Topic 2
Atoms, Elements, Molecules,
Ions, and Compounds
Early in the 19th century John Dalton developed atomic theory.
His theory explained the best available experimental data at
that time. His theory has been modified since then with the
discovery of other data, but his work was the initial ground
1
work that we will examine first.
Atomic Theory of Matter
Protons, neutrons, electrons
• Postulates of Dalton’s Atomic Theory
Later found indivisible
to be untrue.
1.)All matter is composed of indivisible atoms. An
atom is an extremely small particle of matter that
retains its identity during chemical reactions.
2.)An element is a type of matter composed of only
one kind of atom, each atom of a given element
having the same properties. Mass is one such
property. Thus the atoms of a given element have a
Later found all atoms of the same element
characteristic mass.
does not have to have the same mass.
Atoms have different isotopes that have the same #
protons but different # neutrons and hence different
mass. Note: #protons gives identity of atom.
2
Atomic Theory of Matter
• Postulates of Dalton’s Atomic Theory
3.) A compound is a type of matter composed of
atoms of two or more elements chemically combined
in fixed proportions.
– The relative numbers of any two kinds of
atoms in a compound occur in simple ratios.
– Water, for example, consists of hydrogen
and oxygen in a 2 to 1 ratio (2H: 1O) for all
molecules of water.
3
Atomic Theory of Matter
• Postulates of Dalton’s Atomic Theory
i.e., solid sodium
mixed with chlorine
gas forms a new
substance, salt, with
totally different
properties from the
starting materials.
4.) A chemical reaction consists of the
rearrangements of the atoms present in the
reacting substances to give new chemical
combinations present in the substances formed by
the reaction (new chemical with different
properties).
2 Na (s) + Cl2 (g)  2 NaCl (s)
5.) Atoms are not created, destroyed, or broken
into smaller particles by any chemical reaction.
Once again, later
found indivisible
to be untrue.
Protons, neutrons, electrons
4
Atomic Theory of Matter
• The Structure of the Atom
– Although Dalton postulated that atoms were
indivisible, experiments at the beginning of the
1900’s showed that atoms themselves consist of
particles.
– Experiments by Ernest Rutherford in 1910 showed that
the atom was mostly “empty space.”
5
Atomic Theory of Matter
– These experiments showed that the atom consists of two
kinds of particles: a nucleus, the atom’s central core,
which is positively charged and contains most of the
atom’s mass, and one or more electrons.
– Electrons are very light, negatively charged particles
that exist in the region around the atom’s positively
charged nucleus.
Nucleus
(+)
-
e
6
Atomic Theory of Matter
– In 1897, the British physicist J. J. Thompson
conducted a series of experiments that showed
that atoms were not indivisible particles but
instead made of smaller particles.
– From his experiments, Thompson calculated the
ratio of the electron’s mass, me, to its electric
charge, e.
7
Atomic Theory of Matter
– In 1909, U.S. physicist, Robert Millikan obtained the
charge on the electron (1.602 x 10-19 C).
– These two discoveries (Millikan and Thompson)
combined provided us with the electron’s mass of
9.109 x 10-31 kg, which is more than 1800 times
smaller than the mass of the lightest atom
(hydrogen) thereby proving that the atom is made
up of smaller particles.
– These experiments showed that the electron was
indeed a subatomic particle.
8
Atomic Theory of Matter
The nuclear model of the atom.
– Ernest Rutherford, a British physicist, put forth the idea
of the nuclear model of the atom in 1911, based on
experiments done in his laboratory by Hans Geiger and
Ernest Morrison.
– Rutherford’s famous gold leaf experiment gave
credibility to the theory that the majority of the mass of
the atom was concentrated in a very small nucleus.
– Positively charged alpha particles were directed at a
metal foil. Only 1/8000 were deflected indicating that
the nucleus was extremely small and positively
charged. Only those alpha particles that directly hit the
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nucleus were deflected; the rest passed through.
Atomic Theory of Matter
– Most of the mass of an atom is in the nucleus;
however, the nucleus occupies only a very small
portion of the space in the atom.
– The diameter of an atom is approximately 100 pm
while the diameter of the nucleus is approximately
0.001 pm. For comparison, if an atom was 3 miles in
diameter, the nucleus would be the size of a golf ball.
– The nucleus of an atom is composed of two different
kinds of particles: protons (+) and neutrons (neutral).
– An important property of the nucleus is its positive
electric charge.
