Download Lesson 1 - Working With Chemicals

Lesson 1 - Working With Chemicals
 Safety is the number one concern
 WHMIS (Workplace Hazardous Materials
Information System) was established to standardize
information and symbols about chemicals in our lives
o WHMIS informs workers about the chemicals in
three ways
1. Controlled products must have informative labels,
in both English and French on their containers
2. Each controlled product must have a MSDS
(Materials Safety Data Sheet).
3. Workers who handle chemicals must complete an
education program provided through their employer.
http://www.brocku.ca/oehs/safety/whmis_symbols.pdf
 MSDS (Material Safety Data Sheet) was required to
accompany every chemical bought and sold (p.7).
http://www.lindecanada.com/en/msds/linde/Argon__Liquid_EN.pdf
Classifying Matter
Matter is defined as anything that has mass and volume
Matter may be solid, liquid or gas.
1
Matter
Mixture
Pure Substances
Combinations of matter
that can be separated
by physical means
Matter that has a
definite composition.
Do not have a
definite composition
Heterogeneous
Mixtures
(Mechanical
mixtures)
Different
components
of the
mixture are
visible or
layers can be
seen.
Gravel
Homogeneous
Mixtures
Element
Compound
(Solutions)
Cannot be
chemically
broken
down into
simpler
substances
Two or more
elements that
are chemically
combined
Different
components
are not
visible
Milk
Gold
Can be
separated
chemically
into simpler
substances
Water (H2O)
2
Assignment:
1. State whether the following is a pure substance or a
mixture:
a) sea water
b) iron
c) bronze
d) 14k gold
e) table salt
f) oxygen
2. State whether the following mixtures are
homogeneous or heterogeneous:
a) Oil-and-vinegar salad dressing
b) Steel
c) Cranberry juice
d) Sugar dissolve in water
e) Milk
f) Antifreeze
3. State whether the following pure substances is an
element or a compound:
a) Copper
b) Water, H2O
c) Methane, CH4
d) Silver
4. Classify each of the following substances:
a) Graphite, C
b) Shampoo
c) Coffee
d) Motor oil
5. Name the following elements:
3
Au, S, Fe, Hg, W, Cu, At, K, Na, Pb, Zr, Mo, Ag, P,
Ca, Cr, Ac, Ne, Fr, Sc, Ar, N, Mn, Be, Pt, Bi, Kr, C
Hf, Th, Cs, Po, U, He, Y, Ir, In, Rn, Ce, Pu, Sb, Os,
Assignment:
a) p. 11 # 1 – 4
(copy question first)
b) Find the number of protons, electrons, and
neutrons for the following
elements:
Hf, Th, Pt, Po, Au, U, Bi, Cs, Pb, W
Oxidizing material – rusting caused by oxygen (ex.
iron)
4
Developing (History of) Atomic Theories
1. Dalton’s Atom (1766-1844)
Dalton’s Atomic Theory (Pg 12)
- All matter is made up of small particles called
atoms
- Atoms cannot be created, destroyed, or divided
into smaller particles
- All atoms of an element are identical in mass
and size, but they are different in mass and size
from the atoms of other elements
- Compounds are formed when atoms of different
elements combine in fixed proportions
- Chemical reactions change the way atoms are
grouped, but atoms themselves are not changed
in reactions
- “billiard ball” model
http://www.rsc.org/chemsoc/timeline//pages/1803.html
2. J.J. Thomson (1856-1940)
- English physicist
- Atoms contain negatively charged electrons
- Electrons are like raisins in a plum pudding or
“raisin bun” model.
5
3. Ernest Rutherford (1871-1937)
- Atom contains electrons and positively-charged
particles
- Atom composed of
o A nucleus – a central region that is
positively charged and contains most of the
mass
- protons are heavy positive particles within
the nucleus
o Electrons – particles with a negative charge
and are very light (compared to protons).
