Chapter 2 – The Components of Matter
Sample Problem 2.1 p. 39
Distinguishing Elements, Compounds, and Mixtures at the
Atomic Scale
(a)
Mixture - 2
molecular elements
(green & purple), 1
compound (2 red + 1
yellow)
(b)
Element – only one
kind of atom
(c)
Compound molecules that each has two black and six blue
atoms
Follow-up Problem 2.1
Reactants (left) mixture of two elements
Products (right) compound of molecules made up of one red
atom and one blue atom
Sample Problem 2.2 p. 42
Calculating the Mass of an Element in a Compound
102 kg
pitchblende
71.4 kg
uranium
84.2 kg
pitchblende
1000 g
uranium
1 kg uranium
= 86,500 g
uranium
= 8.65 x 104 g Uranium
Follow-up Problem 2.2 p. 42
2.3 t uranium
84.2 t
12.8 t oxygen
pitchblende
71.4 t uranium 84.2 t
pitchblende
= .41 t oxygen
Sample Problem 2.3 p. 44
Visualizing Mass Laws
16 atoms on each
side (9 green + 7
purple)… mass
conservation
The compound
formed has 1 purple
with 2 green… law
of definite
composition
Only one compound
formed so this is
NOT an example of
law of multiple
proportions
Follow-up Problem 2.3 p. 44
ONLY diagram B… it has molecules of 1 green to 1 red and 3
green to 1 red
The Atomic Theory Today
 An atom is an electronically neutral, spherical entity
composed of a positively charged central nucleus
surrounded by one or more negatively charged electrons.
 The atom’s diameter is about 20,000 times the diameter of
the nucleus… (The nucleus would be the size of a pea in an
atom the size of the Houston Astrodome!)
 The nucleus contributes 99.97% of the atom’s mass and
only about 1 quadrillionth of its volume… its density is
about 1014 g/mL!
 The nucleus contains protons and neutrons
o Protons (p+) have a positive charge
o Neutrons (n0) have no charge (neutral)
 Electrons (e-) are found outside the nucleus and have a
negative charge
 An atom is neutral because the number of protons in the
nucleus is equal to the number of electrons surrounding
the nucleus.
Atomic Number, Mass Number, and Atomic Symbol
 The atomic number (Z) of an element equals the number
of protons in the nucleus of each atom
o All atoms of an element have the same atomic number
o The atomic number for each element is unique to that
element
o 117 known elements
 90 occur in nature and 27 have been synthesized
by nuclear scientist
 The mass number (A) is the total number of protons and
neutrons in the nucleus.
o Each proton and neutrons contributes one unit to the
mass number
 A carbon atom with 6 protons and 6 neutrons in
its nucleus has a mass number of 12
 The atomic symbol (or element symbol) of an element is
based on its English, Latin, or Greek name
o Sodium gets its symbol (Na) from its Latin name
natrium
 To calculate the number of neutrons an atom has…
o Number of neutrons = mass number – atomic number
Isotopes
 All atoms of an element have the same atomic number but
not the same mass number.
 Isotopes of an element are atoms that have different
numbers of neutrons and therefore different mass numbers
What is the difference between carbon12 (pictured to the left) and carbon-14?
ANSWER: carbon-12 has 6 protons, 6
neutrons, and 6 electrons… carbon -14
has 6 protons, 8 neutrons, and 6
electrons
What is the difference between oxygen16 and oxygen-18?
ANSWER: oxygen-16 has 8 protons, 8
neutrons, and 8 electrons… oxygen-18
has 8 protons, 10 neutrons, and 8
electrons.
Sample Problem 2.4 p. 49
Determining the Number of Subatomic Particles in the Isotopes
of an Element
28
14
Si – 14p+, 14e-, 14no
29
14
Si – 14p+, 14e-, 15no
30
14
Si - 14p+, 14e-, 16no
Follow-up Problem 2.4 p.49
11
5
Q – 5p+, 5e-, 6no
41
20
R – 20p+, 20e-, 21no
131
53
X – 53p+, 53e-, 78no
(Boron)
(Calcium)
(Iodine)
Atomic Masses of the Elements
 The mass of an atom is measured relative to the mass of an
atomic standard
o The modern standard is the carbon-12 atom which is
defined as exactly 12 atomic mass units
 The atomic mass unit is 1/12 the mass of a
carbon-12 atom
 1 atomic mass unit = 1 dalton (Da)
 Although it is a relative mass it has an
absolute mass of 1.66054 x 10-24 g
 Mass spectrometry – method for measuring the relative
masses and abundances of atomic-scale particles very
precisely.
