Science 1206
Unit 2: Chemistry
Chemicals in Action
Unit 2: Chemistry
CHEMISTRY TERMINOLOGY
Chemistry is the study of matter, its properties, and its changes
or transformations.
Matter is anything that has mass and takes up space. Matter can be
categorized into two distinct types of materials. These are mixtures
and pure substances.
Mixtures
Two or more substances, combined in varying proportions - each
retaining its own specific properties.
The components of a mixture can be separated by physical means,
i.e. without the making and breaking of chemical bonds.
Examples: Air, table salt thoroughly dissolved in water, milk,
wood, and concrete.
There are two types of mixtures.
Heterogeneous Mixture A mixture in which the properties and
composition are not uniform throughout the sample.
A mixture with more than one visible substance.
Examples: Pizza and concrete.
Homogeneous Mixture
Mixture in which properties and composition are
uniform throughout the sample.
A mixture with only one visible substance. Such
mixtures are termed solutions.
Examples: Air and table salt thoroughly dissolved in water.
Pure Substances
A substance with constant composition throughout.
Examples: Table salt (sodium chloride, NaCl), sugar
(sucrose,C12H22O11), water (H2O), iron (Fe), copper
(Cu), and oxygen (O2).
There are two types of pure substances
- Element
- Compound
Element
A substance that cannot be separated into two or more substances by ordinary
chemical (or physical) means.
We use the term ordinary chemical means to exclude nuclear reactions.
Elements are composed of only one kind of atom.
Examples: Iron (Fe), copper (Cu), and oxygen (O2).
Compound
A substance that contains two or more elements, in definite
proportion by weight.
The composition of a pure compound will be invariant, regardless of
the method of preparation. Compounds are composed of more than
one kind of atom. The term molecule is often used for the smallest
unit of a compound that still retains all of the properties of the
compound.
Examples: Table salt (sodium chloride, NaCl), sugar
(sucrose,C12H22O11), and water (H2O).
Read pgs. 172-174
Questions 1-3 pg.175
Chemicals and Chemical Change
Properties of Matter - matter has specific chemical and physical properties.
Physical property a characteristic of a substance.
A substance can have specific physical properties
such as state, colour, texture, shape, mass, odour,
density, taste, etc.
Physical change
a change in the form of a substance.
For instance, from solid to liquid or liquid to gas or solid to
gas, without changing the chemical composition of the
substance. As we will see later, chemical bonds are not
broken in a physical change.
Examples: Boiling of water and the melting of ice.
Chemical property a characteristic behaviour that occurs when a substance
changes to a new substance by reorganization of the atoms.
Chemical change
The change itself involves a chemical reaction.
Example 1
Iron will react with oxygen to produce iron oxide(rust).
Iron + oxygen ------> iron oxide
Fe + O2 ------> Fe2O3
Example 2
Water will decompose, if induced by an electric
current, to produce gaseous hydrogen and gaseous
oxygen.
Water ------> hydrogen + oxygen
H2O ------> H2 + O2
The starting materials in chemical reactions are called reactants,
and the new materials produced are called products.
Evidence of Chemical Change
We use several pieces of evidence to verify if a chemical change has
occurred.
These include:
1. Colour changes
potassium iodide + lead nitrate
KI + Pb(NO3)2
colourless reactants
-----> potassium nitrate + lead iodide
-----> KNO3 + PbI2
-----> coloured products
2. Solid (precipitate) produced
potassium iodide + lead nitrate
KI + Pb(NO3)2
3. Gas produced
zinc + hydrochloric acid
Zn + 2HCl
-----> potassium nitrate + lead iodide
-----> KNO3 + PbI2
-----> lead iodide will ‘settle out’
-----> zinc chloride + hydrogen gas
-----> ZnCl2 + H2
-----> hydrogen gas bubbles out of solution
4. Energy change
Hydrocarbon + oxygen
CH4 + O2
-----> carbon dioxide + water vapour
-----> CO2 + H2O
-----> heat given off
- energy given off (Exothermic)
- energy required (Endothermic)
Chemical Tests (pg. 174)
Chemists can use chemical tests, or distinctive chemical reactions, to identify
unknown gases or other substances.
the presence of oxygen gas,
for example, is tested using a glowing splint.