- Electrons circle around the nucleus
o Empty space surrounding the nucleus is very
large within which electrons move
(planetary model).
o Rutherford also proposed existence of the
neutron to account for the mass difference
between hydrogen and helium
o Neutrons are heavy particles like protons
but have no charge
o Isotopes are atoms of the same element that
differ in mass (but are chemically alike).
(element with different number of
neutrons)
http://www.iun.edu/~cpanhd/C101webnotes/modern-atomic-theory/rutherford-model.html
4. Niels Bohr (1885-1962)
- Electrons exist only in certain energy levels or
orbits around the nucleus
6
- Only a certain number of electrons can exist
in each energy level (orbit).
5. Modern Theory
- Present day models of the atom are much more
complex
- Electron energy levels are divided into
sublevels.
- Neutrons and protons are made of even smaller
particles called quarks.
ATOMIC STRUCTURE

Atom - the smallest part of an element (which
retains the chemical and physical properties of the
element). Atoms are made up of 3 sub-atomic
particles
1. Electron (e- or e)
-smallest particle in an atom
-has a negative charge
-located in the extra nuclear region of the atom
2. Proton (e+ or p) 
-Has a large mass
-Has a positive charge
-Located inside the nucleus
7
3. Neutron (n)
-Same mass as a proton
-Has a neutral charge (no charge)
-Located inside the nucleus
Nuclear Notation
- Atomic number is the number of protons in the
nucleus
- The number of protons equals the number of
electrons in a neutral atom (#p = #e)
Atomic Mass Number is the total number of
protons and neutrons in the nucleus
- The mass number also identifies the particular
isotope.
-
Number of neutrons = mass # - atomic #
Atomic # = #p = #e
Example:
Find the number of protons, electrons and neutrons
for iron and sodium.
8
Fe
Atomic # = 26
Atomic mass = 55.85
p = 26
e = 26
n = 56 – 26 = 30
Note: when finding the number of neutrons we round the
atomic mass to the nearest whole number.
Na
Atomic # = 11
Mass # = 22.99
p = 11
e = 11
n = 23 – 11 = 12
Assignment:
Find the number of protons, electrons, and neutrons of the
elements with atomic numbers 1 to 30 and 40-70.
p. 11 #1-3 (copy questions or complete sentences)
9
Nuclear Notation Continued
1. One way to write isotopes of elements is:
12
6C
where the top number is equal to the
atomic mass, and the bottom number
is equal to the atomic number.
Atomic mass
12
6C
Atomic number
13
6C
#p = 6
#e = 6
#n = 12 – 6 = 6
#n = 13 – 6 = 7
2. Another notation used is: e.g. Lithium–7 or Li
-7
where 7 is equal to the atomic mass.
10
From the table of elements we get the atomic
number (which is 3).
Thus,
#p = 3
#e = 3
#n = 7 – 3 = 4
p.23 #5 to 8
p.24 # 1 to 4
p. 37 (b, c)
sentences)
(copy question first or complete
p. 38 - define the key terms (first column)
11
Bohr’s Model of the Atom
According to Bohr’s model electrons exist only in certain
energy levels or orbits around the nucleus
Only a certain number of electrons can exist in each
energy level or orbit.
The 1st orbit can hold a max. of
The 2nd
a max of
3rd
4th
5th
6th
2 electrons
8
8
18
18
32
When one orbit is filled the remaining electrons go to the
next orbit – you cannot exceed the maximum allowed.
We can draw the Bohr diagram for any element. It must
have a nucleus showing the number of protons and
neutrons and circles outside the nucleus showing the
number of electrons.
Reminder: # of protons = # of electrons = atomic #
e.g. Draw the Bohr model for the following elements:
12
a) Lithium
Step 1 – Look up the atomic number
It’s 3.
So,
# of p = 3
# of e = 3
Step 2 – Look up the atomic mass.
It’s 6.94 = 7 (round to the nearest whole #)
Find the number of neutrons.
Reminder: # of n = atomic mass – atomic #
So, # of n = 7 – 3 = 4
Step 3 – Draw the diagram.
