Download Science 1206 Unit 3 Part 1

Def’n: matter – anything that takes up space, which
have both physical and chemical properties
Def’n: chemistry – the study of matter, its
properties, and its changes or transformations
Q. How do we use chemistry in everyday life?
Why is it important to have safety rules in
a laboratory?
 Textbook page 658 - 660

WHMIS is the Workplace Hazardous Materials
Information System
 It provides information about many hazardous
materials used in the workplace
 Employers are to ensure that their workers are
trained in WHMIS if their workplace contains
any hazardous materials
 There are eight WHMIS symbols:

Class A: Compressed
Gas
This class includes
compressed gases,
dissolved gases, and
gases liquefied by
compression or
refrigeration.
Class B: Flammable
and Combustible
Material
This class includes
solids, liquids, and
gases capable of
catching fire in the
presence of a spark
or open flame under
normal working
conditions.
Class C: Oxidizing
Material
These materials
increase the risk of
fire if they come in
contact with
flammable or
combustible
materials.
CLASS D: POISONOUS AND
INFECTIOUS MATERIAL
Division 1: Materials
Causing Immediate and
Serious Toxic Effects
These materials can cause
death or immediate
injury when a person is
exposed to small
amounts. Examples:
sodium cyanide,
hydrogen sulphide
CLASS D: POISONOUS AND
INFECTIOUS MATERIAL
Division 2: Materials
Causing Other Toxic
EFFECTS
These materials can cause
life-threatening and
serious long-term health
problems as well as less
severe but immediate
reactions in a person
who is repeatedly
exposed to small
amounts.
CLASS D: POISONOUS AND
INFECTIOUS MATERIAL
Division 3: Biohazardous
Infectious MATERIAL
These materials contain
harmful micro-organisms
that have been
classified into Risk Groups
2, 3, and 4 as
determined by the World
Health Organization
(WHO) or the Medical
Research Council of
Canada.
CLASS E: CORROSIVE
MATERIAL
This class includes
caustic and acid
materials that can
destroy the skin or
eat through metals.
Examples: sodium
hydroxide,
hydrochloric acid,
nitric acid
CLASS F: DANGEROUSLY
REACTIVE MATERIAL
These products may selfreact dangerously (for
example, they may
explode) upon
standing or when
exposed to physical
shock or to increased
pressure or
temperature, or they
emit toxic gases when
exposed to water.
MSDS are the Materials Safety Data
Sheets
 They contain important information
about an individual chemical
 A sheet comes with a chemical when
you order them
 There are nine categories on a MSDS:

Product Information: product identifier
(name), manufacturer and suppliers
names, addresses, and emergency
phone numbers
 Hazardous Ingredients
 Physical Data
 Fire or Explosion Hazard Data

Reactivity Data: information on the
chemical instability of a product and the
substances it may react with
 Toxicological Properties: health effects
 Preventive Measures
 First Aid Measures
 Preparation Information: who is
responsible for preparation and date of
preparation of MSDS

Page 173 – 174; 290
 There are several tests used to determine the
presence of certain chemicals:

›
›
›
›
›
›
For oxygen, we have the glowing stick test
For hydrogen, we have the lit splint test
For carbon dioxide, we have the lime water test
For water, we have the cobalt chloride paper test
For acids and bases, we have litmus paper.
For aqueous solutions of salt, we have the
conductivity test
Textbook page 173
 Physical property – a characteristic of a
substance such colour, ability to
conduct heat and/or electricity, lustre,
smell, etc.
 Chemical property – a characteristic
behaviour that occurs when a substance
changes to a new substance such as
ability to burn, reactivity with water, pH,
ability to rust, explosive, etc.

Physical change – a change in the
appearance of a substance without
changing the chemical make up of the
substance such as a phase change,
crumbling, dissolving, bending, etc.
 Chemical change – a change in the
chemical properties such as burning,
rusting, reacting the chemical with
another substance (e.g. vinegar in
baking soda, etc.)


Generally speaking, physical changes
are reversible (in theory) and chemical
changes are not.
Worksheet
 Page 175 #4, 5


Recall that atoms are the basic building
block of matter.
Each atom contains subatomic particles
(protons, neutrons, and electrons)
› Protons are positively charged, neutrons are
neutral, and electrons are negatively
charged.
› Protons and neutrons make up the middle,
or nucleus, of an atom while the electrons
fly around it in various energy levels
›
The outer most level is known as the
valence electron level. It is in this level
that electrons are either gained or lost.
 Atoms want to have full electron levels in
order to be stable.