The glowing splint ‘bursts into flames’ when
placed in a container that has concentrated
oxygen gas present.
the presence of hydrogen gas,
the presence of carbon dioxide gas,
for example, is tested using a flaming
splint. The flaming splint causes the gas to
‘pop’ when placed in a container that has
concentrated hydrogen gas present.
for example, is tested using a limewater
solution. The carbon dioxide gas is bubbled
into the limewater solution. If the solution
turns milky the gas is carbon dioxide.
Questions 4-8 pg.175
Chemicals and Safety (See pg. 176 & 658)
Hazardous Household Product Symbols
(HHPS)
Elements and the Periodic Table
The Periodic Table
a structured arrangement of elements that helps us to explain and
predict physical and chemical properties.
the table was designed by Dmitri Mendeleev
represents on of the most important advancements in science
The periodic table is generally
arranged
• with metals towards the left side of
the table (staircase).
• nonmetals on the right side of the
table (staircase).
• hydrogen the exception - found in
the top left corner but behaves
mostly as a nonmetal
Metals
Elements with the following
properties
Property
Property of Metal
lustre
shiny
malleability
malleable
conductivity
conductors
reactivity with acids
yes
state at room
temperature solids,
except mercury
Non-Metals
Elements with the following
properties
State - the states, at room
temperature, of all the elements
are indicated through a text or
colour code, depending on
individual periodic table.
Property
Property of Metal
lustre
dull
malleability
brittle
conductivity
mostly insulators
reactivity with acids
no
state at room
solids, liquids, gases
Periodic Table Sub Groups
Periods
horizontal rows are called periods
represent the number of energy levels in atoms of an
element
Period 1 - 1 energy level
Period 2 - 2 energy levels and so on
Families
groups of elements in the same vertical column
elements with similar physical and chemical properties
some of these groups include
Family Sub Groups
Alkali Metals
also called Group IA elements
far left column of the periodic table
very reactive
form compounds that are mostly white solids and are very
soluble in water
all shiny, silvery metals
Alkaline Earth
also called Group IIA elements
include magnesium (Mg) and calcium (Ca)
shiny, silvery metals
form compounds often insoluble in water
Noble Gases
far right column of the periodic table
also called Group VIIIA elements
include helium (He) and neon (Ne)
inert - generally do not form compounds
Halogens
second column from the right
nonmetallic elements
also called Group VIIA elements
all poisonous elements
react readily with alkali metals
include chlorine (Cl), bromine (Br)
Hydrogen
a very special element. It is located in Group IA and
Group VIIA. It sometimes behaves as a metal, but
behaves mostly like a nonmetal.
Elements and Atomic Structure
Atoms are the basic building blocks of all forms of matter in our
known universe. Everything we see and touch is made up of atoms.
An atom is the smallest particle of an element that cannot be
further broken down by physical or chemical means. The atom also
has a very particular structure.
Bohr-Rutherford Model of the Atom
There are many models of the atom that have
been proposed throughout history. The BohrRutherford Model of the atom suggests there
are three sub-atomic particles in the atom.
These are located in two regions of the atom,
namely
- the nucleus
- the electron cloud
Atom
neutral in charge
# of protons equal to the atomic number
(ex. Sodium has 11 protons)
# of neutrons equals
the atomic # - atomic mass #
(eg. Sodium has 23 - 11 = 12 neutrons)
the number of electrons is equal to the number of protons
(Ex. Sodium has 11 electrons)
Nucleus
dense positively charged core of the atom
composed of protons and neutrons.
protons are relatively small in volume (size)
protons are heavy (high mass) particles
protons are positively charged particles
neutrons are relatively small in size
neutrons are heavy (high mass) particles
neutrons are neutral particles
the nucleus has an overall net positive charge
the nucleus makes up the vast majority of the mass of the atom
Electron cloud
region surrounding the nucleus where electrons are found
electrons have relatively large volumes(size)
electrons have very little mass
electrons are negatively charged particles
electrons are located in energy levels
first energy level
two electrons
second energy level eight electrons
third energy level
eight electrons
the electron cloud has an overall net negative charge
makes up the vast majority of the volume of the atom
the electrons spin in shells around the nucleus of the atom
Sub-atomic Particles
The number of protons,
neutrons and electrons found in
an atom can be determined
using the periodic table.