10
Atomic Theory of Matter
– A proton is the nuclear particle having a positive
charge equal to that of the electron’s (a “unit”
charge) and a mass more than 1800 times that of
the electron. It is for this reason that we refer to H
as a pure proton.
– The number of protons in the nucleus of an atom is
referred to as its atomic number (Z) and gives the
identity of an element. All species that have same
#p have the same properties.
neutral species: #p = #e-
H Z=1
1p, 1e-
Na Z=11 11p, 11e-
Cl Z=17 17p, 17e-
Na+ Z=11 11p, 10e-
Cl Z=17 17p, 18e-
#p remains constant in species;
#e- can vary and dictates the charge of species
-
+ charge, more p than e- charge, more e- than p
11
Atomic Theory of Matter
– An element is a substance whose atoms all have the
same atomic number (Z). The #protons defines the
identity of an atom and can be found on the periodic
table (large number in top of element box).
– The neutron is a nuclear particle having a mass
almost identical to that of a proton, but no electric
charge. The charge of the nucleus comes from
the #protons. The atoms may have different
masses because of different #neutrons (isotopes).
– Summary of masses and charges of the three
fundamental particles:
particle
mass, kg
charge, C
relative location
charge
electron, e-
9.109 x 10-31
-1.602 x 10-19
-1
outside nucleus
proton, p
1.6726 x 10-27
1.602 x 10-19
+1
nucleus
neutron, n
1.6749 x 10-27
0
0
nucleus
12
The mass number (A) is the total number of protons
and neutrons in a nucleus.
A = #p + #n = Z + #n
How many neutrons does sodium 23 have?
A = 23, Z = 11 (number on periodic table)
#n = 23 - 11 = 12
A = Z + #n
23 = 11 + #n
A nuclide is an atom characterized by a definite atomic
number and mass number.
The shorthand notation for a nuclide consists of its
symbol with the atomic number, Z, as a subscript on the
left and its mass number, A, as a superscript on the left.
A
Z
sodium  23
E
23
11
Na
11p, 12 n, 11e-
23
11
+
Na
11p, 12 n, 10e-
13
What is the nuclide symbol for a nucleus that consists of 17 protons,
18 neutrons, and 17 electrons?
What’s the element? 17 p  atomic number on periodic table for chorine
A = #p + #n = 17 + 18 = 35
35
17
Cl
Note: #e- = #p; therefore, neutral species
How many protons, neutrons, and electrons are in the following
nucleus
80

35 p
45 n 36 eBr
35
A = #p + #n
80 = 35 + #n
#e-
#n = 80 – 35 = 45 n
Note:
> #p;
therefore, negatively charged species
35e- + one additional ebased on -1 charge = 36 e-
HW 9
HW 10
code for
both: proton
14
Atomic Theory of Matter
– Isotopes are atoms whose nuclei have the same
atomic number (Z) but different mass numbers (A);
that is, the nuclei have the same number of protons
but different numbers of neutrons thereby causing
them to have different masses.
– Chlorine, for example, exists as two isotopes:
chlorine-35 and chlorine-37.
35
17
Cl
Frac abund = 75.771%
Mass = 34.97 amu
17p, 18 n, 17e-
37
17
Cl
Frac abund = 24.229%
Mass = 36.97 amu
17p, 20 n, 17e-
– The fractional abundance is the fraction of a sample of
atoms that is composed of a particular isotope.
(0.75771)(34.97 amu) + (0.24229) (36.97 amu) = 35.45 amu
Note: The mixture of isotope masses make up the actual mass of the element given on periodic
15
table. Cl has a mass of 35.45 amu which is based on the two isotopes of Cl-35 and Cl-37.
Atomic Weights
Calculate the atomic weight of boron, B, from
the following data:
ISOTOPE
B-10
B-11
ISOTOPIC MASS (amu)
FRACTIONAL ABUNDANCE
10.013
11.009
0.1978 (19.78%)
0.8022 (80.22%)
Note: fractional
abundances must
add to 1 (100%)
B-10: 10.013 amu x 0.1978 = 1.980
B-11: 11.009 amu x 0.8022 = 8.831
10.811 = 10.811 amu
( = atomic wt.)
Note: mass on periodic table matches
10.811 amu (weighted average of isotopes)
HW 11
code: amu
16
Atomic Weights
Dalton’s Relative Atomic Masses
– Since Dalton could not weigh individual atoms, he
devised experiments to measure their masses
relative to the hydrogen atom.
