Transcript Document

Science 10 Chemistry
Unit A
Study
Complete
homework
Review
Success
Significant Digits
• Any digit from 1-9 is significant
• Sandwich zeros are significant
– E.g. 2.04, 1005.002
• Trailing zeros are significant
– E.g. 6.3800, 12 000
• Leading zeros are not significant
– E.g. 0.0065
• counted objects and constants are not
included
• (+ or -) :
– round to the lowest number of decimal
places
2.35g + 16.77g + 12.1g = 31.22g
31.2g
• (x or  ):
– round to the lowest number of sig digs
100 g  53.29 g/mol = 1.8765 mol
1.88 mol
WHMIS
•Workplace Hazardous Materials Information System
•WHMIS provides guidelines for handling,
storage and disposal of reactive materials
Know your WHIMIS SYMBOLS!
compressed
gas
poisonous and
infectious
material
causing
immediate and
serious toxic
effects
corrosive
poisonous
and
infectious
causing
other toxic
effects
flammable
and
combustible
Dangerously
active material
oxidizing
material
biohazardous
infectious
material
Matter
• It is anything that has mass and
occupies space
• Has properties which describe the
substance
• Two types of properties
– physical
– chemical
• It can be classified as being a mixture
or a pure substance
Substance
(Matter)
Pure Substance
Mixtures
Physical Properties
• Describe its appearance or features you
can measure:
density
color
boiling point
melting point
Chemical Properties
• Describes how it reacts
with others substances
• It cannot be tested without
destroying the substance
E.g. combustion, rusting,
decomposition reactions
Classification of Matter
Substance
(Matter)
Pure Substance
Elements
Metals
Metalloids
Non-metals
Mixtures
Compounds
Ionic
Heterogeneous
Colloids
Homogeneous
Suspension
Alloys
Molecular
Solutions
1. Mixtures
• Are created when 2 or more substances
combine
• There are 2 types:
 Heterogeneous
 Homogenous
A. Heterogeneous Mixture
• 2 or more phases are visible, or two or
more substances are visible
• E.g. chicken soup, pulpy orange juice
• There are two types:
 Suspension
 Colloid
Suspension
• Is a mixture in which the components are
in different states
• E.g. Mud (water & sand)
Colloids
• A mixture in which the components can’t
be easily separated
• E.g. Milk or salad dressing
B. Homogeneous Mixture
• A mixture of 2 or more substances that
appear as one….( uniform properties)
• E.g. kool-aid, coffee, air
• There are 2 types:
 Alloys
 Solutions
Alloys
• A homogenous mixture of 2 or more
metals
• E.g. brass – copper-zinc
Solutions
• A mixture where you can not see the different
parts
2. Pure Substances
• The composition does not change
• They can be:
 Elements
 Compounds
A. Elements
• Cannot be broken down further
• contains only one type of atom
• E.g. Gold
• Have 3 categories:
 Metals
 Nonmetals
 metalloids
Metals
• Are:
 ductile (can be stretched into wire)
 lustrous (shiny)
silver
 conductors
 malleable
 found to the left of the staircase line
 make up approx. 80% of the elements
Non-Metals
• Are:
 non-ductile
 dull
Sulphur
 non-conductors
 brittle
 found to the right of the staircase line
 make up approx. 20% of the elements
Metalloids
• Are elements that have
properties of both metals and
non-metals
arsenic
 E.g. carbon – dull, conducts
silicon – lustrous, nonconductor
silicon
B. Compounds
• Are two or more elements bonded
together in fixed proportions
• E.g. H2O , C6H12O6
The Periodic Table
• It was developed by Dmitri Mendeleev in
the mid 1800’s
Atomic
Number
Ionic
Charge
Element
Symbol
Atomic
Mass
Element
Name
• It follows certain rules (patterns) called
the “Periodic Law”
• It is arranged into groups (families) and
periods
Period
• Are the horizontal rows on the periodic
table
Period 1 - 2 elements
Period 2 - 8 elements
Period 3 - 8 elements
Period 4 - 18 elements
Period 5 - 18 elements
Period 6 – 32 elements
Period 7 – 32 elements
Groups (Families)
•
Vertical columns
•
Have similar properties
(behave in a similar manner)
•
The number indicates the # of electrons
in the last energy level
•
Have 2 numbering systems :
a) Arabic #’s (1 -> 18)
b) Roman numerals & letters (IIA ,VIIB)
Group IA
• Alkali Metals (E.g. Li, Na, K)
• most reactive metals
• never found in pure form in nature
• all have one electron in the last energy
level