#p=3
#e=3
#n=4
P=3
N=4
13
We can draw the orbits using a simplified version.
e.g. Cobalt
P = 27
N = 32
___ 9
___ 8
___ 8
___ 2 e
Assignment: Draw the Bohr model of the atom for the
elements: K, Cl, Mn, Zr, Ne, Ca, Sc, Ti, Fe, Al, Br, Ni,
Cu, S, Ag, He, B, O, Na, Mg, Be, Ar, N, V,
1to10 and 15 to 30.
And p. 38 #1, 5, 6
sentences)
(copy question or complete
14
Periodic Table
Periods are horizontal rows in which elements
increase in atomic mass from left to right
Groups or families are vertical rows made up of
elements with similar properties. There are 4
special named groups.
Group 1 – Alkali Metals
Group 2 – Alkaline Earths
Group 17 – Halogens
Group 18 – Noble gases or Inert gases
- ‘Staircase’ line separates metals from nonmetals
- Metalloids border this line
Francium is the most reactive metal.
Fluorine is the most reactive non-metal.
Valence Electrons
- The outermost occupied energy level (orbit) of
an atom is called its valence energy level.
- The electrons in the valence energy level
(electrons in the last orbit) are
15
called valence electrons.
Electron dot diagrams or Lewis dot are useful
ways to represent an atom.
In an electron dot diagram, the electrons in the last orbit are
shown as dots placed around the symbol.
e.g.
Li
#p = 3
#n = 4
Bohr diagram
e
e
There is 1 valence electron (that is, 1 electron in the last orbit).
•
Li
So, the dot diagram will be
e.g. Oxygen
#p = 8
#e = 8
#n = 8
p=8
n=8
____ 6
____ 2e
16
There are 6 valence electrons.
••
• O ••
•
Electron dot or Lewis dot diagram
Assignment:
Draw electron dot diagrams for the following elements:
K, Cl, Mn, Zr, Ne, Ca, Sc, Ti, Fe, Al, Br, Ni, Cu, S, Ag,
He, B,
and p. 27 #9 to 12
p. 28 complete the table
Science Test Friday
Assignment:
17
Draw electron dot diagrams for:
1. Scandium
2. Fluorine
3. Beryllium
4. Vanadium
5. Gallium
p. 38 – Define the key terms and
p. 38 # 7, 8, 13
Quiz
Draw electron dot diagrams for:
1. Sc
2. Na
3. Chromium
4. Ar
Ions
18
- Ions are atoms, which have gained or lost
electrons, in order to become more stable – it
happens during chemical reactions.
- Ions always have a charge
- Positively charged ions have fewer electrons
than protons – also called cations.
- Most metals form cations – that means they lose
electrons
e.g. Li1+
3
6.94
1+
Li loses an electron
Li1+ Li
Lithium
- Negatively charged ions have more electrons
than protons – also called anions.
- Non-metals that form anions have a name
ending in ‘ide’
e.g. chloride (Cl-), oxide (O-2 or O2-)
19
 All non-metals gain electrons (that is, form
anions).
Compounds
Compounds are formed when two or more elements are
chemically combined.
- Noble gases with their 8 valence electrons are
very stable elements – they usually don’t form
compounds.
- Other atoms have different ways of becoming
stable – they either gain or lose electrons when
they form compounds.
- **Metals give up electrons to other atoms,
forming cations.
- **Non-metals accept electrons, forming
anions.
- **Non-metals may share electrons with other
atoms.
e.g. Sulphur dioxide
non-metal
non-metal
20
Assignment
p.36 #1-3, 5
p.38 #6, 9, 15,16
(copy question or complete sentences)
Compounds
There are two basic types:
1. Ionic
2. Molecular
Ionic Compounds
- Ionic compounds formed from just two elements
are called binary ionic compounds
- A metallic cation is joined to a non-metallic
anion by an ionic bond.
- Ions of an ionic compound are arranged in a
regular repeating pattern called crystal lattice.
