II) oxide d) chromium dioxide

2. What is the correct name of the compound, XS that results from the most
2. What is the name of the compound Al2S3? stable ion for sulfur and the metal ion (X) that contains 24 electrons?

a) dialuminum trisulfide c) aluminum sulfur a) iron(III) sulfide c) nickel(III) sulfate

b) aluminum sulfide d) aluminum trisulfide b) chromium(II) sulfide d) iron(II) sulfide
(2.7) Names of Compounds (2.7) Names of Compounds

Naming of Binary Covalent Compounds (Type III) (Molecular Compounds)

Flow Chart for Naming Binary Compounds
Formed between two nonmetals.
1. The first element in the formula is named first, and the full element name is used.
2. The second element is named as though it were an anion (root of 2nd atom + -ide).
3. Prefixes are used to denote the numbers of atoms present.
4. The prefix mono- is never used for naming the first element.
5. Drop last “a” of the prefix if the name of the nonmetal begins with “o”

Prefixes Used to Indicate Numbers of Atoms

Number Greek Number Greek

Prefix Prefix
1 mono- 6 hexa-
Examples:
2 di- 7 hepta-
CO carbon monoxide
3 tri- 8 octa-
SF6 sulfur hexafluoride
4 tetra- 9 nona-
N2O4 dinitrogen tetroxide
5 penta- 10 deca-

(2.7) Names of Compounds (2.7) Names of Compounds

Naming of Binary covalent compounds (Type III) Naming of Ionic Compounds Containing Polyatomic Ions
Drill Naming ionic compounds containing polyatomic ions follows rules
1. What is the correct name of the compound SeO2? similar to those for binary compounds. The cation is always named first
and the anion second.
a) selenium oxide c) selenium(II) oxide Example: ammonium acetate
b) selenium dioxide d) selenium(IV) dioxide

2. What is the correct name of the compound I2F7?

a) iodine fluoride c) diiodine heptafluoride

b) iodine heptafluoride d) diodine heptafluoride NaOH sodium hydroxide

Mg(NO3)2 magnesium nitrate

3. What is the correct name of the compound P2O5?
(NH4)2SO4 ammonium sulfate
a) diphosphorus pentaoxide c) phosphorus pentoxide
FePO4 iron(III) phosphate
b) phosphorus oxide d) diphosphorus pentoxide
(2.7) Names of Compounds (2.7) Names of Compounds

Naming of Binary ionic compound with Type II metal (Type II) Naming of Acids
• Acids can be recognized by the hydrogen that appears first in the
Drill
formula e.g. HCl (hydrochloric acid)
1. What is the name of the compound KClO3? • Molecule with one or more H+ ions attached to an anion.

a) potassium chlorite c) potassium perchlorate

Rules for Naming Acids
b) potassium chlorate d) potassium carbonate
• If the anion does not contain oxygen or if the name of the anion
ends in –ide, the acid is named with the prefix hydro– and the suffix
2. What is the name of the compound Fe2(CO3)3? –ic attached to the root name for the element.

a) iron carbonate c) iron(III) carbonate

• Examples: anion
b) diiron tricarbonate d) iron(II) carbonate HF(aq) F fluoride hydrofluoric acid
HCN(aq) CN cyanide hydrocyanic acid
3. What is the name of the compound Al(NO3)3?
H2S(aq) S2 sulfide hydrosulfuric acid
a) aluminum nitrate c) aluminum(III) trinitrate
b) aluminum(III) nitrate d) aluminum trinitrate

(2.7) Names of Compounds (2.7) Names of Compounds

Naming of Oxyacids
Drill
• If the anion contains oxygen (oxyanion):
Examine the following table of formulas and names.
Which of the compounds are named correctly?
a. if the anion name ends in –ate, the suffix –ic is added to the root name.
Formula Name anion acid
Examples: HNO3(aq) 
NO3 nitrate nitric acid
I P2O5 Diphosphorus pentoxide 9 H2SO4(aq) SO42 sulfate sulfuric acid
a) I, II
II ClO2 Chlorine oxide 8 chlorine dioxide

III PbI4 Lead iodide8 lead(IV) iodide

HC2H3O2(aq) C2H3O2 acetate acetic acid
b) I, III, IV
IV CuSO4 Copper(I) sulfate8 copper(II) sulfate
c) I, IV b. if the anion name ends in –ite, the suffix –ous is added to the root name.
d) I only anion acid
Examples: HNO2(aq) NO2 nitrite nitrous acid
H2SO3(aq) SO32 sulfite sulfurous acid

HClO(aq) ClO hypochlorite hypochlorous acid

(2.7) Names of Compounds (2.7) Names of Compounds
Drill. Name the acids.
Flowchart for Naming Acids 1. H2SO4 (aq) anion: SO42 name of anion: sulfate
name of acid: sulfuric acid