Group IIA
• Alkaline Earths (E.g. Mg, Ca, Ba)
• very reactive metals
• all have 2 electrons in the last energy
level
Group VII A
• Halogens E.g. (Cl2, Br2, )
• most reactive non-metals
• all are diatomic elements E.g.( F2, I2, At2 )
• are one electron short of a full energy level
Group VIII A
• Are Noble Gases
• non-reactive elements (inert gases)
• all have full energy levels
Series
•
Are at the bottom of the periodic table
•
Lanthanum (Lathanides)
 are rare earth metals
 located 57-71
•
Actinium (Actinides)
 Are located 89 to 103
Transition Elements
• Belong to Group B
• Are located in the middle of the table
magnetic elements are found here
Trends of Transition Metals
• metals:
– get more reactive as you move
(R
L) and
– The most reactive metal is FRANCIUM
• Non-metals :
– get more reactive as you move
(L
R) and
– this excludes noble gases
Atomic Structure
• Describes what an atom is made up of
• http://www.youtube.com/watch?v=07yDiEL
e83Y&feature=related
• A Theory was proposed by Dalton in
1808
Dalton’s Theory
1. All matter is composed of tiny,
indivisible particles called atoms.
2. Atoms of an element have
identical properties
3. Atoms of different elements have
different properties
4. Atoms of two or more elements
can combine in constant ratio to
form new substances.
Know this: Page 22 Textbook
J.J Thompson: 1897
• credited with discovery of
electrons
• “raisin bun” model or “plum
pudding” model
• atom is a sphere which is
positive, with negative
electrons embedded in it like
raisins in a bun
• most of the mass is
associated with the positive
charge
Ernest Rutherford: 1911
• atoms have a nucleus
which is positive and has
most of the mass
• most of the atom is
empty space occupied by
the moving negatively
charged electrons
• proposed the existence
of protons
http://www.learnerstv.com/animation/animati
on.php?ani=121&cat=chemistry
Neils Bohr: 1913
• electrons move in circular orbits around the
nucleus
• cannot exist between orbits
James Chadwick: 1932
•
showed that the nucleus must contain heavy
neutral particles to account for all of the atom’s
mass (neutrons)
Schrodinger/de Broglie: 1930
• quantum mechanical model
• electrons have distinct energy levels
• exact locations of electrons are not defined, but
the probable location in a region of space can be
predicted
Atoms
• Are the smallest part of an element
which retains the chemical and physical
properties of an element.
• They are neutral
• It is made up of 3 sub-atomic particles:
–Protons
–Neutrons
–Electrons
Protons (P+)
• Are large positively charged particle in the
atom’s nucleus
• They make up 99% of the mass of the atom
• the number of protons determines the
atomic number
• E.g. Cu has 29 protons… atomic number =29
Neutrons (N0)
• Are large particles in the atom’s
nucleus
• They hold the nucleus together
• They have no charge
Electrons (e-)
• Are the smallest particle in an
atom
• They have a negative charge
• Are located outside the nucleus
of the atom
• They take up most of the space
• They are arranged in energy
levels
• The # of electrons in each level is :
• Level # of e• 1
2
• 2
8
• 3
8
• 4
18
You MUST
MEMORIZE
THIS!!!!!
Atomic Number
• Is the # of protons in an atom
E.g. oxygen:
atomic # = 8  it has 8 protons
Atomic
Number
Ionic
Charge
Element
Symbol
Atomic
Mass
Element
Name
Mass Number
• Also called the atomic mass
• its the sum of the protons and neutrons
(These are averages)
Atomic
Number
Ionic
Charge
Element
Symbol
Atomic
Mass
Element
Name
• E.g. Lithium
 Atomic Number
 Atomic Mass
3
6.9
(so it has 3 Neutrons)
(round it)
• Try: Mercury
 Atomic Number
 Atomic Mass
___ (so it has _ Neutrons)
___
(round it)
• Try: Chlorine
 Atomic Number
 Atomic Mass
___ (so it has _ Neutrons)
___
(round it)
Finding The Number of Neutrons
• Use the Formula:
 Neutrons = Atomic Mass – Protons
• E.g. Oxygen
 Mass =
 Protons =
____
____
 Neutrons = Atomic Mass – Protons
 Neutrons =
_____ - ____
 Neutrons =
_____
Try the Following
• E.g. Calcium
 Mass =
____
 Protons = ____
• Neutrons = Atomic Mass – Protons
• Neutrons =
_____ - ____
• Neutrons =
_____
Isotopes
• Have a different atomic mass
• They have a different number of
neutrons…. but same number of protons
• E.g.