# protons
atomic # from periodic table
eg. Sodium(Na) atoms have 11 protons
# neutrons
atomic mass # - atomic #
eg. Sodium atoms have 23 - 11 = 12 neutrons
# electrons number of protons
eg. Sodium atoms have 11 electrons
Complete the following chart, for the first 18
elements, by listing the required information about
individual atoms.
Electron Energy Level Diagrams (Bohr
Diagrams) for Atoms
These diagrams indicate the main features of the atom
of an element. Bohr diagrams are used to represent
the number of protons and neutrons found in the
nucleus of an atom and the number of electrons in the
electron cloud.
They also represent the arrangement of electrons in
various orbits, or energy levels, around the nucleus.
Each electron energy level has a specific and definite
number of electrons.
Energy Level
Maximum # of Electrons
1
2
3
2
8
8
The outer-most energy level that contains electrons is called the
valence energy level, or valence shell and electrons located here
are called valence electrons.
The Bohr diagram also show the name and symbol of the atom. As
well, atoms are electrically neutral since the number of positively
charged protons is equal to the number of negatively charged
electrons.
Electron Energy Level Diagrams - Atoms:
Chlorine Atom
Nitrogen Atom
Argon Atom
Aluminum Atom
Lithium Atom
Beryllium
Read pgs. 184-187
Questions 1-3 pg. 187
Complete the chart on page 21 of you notes by filling in the missing
information.
Compounds and Bonding
Atoms to Ions
It is important to note that atoms tend to be unstable structures.
They are generally unstable because the outermost energy levels
that contain electrons are not filled.
Atoms want to gain or lose electrons to ensure that their outer
energy levels are filled with electrons, and to ensure they are
stable structures.. However, when atoms gain or lose electrons they
become ions.
Atoms become ions to attain an electron energy level arrangement like the nearest
noble gas. The noble gases have stable atoms because their outermost energy
levels are filled with electrons.
In order to attain this stable electron structure metallic atoms, which have weak
attractions for valence electrons, will lose electrons to become like their nearest
noble gas. This will result in the metallic ions having a positive charge.
The non-metallic atoms, which have strong attractions for valence electrons, will
gain electrons to become stable like their nearest noble gas. This will result in the
non-metallic ions having a negative charge.
Of course, this transfer of electrons will
only occur if there are both metallic and
non-metallic atoms in close proximity to
each other.
Electron Energy Level Diagrams for Ions:
Aluminum Ion
Aluminum will give
up three electrons
to become an ion,
and to be stable like
its nearest noble gas
- neon
Because aluminum
loses three electrons
to become an ion it
attains an overall
charge of 3+.
Note that the
symbol changes, as
well, to represent
the charge of the
ion.
Electron Energy Level Diagrams for Ions:
Oxygen Ion
Oxygen will gain two
electrons to become
an ion, and to be
stable like its
nearest noble gas neon.
Because oxygen
gains two electrons
to become an ion it
attains an overall
net charge of 2-.
When non-metallic
atoms gain, or lose,
electrons to become
ions their name
changes. Oxygen
atom, for example,
changes to become
oxide ion.
Electron Energy Level Diagrams for Ions:
Beryllium Ion
Beryllium will give
up two electrons to
become an ion and
to be stable like its
nearest noble gas helium.
Because beryllium
loses two electrons
to become an ion it
attains
Complete the charts and questions on pages 24-27 of your notes
• Remember that families of elements have similar chemical and
physical properties. These families of elements will gain, or lose,
specific numbers of electrons to attain a stable ‘noble gas like’
electron arrangement. All elements in group IA, for example will
lose one electron to be like the nearest noble gas. The other
families are as follows.
Group
Gain or Lose e- to become ion
Ion Charge
Group IA (alkali metals)
lose of one electron to become an ion
+1
Group IIA (alkaline earth)
lose two electrons to become an ion
+2
Group IIIA
lose three electrons to become an ion
+3
Group VA
gain three electrons to become an ion
-3
Group VIIA (halogens)
gain two electrons to become an ion
-1
Group VIIIA (noble gases)
gain one electron to become an ion
0
Read pgs. 184-187
Questions 4-8 pg. 187
Bonding: How Elements Form Compounds
Bonding occurs between elements when conditions are such that
there are two, or more, elements present in close proximity to each
other. A force of attraction between these elements results in the
formation of a bond, and a compound is created.