2. H2S (aq) anion: S2 name of anion: sulfide

name of acid: hydrosulfuric acid

3. H2C2O4 (aq) anion: C2O42 name of anion: oxalate

name of acid: oxalic acid

Drill. Which of the following compounds is named incorrectly?

a) KNO3 potassium nitrate
b) TiO2 titanium(II) oxide
c) Sn(OH)4 tin(IV) hydroxide
d) PBr5 phosphorus pentabromide
e) H2SO3(aq) sulfurous acid

(2.7) Names of Compounds (2.7) Names of Compounds

Naming of Acids Writing of Chemical Formulas from the Name of the Compound
Examples • Sodium hydroxide = NaOH
Practice: Name the Following
• Potassium carbonate = K2CO3
1. HBrO3 (aq)
Solution: anion: BrO3 • Sulfuric acid = H2SO4(aq)
name of anion: bromate (change ate to ic and add acid)
name of acid: bromic acid • Dinitrogen pentoxide = N2O5

• Cobalt(III) nitrate = Co(NO3)3

2. HBrO2 (aq)
Solution: anion: BrO2
name of anion: bromite (change ite to ous and add acid) Drill. A compound has the formula XCl3 where X could represent a
name of acid: bromous acid metal or nonmetal. What could the name of this compound be?

3. HBr (aq) a) phosphorus trichloride

Solution: anion: Br
b) carbon monochloride
name of anion: bromide o ends in ide
(change ide to ic; use prefix “hydro” + brom + ic and add acid) c) tin(IV) chloride
name of acid: hydrobromic acid
d) magnesium chloride
(2.7) Names of Compounds (2.7) Names of Compounds

Writing Formulas for Acids Given Their Name NAMING COMPOUNDS: SUMMARY DRILL
• When name ends in acid, formulas starts with H.
Identify the compound INCORRECTLY named and modify the name
• Write formulas as if ionic, even though it is molecular.
to make it correct.
• Hydro- prefix means it is binary acid, no prefix means it is an oxyacid.
• For an oxyacid, if ending is –ic, polyatomic ion ends in –ate; if ending is – 1. NO2 – mononitrogen dioxide = nitrogen dioxide
ous, polyatomic ion ends in –ite.
mono is not used in the beginning of the name of a
covalent compound

DRILL: Write the formula of the following acids. (Hint: Identify the anion first)
ions formula 9 2. SiO2 – silicon dioxide
1. Chlorous acid H+ with ClO2– HClO2(aq)

2. Phosphoric acid H+ with PO43– H3PO4(aq)

3. CaO – calcium monoxide = calcium oxide
3. Hydrofluoric acid H+ with F– HF(aq) This is an ionic compound!
Prefixes mono, di, etc … are NOT used in ionic compounds!

(2.7) Names of Compounds (2.7) Names of Compounds

Writing Formulas for Compounds Given Their Name NAMING COMPOUNDS: SUMMARY DRILL
Drill A. Give the formula for the following Type III binary compounds Identify the compound INCORRECTLY named and modify the name
containing only nonmetals (molecular compounds). to make it correct.
Name Formula
4. N2O4 – dinitrogen tetraoxide = dinitrogen tetroxide
1. sulfur hexafluoride SF6

2. radon dibromide RnBr2 “a” in tetra, penta, hexa, hepta, octa are dropped if the last
atom to be named starts with “o”
3. dinitrogen pentoxide N2O5
95. Sr3(PO4)2 – strontium phosphate
Drill B. Give the formula for the following ionic compounds.
Name Formula
6. Ba(NO3)2 – barium dinitrate = barium nitrate
1. zinc iodide ZnI2
2. iron(III) selenide Fe2Se3 This is an ionic compound!
Prefixes mono, di, etc … are NOT used in ionic compounds!
3. lead(IV) telluride PbTe2

4. lithium sulfate Li2SO4

5. ammonium carbonate (NH4)2CO3

(2.7) Names of Compounds
Identify the compound INCORRECTLY named and modify the name
to make it correct.
7. ZnCO3 – zinc(II) carbonate = zinc carbonate

Zinc is a Type I cation. It forms only one type of charge

(+2), thus the charge need NOT be indicated in the name.

8. HBrO – hypobromic acid = hypobromous acid

The name of the anion, BrO is hypobromite, therefore

–ite is replaced by ous.

99. NH3 – ammonia

910. H O – water
2

9 11. CH – methane
4

Ammonia (NH3), water (H2O) and methane (CH4) are common names used
for these corresponding compounds. They are NOT named using the
prefixes for covalent compounds.