Ionic compound – metal and non-metal joined
chemically.
21
In ionic compounds electrons are traded.
e.g.
NaCl
(p.30 – sketch Fig. 1.27 here)
Molecular Compounds
- Atoms which share electrons to become stable
form molecular compounds (see p.32)
- These groups of atoms are called molecules
- Atoms in molecules are joined by covalent
bonds.
- All atoms in molecular compounds are nonmetals.
Molecular compounds – non-metal and non-metal
joined chemically.
e.g. CO2
(p. 32 – sketch Fig. 1.29 here)
Assignment:
p.37 (i,j,k)
p. 38 #13,15,16,17,18
(copy question)
22
Investigation 1-A Pg 33
Check Your Understanding Pg 36
Read Pg 37
Review Pg 38
23
Naming and Writing Binary Molecular
Compounds
- When two (binary) non-metallic atoms join by a
covalent bond we have a molecular compound.
e.g. Carbon dioxide
Rules for naming
1. The first element in the compound is the one most
left on the periodic table.
2. The suffix ‘ide’ is attached to the name of the
second element.
3. Prefixes are used to indicate how many atoms of
each type are present in one molecule of the
compound.
Prefixes:
1 = mono
2 = di
3 = tri
6 = hexa
7 = hepta
8 = octa
MEMORIZE
24
4 = tetra
5 = penta
9 = nona
10 = deca
No “mono” is used with the first element.
e.g. Give the name or formula for each compound:
NO2 – Nitrogen dioxide
N2O – Dinitrogen monoxide
N2O4 – Dinitrogen tetraoxide
Nitrogen monoxide - NO
Dinitrogen pentaoxide – N2O5
Carbon dioxide – CO2
Assignment:
Name or give the formula:
1. Silicon dioxide
2. Sulphur monoxide
3. OF2
4. SiBr4
25
5. PH3
6. N2O
7. CO
8. NBr3
9. P2I3
10. SO3
11. N2O4
12. Tetraphosphorous hexaoxide
13. Dinitrogen tetraoxide
14. Heptasilicon monobromide
15. Octaboron decaiodide
16. B2O3
17. BrF7
18. N3O6
19. H2Cl5
20. Triselenium diastatide
21. Diarsenic pentaoxide
22. Sulphur trioxide
23. C3O2
24. C2H6
25. As3Br7
26. SO2
27. Selenium monoxide
28. Diboron trioxide
29. PF3
30. P2O5
31. P4O10
32. Arsenic trifluoride
26
33.
34.
35.
BrF7
Hydrogen chloride
N2O
And p. 44 #1- 4, p. 62 #1
(copy question first)
Binary Ionic Compounds
- Are composed of ions of one metal element and
ions of one non-metal element joined by ionic
bonds
Rules for naming
1. The first element in the name of the formula is
the metal
27
2. The second element, the non-metal, is named as
an ion. The suffix ‘ide’must be present.
3. No prefixes are used.
e.g.
Fe2O3 – Iron oxide
CuS – Copper sulfide
KCl – Potassium chloride
p. 45 #5 - 7
p. 46 #9, 10
p. 47 #12
p. 48 #3, 5
(copy question)
Writing Formulas for Binary Ionic Compounds
In an ionic compound the total number of positive
charges must equal the total negative charges – the
compound must be electrically neutral.
28
This fact tells us how many of each atom is necessary to
form a compound.
e.g. sodium chloride
Step 1 – use the table to find the charges on each ion
(element)
Na1+
Cl1-
Step 2 – bring the two ions close together and see what the
net charge is.
Na+Cl- the two charges are equal so the formula is
NaCl.
Magnesium chloride
Mg2+Cl1Question: how many of each ion is needed so that the
molecule is neutral.
Cl1-
Mg2+
29
Cl1-
Therefore the formula is
Chromium oxide
3+
Cr O
MgCl2
Cr3+O2-
2-
to balance the charges we use a
shortcut method – charges are “traded”
across.