There are many different types of bonding that can occur between
atoms of elements.
In Science 1206, we focus on two main types, namely;
1. Ionic bonding
2. Covalent (molecular) bonding
The type of bonding that will actually occur depends on the types of
atoms that are in close proximity to each other. In order for these
atoms to create a bond there must be a force of attraction between
the them.
In order for sodium chloride to be formed there
must be sodium and chlorine atoms in close
proximity to each other, and there must be a
force of attraction between them. This force of
attraction will create sodium chloride, an ionic
compound.
Ionic Compounds
made up of negative and positive ions that have resulted from the
transfer of electrons from a metal to a nonmetal.
electrons are transferred from metals (positive ions) to non-metals
(negative ions)
the force of attraction exists between oppositely charged ions
(opposites attract)
dissolve in water and conduct electricity (electrolytes)
Molecular Compounds
formed when non-metals combine with non-metals
a force of attraction exists between the positive nucleus of one
atom and the negatively charged electrons of the other atom(s).
electrons are shared between the two non-metals
do not conduct electricity in solution (non-electrolytes)
Chemical Formula: Ratios of Atoms
The compounds that are created will have specific ratios of one
atom(element) to another. This ratio is represented using a
chemical formula. This chemical formula contains a combination of
symbols that represent a particular compound and denotes the
relative numbers of each element in the compound.
i.e.: NaCl
This chemical formula indicates that there is one sodium for every
one chlorine ion(chloride) in the compound
i.e.: CaCl2
This chemical formula indicates that there is one calcium ion for
every two chlorine ions(chloride) in the compound
I.e.: P4O10
This chemical formula indicates that there are four phosphorous
ions for every ten oxygen ions(oxide) in the compound
Read pgs. 188-189
Questions 1,2 & 4 pg. 180/ 1 & 2 pg. 191
Ionic Bonding and Ionic Compounds
Ionic bonding occurs when atoms(ions) with opposite charges are in
close proximity to each other. Atoms of elements attain charges by
gaining, or losing, electrons to form ions. The atoms will gain or
lose electrons if there are elements present that have different
attractions for electrons.
Metallic atoms, for example, have weak attractions for electrons
while non-metallic atoms have strong attractions for electrons. As a
result, when metallic and non-metallic atoms are in close proximity
to each other the non-metallic atoms steal the some of the
electrons away from the metallic atom. This results in a positively
charged metallic ion, and a negatively charged non-metallic ion.
The number of valence electrons gained, or lost, will depend on the
atoms present and on the number of electrons they need to transfer
to obtain a noble gas like electron structure - to become stable
ions.
Sodium, for example has a weak attraction for electrons while
oxygen has a strong attraction for electrons. When these elements
are in close proximity to each other there will be a transfer of
electrons from the sodium to the oxygen. The sodium atoms will
end up with a positive charge and the oxygen will end up with a
negative charge.
These ions, with opposite charges, are then attracted to each
other(opposites attract), creating the compound Na2O. Opposites
attract in ionic bonding.
Remember that valence electrons are those found in the outermost
energy level of an atom, or ion.
Writing Chemical Formulas for Ionic
Compounds
There are four distinct combinations of ions that determine the formation of an
ionic compound. These criteria are dependent on the types of atoms, or ions,
that are present to form the compound. These combinations of ions are:
Type 1. metal ion + non-metal ion
Type 2. metal ion + negatively charged complex ion(polyatomic ion)
Type 3. positively charged complex ion + negatively charged complex ion
Type 4. positively charged complex ion + non-metal
Rules for Determining Chemical Formula of
Ionic Compounds.
Rule 1
Rule 2
Write the symbols, with the cation(+) first
Write the Ionic charge above each symbol to indicate
the stable ion each element forms
Rule 3
Choose the number of ions to balance each charge.
Use the cross-hatch method. Let the number that
represents the ionic charge be the subscript for the
oppositely charged ion. This creates a neutrally
charged compound.