Copper has 35 neutrons ..but an
isotope of copper can have 36
Isotope Notation
x
A
z
A = symbol
x = mass # (p+ + n°)
z = atomic # (p+)
Try copper – 64
(together)
Isotope Notation x
z
A
A = symbol
x = mass #
z = atomic #
(#p+ + n°)
(#p+)
copper - 64
__
Cu
__
Try copper - 62
# p+ = ____
#n° = ___ - ___ =___
# e- = ____
# p+ = _____
#n° = _____
# e- = _____
Try the Following
• Write the isotope notation for the following:
 Cobalt - 61
 Cobalt – 60
EELR - Electron Energy Level
Representations
• It is a diagram that shows the following:
– The Nucleus
– Energy Levels
– Valence electrons
• Energy Levels:
 show the number of electrons in each
level
 number of levels = the period #
• Valence Electrons:
 Are electrons in the last energy level
 It is the group #
E.g. Sodium
# p+ =
# e- =
#n° =
- atomic #
- mass #
___
___
____
____
____ - ____ = _____ (round)
__ e__ e-
____
____ e-
__ ep+ = ___
n° = ___
Level 1 = 2
Level 2 = 8
Level 3 = 8
Level 4 = 18
E.g. Sodium
# p+ =
# e- =
#n° =
- atomic #
- mass #
11
11
22.99 - 11 = 11.99
1 e8 e-
11
22.99
12
11 e-
2 ep+ = 11
n° = 12
Level 1 = 2
Level 2 = 8
Level 3 = 8
Level 4 = 18
Krypton
# p+ =
# e- =
#n° =
- atomic #
- mass #
_____
_____
___
___
___ - ____ = ____ round
___
___ e___ e___ e___ ep+
= ___
n° = ___
____ e-
Level 1 = 2
Level 2 = 8
Level 3 = 8
Level 4 = 18
Krypton
# p+ =
# e- =
#n° =
- atomic #
- mass #
36
83.80
36
36
83.80 - 36 = 47.80
48
18 e8 e8 e2 ep+
= 36
n° = 48
36 e-
Level 1 = 2
Level 2 = 8
Level 3 = 8
Level 4 = 18
Assignment
Draw an EELR for each of the following atoms
(a)
(c)
(e)
(f)
berylium
chlorine
calcium
magnesium
Ions
• Are atoms or group of atoms that have:
– become electrically charged
– gained or lost electrons to be isoelectronic with
a noble gas (same configuration)
• They can be:
– Monatomic
– Polyatomic
Their outer orbital is filled
A. Monatomic Ions
• Are made of one element
• Can be:
– Cations
– Anions
Cations
• Are positively charged
• Are METALS
• lose electrons …….to obtain a stable
electron configuration (like a noble gas)
• E.g. Sodium ( Na+1 )
•
•
•
•
E.g. Sodium ion
Losses 1 electron
It tries to have the same configuration as Neon
Na+1 has 10 electrons
•
•
•
•
Try: Calcium Ion
Losses _____ electron(s)
It tries to have the same configuration as _____
Ca 2+ has ________ electrons
Anions
• Are negatively charged
• They are NON-METALS
• gain electrons….. to obtain a stable
electron configuration (like a noble gas)
• They have an “ide” ending
• E.g. Sulfide (S 2-)
•
•
•
•
E.g. Fluorine ion
Gains 1 electron
It tries to have the same configuration as Neon
F-1 has 10 electrons
• Try: Oxygen Ion
• Gains_____ electron(s)
• It tries to have the same configuration as _____
• O2- has ________ electrons
B. Polyatomic Ions
• Are a group of atoms that act like
anions. (Except: ammonium…it acts like a cation)
• They are located at the top of the periodic
table
They don’t use bi anymore… use
hydrogen instead
Different Ways of Writing Acetate
−
CH3COO or C2H3O
−
EELR of Ions
• It is a diagram that shows the following:
– The Nucleus
– Energy Levels
– Valence electrons
• The only difference is the electrons!!!