Cr2O3
Calcium oxide
Ca2+O2-
Ca2O2 CaO
Multivalent Cations (metals)
- Some atoms are able to form more then one
cation. Ex. Ni2+ or Ni3+
- In the Stock system, the charge on the cation is
written in brackets, as a Roman numeral after
the name of the metal
Example
30
Copper (II) oxide
Cu2+O2-
Tin (IV) fluoride
Sn4+F1- SnF4
PbI2
CuO
Lead (II) iodide
Pb2+ I1-
Cr2S3
Chromium (III) sulfide
Cr3+S2Is this formula correct LiO
Li1+O2-
No – correct formula is Li2O
p. 47 #11
p. 48 #4, 5
p. 49 #7, 9
31
Polyatomic Ions
- Consist of two or more different atoms
(covalently bonded) containing an overall
charge. e.g. NO3- Found in the box at the top of the table.
- All are negatively charged, except ammonium
ion, and most names end in ‘ate’
- All act as non-metals except ammonium ion,
NH4+, which acts as a metal in compounds.
- The name of the cation (metal) is followed by
the name of the anion (non-metal – negatively
charged).
- When writing formulas, brackets must
surround the polyatomic ion (when more than
one is present – i.e. subscript is not 1).
Examples:
1. Potassium sulphate
K2(SO4)
K1+(SO4)2- “trade”
charges
or K2SO4
32
NH4NO3 Ammonium nitrate
Al(NO3)3 Aluminum nitrate
Sodium sulfate
Na1+(SO4)2-
Na2SO4
Na1+
SO42Na1+
Ammonium phosphate
(NH4)1+(PO4)3-
(NH4)3PO4
33
Gallium hydrogen carbonate
Ga3+(HCO3)1-
Ga(HCO3)3
Assignment:
Practice Problems p. 52 #13-16
Practice Problems p. 53 #17-18
p. 55 #1-3
Properties of Ionic Compounds
-
In the solid state ionic compounds are crystalline
Ionic compounds have fairly high melting points
In the solid form they do not conduct electricity
In the aqueous (dissolved in water) form ionic
compounds are electrolytes – they conduct
electricity
34
Properties of Molecular Compounds
- Most molecular compounds have fairly low
melting points – weak intermolecular bonds
- Non-electrolytes – do not conduct electricity
- When dissolved in water most do not conduct
electricity (some do)
p. 55 #4, 5
p. 80 #1 - 3, 6, 17,18, 20
(copy question)
Special Compounds and Elements
Special compounds – these compounds have special
names, which do not follow the rules for naming.
Water
H2O
Ozone
Ammonia
Hydrogen Peroxide
Methanol
Ethanol
O3
NH3
H2O2
CH3OH
C2H5OH
or
HOH
35
Sucrose
Glucose
Methane
C12H22O11
C6H12O6
CH4
Diatomic and Polyatomic Elements
If these elements are FREE, that is ALL ALONE, they are
written as:
H2
N2
O2
F2
Cl2
Br2
I2
At2
P4
S8
MEMORIZE
For example, hydrogen has one electron and thus it wants
to fill that orbit in order to become stable - so it will pair
up with another hydrogen atom and they will share the two
electrons – covalent bonding.
H••H
Thus, hydrogen when it’s not in a compound but
all-alone is written as H2.
36
Assignment:
p. 79 (a, b, c, e,)
p. 80 – key terms
p. 81 #25
Test tomorrow.
Properties of Acids and Bases
1. Acids
- a substance that reacts and releases hydrogen
ions, H+(aq),in a water solution
- taste sour
- form colourless solutions
- conducts electricity
- formula starts with Hydrogen
e.g. HCl
H2SO4
-
Hydrochloric acid
Sulphuric acid
37
2. Bases
- a substance that dissolves in water and releases
hydroxide ions, OH- bitter tasting
- feel slippery
- form colourless solutions
- conducts electricity
e.g. NaOH
-
Sodium hydroxide
Indicators and pH
- an indicator is a chemical that changes a
different colour in an acid vs a base
- litmus is red in acids and blue in bases
- phenolphthalein is colourless in acids but pink in
bases
- pH is a scale used to indicate the strength of the
acid or base
- pH scale ranges from 0 – 14,
- pH of 7 is neutral – pure water
- pH less than 7 – acid
- pH greater than 7 – base
0
acid
7
base
14
38
p.80 #12,13
p.81 #23,25
(copy questions)
Naming Acids
All acids start with hydrogen. Acids have special names,
which derive from the following rules.