Write the formula using subscripts
Rule 4
Type 1: Metal Ion(cation) + Nonmetal
Ion(anion)
Ex. A compound formed by Aluminum and Bromine
- note the overall charge on the compound
- do several more examples
Complete the chart on page 33-39
strontium and nitrogen
aluminum and oxygen
magnesium and sulfur
sodium and chlorine
germanium and selenium
zinc and phosphorous
aluminum and arsenic
barium and phosphorous
Type 1: (continued)
Multivalent Metal Ion(cation) + Nonmetal
Ion(anion)
Note: Some metallic atoms are capable of forming ions with more
than one possible ionic charge. These elements are referred
to as multivalent.
If the charge of the multivalent ion is not given then
you must always choose the top one as this is the one
that occurs most often.
Ex. A compound formed by Iron (II) and Nitrate
- note the overall charge on the compound
- do several more examples
Complete the chart on page 39-40
titanium (IV) & nitrogen
vanadium and oxygen
nickel (III) and sulfur
iron (III) and sulfur
cobalt (II) and nitrogen
lead (IV) and selenium
bismuth (V) and astatine
gold (III) phosphorous
Type 2: Metal(cation) + Negative
Complex(Polyatomic) Ion(anion)
The complex ions are a special group of ions that have formed when
a group of atoms have combined to produce a single, stable
complex ion with an overall charge.
The elements boron and oxygen can form a single stable ion called
borate. This complex ion has an over charge of 3- and will have the
formula BO3-. It is also located in the table of complex ions. Note
the names of the complex ions ending with the suffixes ate and ite.
Polyatomic Ion Formula
Name
chlorite
ClO2chromate
CrO42-
• There are three complex ions
that do end in ide, namely
cyanide, hydroxide and
bisulfide, and these must be
remembered.
As well, it is important to
note that if there is more
than one complex ion in the
chemical formula of the
compound, brackets must
be used to obtain the
proper ratio for the
complex ion.
Ex.
A compound formed by
Indium and Nitrate
Magnesium & chlorite
Potassium & chromate
Nickel III & carbonate
Strontium & cyanide
Bismuth III & thiocyanate
Calcium & nitrite
Titanium III &
permanganate
Type 3: Positive Complex Ion + Negative
Complex Ion
Ex. A compound formed by Ammonium and Carbonate
Ammonium & iodate
Ammonium & cyanide
Ammonium & stearate
Ammonium & sulfate
Ammonium & bisufite
Ammonium & hydroxide
Ammonium & carbonate
Ammonium & benzoate
Type 4: Positive Complex Ion + Nom-metallic
Ion
Ex. A compound formed by ammonium and oxygen.
Ammonium & oxygen
Ammonium & fluorine
Ammonium & nitrogen
Ammonium & sulfur
Ammonium &
phosphorous
Ammonium & chlorine
Ammonium & tellurium
Ammonium & iodine
Rules for Writing Compounds Names of Ionic
Compounds.
Rule 1:
- Name the positive ion(cation) first.
- If the ion is milti-valent then use roman numerals to
indicate the charge on the ion.
Rule 2:
- Name the anion second.
- If the anion is a single element (non-metal) remove
the suffix and append "-ide".
- If the anion is a complex ion write as given in the
table of complex ions.
Example 1:
A compound formed by aluminum and bromine
The cation is aluminum so it gets written in full.
The anion is bromine so we drop the ine and add ide.
Bromine becomes bromide.
Thus the compound name is aluminum bromide.
• Example 2:
A compound formed by magnesium and oxygen
The cation is magnesium so it gets written in
full.
The anion is oxygen so we drop the ygen and
add ide. Oxygen becomes oxide.
Thus the compound name is magnesium oxide
Example 3:
A compound is formed from copper and chlorine. Note that copper
is multivalent and if you do not know the charge on the ion then
choose the ion at the top. In this case it is Cu2+. The name of the
compound must now indicate this charge.
The cation is copper so it gets written in full, with the charge
indicated in roman numerals - copper (II)
The anion is chlorine so we drop the ine and add ide. Chlorine
becomes chloride
Thus the compound name is copper (II) chloride.
Complete chart on pg. 45
Questions 1,3-6 pg. 195
Multivalent Ion Note:
If you are given the chemical formula for the
compound then the charge on the multivalent ion
can be determined by multiplying the charge on the
negatively charged anion with the subscript denoting
the ratio of the anion to the cation. Use the cross
method to determine the charge on the multivalent
ion.