Example…
Mass of 28Si atom
= 2.331411
Mass of 12C standard
To calculate the isotopic mass of 28Si
Isotopic mass of 28Si = measured mass ratio x mass of 12C
Isotopic mass of 28Si = 2.331411 x 12 amu
Isotopic mass of 28Si = 27.97693 amu
 Atomic Mass (also called atomic weight) – the average of
the masses of its natural occurring isotopes
To calculate the atomic mass of an element…
Isotope
28
Si
29
Si
30
Si
Mass (amu)
27.97693
28.976495
29.973770
Abundance %
92.23 % (.9223)
4.67% (.0467)
3.10% (.0310)
Atomic Mass
25.8031 amu
1.3532 amu
0.9292 amu
28.09 amu
Sample Problem 2.5 p. 52
Calculating the Atomic Mass of an Element
Isotope Mass (amu) Abundance
Ag - 107 106.90509
.5184
Ag - 109 108.90476
.4816
Atomic Mass
Follow-up Problem 2.5 p. 52
10.0129x + [11.0093 (1- x)] = 10.81
x = .2000 (10B)
1- x = .8000 (11B)
55.42 amu
52.45 amu
107.87
Periodic Table
 Dmitri Mendeleev (Russian chemist) published the most
successful organization of the elements in 1871 (he
arranged the elements in order by mass and arranged
elements with similar properties in the same column)
 The modern periodic table of elements is based on
Mendeleev’s version except elements are arranged by
atomic number
(1) Elements are arranged in increasing order by atomic
number (number of protons in the nucleus)
(2) Horizontal rows are called periods
(3) Vertical columns are called groups (or families)
(4) The eight A groups contain the main-group elements
(5) The ten B groups contain the transition elements
(6) Inner transition elements include the lanthanides and
actinides (pulled out of the main portion of the table at
the bottom)
(7) Metals (left of the “staircase”)… shiny, solids at room
temperature (mercury is the only liquid), conduct heat
and electricity well, malleable (they can be shaped with
tools), ductile (can be drawn into wires)
(8) Nonmetals (upper right corner above “staircase”)…
gases or dull brittle solids at room temperature (bromine
is the only liquid), poor conductors of heat and electricity
(9) Metalloids (lay along the “staircase”)… have properties
between those of metals and nonmetals
(10) Alkali metals – Group 1A (except hydrogen)
(11) Alkaline earth metals – Group 2A
(12) Oxygen family – Group 6A
(13) Halogens – Group 7A
(14) Noble gases – Group 8A
Compounds : Introduction to Bonding
 With a few exceptions, most elements are found in
compounds combined with other elements
 Elements combine in two general ways
o Ionic compounds – transfer electrons from one
element to another
o Covalent compounds – atoms sharing electrons
 Both processes generate chemical bonds
(Forces that hold atoms together)
Formation of Ionic Compounds
 Ionic compounds are composed of ions
o Charged particles that form when at atom (or small
group of atoms) gains or loses one or more electrons
o Simplest type of ionic compound is a binary ionic
compound
 Typically forms when a metal reacts with a
nonmetal
 Metal atoms lose one or more electrons and
form a cation
o Positively charged ion
 Nonmetal atoms gain one or more electrons
and form an anion
o Negatively charged ion
 Metal atoms transfer electrons to nonmetals
 A cation or anion derived from a single atom is
called a monatomic ion
The Case of Sodium Chloride
Predicting the Number of Electrons Lost or Gained
 Ionic compounds are neutral because they contain equal
numbers of positive and negative charges (In the case of
sodium chloride there are equal numbers of Na+ and Clbecause both are singly charged
o In sodium oxide (where we have Na+ and O2- ions), we
have twice as many Na+ as we have O2 To predict the number of electrons that a given atom will
gain or lose…
o For A group elements
 Metals lose electrons to form an ion that has the
same number of electrons as the nearest noble gas
 Group 1A – lose one electron form 1+ ions
 Group 2A – lose two electrons form 2+ ions
 Aluminum in Group 3A – loses two
electrons form 3+ ion
 Nonmetals gain electrons to form an ion that has
the same number of electrons as the nearest noble
gas
 Group 7A gains one electron forms 1- ions
 Group 6A gains two electrons forms 2- ions
 Nitrogen in Group 5A gains three electrons
forms 3- ions
Sample Problem 2.6 p. 57
Predicting the Ion an Element Forms
(a) I- [Iodine is in group 7A (17)… so it wants to gain (it is
a nonmetal) one more electron to have the same number
of electrons as Xe. When it gains one more electron, it
now has 17p+ and 18e-. Its net charge is now 1-.]