› Not the Noble gases have full valence levels
and, therefore, do not form ions.
› In order to become stable atoms will lose or
gain electrons (whichever is least) to form
simple ions (only one charged atom).

Metals (the left side of the periodic
table) always lose electrons to obtain
the nearest noble gas configuration.
› These form cations
› Since they lose electrons, these atoms now
have more protons than they do electrons
and are, therefore, positively charged
› The group number (A’s) is the number of
electrons which the atom will lose

Non-metals (right side of the periodic
table) always gain electrons to obtain
the nearest noble gas configuration.
› These form anions
› Since these gain electrons, these atoms now
have more electrons than protons and are,
therefore, negatively charged

Worksheet #1
Recall from last year that you can draw
the electron levels in a Bohr diagram.
 There are two ways to do this, either with
circles (messy) or with lines
 The number of energy levels in a Bohr
diagram is equal to the row number in
which the element is found.

Ex. Draw a Bohr diagram for each of the
following:
1. Magnesium
2. Chlorine


When drawing a Bohr diagram for an
ion, we need to think about how an
atom becomes an ion (by losing or
gaining electrons).
› We need to show this on the digram.
› Note that all the Bohr diagrams for metallic
ions will look like the diagram for the Nobel
Gas that proceeds it, while the Bohr
diagrams for non-metallic ions look like the
diagrams for the Noble Gas which comes
after it.
Ex. Draw a Bohr diagram for each ion.
1. Aluminium ion
2. Fluoride ion

*Bohr diagram worksheet

Page 189; 201 - 202
Binary Molecular Compounds
 Involve the sharing of electrons resulting in
covalent bonding
 Composed of two non-metals
 Exist as individual molecules
 May or may not be in lowest whole number
ratios
Covalent Bonding (Molecular Compounds)
 Sharing of electrons
› the two non-metals both want to gain
electrons (to be like the noble gases), since
both cannot gain electrons at the same time
they share.
 The unit formed by a covalent bond is called a
molecule (it is neutral it has no charge)

Mono-atomic elements
› Mono means one so these are elements that exist on their
own in nature.
› Namely the noble gases:






He – helium
Ne – neon
Ar – argon
Kr – krypton
Xe – xenon
Rn – radon
No need to memorize these as they
are listed in the periodic table

Diatomic molecular compounds
› These are non-metals that exist in nature as two atoms
joined together.
› They are:







H2 – hydrogen
O2 – oxygen
N2 – nitrogen
Cl2 – chlorine
F2 – flourine
Br2 – bromine
I2 - iodine

Polyatomic molecular elements
› These are non-metals that exist in nature with many
atoms joined together
› These are:




S8 – sulfur
O3 – ozone
P4 – phosphorus (red)
P10 - phosphorus (white)
Molecule – a neutral particle that is
made up of two or more atoms that are
joined together by covalent bonds
 Molecular formula – chemical formula
which denotes the number and type of
different atoms in a molecule
 Empirical formula – the simplest chemical
formula that can be written for a
compound (smallest whole number ratio
of atoms

For example, water has the molecular
formula H2O and the empirical formula
H2O since the atoms are already in the
simplest form. Whereas hydrogen
peroxide has the molecular formula H2O2
and the empirical formula HO.
 Polyatomic ion – an ion that consists of
two or more different non-metal atoms
that are joined by covalent bonds
Simple ion – an atom that carries an
electrical charge; positively charged
ions are called cations and negatively
charges ions are called anions
 Formula unit - the chemical formula with
the least number of elements out of the
set of empirical formulas having the
same proportion of ions as elements. Ex.
NaCl is the formula unit for the ionic
compound sodium chloride.