Example 4: TiO2 - the titanium is multivalent
In this case, you can determine the charge on the titanium by
multiplying the charge on the anion by the ratio number of the ion 2times 2 = 4-.
Since the compound must be neutral then there must be a net positive
charge of 4+, and since there is only one titanium ion it must have a
charge of 4+.
Name: Titanium (IV) oxide
Example 5: SnCl2 - the tin is multivalent
In this case, you can determine the charge on the tin by
multiplying
the charge on the anion by the ratio number of the ion 1- times 2 = 2 -.
Since the compound must be neutral then there must be a net positive
charge of 2+, and since there is only one tin ion it must have a charge of
4+.
Name: Tin(II) chloride
Complete chart on pg. 47
Questions 7-9 pg. 47
Example 6:
A compound is formed from calcium and sulfate. Note
name ending of sulfate as ‘ate’. This indicates this ion
is a complex ion. Ions ending in ‘ite’ are also complex
ions.
The cation is calcium so it gets written in full.
The anion is sulfate, a complex ion, so it gets written
as in the table of complex ions.
Thus the compound name is calcium sulfate.
Example 7:
A compound is formed from aluminum and sulfite.
Note name ending of sulfite as ‘ite’. This indicates this
ion is a complex ion. Ions ending in ‘ate’ are also
complex ions.
The cation is aluminum so it gets written in full.
The anion is sulfite, a complex ion, so it gets written
as in the table of complex ions.
Thus the compound name is aluminum sulfite.
Complete chart on pg. 48
Questions 1-4, 6 & 7 pg. 49
Complete charts on pgs. 50-52
Hydrates
Hydrated substances are ionic compounds that contain water as part of
their crystalline structure.
Cobalt chloride, for example, can incorporate a water component
depending on how the compound is formed. The hydrated ionic compound
produced is called cobalt (II) chloride dihydrate. The ‘di’ prefix is used to
indicate the amount, or number, of water molecules contained in the
compound.
The chemical formula for this ionic compound will be CoCl2 * 2H2O.
When writing the compound name for a hydrated compound it is
necessary to indicate the amount, or number, of water molecules
contained in the compound by using the a prefix system.
Example:
NaCl * 3H2O Note that there are three water molecules that are
contained in this compound. The prefix for three is ‘tri’, so the
name of this compound will be sodium chloride trihydrate.
Number
Prefix
1
2
3
4
5
6
7
8
9
10
mono
Di
Tri
Tetra
Penta
Hexa
Hepta
Octa
Nona
Deca
Complete chart on pg. 53
Molecular Compounds
Many compounds are formed when non-metallic atoms are linked
together to form electrically neutral compounds called molecules. These
compounds are formed when one or more non-metallic elements combine
through a process called covalent bonding. Molecular compounds share
electrons to become stable, like its nearest noble gas, and to attain an
overall neutral electric charge.
In molecules, the electrons of one atom are attracted to the nucleus of
its neighbouring atom, and vive-versa. Carbon dioxide (CO2), for
example, is a molecular compound since both carbon and oxygen are nonmetallic elements. Carbon tetrafluoride is another molecular compound.
Its formula will be CCl4.
These compounds may contain just two atoms of the same element
(diatomic molecules), such as hydrogen gas (H2) and oxygen gas (O2).
Most molecules are complex and will contain many more atoms such as
sugar (C12H22O11). A compound in blood, for example, has the chemical
formula C3032H4816O872N780S8F4.
Molecular Bonding( Covalent bonds)
• bonds in which electrons are shared between the nuclei of atoms.
• molecules are electrically neutral
• covalent bonds form to complete the outer energy level of atoms
and allow the atoms to have a noble gas like electron structure
Rules for Naming Molecular Compounds
•
•
•
•
•
The first element in the formula is written with its full name.
The second element is written as if it were an ion using “ide”
Prefixes are used to denote the numbers of atoms present.
The prefix mono- is only used to name the second element.
For example, CO is carbon monoxide, not monocarbon monoxide.
Sometimes, however, the prefix mono is not used in naming the second
element present. If there is no prefix the assumption is that mono should be
there.