Electrons = # of protons - the charge
Nitride
N3-
# p+ =
# e- =
#n° =
- atomic #
- mass #
_____
_____
____
____ - (-___) = ___
_____ - ___ = ____
___ e___ ep+ = ___
n° = ___
____ e-
______
Nitride
N3-
# p+ =
# e- =
#n° =
- atomic #
- mass #
7
14.01
7
7 - (-3) = 10
14.01 - 7 = 7.01
8 e2 ep+ = 7
n° = 7
7
10 e-
Cadmium ion
Cd 2+
# p+ =
# e- =
#n° =
- atomic #
- mass #
____
____
____
____ - (+___) = ___
____ - ____ = _____
___ e___e___e___e___ ep+ = __
n° = __
___ e-
___
Cadmium ion
Cd 2+
# p+ =
# e- =
#n° =
- atomic #
- mass #
48
48 - (+2) = 46
112.41 - 48 = 64.41
10 e18 e8 e8 e2 ep+ = 48
n° = 64
46 e-
48
112.41
64
Iron (III) ion - atomic #
- mass #
Fe3+
# p+ =
# e- =
#n° =
___
___
____
____ - (+___) = ___
____ - ___ = _____ ==>
__e__ e__ e___ ep+ = __
n° =___
____ e-
____
Iron (III) ion - atomic #
- mass #
Fe3+
# p+ =
# e- =
#n° =
26
55.85
26
26 - (+3) = 23
55.85 - 26 = 29.85 ==> 30
5e8 e8 e2 ep+ = 26
n° =30
23 e-
Draw an EELR for each of the following ions:
1.
2.
3.
4.
5.
Sulfide
strontium ion
Copper (II) ion
Iodide
Vanadium (V) ion
Elements
• Metallic elements:
– exist as single atoms (monatomic)
– The formula is the symbol followed by
the state
E.g.
[sodium]
Na (s)
• Non-metals:
– Excluding noble gases don’t exist as
single atoms… they are diatomic or
polyatomic
H2
N2
O2 F2
P4
S8 Cl2
Br2
polyatomic
I2
Chemical Bonds
• Are interactions that occur between
atoms or molecules
• There are two types:
 Covalent
 Ionic
Covalent
• Are formed when electrons are shared
between atoms
• It occurs between non-metals
Each needs one electron to complete
it’s 1st orbital
1+
1+
Hydrogen
Ions
Single Bond
1+
1+
Hydrogen Molecule
H -- H
H2
Ionic
• Are formed when
electrons TRANSFER
from one atom to another
• It occurs between a metal
and a nonmetal.
Compounds
• are two or more substances that are
held together by chemical bonds
• There are two types:
– Molecular
– Ionic
Molecular
• Are solid, liquid or gas at room
temperature
• Contain only non- metals
• Have a covalent bond
• Don’t conduct electricity
• may dissolve in water to produce either
(a) neutral molecular solution
(b) acidic solution
Ionic Compounds
• Are crystal solids at room temperature
• occur because of a force of a attraction
between the + and - ions
• contain a metal and a non metal
• have an ionic bond
• Have high melting point and boiling
points
• Conduct electricity in water
• May dissolve in water to produce either a
(a)neutral ionic solution
(b)basic solution.
Ionic & Molecular Compounds
Ionic
Molecular
cation + anion
(metal + non-metal)
all elements are non metals
NO PREFIXES
PREFIXES
ALL SOLIDS AT R.T.