Aqueous solution – a solution in which
water is the solvent
 Electrolyte – a substance that dissolves in
water, producing a solution that is able
to conduct electricity
 Nonelectrolyte – a substance that
dissolves in water and does not produce
a solution that conducts electricity

Page 203 - 204
 Trivial names

› These are names on compounds that do not
follow the guidelines put out by IUPAC.
› These are:




H2O – water
H2O2 – hydrogen peroxide
NH3 – ammonia
C12H22O11 – sucrose

For binary molecular compounds, we use prefixes to
indicate the number of atoms that are present in the
compound. These are:
›
›
›
›
›
›
›
›
›
›
Mono – one
Di – two
Tri – three
Tetra – four
Penta – five
Hexa – six
Hepta – seven
Octa – eight
Nona – nine
Deca - ten
When writing the name of a molecular
compound start with the first element, not
changing the name except for the prefix. Then
write the name of the second element with the
prefix and changing the ending of the name to –
ide.
 Note: the prefix mono on the first element only is
optional.

For example, name the compound CCl4
The first element is carbon and there is only one of
them.
The second element is chlorine and there are 4 of
them. Therefore the name would be
monocarbon tetrachloride
or
carbon tetrachloride

Ex 2. Write the IUPAC name for P4F6.
Sol’n: The first element is phosphorus and there
are four atoms.
The second element is flourine and there are six
atoms. Therefore the name of this compound is:
tetraphosphorus hexafluoride
Your turn. Name each of the following.
1. P4F5
2. C2O4
3. NO3
4. Cl2O
5. N3Br2
6. O6I2
When writing molecular formulas you look up
each element on the periodic table to determine
the chemical symbol for the element. (Don’t
forget that the ending on the second element has
been changed to –ide)
 Use the prefix on each element to determine the
number of atoms present. Write this number to
the right of the symbol and as a subscript.

Ex. Write the formula for diboron hexahydride.
The first element is boron which is B and di
means 2.
The second element is hydride (or hydrogen)
which is H and hexa means 6.
 We have B2H6
Ex. Write the formula for nitrogen triiodide.
The first element is nitrogen which is N and since
there is no prefix there is only one.
The second element is iodide (or iodine) and tri
means 3.
 We have N1I3 or NI3.
Your turn. Write the formulas for each of the
following.
a. Carbon disulfide
b. Dinitrogen pentabromide
c. Silicon trifluoride
d. Hexaphosphorus pentachloride
Worksheet #2
Ionic Compounds:
 Involve the transfer of electron(s), i.e. gaining
and losing electrons, resulting in ionic bonding
 Made up of two oppositely charged ions (metal
and non-metal, or combination involving a
polyatomic ion)
 Exist in the form of an ionic crystal lattice (not
individual molecules)
 Are always written as empirical formulas (lowest
whole number ratio)


Page 192 - 194
There are three categories of ionic compounds
that we will deal with.
1. Binary ionic
a) Simple ions (only single charges)
b) Multivalent ions (more than one charge)
2. Polyatomic ions (complex ions)
3. Hydrates

Binary ionic compounds are composed of a metal
ion (+) and non-metal ion (-).
› Binary simply means that only 2 ions are involved.
Ex. NaCl – the combination of a sodium ion
and a chloride ion.
Rules for naming simple binary ionic compounds
(meaning a compound with only two elements:
1. Name the cation (+) by writing the full name of
the metal.
2. Name the anion (-) by shortening the name of
the atom and adding the –ide ending.
Ex. NaCl  sodium chloride
CaF2  calcium fluoride
K2O  potassium oxide
**Note: Do NOT use prefixes – they are for
molecular compounds only (i.e. two non-metals)
Your turn. Name the following:
1. LiBr
2. AlCl3
3. Rb2S
4. Mg3P2
5. CaO
Rules for writing binary ionic formulas:
1. Write down the symbols of the ions involved.
2. Cross over the charges and write as subscripts.
3. Determine the lowest whole number ratio of
ions that will give a net charge of zero.
4. You do not need to write 1’s.
Ex. 1) Write the chemical formula for
potassium bromide.
2) Write the chemical formula for calcium
oxide.
3) Write the chemical formula for
magnesium iodide.
Your turn. Write the chemical formula for
each compound:
1. Sodium sulfide
2. Aluminum bromide
3. Barium iodide
4. Magnesium nitride
5. Aluminum nitride
Worksheet #3
Page 195
 Ions of certain elements can have more than one
possible charge. Such elements are called
multivalent species.
Ex. Copper is multivalent  its ions can have
either a 1+ or 2+ ion charge (Cu+ or Cu2+)