– Hydrogen was chosen as it was believed to be the
lightest element. Daltons assigned hydrogen a
mass of 1 (1 Dalton = mass of H).
– For example, he found that carbon weighed 12
times more than hydrogen. He therefore assigned
carbon a mass of 12 ( mass of carbon = 12
Daltons).
17
Atomic Weights
Dalton’s Relative Atomic Masses
– Dalton’s atomic weight scale was eventually replaced
in 1961, by the present carbon–12 mass scale.
– One atomic mass unit (amu) is, therefore, a mass unit
equal to exactly 1/12 the mass of a carbon–12 atom.
– On this modern scale, the atomic weight of an element
is the average atomic mass for the naturally occurring
element, expressed in atomic mass units. Periodic table
is based on atom mass with units of amu.
Na - 23.1 amu  mass of 1 atom of sodium
18
The Periodic Table
In 1869, Dmitri Mendeleev discovered that if the known
elements were arranged in order of atomic mass (A),
they could be placed in horizontal rows such that the
elements in the vertical columns had similar properties.
– periodic table - tabular arrangement of elements in
rows and columns, highlighting the regular repetition of
properties of the elements.
– periodic law – states that certain sets of physical and
chemical properties recur at regular intervals
(periodically) when the elements are arranged
according to increasing atomic number (Z).
– Note: eventually changed from atomic mass to atomic
number because of a couple of anomalies.
19
Figure: A modern form of the periodic table.
anomalies
20
The Periodic Table
Periods and Groups
– A period consists of the elements in any one
horizontal row of the periodic table.
– A group consists of the elements in any one column
of the periodic table (similar properties/structure).
– The groups are usually numbered (North American
uses roman numbers and A/B; IUPAC 1-18).
– The eight “A” groups are called main group (or
representative) elements.
21
The Periodic Table
• Periods and Groups
– The “B” groups are called transition elements.
– The two rows of elements at the bottom of the
table are called inner transition elements.
– Elements in any one group have similar properties
because their outer shells have the same number
of valence electron (discuss in later sections).
22
The Periodic Table
• Periods and Groups
– The elements in group IA (except H) - alkali metals
– The elements in group IIA - alkaline earth metals,
– The group VIIA elements - halogens
– The group VIIIA elements – noble gases (monoatomic)
– Diatomic elements – H2, N2, O2, F2, Cl2, Br2, I2
– Most species are solids at room temperature; H2,
N2, O2, F2, Cl2, and noble gases are gases; Br2
and Hg are liquids.
23
Metals, Nonmetals, and Metalloids – generally,
left of staircase are metals, touching staircase
are metalloids, right of staircase are nonmetals.
This is important for determining bond type,
using proper terminology, and making decisions.
Metallic character
HW 12
nonmetals
Metallic character
metals
code:
table
24
Chemical Formulas; Molecular and
Ionic Substances
The chemical formula of a substance is a
notation using atomic symbols with subscripts
to convey the relative proportions of atoms of
the different elements in a substance.
– aluminum oxide, Al2O3
2Al:3O ratio
– sodium chloride, NaCl
1Na:1Cl ratio
– calcium nitrate, Ca(NO3)2
1Ca:2NO3- ratio
or 1Ca:2N:6O ratio
25
Chemical Formulas; Molecular and
involves covalent bond – share electrons
Ionic Substances
between atoms – typically nonmetal/nonmetal
involves ionic bond – transfer electrons between
atoms – attraction between charged particles –
+ Na Cl
typically metal/nonmetal or polyatomic ions
C:C
• Molecular substances
– A molecule is a definite group of atoms that are
chemically bonded together through sharing of
electrons (covalent bonding, generally nonmetalnonmetal including H).
– A molecular substance is a substance that is
composed of molecules, all of which are alike.
– A molecular formula gives the exact number of
atoms of elements in a molecule (i.e. C2H6O).
– Structural formulas show how the atoms are bonded
to one another in a molecule.
i.e. ethanol (C2H6O) has a structural formula of CH3CH2OH
26
Ionic substances
– Although many substances are molecular, others are
composed of ions (charged particles) that have
transferred electrons and have ionic bonding; occurs
generally with metal-nonmetal interactions.
– An ion is an electrically charged particle obtained from
an atom or chemically bonded group of atoms by
adding or removing electrons.