SOLIDS, LIQUIDS & GAS
Solutions conduct
electricity
solutions do not conduct
electricity
solutions are basic or
neutral
solutions are acidic or
neutral
ex. NaCl
ex. OF2
Identify if it is M or I
• A yellow gas forms a neutral
solution_____
• A purple solid dissolves in water to
produce a conducting solution
_____
• A white solid dissolves in water and turns
red litmus paper blue ______
(solutions)
• A yellow gas forms a neutral
solution____M_
• A purple solid dissolves in water to
produce a conducting solution
__I___
• A white solid dissolves in water and turns
red litmus paper blue __I____
Naming
&
Writing Formulas
For
Molecular & Ionic Compounds
Naming Molecular Compounds
Non-metal + Non-metal
• Step 1:
 write the element name for the first
non-metal
 add a prefix ….(if there is more than
one atom)
• Step 2:
 Write the second element name and
add an ide ending
 Add a prefix
Exceptions
• DO NOT use any prefixes at all if the first
element is hydrogen ……these are acids
Prefixes
6=
hexa
7=
hepta
3=
mono
di
tri
8=
octa
4=
tetra
9=
nona
5=
penta
10 =
deca
1=
2=
Examples
• P4O10
tetraphosphorus decaoxide
• BrH7
bromine heptahydride
• H 2S
hydrogen sulfide
Try the Following
• CO(g)
carbon monoxide
• CO2(g)
carbon dioxide
• N3F8
Trinitrogen octafluoride
Writing Formulas for Molecular
Compounds
• Steps:
 Write each elements symbol
 Write the subscript number (the prefix)
• E.g.
 dinitrogen oxide
 disulfur oxide
N2O
S20
Try the Following
• oxygen dibromide
OBr2
• diphosphorus pentasulphide
P2S5
• carbon tetraiodide
CI4
• phosphorus pentachloride
PCl5
Molecular Compounds that Must be memorized !!!
NH3 ( g) = ammonia
H2O ( l) = water
H2S ( g) = hydrogen sulphide
CH4 ( g) = methane
CH3OH ( l) = methanol
C2H6 ( g) = ethane
C2H5OH ( l) = ethanol
C6H12O6 ( s) = glucose
C12H22O11 ( s) = sucrose
O3 ( g) = ozone
H2O2 ( l) = hydrogen peroxide
Naming Binary Ionic Compounds
Metal + non-metal
• DO NOT USE PREFIXES
• Steps
1.Write the metal 1st
2. Write the non-metal 2nd with an ide
ending
E.g.
NaF sodium fluoride
Na2S sodium sulphide
two sodium ions are bonded with
one sulphide ion… this doesn’t
matter for naming ionic compound
Try the Following
•
•
•
•
•
•
•
LiF
KCl
BeS
Rb3P
MgF2
Na2O
CsBr
lithium fluoride
potassium chloride
beryllium sulphide
rubidium phosphide
magnesium fluoride
sodium oxide
cesium bromide
Try the Following
• KCl
potassium chloride
• MgBr2
magnesium bromide
• Ba3N2
barium nitride
• ScP
scandium phosphide
Writing Formulas for Binary Ionic
Compounds
• Steps
1.Look up the symbol for each…& write the
metal first
2.Balance the charges
(total + charges = total – charges)
3.Use subscripts to show the # of each
element
sodium oxide
1+
2  Charge
Charge
1+  2 = 2 +
2  1 = 2
Na2O
calcium phosphide
2+
Charge
2+
3=
3  Charge
32=6
6+
Ca3P2
Try the Following
• magnesium chloride
MgCl2
• calcium chloride
CaCl2
• zinc sulphide
ZnS
• silver sulphide
Ag2S
• germanium oxide
GeO2
• calcium arsenide
Ca3As2
• magnesium nitride
Mg3N2
Try the Following
• lithium iodide
LiI
• zinc fluoride
ZnF2
• strontium phosphide
Sr3P2
• silver oxide
• germanium arsenide
Ag2O
Ge3As4
Naming Multivalent Ionic
Compounds
transition metal + non-metal
• Transition metal ions have more than one
possible charge
Cu2+, Cu+, Fe3+, Fe2+
• Steps
1. Write metal 1st with the charge in roman
numerals
Roman Numerals (I,II,III,IV,V,VI,VII)
2.Write non-metal second
remember the charges have to balance
Examples
uranium (VI) fluoride
U6+
F–
UF6
chromium (III) nitride
Cr3+
N3-
CrN
cobalt (II) chloride
Co2+
ClCoCl2
Try the Following
• AuBr
gold (I) bromide
• CrCl2
chromium (II) chloride
• Co2O3
cobalt (III) oxide
• VS2
vanadium (IV) sulphide
• PuN2
plutonium (VI) nitride
Naming Complex Ions
Metal + complex ion
• Steps:
1. Name the metal ion
2. Name the complex ion
E.g.) PO43
Note: NH4+ (ammonium ion) is the
only positive complex ion…it
will take the place of a metal
Examples
• CaCO3
• Ba(OH)2
• (NH4)3N
Solutions
• CaCO3
Calcium carbonate
• Ba(OH)2
Barium hydroxide
• (NH4)3N
Ammonium nitride
Try the Following
• KIO3
potassium iodate
• NaCH3COO
sodium acetate
• MgSO3
magnesium sulphite
• NH4NO3
ammonium nitrate
• Ca3(PO4)2
calcium phosphate
Writing Formulas For Complex Ions
• Steps:
1. Look up the symbol for each ion
2. Balance the charges
• Note: if you need more than 1 complex ion
to balance the charges use brackets
2+
1-
• E.g. Ca(CH3COO)2
Try the Following
• aluminum phosphate
AlPO4
• calcium sulphite
CaSO3
• scandium acetate
Sc(CH3COO)3
• ammonium sulphate
• nickel (II) phosphate
• aluminum chlorate
(NH4)2SO4
Ni3(PO4)
Al(ClO3)3
2
Solubility
Will the compound dissolve in water?