- these charges are provided on the periodic table

These multivalent elements are all transition
metals

When we name these compounds we need to
indicate which ion we are using and we do this
by using Roman numerals to represent the ionic
charge. (Do worry if you don’t know your
Roman numerals as they are written on your
periodic table of ions, but if you plan on doing
Chemistry 2202/3202 you will need to know
them)

Roman numerals:
› One – I
› Two – II
› Three – III
› Four – IV
› Five – V
› Six – VI
Rules for naming multivalent binary ionic
compounds:
1. Determine if the metal has more than one
possible charge (i.e. multivalent)
 consult periodic table
2. “Uncross” the subscripts and write them
as ionic charges. I.e. work backwards.
3.
4.
See if that charge is there on the periodic table
for that element. If not then you will have to
look up the charge on the anion (i.e. the non
metal) and see what scale factor to use.
Write the name of the compound. Be sure to
indicate the identity of the metal ion with
Roman numerals.
Ex 1. Write the IUPAC name for SnCl4.
**See other board for solution.
Ex 2. Write the IUPAC name for CrBr3.
**See other board for solution.
Your turn. Write the names for the following
compounds:
1. TiO2
2. AuCl3
3. Fe2O3
4. AgI
These are done in the same manner as binary
ionic compounds.
 Use the Roman numeral to determine the ionic
charge on the cation (i.e. the positive ion)
Ex. Write the chemical formula for iron (II)
chloride.
**See other board for solution.

Your turn. Write chemical formula for the
following compounds:
1. Titanium (III) fluoride
2. Titanium (IV) fluoride
3. Nickel (II) oxide
4. Lead (IV) sulfide
5. Vanadium (V) oxide
Worksheet #4
Page 196 - 198
 A polyatomic ion is two or more atoms
covalently bonded together that carry an overall
charge. Since they have a charge they can gain
or lose electrons to form ionic compounds.
 These ions are listed on your periodic table of
ions but this list is not exhausted, i.e. There are
more polyatomic ions than what’s listed.

To identify these compounds look at the elements
involve. If you have more than two elements in
the compound and the compound is not in your
trivial names list than it contains a polyatomic
ion.
 You name these the same way you name ionic
compounds: look up the element/polyatomic ion
and write the names down not changing the
second name to have the ending –ide.

Ex 1. (NH4)3PO4
**See other board for solution
Ex 2. Cr(NO3)3
**See other board for solution
Your turn. Write the name of the following
compounds.
1) Zn(OH)2
2) Pb(NO3)2
3) Mg(CH3COO)2
4) Na3BO3
5) K2Cr2O7
When writing the formulas for compounds
containing a polyatomic ion(s), you treat the ion
as a single unit, not as individual elements.
 If you need to write a subscript for a polyatomic
ion, then you must enclose the ion in brackets.
 Remember to use the charges on the polyatomic
ions to determine how many you need to have a
balanced compound, NOT the subscripts!!!

Ex 1. sodium chlorite
**See other board for solution
Ex 2. iron (III) sulfate
**See other board for solution
Ex 3. ammonium permanganate
**See other board for solution
Your turn. Write the formula for each of the
following.
1. Sodium hydroxide
2. Potassium bicarbonate
3. Potassium carbonate
4. Magnesium hydroxide
5. Ammonium sulfate
Worksheet #5
Page 236
Def”n: ionic hydrates – a compound that has water
associated with it. Water is part of its crystalline
structure.
Ex. Bluestone (CuSO45H2O) contains five water
molecules per copper (II) sulfate molecule in the
crystal.
Def’n: anhydrous – without water.
Ex. Anhydrous bluestone is CuSO4
When writing formulas for ionic hydrates, you
write the formula for the ionic compound as
already learned and add on H2O.
 Hydrates use the same prefixes as molecular
compounds to indicate how many water
molecules are associated with compound.
 To tell whether or not you have a hydrate or not,
you look for the word hydrate.