+
Na  1e  Cl
– Sodium chloride is a substance made up of ions.
Na+ Cl-
27
Chemical Formulas; Molecular
and Ionic Substances
Ionic substances
– The formula of an ionic compound is written by giving
the smallest possible whole-number ratio of different
ions in the substance.
– The formula unit of the substance is the group of
atoms or ions explicitly symbolized by its formula.
C : O
Covalent bond (share e-)
Ionic bond (transfer e-/
attraction charged particles
nm –nm
m – nm and charged ions
Molecules
Formula unit
Molecular substance
Ionic substance
Molecular formula
formula
+ -
Na Cl
28
Ionic substances
– When an atom gains extra electrons, it becomes a
negatively charged ion, called an anion (more
electrons than protons). i.e, Cl– An atom that loses electrons becomes a positively
charged ion, called a cation (more protons than
electrons). i.e., Na+
– An ionic compound is a compound composed of
cations and anions.
Answer the following questions for species below:
ionic or molecular substance; formula unit or molecule; ionic or covalent bonds involved?
NaCl
CaBr2
Na2SO4
CO2
ionic substance; formula unit; ionic bond
ionic substance; formula unit; ionic bonds
ionic substance; formula unit; ionic and covalent bonds in SO42molecular substance; molecule; covalent bonds
29
Ions in Aqueous Solution
Many (not all) ionic compounds (ionic
bond/m-nm) dissociate into independent
ions when dissolved in water
NaCl (s)  Na+(aq) + Cl-(aq)
Soluble salt
charges particles
Soluble ionic compounds dissociate 100%
- referred to as strong electrolytes –
breaks into charged particles until reaches
saturation point.
30
Ions in Aqueous Solution
Most molecular (covalent bond/nm-nm)
compounds dissolve but do not dissociate into
ions, exception acids.
C6H12O6 (s)  C6H12O6 (aq)
no charges particles; remains whole
These compounds are referred to as
nonelectrolytes; no charged particles; soluble
to saturation point but no ions formed.
How would sodium sulfate dissolve based on bonding?
Na2SO4 (s)  2Na+(aq) + SO42-(aq)
ionic bond dissociates while covalent bonds in
sulfate remain intact
31
Chemical Substances;
Formulas and Names
Ionic compounds
– Most ionic compounds contain metal and nonmetal
atoms (as well as polyatomic ions); for example, NaCl.
– You name an ionic compound by giving the name of
the cation followed by the name of the anion with -ide.
Sodium chloride, NaCl
Calcium Iodide, CaI2
Potassium Bromide, KBr
– We give the monatomic ion name for the cations and
anions when naming compounds. A monatomic ion is
32
an ion formed from a single atom.
How do we get the charge for ions?
Rules for predicting charges on monatomic ions
– Most of the main group metals form cations with the
charge equal to their roman group number.
– The charge on a monatomic anion for a nonmetal
equals the roman group number minus 8.
– Most transition elements form more than one ion, each
with a different charge (exceptions Cd2+, Zn2+, Ag+).
– Other important elements with variable charge
Pb4+, Pb2+ Sn4+, Sn2+ As5+, As3+ Sb5+, Sb3+
0
43+ 4+ 3- 2- 1-
1+
2+
varies
33
Rules for naming monatomic ions
– Monatomic cations are named after the element. For
example, Al3+ is called the aluminum ion.
– If there is more than one cation of an element (charge),
a Roman numeral in parentheses denoting the charge
on the ion is used. This often occurs with transition
elements.
Na+ sodium ion
Ca2+ calcium ion
Fe2+ iron (II) ion
Fe3+ iron (III) ion
Older name: higher ox state (charge) – ic, / lower, -ous
Fe3+ ferric ion
Fe2+ ferrous ion
Cu2+ cupric ion
Cu+ cuprous ion
Hg2+ mercuric ion Hg22+ mercurous ion
also done with Pb4+, Pb2+ ; Sn4+, Sn2+ ; As5+, As3+ ; Sb5+, Sb3+.
For the names of the monatomic anions, use the stem
name of the element followed by the suffix – ide. For
34
example bromine, the anion is called bromide ion, Br .