Chemical name
Hydrogen _______ide becomes
e.g. HCl
Hydrogen chloride
Acid name
Hydro______ic acid
Hydrochloric acid
39
H2S Hydrogen sulfide
Hydrosulfuric acid
Hydrogen _______ate
_______ic acid
H2SO4
Hydrogen sulfate
HClO
Hydrogen chlorate
Hydrogen ______ite
H2SO3
Hydrogen sulfite
HClO2 Hydrogen chlorite
Sulfuric acid
Chloric acid
_________ous acid
Sulfurous acid
Chlorous acid
p. 70
#20 (a, b, c) #21 (a, b, c)
#22,23
p. 71 #6
p. 79 (e, k) and p. 135 #26-28
40
Water
- the shape of the water molecule is
Oxygen end – slightly negative.
1050
Hydrogen end – slightly positive.
- has two covalent bonds but the electrons shared
in these bonds are not shared equally
41
- oxygen attracts the pairs of electrons closer to it
- this creates an uneven distribution of charges or
partial charges
- the result is a polar molecule or dipole
- the negative end or oxygen of one water attracts
the positive end or hydrogen of another –
hydrogen bond
- hydrogen bonds are one kind of intermolecular
force
- intermolecular forces are attractions between
molecules
- intramolecular forces are attractions within
molecules
Properties of Water
- the boiling point and melting points are higher in
water than other similar substances – the need to
break the hydrogen bonds
- it requires a great deal of energy to raise the
temperature of water – strong intermolecular
forces
- has a concave meniscus and shows capillary
action – strong force of attraction between water
and other molecules
42
- ice floats in liquid water – due to the
rearrangement of the hydrogen bonds in the
solid creating a greater volume and lower
density
- has a high surface tension – again due to the
hydrogen bonds
Chemical Reactions
- chemical reactions occur when one or more
substances change to form new substances
- also called a chemical change
- the substances that change are called the
reactants
- the substances formed are called products
- evidence that a chemical change has occurred
could involve one or more of the following
o energy change – heat and/or light
 exothermic – release energy
 endothermic – absorb energy
o odour change
o colour change
o formation of a gas – bubbling
o formation of a solid – precipitate
43
Predicting Solubility
- some ionic compounds are highly soluble in
water while others have a very low solubility
- we use a solubility table to help determine
whether a substance is soluble or not – back of
table.
Step 1 – Locate one of the ions in the compound in the
boxes across the top.
Step 2 – Look for the other ion in the two vertical boxes
below.
If it is soluble write (aq) behind the compound to show
that it is aqueous – it dissolves.
If it is slightly (low) soluble show that it does not dissolve
by writing (s) behind the compound so that it is solid.
e.g. Determine if the following compounds are soluble or
not by using the appropriate notation.
NaCl(aq)
Look for Na1+ or Cl1- across the top horizontal row.
44
PbI2(s)
NH4OH(aq)
CuSO4(aq)
p.90 #1, 2
p. 93 #1- 4
p.128 # 9
Law of Conservation of Energy
-Energy can be converted from one form to another,
but the total energy of the universe remains constant
(energy cannot be created nor destroyed).
- Breaking chemical bonds is an endothermic
process (energy is used).