Soluble
• Refers to whether or not the compound
dissolves in water
• If it is…. the compound is aqueous (aq)
• All acids are soluble
• Some ionic compounds are soluble… the
rest are solids
Is It soluble?
• This will apply to ionic compounds (only)
• Steps
1. Find each ion in the boxes across the
top
2. if it is soluble it will have (aq) aqueous
3. If it does not dissolve it will have (s)
solid.
Determine if the following compounds are
soluble in water. Use the proper subscript
to indicate the state.
•
•
•
•
•
•
•
•
AgCl
BaCO3
LiOH
Ca2(PO4)3
NaCl
CaI2
Pb(NO3)2
HMnO4
• AgCl
(s)
• BaCO3
(s)
• LiOH
(aq)
• Ca2(PO4)3 (s)
• NaCl
(aq)
• CaI2
(aq)
• Pb(NO3)2
(aq)
• HMnO4
(aq)
Acids & Bases
Acids
• They are always soluble in water
• Conduct electricity
• Taste sour
• React with metals to produce hydrogen
gas (H2(g))
• Neutralize a base
• they ALWAYS have hydrogen ….usually
as the first element
• E.g. HCl(aq) , H3PO4(aq)
• There are three types:
 Binary
 Oxo
 Organic
Binary Acids
• contain only H and one other element (Cl,
Br, etc.)
• E.g. HCl (aq)
Oxo Acids
• Contain H and Oxygen
• E.g. H3PO4(aq)
Organic Acids
• Contain C, H, & O
• The H is written at the end
• All have COO• E.g. CH3COOH(aq) – acetic acid
C6H5COOH(aq) - benzoic acid
HOOCCOOH(aq) – oxalic acid
Acid Indicators
• Turns blue litmus paper red
• Able to turn bromothymol blue to yellow
• Phenolphthalein remains
colorless
• E.g. lemon juice
Bases
• Are usually soluble in water
• Conduct electricity (not weak ones)
• Neutralize acids
• Taste bitter
• Usually solids
• Feel slippery
Base Indicators
• Turns red litmus paper blue
• Bromothymol blue remains blue
• Turns phenolphthalein pink
• E.g. baking soda, Rolaids,
soap, Draino crystals
Naming Acids
• Steps:
 Hydrogen ____ide becomes hydro____ic acid
 Hydrogen ____ate becomes _______ic acid
 Hydrogen ____ ite becomes _______ ous acid
Examples
• HF (aq)
hydrogen fluoride = hydrofluoric acid
• H2SO3 (aq)
hydrogen sulphite = sulphurous acid
• H3BO3 (aq)
hydrogen borate = boric acid
• HCl (g)
Hydrogen chloride (not an acid)
Try the Following
hydrosulphuric acid
phosphorus acid
carbonic acid
hydrogen sulphide
1+
2-
H2S (aq)
hydrogen phosphite
1+
3-
H3PO3 (aq)
hydrogen carbonate H2CO3(aq)
1+
2-
Writing Acid Formulas
• Steps:
1. Use the naming rules in the opposite
direction
• Example:
hydrosulphuric acid hydrogen sulphide H2S(aq)
Try the Following
• carbonic acid hydrogen carbonate H CO
2
3(aq)
• chlorous acid hydrogen chlorite
HClO2(aq)
Naming Bases
• Steps:
1. Write the metal name 1st
2. Write hydroxide or bicarbonate
E. g.
NaOH
sodium hydroxide
Try the Following
• KOH
• Ba(OH)2
• NaHCO3
potassium hydroxide
barium hydroxide
sodium bicarbonate
Chemical Reactions
• Can cause a physical or a chemical
change
• Always results in the formation of a new
substance
• Evidence:
1. Temperature change
2. Formation of a precipitate
3. Colour change
4. Gas produces
Reactants
Products
balancing
1 Zn(s) + 2 HCl(aq)
states
1 ZnCl2(aq) + 1 H2 (g)
states
Energy Changes
• Can occur in the form of heat, light,
electrical, or mechanical
• There are two types:
 Endothermic
 Exothermic
• Endothermic
 Energy is absorbed (enters)
 Reactants + Energy
products
• Exothermic
 Energy is released (leaves)
 Reactants
product + energy
Balancing Equations
• There must be equal numbers of each
element on both sides of the equation
Use lowest numbers
Example
____Mg(s) + ___ O2(g)  ___ MgO(s)
____ H2O(l)  ____H2(g) + ____ O2(g)
• When chemicals react they follow the Law
of Conservation of Matter:
 Matter can not be created or destroyed
it only changes form
• Mass of reactants = mass of products
Counting Practice
• How many of each element are in the
following compounds?
1. NaCl
5. NH4CH3COO
2. BaBr2
6. 3 (NH4)2S
3. (NH4)3P
7. 2 CaCl2
4. Ba(OH)2
8. 8 PbI2
9. 4 Zn(CH3COO)2
Balancing Practice
____ Cu(s) + ____ AgNO3(aq)  ____ Ag(s) + ___ Cu(NO3)2(g)
___ Cl2(g) + ____ NaBr(aq)  ____ Br2(l) + ____ NaCl(aq)
____ KI(aq) + ____ Pb(NO3)2(aq)  ___ PbI2(s) + ___NO3(aq)
____ CH4(g) + ____ O2(g)  ____ CO2(g) + ____ H2O(g)
Types of Reactions
• There are 5 types of reactions:
 Simple Composition
 Simple decomposition
 Single replacement
 Double replacement
 Hydrocarbon combustion
Composition/ Formation Reactions
• Elements combine to form a compound
Element + element  compound (s)
• These are usually exothermic
E.g. 2 Mg (s)
+
O2 (g)
 2 MgO(s)
Try the Following
Mg(s) + Cl2(g) 
2+
14 Fe(s) + 3 O2(g) 
3+
22 Zn(s) +
2+
O2 (g) 
2-
MgCl2(s)
2 Fe2O3(s)
2 ZnO(s)
Simple Decomposition
• Compound decomposes into its elements
Compound  element + element
• These are usually Endothermic
E.g. 2 H2O(l)
 2 H2(g)
+
O2(g)
Try the Following
2 HMnO4(s) 
2 HCl(g)

Pb(NO3)2(s) 
H2(g) + 2 Mn(s) + 4O2(g)
H2(g) + Cl2(g)
Pb(s) + N2(g) + 3 O2(g)
Single Replacement
• An element reacts with an ionic compound
to form a different element and compound
element + compound
 element + compound
• This occurs in water (use solubility chart)
E.g.