The formula of an ionic compound is written by giving the
smallest possible whole-number ratio of different ions in
Based on the
the substance.
ions and charges formula
charge of the ions
Sodium chloride
Na+ ClNaCl
and balancing the
1Na+ = 1+ charge
1Cl- = 1- charge
balanced
Roman number tells
charge of transition metal
overall charge on
the compound by
adjusting the
Iron (III) sulfate
Fe3+ SO42Fe
(SO
)
2
4 3 number of ions, a
2Fe3+ = 6+ charge
2
3
formula is written.
3SO42- = 6- charge
balanced
Note the sum of all
the charges must
Cr2O3
Chromium (III) oxide Cr3+ O2equal zero, and you
do not display the
charges in the final
formula.
Calcium nitrate
Ca2+ NO3Ca(NO3)2 Generally, you can
crisscross the
charge of one ion
as the subscript on
+
3Na3PO4 the second ion,
Sodium phosphate
Na PO4
reducing when
HW 13 & 14
code for both: formula
possible.
Strontium oxide
Sr2+ O2-
SrO
35
Naming Ionic Binary Compounds
NaF
-
LiCl
-
MgO
-
The charge on Mn must be 2+
to balance out the 2Br- charges.
MnBr2
-
To name a compound, you
must know if it is a molecular
sodium fluoride or ionic compound so that you
know which rules to follow. If
you have a metal-nonmetal (or
ion), it is an ionic
lithium chloride polyatomic
compound where you name
the metal first then the
nonmetal with changing the
magnesium oxide ending to –ide. If it is a
transition metal, you must
include the charge of the metal
(Roman numbers).
manganese (II) bromide
The Roman number 3 tells us the charge on
Co is 3+ which helps us determine the
formula knowing that O is 2-.
If you have nonmetalnonmetal, it is a molecular
compound which we haven’t
discussed yet.
Co2O3
-
cobalt (III) oxide
CuCl2
-
copper (II) chloride or cupric chloride36
Chemical Substances; Formulas and Names
Polyatomic ions
– A polyatomic ion is an ion consisting of two or more
atoms chemically bonded together and carrying a net
electric charge. We name the compounds the same
way we just discussed except each polyatomic ion has
a particular name.
– Books typically have a table that lists common
polyatomic ions. Most are oxo anions – consists of
oxygen with another element (central element).
NO3- nitrate
SO42- sulfate
NO2- nitrite
SO32- sulfite
Most groups have –ate, -ite endings and differ by #O.
Mn, Br, Cl, I have per- -ate, -ate, -ite, hypo- -ite.
37
Ions You Should Know
Polyatomic ions
• NH4+ - Ammonium
• OH- - Hydroxide
• CN- - Cyanide
• SO42- - Sulfate
• SO32- - Sulfite
• ClO4- - perchlorate
• ClO3- - chlorate
• ClO2- - chlorite
• ClO- - hypochlorite
• Hg22+ - mercury (I) or
mecurous
• S2O32- - thiosulfate
• SCN- - thiocyanate
• CNO- - cyanate
• MnO4- - permanganate
•
•
•
•
•
•
•
•
•
•
•
•
•
•
O22- - Peroxide
PO43- - Phosphate
PO33- - Phosphite
CO32- - Carbonate
HCO3- - Bicarbonate or
Hydrogen Carbonate
N3- - azide
NO3- - nitrate
NO2- - nitrite
C2H3O2- or CH3COO - acetate
Cr2O72- - dichromate
CrO42- - chromate
C2O42- - oxalate
HSO4- - bisulfate or hydrogen
sulfate
38
H2PO4- - dihydrogen
phosphate
SnSO4
Na2SO3
tin (II) sulfate or stannous sulfate
sodium sulfite
Ca(ClO)2
calcium hypochlorite
Ba(OH)2
barium hydroxide
KClO4
potassium perchlorate
Cr2(SO4)3
chromium (III) sulfate
Note: Not a polyatomic ion;
monoatomic anion of N.
Mg3N2
magnesium nitride
Fe3(PO4)2
iron (II) phosphate or ferrous phosphate
Ti(NO3)4
titanium (IV) nitrate
39
Chemical Substances;
Formulas and Names
Molecular compounds
– Binary compounds composed of two nonmetals are
usually molecular and are named using a prefix
system (name same as ionic except must indicate
how many atoms are present using mono, di, tri,
etc.). No charges (share electrons) involved with
molecular compounds, but we typically put more
metallic compound first.
Which way is the correct way to write the following formula
based on putting the more metallic compound first?
NF3
F3N
40
Chemical Substances;
Formulas and Names
Binary molecular compounds
– The name of the compound has the elements in the
order given in the formula.