(b) Ca2+ [Calcium is in group 2A (2)… so it wants to lose
(it is a metal) 2 electrons so that it has the same number
of electrons as Ar. When it loses 2 electrons, it now has
20p+ and 18e-. Its net charge is now 2+.]
(c) Al3+ [Aluminum is in group 3A (13)… so it wants to
lose (it is a metal) 3 electrons so that it has the same
number of electrons as Ne. When it loses 3 electrons, it
now has 13p+ and 10e-. Its net charge is 3+.]
Follow-up Problem 2.6 p. 57
(a) S2- [Sulfur is in group 6A (16)… so it wants to gain (it
is a nonmetal) 2 more electrons so it has the same
number of electrons as Ne. When it gains 2 electrons, it
now has 16p+ and 18e-. Its net charge is 2-.]
(b) Rb+ [Rubidium is in group 1A (1)… so it wants to lose
(it is a metal) 1 electron so it has the same number of
electrons as Kr. When it loses 1 electron, it now has
37p+ and 36e-. Its net charge is 1+.]
(c) Ba2+ [Barium is in group 2A (2)… so it wants to lose (it
is a metal) 2 electrons so that it has the same number of
electrons as Xe. When it loses 2 electrons, it now has
56p+ and 54e-. Its net charge is 2+.]
Formation of Covalent Bonds
 Covalent compounds form when elements share electrons
o Usually occurs between nonmetals
 Covalent bond – a pair of electrons mutually attracted by
two nuclei
Covalent and Ionic Substances
 Most covalent substances consists of molecules
o A cup of water consists of individual water molecules
lying near each other.
 There are no molecules in ionic compounds
o Covalent bonding involves the mutual attraction
between two (positively charged) nuclei and the two
(negatively charged) electrons.
 An ionic substance such as sodium chloride is a continuous
array in three dimensions of oppositely charged sodium and
chloride ions, not a collection of individual sodium chloride
“molecules.”
o Ionic bonding involves the mutual attraction between
positive and negative ions.
Polyatomic Ions: Covalent Bonds Within Ions
 Many ionic compounds contain polyatomic ions
o An ion that consists of two or more atoms bonded
together covalently and has a net positive or negative
charge.
Formulas, Names, and Masses of Compounds
Chemical formula – uses element symbols and often numerical
subscripts to show the type and number of each atom in the
smallest unit of the substance.
 Represents one molecule if it is a molecular substance
 Represents one formula unit if it is an ionic compound
Binary Ionic Compounds
Two general rules…
(1) For all ionic compounds, names and formulas give the
positive ion (cation) first and the negative ion (anion)
second.
(2) For all binary ionic compounds, the name of the cation
is the name of the metal, and the name of the anion has
the suffix –ide added to the root of the name of the
nonmetal.
Example… calcium (metal) + bromine (nonmetal)
bromine become bromide
calcium bromide
Compounds of Elements That Form One Ion
The periodic table presents some key points about the formulas
of main-group monatomic ions.
 Monatomic ions of main group elements have the same
ionic charge
o The alkali metals [1A (1)] form ions with a 1+ charge
o The halogens [7A (17)] form ions with a 1- charge
 For cations (positive ions), the ion charge equals A-group
number (exceptions are Sn2+ and Pb2+)
 For anions (negative ions), the ion charge equals A-group
number minus 8
o Example… sulfur is in family 6A. (6 – 8 = -2) Sulfur
forms S2 Main group elements have the same number of electrons as
the nearest noble gas (Aluminum has 13p+ and 10e- and
thus forms a Al3+ ion)
Formula unit
 Ionic compounds consists of an array of ions rather than
separate molecules
o The formula of an ionic compound represents the
formula unit
 A formula unit is the relative numbers of cations
and anions in the compound. The compound has
a zero net charge, so the positive charges of the
cations balance the negative charges of the
anions.