- Forming new chemical bonds is an exothermic
process
1. When more energy is required to break bonds than
is released when new bonds form, the reaction is
endothermic
e.g. energy + water
hydrogen + oxygen
45
2. When less energy is required to break bonds than
is released when new bonds form, the reaction is
exothermic
Example: hydrogen + oxygen → water + energy
Lavoisier’s Law of Conservation of Mass
During a chemical reaction, the total mass of the reacting
substances (reactants) is always equal to the total mass of
the resulting substances (products).
Balanced Chemical Equations
- a balanced chemical equation shows that
atoms are conserved in a chemical reaction
(that is, the numbers of each atom must be
equal on both sides of the equation).
- reactants are on the left side of the equation and
products are on the right side
46
- coefficients are used to balance a chemical
equation (tells us how many molecules or atoms
are needed in the reaction).
e.g.
H2 + O2 → H2O (this is called a
skeleton equation)
2H2 +1O2 → 2H2O
H +
H
H
H
O →
O
(this is a balanced equation)
O
H H
+
O
H H
47
http://funbasedlearning.com/chemistry/chemBalancer/
Assignment:
Model Problem 1 p. 99
Model Problem 2 & 3 p. 100
p. 101 #5
p. 102 #5 (copy the EQUATIONS)
Types of Chemical Reactions
1. Formation Reactions or Simple Composition
- Two or more elements combine to form a new
compound
Element + Element → Compound
X +Y →
XY
48

The reactions must be balanced.
e.g. Iron combines with oxygen to form iron (III) oxide.
4 Fe + 3 O2 → 2 Fe2O3
Copper reacts with chlorine to form copper (I) chloride.
2 Cu + 1Cl2 → 2 CuCl
p. 114 #3
p. 115 #6
(Copy and balance)
2. Decomposition Reaction or Simple Decomposition
One compound breaks down into two or more
elements
49
compound → element + element + …
XY →
X +Y
XYZ → X + Y + Z
e.g.
2HCl → 1H2 + 1Cl2
2K2IO3 → 4K + 1I2 + 3O2
p. 127 (a-i,l,m) → complete sentences or copy the
question.
p. 128 #2, 4, 9
50
3. Single-Replacement Reactions

one element takes the place of another element in
a compound
- many involve the reaction between a metal and a
compound
element + compound → new element + new compound
A + BX →
AX + B
AX + Y
AY + X
→
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Cu + 2AgNO3 → 2Ag + Cu(NO3)2
2NaBr + Cl2 → 2NaCl + Br2
Note: Metal replaces (switches with) a metal.
Non-metal replaces a non-metal.
Mg + CuSO4 → Cu + MgSO4
When a metal reacts with water, the water formula is
written as HOH (first H “acts” as a metal).
2Na + 2HOH → H2 + 2NaOH
OH1- (hydroxide ion)
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4. Double-Replacement Reactions
- Two different compounds react, forming two
new compounds
compound + compound → new
+ new
compound
compound
AX + BY →
BX + AY
Note: metal switches with a metal and a non-metal
with a non-metal.
- a special kind of double-replacement reaction
called neutralization is between an acid and a
base
NaOH + HCl → NaCl + HOH
Base
Acid
Ba(OH)2 + Na2SO4 → 2NaOH + BaSO4
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p. 114 #2,3,7
p.134 #21
p.136 #40
5. Hydrocarbon Combustion
A hydrocarbon is an organic compound containing
carbon and hydrogen (sometimes oxygen also)
When hydrocarbons are burnt in a plentiful supply
of oxygen complete combustion occurs
- The two products are always carbon dioxide and
water vapour
Hydrocarbon + oxygen → carbon dioxide + water
Hydrocarbon + O2 → CO2 + H2O
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When hydrocarbons are burnt in a poor supply of
oxygen incomplete combustion occurs.
The products of this reaction are; carbon dioxide,
water, carbon (soot) and carbon monoxide.
Carbon monoxide is an odourless, colourless and
highly toxic gas.
- CO binds 200x more strongly to hemoglobin in
the red blood cells than does O2
e.g.