• Common names are preferred where compounds have common names
Number
Prefix
1
2
3
4
5
6
7
8
9
10
mono
Di
Tri
Tetra
Penta
Hexa
Hepta
Octa
Nona
Deca
Example 1: P4O10
Note the prefixes for both the phosphorous and oxygen atoms. The
prefix 4 is tetra and 10 is deca. These will now be used to name the
compound. As well, the name of the second atom must be changed
to end in ‘ide’.
Name: tetraphosphorous decaoxide
• Example 2: CF4
• Note the prefixes for both the carbon and fluorine atoms. The
prefix 1 is mono and 4 is tetra. These will now be used to name
the compound. Once again, the name of the second atom must be
changed to end in ‘ide’. Remember that the rules indicate the
prefix ‘mono’ does not get used with the first atom.
• Name: Monocarbon tetrafluoride becomes carbon tetrafluoride
#
Compound Name
1
Nitrogen monoxide
2
3
N2O
Nitrogen dioxide
4
5
Compound Formula
N 2 O3
Dinitrogen tetraoxide
Complete chart on pg. 56-57
Compounds with Common Names.
There is a large group of molecular compounds that have been used
over a long period of time and, as a result, have common names.
These common names are the ones that must be used when writing
the compound name for these substances. These compounds include
the following;
#
Compound Name
Compound Formula
1
Water
H20
2
Ozone
O3
3
Ammonia
NH3
4
Methane
CH4
5
Sucrose
C12H22O11
6
Methanol
CH3OH
7
Ethanol
C2H5OH
8
Hydrogen peroxide
H2O2
Diatomic Molecules
Other non-metallic atoms are capable of forming compounds by sharing
electrons with other atoms of the same element. These molecular
compounds are referred to as diatomic molecules, and they exist in
nature as bonded pairs of the same element. The diatomic molecules are;
#
Compound Name
Compound Formula
1
Hydrogen
H2
2
Nitrogen
N2
3
Oxygen
O2
4
Fluorine
F2
5
Chlorine
Cl2
6
Bromine
Br2
7
Iodine
I2
8
astatine
At2
Other unique Molecules
Atoms of the elements phosphorous and sulfur will also combine
with each other to form molecular compounds.
#
Compound Name
Compound Formula
1
Sulfur
S8
2
Phosphorus
P4
Monatomic Elements
The noble gases are monatomic elements meaning they form
molecular compounds as single atoms.
#
Compound Name
Compound Formula
1
Helium
He
2
Neon
Ne
3
Argon
Ar
4
Krypton
Kr
5
Xenon
Xe
6
Radon
Rn
Complete chart on pg. 62 & 63
Hydrogen Compounds
Hydrogen compounds are a special group of molecular compounds
that are composed of a hydrogen atom combined with another nonmetal atom or with a negatively charged complex ion.
Hydrogen compounds are often said to have an ‘identity crisis’.
They are molecular, covalently bonded compounds that behave as
ionic compounds.
Hydrogen compounds
- may be solid, liquid or gas at room temperature.
- are soluble in water
- form coloured and colourless solutions
- form solutions that conduct electricity (electrolytes)
- form solutions which turn blue litmus red
Most hydrogen compounds are named as acids. A couple of
exceptions include HCl(g), H2S(g) and HCN(g). These are named as if
they were ionic.
Acids
When hydrogen compounds are placed in water they ionize (break
apart into individual ions) to become acids. Because the acids are
ionized in water they are given the state designation aqueous(aq) meaning dissolved in water. All acids must have this (aq)
designation.
#
Compound Formula
Compound Name
1
HCl(g)
Hydrogen chloride
2
HCl(aq)
Hydrochloric acid
3
H2SO4(g)
Hydrogen sulfate
4
H2SO4(aq)
Sulfuric acid
• Any compound
whose chemical
formula begins
with a hydrogen
cation (H+) is an
acid . There are
two different
types of acids.
They are;
Type
Explanation
Example
Binary
Molecule contains hydrogen and one other
non-metal
HF(aq) and HCl(aq)
OXO
Molecule contains hydrogen, oxygen and
another non-metal.