 Example… in calcium bromide we have
twice as many Br - as Ca2+ [ CaBr2 ]
 Reduce the subscripts to the smallest whole
numbers… Ca2O2 would be CaO
Sample Problem 2.7 p. 61
Naming Binary Ionic Compounds
(a) magnesium nitride (magnesium is the metal and
nitrogen is the nonmetal… so magnesium is written first
and nitrogen gets an ide ending)
(b) cadmium iodide (cadmium is the metal and iodine is
the nonmetal… so cadmium is written first and iodine
gets an ide ending)
(c) strontium fluoride (strontium is the metal and fluorine
is the nonmetal… so strontium is written first and
fluorine gets an ide ending)
(d) cesium sulfide (cesium is the metal and sulfur is the
nonmetal… so cesium is written first and sulfur gets an
ide ending)
Follow-Up Problem 2.7 p. 61
(a)
(b)
(c)
(d)
zinc [Group 2B (12)], oxygen [Group 6A (16)]
silver [Group 1B (11)], bromine [Group 7A (17)]
lithium [Group 1A (1)], chlorine [Group 7A (17)]
aluminum [Group 3A (13)], sulfur [Group 6A (16)]
Sample Problem 2.8 p. 61
(a) Mg2+ and N3- (you need 3 Mg2+ to get 6+ and 2 N3- to
get 6- so that you have a neutral compound) Formula
is… Mg3N2
(b) Cd2+ and I - (you need 1 Cd2+ to get 2+ and 2 I - to get 2so that you have a neutral compound) Formula is…
CdI2
(c) Sr2+ and F - (you need 1 Sr2+ to get 2+ and 2 F – to get 2so that you have a neutral compound) Formula is…
SrF2
(d) Cs+ and S2- (you need 2 Cs+ to get 2+ and 1 S2- to get 2so that you have a neutral compound) Formula is…
Cs2S
Follow-Up Problem 2.8 p. 61
(a) Zinc is Zn2+ because it is in Group 2B and oxygen is O2because it is in Group 6A… You need 1 Zn2+ to get 2+
and you need 1 O2- to get to 2- so that you have a neutral
compound… Formula is ZnO
(b) Silver is Ag+ because it is in Group 1B and bromine is
Br – because it is in Group 7A… You need 1 Ag+ to get
1+ and 1 Br – to get 1- so that you have a neutral
compound… Formula is AgBr
(c) Lithium is Li+ because it is Group 1A (1) and chlorine is
Cl – because it is in Group 7A (17)… You need 1 Li+ to
get 1+ and 1 Cl – to get 1- so that you have a neutral
compound… Formula is LiCl
(d) Aluminum is Al3+ because it is in Group 3A (13) and
sulfur is S2- because it is in Group 6A (16)… You need
2 Al3+ to get 6+ and 3 S2- to get 6- so that you have a
neutral compound… Formula is Al2S3
Compounds with Metals That Form More Than One Ion
Many metals, especially the transition metals (B groups), can
form more than one ion…
Names of compounds containing these elements (and others)
that form more than one ion include a Roman numeral within
parentheses immediately after the metal ion’s name to indicate
its ionic charge.
Examples – FeCl2 is named iron (II) chloride and FeCl3 is
named iron (III) chloride
Trivial name rules…
(a) The suffix –ous for the ion with the lower charge
(b) The suffix –ic for the ion with the higher charge
Examples – iron (II) chloride is also called ferrous chloride
and iron (II) chloride is also called ferric chloride
Sample Problem 2.9 p. 62
Determining Names and Formulas of Ionic Compounds of
Metals That Form More Than One Ion
(a) Sn2+ because the name told you tin (II) and F – because F
is in Group 7A (17)… the formula is SnF2 (Because
you need 1 Sn2+ to get 2+ and 2 F – to get 2- so that you
have a neutral compound)
(b) You determine Cr is Cr3+ because there are 3I – in the
formula so you have 3-… there is only 1 Cr in the
formula so you know it has to be 3+ so that you have a
neutral compound. The name is chromium (III) iodide
(the Roman numeral in parentheses is necessary because
chromium forms more than one ion… Cr2+ and Cr3+)
(c) Ferric tells you it is Fe3+ and oxygen is O2- because it is
in Group 6A (16)… so you would need 2 Fe3+ to get 6+
and you need 3 O2- to get 6-… the formula is Fe2O3
(d) You determine Co is Co2+ because there is 1 S2- in the
formula so you have 2-… there is 1 Co in the formula so
you know it has to be 2+ so that you have a neutral
compound. The name is cobalt (II) sulfide (the Roman