Ex. Write the chemical formula for each hydrate:
a) Sodium thiosulfate pentahydrate
**See other board for solution
b)
Copper (II) sulfate pentahydrate
**See other board for solution.
Your turn. Write the formulas for each of the
following.
1. Zinc sulfate heptahydrate
2. Potassium sulfate decahydrate
3. Cadmium nitrate tetrahydrate
Name the ionic compound as already learned and
add on hydrate with the appropriate prefix.
Ex. Write the name for:
1. Ni3(PO4)28H2O
**See other board for solution
2. Fe(OH)33H2O
**See other board for solution

Your turn. Name the following compounds.
1. CuSO4∙5H2O
2. NiCl2∙6H2O
3. CoCl2∙5H2O
Worksheet #6
Complete the Pre-lab Activity
 Complete the lab activity and prepare a
lab report following the handout.
 Attach answers to questions to the back
of the report.
 Go back and complete the chart –
Properties of Ionic and Molecular
Compounds p. 23

Page 288 – 290
 Page 293 – 294
 Page 295

Def’n: acids – molecules that ionize in water to
produce hydrogen ions (H+)
The properties of acids include:
 Turn blue litmus paper red
 React with metals to produce hydrogen gas
 Neutralize bases
 Have low pH (<7)
 Taste sour
Def’n: base – ionic compounds that contain the
hydroxide ion (OH¯)
The properties of bases include:
 Turn red litmus blue
 Neutralize acids
 High pH (>7)
 Form slippery solutions
 Bitter taste

Salts are formed as a result of the reaction
between an acid and a base
› Salts form electrolytic solutions when dissolved in
water
Acid + Base  Salt + Water
Ex. HCl(aq) + NaOH(aq)  H2O(aq) + NaCl(aq)
Note: There are more salts than just table salt
(NaCl)
To distinguish acids and bases from other types
of compounds look for (aq) next to the chemical
formula which means that the compound is
dissolved in water which acids and bases need to
be.
 Name bases the same way you name any ionic
compound.


Naming acids depends on the anion (negative
ion)
› If the anion is an element or a polyatomic ion ending
with –ide then
1. Drop the –gen ending of hydrogen
2. Replace the –ide ending of the anion with –ic
3. Add the word acid
Ex. HCl(aq)  hydrochloric acid
› If the anion is polyatomic and ends with –ate
1. Drop the name hydrogen
2. Replace the –ate ending of the anion with –ic
3. Add the word acid
Ex. HClO3(aq)  chloric acid
› If the anion is polyatomic and ends with –ite
1. Drop the name hydrogen
2. Replace the –ite ending of the anion with –
ous
3. Add the word acid
Ex. HClO2(aq)  chlorous acid
Note: With sulfur leave the –ur, for example,
H2SO3(aq)  sulfurous acid.
With phosphorus leave the –or , for example,
H3PO4(aq)  phosphoric acid
Your turn. Name the following acids:
1. HBr(aq)
2. HNO3(aq)
3. HNO2(aq)
4. HCN(aq)
5. H2CrO4(aq)
6. HClO(aq)
To write bases, you do the same as with any ionic
compound.
 To write acids, first convert the acid name to the
asscotiated ionic name

Hydro________ic acid  hydrogen _______ide
________ic acid  hydrogen _______ate
_________ous acid  hydrogen _________ite

Then write the formula as you learned
previously.

Ex 1. hydroiodic acid

Ex 2. boric acid

Ex 3. Nitrous acid
Your turn. Write the formula for each of the
following.
1. Hydrofluoric acid
2. Carbonic acid
3. Sulfurous acid
4. Hydrosulfuric acid
5. Perchloric acid
6. Silicic acid
Worksheet #7
Page 314
 Page 317 – 319
 Neutralizations reaction – mixing an acid
with a base (or vice versa) to temper
(reduce) the effects of one or the other

› Produces water and a salt

Salt – an ionic compound that will
conduct electricity when dissolved in
water (aqueous)
› Salts do NOT change the colour of litmus
paper
Generalized neutralization reaction:
Acid + Base → Water + Salt
 For example, hydrochloric acid + sodium
hydroxide → water + sodium chloride


Applications of neutralization reactions
› Using lemon juice (acidic) to eliminate fish
odors (base)
› Using TUMS (basic) for heartburn (acidic)
› Baking with baking soda (basic) and an acid
(like lemon juice or buttermilk) to make your
cakes rise
› Cleaning up acid spills in the lab with baking
soda (base)
Page 296
 The pH scale is a measure of how acidic or basic
a solution is.
 It ranges from 0 – 14, with 7 being neutral, less
than 7 being acidic and greater than 7 being
basic.
 The further from 7 you get in either direction
results in a stronger acid/base.
 Indicators are used to test for pH.