– You name the first element using the exact element
name.
– Name the second element by writing the stem name of
the element with the suffix “–ide.”
– If there is more than one atom of any given element,
you add a prefix (di, tri, tetra, penta, hexa, hepta, octa,
etc.)
41
• Binary molecular compounds
– N2O3
dinitrogen trioxide
– SF4
sulfur tetrafluoride
– ClO2
chlorine dioxide
– SF6
sulfur hexafluoride
– Cl2O7
Drop the “a” on prefix if you encounter
double vowel in name.
dichlorine heptoxide
Since this is a gas, we name using molecular
rules; however, if acid we have other rules.
To name a compound, you
must know if it is a molecular
or ionic compound so that you
know which rules to follow. If
you have a metal-nonmetal (or
polyatomic ion), it is an ionic
compound where you name
the metal first then the
nonmetal with changing the
ending to –ide. If you have
nonmetal-nonmetal, it is a
molecular compound which
you do similarly as the ionic
compound except that you
must use prefixes to indicate
the number of atoms.
– HCl (g)
hydrogen chloride
Name this compound but think about bonding:
magnesium chloride; ionic comp, no prefix
MgCl2
Older names: water - H2O, ammonia – NH3,
hydrogen sulfide – H2S, nitric oxide – NO, hydrazine – N2H4
42
Chemical Substances;
Formulas and Names
Acids
– Acids are traditionally defined as compounds with a
potential H+ as the cation.
– Binary acids consist of a hydrogen ion and any single
anion in aqueous solution. For example, HCl (aq) is
hydrochloric acid. Binary acid: hydrostemic acid
– An oxoacid is an acid containing hydrogen, oxygen,
and another element. An example is HNO3, nitric acid.
The oxoacids are a derivation of the oxoanions we
discussed earlier.
43
oxoacids
Anion prefix/suffix
per- -ate ion
-ate ion
-ite ion
hypo- -ite ion
NO3- nitrate ion
NO2- nitrite ion
ClO4- perchlorate ion
If you learn the oxoanions, you can
easily adapt to naming the oxoacids:
-ate  -ic and –ite  -ous
acid prefix/suffic
per- -ic acid
-ic acid
-ous acid
hypo- -ous acid
HNO3 nitric acid
HNO2 nitrous acid
HClO4 perchloric acid
For some species there is a change in spelling in the name.
SO42- sulfate ion
H2SO4 sulfuric acid
PO43- phosphate ion
H3PO4 phosphoric acid
44
Chemical Substances; Formulas and
Names
Hydrates
– A hydrate is a compound that contains water
molecules weakly bound in its crystals.
– Hydrates are named from the anhydrous (dry)
compound, followed by the word “hydrate” with a prefix
to indicate the number of water molecules per formula
unit of the compound.
CuSO4. 5H2O
copper(II)sulfate pentahydrate
Magnesium sulfate heptahydrate MgSO4 . 7H2O
HW 15 - 18
code for all: names
45
Chemical Substances; Formulas and
Names
Naming simple compounds
– Chemical compounds are classified as organic or
inorganic.
– Organic compounds are compounds that contain
carbon combined with other elements, such as
hydrogen, oxygen, and nitrogen.
– Inorganic compounds are compounds composed of
elements other than carbon.
46
Chemical Formulas; Molecular and
Ionic Substances
Organic compounds
– An important class of molecular substances that contain
carbon is the organic compounds.
– Organic compounds make up the majority of all known
compounds.
– The simplest organic compounds are hydrocarbons compounds containing only hydrogen and carbon.
– Common examples include methane, CH4, ethane, C2H6,
and propane, C3H8.