numeral in parentheses is necessary because cobalt forms
more than one ion… Co2+ and Co3+)
Follow-up Problem 2.9 p. 62
(a) Lead (IV) tells you Pb4+ and you know that oxygen is O2because it is in Group 6A (16). You need 1 Pb4+ to have
4+ and you would need 2 O2- to have 4- so that you have
a neutral compound. The formula is PbO2
(b) The formula tells you that you have 1 S2-. [we know that
sulfur is 2- because it is in Group 6A (16)] Since the
formula tells us that we have 2 Cu, Cu must be Cu+1. 1
S2- gives us 2- and 2 Cu+ give us 2+ so that we have a
neutral compound. The name of the compound is
copper (I) sulfide (the Roman numeral in parentheses is
necessary because copper forms more than one ion…
Cu+ and Cu2+)
(c) The formula tells us that we have 2 bromines. Each
bromine is Br -. We know that bromine is 1- because
bromine is in Group 7A (17). So from the bromines we
have 2-. Since we have only 1 Fe in the formula, that
means that Fe is 2+. The name of the compound is iron
(II) bromide (the Roman numeral in parentheses is
necessary because iron forms more than one ion… Fe2+
and Fe3+)
(d) Mercuric tells us that it is Hg2+ and we know that
chlorine is Cl – because it is in Group 7A (17). We need
1 Hg2+ to give us 2+ and 2 Cl – to give us 2- so that we
have a neutral compound. So the formula is HgCl2
Compounds That Contain Polyatomic Ions
Many ionic compounds contain polyatomic ions. The table
below lists the common polyatomic ions…
Formula for potassium nitrate is
KNO3 because potassium is K+
and nitrate is NO3- (1 of each ion
makes a neutral compound)
Formula for sodium carbonate is
Na2CO3 because sodium is Na+
and carbonate is CO32- (2 Na+
combine with 1 CO32- to make a
neutral compound)
Formula for calcium nitrate is
Ca(NO3)2 because calcium is Ca2+
and nitrate is NO3 – (1 Ca2+ combines with 2 NO3 – to make a
neutral compound. Note when you use more than one
polyatomic ion in a formula, you need to place it in parentheses
with a subscript noting the number of ions needed to make a
neutral compound.
Families of Oxoanions
Most polyatomic ions are oxoanions, or oxyanions. (Those in
which an element, usually a nonmetal is bonded to one or more
oxygen atoms) There are several families of two or four
oxoanions that differ only in the number of oxygen atoms.
 Rules when we have two oxoanions in a family
o The ion with more oxygen atoms takes the nonmetal
root and the suffix –ate.
o The ion with less oxygen atoms takes on the nonmetal
root and the suffix –ite.
 Examples…
 SO42- is sulfate and SO32- is sulfite
 NO3 – is nitrate and NO2 – is nitrite
 Rules when you have four oxoanions in a family
o The ion with the most oxygen atoms has the prefix
per, the nonmetal root, and the suffix –ate.
o The ion with one fewer oxygen atoms has just the root
and the suffix –ate.
o The ion with two fewer oxygen atoms has just the root
and the suffix –ite.
o The ion with the least (three fewer) oxygen atoms has
the prefix hypo, the root, and the suffix –ite.
 Examples…
 Chlorine oxoanions
o ClO4 – is perchlorate
o ClO3 – is chlorate
o ClO2 – is chlorite
o ClO – is hypochlorite
 Bromine oxoanions
o BrO4 – is perbromate
o BrO3 – is bromate
o BrO2 – is bromite
o BrO – is hypobromite
 Iodine oxoanions
o IO4 – is periodate
o IO3 – is iodate
o IO2 – is iodite
o IO – is hypoiodite
Hydrated Ionic Compounds
Ionic compounds called hydrates have a specific number of
water molecules in each formula unit, which is shown after a
centered dot in the formula and noted in the name by a Greek
numerical prefix and the word hydrate.
The Greek prefixes are…
Example… MgSO4 7H2O would be
named magnesium sulfate heptahydrate
Example… CaSO4 2H2O would be
named calcium sulfate dihydrate
Example… CuSO4 5H2O would be named copper (II) sulfate
pentahydrate
Sample Problem 2.10 p. 63
Determining Names and Formulas of Ionic Compounds
Containing Polyatomic Ions
(a) ClO4 – is perchlorate, which has a 1- charge. Since there
are two perchlorates in the formula, this is a total of 2-.