CH4 + 2O2 → CO2 + 2H2O
2C2H6 + 7O2 → 4CO2 + 6H2O
p. 114 #1,4
p. 135 #27-29
Ammonia
NH3
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Ethane
C2H6
Glucose
C6H12O6
The Mole
The mole is defined as the amount of substance that
contains as many elementary entities (atoms, molecules, or
formula units) as exactly 12 g of carbon-12, the most
common isotope of carbon.
One mole of a substance has been determined to
contain 6.02 x 1023 elementary entities of a substance
(atoms, molecules). This number is called Avogadro’s
number.
(Similar to dozen) dozen = 12
mole = 6.02 x 1023
Atomic Molar Mass
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- is a weighted average of the mass of 1 mol of all
of the naturally occurring isotopes of the
element
- listed for each element on the periodic table
- example 1 mol of iron = 55.85 g/mol
1 mol of zinc = 65.39 g/mol
- some elements exist as molecules such a
nitrogen gas
1mol N2 = 2 x 14.01g/mol = 28.02g/mol
Molar Mass of a Compound (M)
- refers to the mass of 1 mol of any pure
substance.
- to find the molar mass of a compound use the
chemical formula
e.g.
CO2 contains 1 carbon and 2 oxygen
1C = 1 x 12.01g/mol = 12.01
2 O = 2 x 16.00g/mol = 32.00
M = 44.01 g/mol
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H2O
Ca(OH)2
2 x 1.01 = 2.02
1 x 16.00 = 16.00
M = 18.02 g/mol
1 x 40.08 = 40.08
2 x 16.00 = 32.00
2 x 1.01 = 2.02
M = 74.10 g/mol
p. 120 #9, 10
p. 123 #15, 16, 20
p. 125 #3
Calculate the molar mass of the following compounds:
1. PbI2
2. NH4OH
4. CaPO4
5. Mn(NO3)5
7. NH3
8. S2N4
10. C6H12O6
11. NH4HS
13. CoCl2
14. Cobalt(III) silicate
15. Potassium phosphate
16. Polonium (II) oxide
17. Mercury (II) sulfide
18. Fe2(OOCCOO)3
19. Zn(OH)2
20.Cu(NO2)2
21.Co2(Cr2O7)3
22. MgHPO4
3. CuSO4
6. Fe(OH)3
9. BaSO4
12. GaI3
58
Calculating mass of a sample (m)
Molar mass (M) is equal to the mass of one mole of a
compound.
For example the molar mass of water is 18.02 g/mol.
What if we have 2 moles of water?
Then the mass of the water would be 2 x 18.02 = 36.04 g.
We use the following formula:
m = nM
n = # of moles (mol)
m = mass (g)
M = molar mass (g/mol)
How many grams are there in 3.5 moles of francium
nitride?
59
Step 1 – Write the formula and find the molar mass.
Fr3N
3 Fr – 3 x 223.00 = 669.00
1 N – 1 x 14.01 = 14.01
M = 683.01 g/mol
Step 2 – List what’s given and apply the formula.
n = 3.5 moles
M = 683.01 g/mol
m=?
m = nM
= (3.5)(683.01)
m = 2390.54 g
Mass of a substance to moles
If the mass of the sample is given rearrange the formula
for “n”
n=m
M
e.g.
60
How many moles are there in a 16 g sample of carbon
dioxide.
CO2
1 x 12.01 = 12.01
2 x 16.00 = 32.00
M = 44.01 g/mol
m = 16 g
M = 44.01 g/mol
n=?
n = m/M
= 16/44.01
n = 0.36 moles
p. 122 #11-14
p. 123 #17,18
p. 125 #5,6
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Moles summary
1. Molar mass (M) – must be calculated using the table.
2. Mass (m) – use the formula
3. Number of moles (n) – use
m = nM
n=m
M
p. 135 #29 - 32
p. 136 #41 - 43
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Practice Problems Pg 122
Practice Problems Pg123
Check Your Understanding Pg 125
Read Pg 127
Chapter 3 Review Pg 128
Unit 1 Review Pg 134
TEST
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