Cu(s) + 2 AgNO3(aq)  2 Ag(s) + Cu(NO3)2(aq)
Example
Cl2(g) + 2 NaBr(aq)  Br2(l) + 2 NaCl(aq)
Try the Following
Zn (s) + 2 NaCl (aq) ZnCl2 (aq) + Na (s)
2
Pb (s) + Cu(NO3)2 (aq)  Pb(NO3)2 (aq) + Cu (s)
Mg (s) + 2 HOH (l)
(g)
+
H
(s)
Mg(OH)
2

2
2 KCl (aq) + Br2 (l)  2 KBr (aq) + Cl2 (g)
Double Replacement
• Two ionic compounds react to form two
different ionic compounds
compound + compound
 compound + compound
• It occurs in water (solubility chart)
• There are two types of reactions:
 Precipitation
 Neutralization
a. Precipitation Reaction
• One product that is formed is insoluble
compound + compound
 insoluble + compound
compound
E.g. Pb(NO3)2(aq) + 2 KI(aq)  PbI2(s) + 2 KNO3(aq)
solid
b. Neutralization
• An acid reacts with a base to form water
and salt
acid + base  water + salt
E.g. 1 HCl(aq) + 1 NaOH(aq)  1 H2O(l) + 1 NaCl(aq)
Hydrocarbon Combustion
• A hydrocarbon is made up of C &H (E.g. CH4)
• Occurs when a hydrocarbon burns in the
presence of oxygen
• it always produces CO2 (g) + H2O (g)
C?H? + O2(g)
 CO2(g)
hydrocarbon
E. g. CH4(g) + 2 O2(g)
+
H2O(g)
 CO2(g) + 2 H2O(g)
Try the Following
6 CO2(g) +
C6H12O6(s) + O
62(g) 
C2H6 (g) + O2(g) 
2 CO2(g) + 3 H2O(g)
2 C2H6 (g) + O
7 2(g) 
2
H
6 2O(g)
C10H22 (l) + 31O2(g) 
4 CO2(g) + 6 H2O(g)
20 CO2(g) + 22H2O(g)
Other
• any reaction that does not follow any of the
above patterns
E.g. Ca(OH)2(aq) + CO2(g)  CaCO3(s) + H2O(g)
HBrO3(aq) +
5 
HBr
HOH
3 (l) + Br2(l) 3
Mole
• is a quantity equal to 6.02 x1023
atoms, ions, molecules, etc…. Which is
Avogadro’s number
• originally defined as the number of
atoms in exactly 12 g of carbon-12
• it is used to determine the mass of all
other elements
Molar Mass
• Is the mass of 1 mole of a substance
Symbol – M
Units – g/mol
Steps:
1. Write the correct chemical formula.
2. List all elements present
3. Determine how many of each element are
present
4. Multiply by the atomic molar mass of that
element.
5. Find the sum.
Example 1:
Determine the molar mass of barium
hydroxide
Ba(OH)2
Ba 1 x 137.33 g/mol= 137.33 g/mol
O
2 x 16.00 g/mol = 32.00 g/mol
H
2 x 1.01 g/mol = 2.02 g/mol
171.35 g/mol
Example 2:
Determine the molar mass of magnesium
nitrate
Mg(NO3)2
Mg 1
2
N
6
O
x 24.31 g/mol
x 14.01 g/mol
x 16.00 g/mol
= 24.31 g/mol
= 28.02 g/mol
= 96.00 g/mol
148.33 g/mol
Finding the # of Moles
m = nM
n = moles (mol)
n =m
M
m = mass (g)
M = molar mass (g/mol)
E.g.1 How many moles are there in a 25.0 g sample of
lithium sulfate?
Li2SO4 =
Li – 2 x 6.94 g/mol
S - 1 x 32.06 g/mol
O – 4 x 16.00 g/mol
109.94 g/mol
n = m/M
= 25.0 g
109.94 g/mol
= 0.227 mol
E.g.2 How many moles are there in a 32.8 g sample of
potassium permanganate?
KMnO4
=
K – 1 x 39.10 g/mol
Mn - 1 x 54.94 g/mol
O – 4 x 16.00 g/mol
159.04 g/mol
n = m/M
= 32.8 g
159.04 g/mol
= 0.20754… mol
= 0.208 mol
Example 3
What mass of copper (II) sulfate is present in a
0.3750 mol sample?
CuSO4
m=?
n = 0.3750 mol
M=
Cu – 1 x 63.55 g/mol
S - 1 x 32.06 g/mol
O - 4 x 16.00 g/mol
159.61 g/mol
m = nM
= (0.3750 mol)(159.61 g/mol)
= 59.85 g
Example 4
What mass of sodium hydrogen carbonate is
present in a 2.50 mol sample?
NaHCO3
m=?
n = 2.50 mol
M=
Na - 1 x 22.99 g/mol
H - 1 x 1.01 g/mol
C - 1 x 12.01 g/mol
O - 3 x 16.00 g/mol
84.01 g/mol
m = nM
= (2.50 mol)(84.01 g/mol)
= 210g