Since there is only one Fe in the formula, it must be Fe2+
because that would give us a neutral compound (2- and
2+) The name would be iron (II) perchlorate
(b) Sodium is Na+ and sulfite is SO32-. You would need two
Na+ and one sulfite to make a neutral compound, so the
formula would be Na2SO3
(c) Ba2+ is barium and OH – is hydroxide. There are 8 water
molecules for each formula unit. The name would be
barium hydroxide octahydrate
Follow-Up Problem 2.10 p. 63
(a) Cupric tells me Cu2+ and nitrate tells me NO3 - . You need
2 NO3 – for each Cu2+ in the formula to create a neutral
compound. Trihydrate tells me that I have three waters.
The formula would be Cu(NO3)2 3H2O
(b) Zinc is Zn2+ and hydroxide is OH- , so you need 2 OHand 1 Zn2+ to make a neutral compound. The formula
would be Zn(OH)2
(c) Li+ is lithium and CN- is cyanide. The name is lithium
cyanide
Sample Problem 2.11 p. 64
Recognizing Incorrect Names and Formulas of Ionic
Compounds
(a) Prefixes are not used in naming ionic compounds. The
name would simply be barium acetate
(b) You only place parentheses around polyatomic ions
when you are using more than one. Sulfide is S2- not
SO32-. The correct formula would be Na2S
(c) Iron II would be Fe2+ and sulfate is SO42- so they would
simply combine in a 1:1 ratio. The correct formula
would be FeSO4
(d) Cesium is Cs+ and carbonate is CO32-. You would need 2
Cs+ and 1 CO32- to make a neutral compound. The
correct formula would be Cs2CO3.
Follow-Up Problem 2.11 p. 64
(a) Ammonium is NH4+ and phosphate is PO43-. You would
need 3 NH4+ and 1 PO43- to make a neutral compound.
The correct formula would be (NH4)3PO4
(b) When you use more than one polyatomic ion, it needs to
go in parentheses with a subscript. The correct formula
would be Al(OH)3
(c) Mg is magnesium (not manganese) and it is Mg2+ as an
ion (it is in Group 2A) and HCO3- is bicarbonate. The
correct name is magnesium bicarbonate
(d) Do not use the –ic ending with Roman numerals and
NO3- is nitrate. The correct name would be chromium
(III) nitrate or chromic nitrate
(e) Ca is calcium and NO2- is nitrite. The correct name
would be calcium nitrite
Acid Names from Anion Names
When naming acids and writing the formulas for acids we
consider acids as anions that are connected to the number of
hydrogen ions (H+) needed for charge neutrality. The two
common types of acids are binary acids and oxoacids.
 Binary acids form when certain gaseous compounds
dissolve in water and form an aqueous solution. (Hydrogen
plus one other element… two (bi) total elements or an
anion that ends in –ide)
o When naming binary acids, you add the prefix hydro
and the suffix –ic to the nonmetal root and then the
word acid
 HCl (aq) would be hydrochloric acid
 HBr (aq) would be hydrobromic acid
 HI (aq) would be hydroiodic acid
 Oxoacids names are similar to those of the oxoanions,
except for two suffix changes
o Oxoanions that end in –ate become –ic in the acid
 BrO3 – is bromate, so HBrO3 (aq) is bromic acid
o Oxoanions that end in –ite become –ous in the acid
 IO2 – is iodite, so HIO2 (aq) is iodous acid
Sample Problem 2.12 p. 65
Determining Names and Formulas of Anions and Acids
(a) Br – will form a binary acid with H+. The correct name
would be hydrobromic acid / HBr
(b) IO3 – will form an oxoacid with H+. Since IO3 – is iodate,
the acid would be named iodic acid / HIO3
(c) CN – is an anion that ends in –ide (cyanide), so it follows
the binary acid rules. The correct name would be
hydrocyanic acid / HCN
(d) SO42 – is the sulfate ion and will form an oxoacid with
H+. Since sulfate ends in –ate the acid’s name would be
sulfuric acid / H2SO4.
(e) NO2 – is the nitrite ion and will form an oxoacid with H+.
Since nitrite ends in –ite the acid’s name will end in –
ous. The correct name of the acid would be nitrous
acid / HNO2
Follow-Up Problem 2.12 p. 65
(a) Chloric has an –ic ending so that means it has the
chlorate ion… The chlorate ion is ClO3 – and hydrogen
is H+ so the formula would be HClO3.