47
Classifying Compounds
Organic vs. Inorganic
• in the 18th century, compounds from living
things were called organic; compounds from
the nonliving environment were called
inorganic
• organic compounds easily decomposed and
could not be made in 18th century lab
• inorganic compounds very difficult to
decompose, but able to be synthesized
48
Modern Classifying Compounds
Organic vs. Inorganic
• today we commonly make organic compounds
in the lab and find them all around us
• organic compounds are mainly made of C and
H, sometimes with O, N, P, S, and trace
amounts of other elements
• the main element that is the focus of organic
chemistry is carbon
49
Carbon Bonding
• carbon atoms bond almost exclusively covalently
– compounds with ionic bonding C are generally
inorganic
• when C bonds, it forms 4 covalent bonds
– 4 single, 1 double + 2 singles, 2 double, or 1 triple + 1 single
– carbon is unique in that it can form limitless chains of
C atoms, both straight and branched, and rings of C
atoms
50
Examples of Carbon Compounds
51
Classifying Organic Compounds
• there are two main
categories of organic
compounds, hydrocarbons
and functionalized
hydrocarbons
• hydrocarbons contain
only C and H
• most fuels are mixtures of
hydrocarbons
52
Classifying Hydrocarbons
• hydrocarbons containing only single bonds
are called alkanes
• hydrocarbons containing one or more C=C
double bonds are called alkenes
• hydrocarbons containing one or more CC
triple bonds are called alkynes
• hydrocarbons containing C6 “benzene” ring
are called aromatic
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54
Naming Straight Chain Hydrocarbons
• consists of a base name to indicate the number of
carbons in the chain, with a suffix to indicate the
class and position of multiple bonds
– suffix –ane for alkane, –ene for alkene, –yne for alkyne
Base Name
No. of C
Base Name
No. of C
meth-
1
hex-
6
eth-
2
hept-
7
prop-
3
oct-
8
but-
4
non-
9
pent-
5
dec-
10
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Functionalized Hydrocarbons
• functional groups are non-carbon groups that
are on the molecule
• substitute one or more functional groups
replacing H’s on the hydrocarbon chain
• generally, the chemical reactions of the
compound are determined by the kinds of
functional groups on the molecule
56
Functional Groups
57
Chemical Reactions: Equations
Writing chemical equations
– A chemical equation is the symbolic representation of a
chemical reaction in terms of chemical formulas.
– For example, the burning of sodium and chlorine to
produce sodium chloride is written
2Na  Cl 2  2NaCl
Reactants (consumed)

Products (produced)
– The reactants (consumed; left side of reaction) are
starting substances in a chemical reaction. The arrow
means “yields.” The formulas on the right side of the
arrow represent the products (produced).
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Chemical Reactions: Equations
Writing chemical equations
– In many cases, it is useful to indicate the states of the
substances in the equation (s, g, l, aq).
– When you use these labels, the previous equation
becomes
2Na(s )  Cl 2 ( g )  2NaCl(s )
We write above the arrow any conditions for the reaction such as
pressure, catalyst, heat, etc. A reaction gives a recipe for the amount
of reactants needed to produce the amount of products. Species with
no coefficient have an understood coefficient of 1.
59
Chemical Reactions: Equations
Writing chemical equations
– The law of conservation of mass dictates that the
total number of atoms of each element on both sides of
a chemical equation must match. The equation is then
said to be balanced.
CH 4  2 O 2  CO 2  2 H 2O
We must have the same number of atoms on both sides for a reaction to be considered
balanced and obeying the law of conservation of mass. To balance a reaction:
1. First, balance the atoms for elements that occur in only one substance on each side of the
reaction. In this problem, O is involved with two substances on the product side;
therefore, I will wait on balancing O until later. C & H are only in one species on both
sides so I will balance them first. C needs no changes because there are one on each
side, but H needs a 2 in front of H2O to balance the 4H on the reactants side.
2. Now that we have changed the coefficient of one of the O on the product side, it is easier
to balance the O. We determine that we need a 2 coefficient on the O2 to balance the O
on both sides at 4. Now the equation is balanced with 1C, 4O, and 4H on both sides.60
Chemical Reactions: Equations
Caution: For formulas that have subscripts, you
must account for all atoms especially when
dealing with parentheses for polyatomic species.
For example,
Fe2(SO4)3:
has 2-Fe, 3x1 = 3-S, 3x4 = 12-O
Caution: Remember that you can’t change the
subscripts in formulas to balance equations; you
may only change coefficients. If you change the
subscripts, you are changing the substance.
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Chemical Reactions: Equations
O2 
P4 
2 PCl 3  2 POCl
6 N 2O 
P4 O 6 
3
6N2
Technique to handle odd numbers: determine number needed and divide by subscript of
species. Next, you multiple the entire equation by the subscript to obtain whole numbers.
[
2 As 2 S 3 
2 As 2 S 3 
Ba(NO
HW 19
3
O2 
As 2 O 3 
9 O2 
)2 
2
NaCl
3 SO 2
2 As 2 O 3 

code: balance
2 NaNO
3
]
6 SO 2

BaCl
2
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