These compounds form when hydrogen
combines with polyatomic ions that contain
oxygen
H2SO4(aq) and HNO2(aq)
Rules for Naming, and Writing Chemical
Formulas, for Acids
Because of the nature of acids, being molecular compounds that behave as
if they were ionic compounds, a special naming system was developed for
these molecules. A special chart, called Naming Acids, is present in the
periodic table to help accommodate this unique naming system.
The general rule is to create an intermediate ionic name to represent the
molecules chemical formula or compound name. This intermediate ionic
name is then used in conjunction with the naming acids table in the
periodic table.
Rule #1: Naming:
Acids with intermediate ionic names ending with ‘ide’. The root of the anion
name is combined with ‘hydro’ and ‘ic’ to create the name of the acid.
HCl(aq) has the intermediate ionic name hydrogen chloride - ending in ide. As a
result the root of the anion chloride is combined with hydro and ic to produce
the acid name hydrochloric acid.
HCN(aq) has the intermediate ionic name hydrogen cyanide - ending in ide. As a
result the root of the anion cyanide is combined with hydro and ic to produce
the acid name hydrocyanic acid.
Rule #1: Chemical Formulas:
If the acid name begins with ‘hydro’ and ends in ‘ic’, then the
intermediate ionic name would have been hydrogen ‘root’ide. The
ions are then balanced as if they were ionic.
Hydrobromic acid, for example, has an acid name beginning with
‘hydro’ and ending in ‘ic’. Its intermediate ionic name will be
hydrogen bromide. The ions present will be H+ and Br -. Using the
cross-hatch we can predict the formula to be HBr(aq).
Rule #2: Naming:
Acids with intermediate ionic names ending with ‘ate’. The root of the
anion name is combined with ic to create the name of the acid.
H2CO3(aq) has the intermediate ionic name hydrogen carbonate - ending in
‘ate’. As a result the root of the anion carbonate is combined with ‘ic’ to
produce the acid name carbonic acid.
H2SO4(aq) has the intermediate ionic name hydrogen sulfate - ending in
‘ate’. As a result the root of the anion sulfate is combined with ‘ic’ to
produce the acid name sulfic acid.
The anions sulfur and phosphorous retain their ur and or designation when
named as acids. The acid name then is sulfuric acid.
Rule #2: Chemical Formulas:
If the acid name ends with ‘ic’, then the intermediate ionic name
would have been hydrogen ‘root’ate. The ions are then balanced as
if they were ionic.
Chloric acid, for example, has an acid name ending in ‘ic’. Its
intermediate ionic name will be hydrogen chlorate. The ions
present will be H+ and ClO3-. Using the cross-hatch we can predict
the formula to be HClO3(aq).
Rule #3: Naming:
Acids with intermediate ionic names ending with ‘ite’. The root of the
anion name is combined with ‘ous’ to create the name of the acid.
HNO2(aq) has the intermediate ionic name hydrogen nitrite - ending in ite.
As a result the root of the anion nitrite is combined with ‘ous’ to produce
the acid name nitrous acid.
H2SO3(aq) has the intermediate ionic name hydrogen sulfite - ending in ite.
As a result the root of the anion sulfite is combined with ‘ous’ to produce
the acid name sulfous acid.
The anions sulfur and phosphorous retain their ‘ur‘ and ‘or’ designation
when named as acids. The acid name then is sulfurous acid.
Rule #3: Chemical Formula:
If the acid name ends with ‘ous’, then the intermediate ionic name
would have been hydrogen ‘root’ite. The ions are then balanced as
if they were ionic.
Sulfurous acid, for example, has an acid name ending in ‘ous’. Its
intermediate ionic name will be hydrogen sulfite. The ions present
will be H+ and SO3 2-. Using the cross-hatch we can predict the
formula to be H2SO3(aq).
Acids with COO- groups.
Any acid with a COO- group at the end, or beginning, of the its
chemical formula will need to have the Hydrogen(H+ ) cation placed
next to the COO- group.
Acetic acid, for example, is composed of the cation H+ and the
anion acetate CH3COO-. These two ions will combine to create an
acid with the chemical formula the HCH3COO(aq). Note, however,
that the anion has a COO- group, and recall that the H+ ion has to
be placed next to this COO- polyatomic ion group. This will result in
the chemical formula, for this compound, being CH3COOH(aq)