(b) HF is a binary acid (only two elements) so it follows
those rules. The name is hydrofluoric acid.
(c) Acetic means it has the acetate ion which is C2H3O2 – or
CH3COO –. So the formula would be HC2H3O2 or
CH3COOH
(d) Sulfurous tells us that it has the sulfite ion which is SO32So the correct formula would be H2SO3
(e) BrO – is the hypobromite ion. The correct name would
be hypobromous acid.
Binary Covalent Compounds
Binary covalent compounds are typically formed by the
combination of two nonmetals. Some are so familiar that we use
their common names, such as ammonia (NH3), methane (CH4),
and water (H2O), but most are named systematically.
Rules For Naming Binary Covalent Compounds
 The element with the lower group number in the periodic
table comes first in the name. The element with the higher
group number comes second and is named with its root and
the suffix –ide.
o Exception… When the compound contains oxygen
and any of the halogens chlorine, bromine, or iodine,
the halogen is named first.
o If both elements are in the same group in the periodic
table, the one with the higher period number is named
first.
 Covalent compounds use Greek numerical prefixes to
indicate the number of atoms of each
element. The first element in the
name has a prefix only when more
than one atom of it is present; the
second usually has a prefix.
Example… NF3 is nitrogen
trifluoride
Example… H2O is dihydrogen monoxide
Example… CO2 is carbon dioxide
Example… CCl4 is carbon tetrachloride
Sample Problem 2.13 p. 65
Determining Names and Formulas of Binary Covalent
Compounds
(a) CS2 (1 carbon and 2 sulfurs)
(b) phosphorus pentachloride (1 phosphorus and 5
chlorines)
(c) N2O4 (dinitrogen tetraoxide)
Follow-Up Problem 2.13 p. 65
(a)
(b)
(c)
(d)
sulfur trioxide
silicon dioxide
N2O
SeF6
Sample Problem 2.14 p. 66
Recognizing Incorrect Names and Formulas of Binary Covalent
Compounds
(a) SF4 is sulfur tetrafluoride
(b) dichlorine heptaoxide is Cl2O7
(c) dinitrogen trioxide
Follow-Up Problem 2.14 p. 66
(a) disulfur dichloride
(b) NO
(c) Bromine trichloride
Molecular Masses from Chemical Formulas
 Molecular Mass = sum of atomic masses
o Example… H2O
 Molecular mass = (2 x atomic mass of hydrogen)
+ (1 x atomic mass of oxygen)
 Molecular mass = (2 x 1.008 amu) + (1 x 16.00)
 Molecular mass = 18.02 amu
Formula Mass from Chemical Formulas
 Ionic compounds do not consists of molecules, so the mass
of a formula unit is termed the formula mass
o Example… Ba(NO3)2
 Formula mass = sum of atomic masses
 Formula mass = (1 x atomic mass of barium) + (2
x atomic mass of nitrogen) + (6 x atomic mass of
oxygen)
 Formula mass = (1 x 137.3 amu) + (2 x 14.01
amu) + (6 x 16.00 amu)
 Formula mass = 261.3 amu
Sample Problem 2.15 p. 67
Calculating the Molecular Mass of a Compounds
 Molecular mass P4S3 = sum of atomic masses
o Molecular mass P4S3 = (4 x atomic mass of
phosphorus) + (3 x atomic mass of sulfur)
o Molecular mass P4S3 = (4 x 30.97 amu) + (3 x 32.07
amu)
o Molecular mass P4S3 = 220.09 amu
 Molecular mass NH4NO3 = sum of atomic masses
o Molecular mass NH4NO3 = (2 x atomic mass of N) +
(4 x atomic mass of hydrogen) + (3 x atomic mass of
oxygen)
o Molecular mass NH4NO3 = (2 x 14.01 amu) + (4 x
1.008 amu) + (3 x 16.00 amu)
o Molecular mass NH4NO3 = 80.05 amu
Follow-Up Problem 2.15 p. 67
 Molecular mass of H2O2 = sum of atomic masses
o Molecular mass of H2O2 = (2 x atomic mass of
hydrogen) + (2 x atomic mass of oxygen)
o Molecular mass of H2O2 = (2 x 1.008 amu) + (2 x
16.00 amu)
o Molecular mass of H2O2 = 34.02 amu
 Molecular mass of CsCl = sum of atomic masses
o Molecular mass of CsCl = (1 x atomic mass of
cesium) + (1 x atomic mass chlorine)
o Molecular mass of CsCl = (1 x 132.9 amu) + (1 x
35.45 amu)
o Molecular mass of CsCl = 168.4 amu
 Molecular mass of H2SO4 = sum of atomic masses
o Molecular mass of H2SO4 = (2 x atomic mass of
hydrogen) + (1 x atomic mass of sulfur) + (4 x atomic
mass of oxygen)
o Molecular mass of H2SO4 = (2 x 1.008 amu) + (1 x
32.07 amu) + (4 x 16.00 amu)
o Molecular mass of H2SO4 = 98.09 amu
 Formula mass of K2SO4 = sum of atomic masses
o Formula mass of K2SO4 = (2 x atomic mass of
potassium) + (1 x atomic mass of sulfur) + (4 x atomic
mass of oxygen)
o Formula mass of K2SO4 = (2 x 39.10 amu) + (1 x
32.07 amu) + (4 x 16.00 amu)
o Formula mass of K2SO4 = 174.27 amu
Sample Problem 2.16 p. 67
Using Molecular Depictions to Determine Formula, Name, and
Mass
(a) NaF [this is an ionic compound… both the picture and
the fact that sodium is a metal and fluorine is a nonmetal
makes this an ionic compound. Sodium is Na+ because
sodium is in Group 1A and fluorine is F – because it is in
Group 7A.] Name is sodium fluoride
Formula mass of NaF = sum of atomic masses
Formula mass of NaF = (1 x atomic mass of sodium) +
(1 x mass of fluorine)
Formula mass of NaF = (1 x 22.99 amu) + (1 x 19.00
amu)
Formula mass of NaF = 41.99 amu
(b) NF3 [this is a molecule… both nitrogen and fluorine are
nonmetals so you need to follow the rules for naming
binary covalent compounds]
Name is nitrogen trifluoride
Molecular mass of NF3 = sum of atomic masses
Molecular mass of NF3 = (1 x atomic mass of nitrogen) +
(3 x atomic mass of fluorine)
Molecular mass of NF3 = (1 x 14.01 amu) + (3 x 19.00
amu)
Molecular mass of NF3 = 71.01 amu
Follow-Up Problem 2.16 p. 68
(a) Formula is Na2O [This is an ionic compound… sodium
is a metal and oxygen is a nonmetal]
Name is sodium oxide
Formula mass of Na2O = sum of atomic masses
Formula mass of Na2O = (2 x atomic mass of sodium) +
(1 x atomic mass of oxygen)
Formula mass of Na2O = (2 x 22.99 amu) + (1 x 16.00
amu)
Formula mass of Na2O = 61.98 amu
(b) Formula is NO2 [This is a binary covalent compound…
both nitrogen and oxygen are nonmetals]
Name is nitrogen dioxide
Molecular mass of NO2 = sum of atomic masses
Molecular mass of NO2 = (1 x atomic mass of nitrogen)
+ (2 x atomic mass of oxygen)
Molecular mass of NO2 = (1 x 14.01 amu) + (2 x 16.00
amu)
Molecular mass of NO2 = 46.01 amu
Representing Molecules with Formulas and Models
Mixtures: Classification and Separation
In the natural world, matter usually occurs as mixtures…
 Sample of clean air consists of many elements and
compounds physically mixed together
o O2, N2, CO2, the noble gases [Group 8A (18)], and
water vapor
 Oceans are complex mixtures of dissolved ions and
covalent substances including
o Na+, Mg2+, Cl –, SO42-, O2, CO2, and of course H2O
 Rocks and soils are mixtures of numerous compounds,
including calcium carbonate, silicon dioxide, aluminum
oxide, and iron (III) oxide.
 Living things contain thousands of substances –
carbohydrates, lipids, proteins, nucleic acids, and many
simpler ionic and covalent compounds.
There are two broad classes of mixtures…
 Heterogeneous mixtures
o Its composition is not uniform (varies from one region
to another)
 Homogeneous mixtures
o Its composition is uniform (the same throughout the
entire sample)
Aqueous solutions – solutions in water
Mixtures differ from compounds…
(1) The proportions of the components can vary in a
mixture.
(2) The individual properties of the components of a mixture
are observable.
(3) The components of a mixture can be separated by
physical means.