Transcript Unit II - Corner Brook Regional High

Unit II
Physical Science I
Chemistry
Chemistry Terms


Matter - anything that has mass
and takes up space
Chemistry - is the study of
matter, its properties and the
changes or chemical reactions
that matter can undergo.
• Ex. rusting, combustion of fuel/candle
wax, explosion of TNT, vinegar and
baking soda


Pure Chemistry - describing known
substances and discovering new
compounds for research purposes.
Applied Chemistry – the search for
uses for chemical substances.
Modern society demands chemistry
understanding.
• Ex. technology; government.


Mass- The amount of matter an object
contains in grams (g).
States of Matter– 3 physical states:
1) Solid
2) Liquid
3) Gas
definite volume/shape
definite volume/indefinite shape
indefinite volume/shape
(aq)=aqueous; dissolved in H2O

Physical Property- a characteristic
of a substance; can be observed
without changing into a new
substance.
• Ex. State of matter, hardness,
melting point, boiling point, odour,
solubility colour, malleability,
ductility, brittleness, conductivity.

Physical Change- a change in state
of a substance (no new substance
formed).
• Ex. melting (s to l);
evaporation/boiling (l to g);
condensation (g to l);
sublimation (s to g)
H2O
heat
(s)
H2O
(l)

Chemical Property- a characteristic
behaviour of a substance that occurs
when a substance changes into a
new substance.
• Ex.
2 Mg(s) +O2(g)
+ light
energy
2 MgO(s)
Mg ribbon

Chemical Change- a change in which
one or more NEW substances are
formed.
• Ex.
coal combustion
C(s) + O2(g)
2 CO2(g) + heat
rusting
4 Fe(s) +
3 O2(g)
Fe2O3(g)

Indicators of Chemical Change –
new colour
 heat/light given off
 Bubbles of gas
 Precipitate (solid) formation


Change is difficult to reverse


Mixture – contains 2 or more pure
substances.
Two Types:
• Homogeneous Mixture- aka
“solution”- have only one visible
phase throughout.

ex. air, apple juice, salt water
• Heterogeneous Mixture- contain 2 or
more visible components or phases

ex. soil, soup

Pure Substance – made up of only one
type of atom or atom combination.
(Ex. O2, H2O)
Stays the same in response to physical
change.
Two Types: Compounds and Elements

Compounds - pure substances that
contain two or more different
elements in a fixed proportion
ie., CO2, H2O, C6H12O6 and NaCl
Can be broken down to elements via
chemical means.
Ex. 2 NaCl
electricity
(l)
2 Na
(l)
+ Cl2
(g)

Elements - pure substances that
CANNOT be broken down into simpler
substances by regular laboratory
conditions; made up of 1 type of atom.
• ie., oxygen, nitrogen, carbon and
phosphorus

Element Symbols are always written
with the first letter uppercase and the
second letter lowercase.
• Ex. Au, Mg, Ar

Element Names are always written in
lowercase letters.

Diatomic Molecules – There are 7
elements that are diatomic gases in
their natural state.

These are: H2 O2 F2 Br2 I2 N2 Cl2
Also P4 and S8

How can we remember these?

• HOFBrINCl PS! (or an upside down “L”
on the periodic table).
Matter Flow Chart
Matter
Pure Substance
Compound
Mixture
(Ionic, Molecular,
Acids)
Homogeneous
Element
atom
(Metals,
Nonmetals)
(solutions)
Heterogeneous

Reactants – starting materials.

Products – new substances formed.

Chemical Reaction
Reactants
“go to form”
Products
Periodic Table


Mid 1800s- 65 known elements.
Began to recognize patterns after
recording reactivity, masses, etc.
Dmitri Mendeleev (1834-1907)
 Wrote out elements in order of
increasing atomic mass, result
was a table.
 “periodic” table- “periodic” meaning
repeating patterns and properties.


We now organize the periodic table
according to atomic number
Also organized according to # of
electrons (e-) in atoms of each
element.
Periodic Table- A Review



It is designed to arrange elements in
a pattern that helps us predict
properties and bonding patterns of
elements.
Elements are organized by Atomic
Number and Number of Electrons.
The periodic Table is arranged in
rows and columns.



Period: horizontal row (7 in total)
: atomic mass and atomic number
increase( )from L to R.
Group/Family: vertical columns (18)
:Elements of the same group
have similar but not identical properties.
Some groups have species names:
Group 1- Alkali Metals
Group 17- Halogens
Group 2- Alkaline Earths
Group 18- Noble Gases
• Lanthanides (rare earth)
• Actinides

Groups have 2 numbering systems:
• New Group 1-18
• Old Roman Numerals/Letters
 IA-VIIIA - Representative Elements
 IB-VIIIB - Transition Elements.
Metals and Nonmetals
1
2
3
H
He
1
2
Li
Be
B
C
3
4
5
Na Mg
11
4
K
19
5
7
Ca Sc
O
F
Ne
6
7
8
9
10
Al
Si
P
S
Cl
Ar
13
14
15
16
17
18
Ti
V
Cr Mn Fe Co Ni Cu Zn Ga Ge As Se Br
Kr
23
24
35
36
I
Xe
53
54
20
21
22
Rb Sr
Y
Zr Nb Mo Tc Ru Rh Pd Ag Cd
In
39
40
41
42
49
Hf
Ta
W
72
73
74
37
6
12
N
38
Cs Ba
55
56
Fr
Ra
87
88
*
W
Nonmetals
25
26
27
28
29
30
METALS
43
44
Re Os
75
76
47
45
46
Ir
Pt Au Hg
Tl
77
78
81
79
48
31
80
32
33
34
Sn Sb Te
50
51
Pb Bi
82
83
52
Po At Rn
84
85
86
Rf Db Sg Bh Hs Mt
104
105
106
107
108
Metalloids
109
La Ce Pr Nd Pm Sm Eu Gd Tb Dy Ho Er Tm Yb Lu
57
58
59
Ac Th Pa
89
90
91
60
U
92
61
62
63
64
65
66
Np Pu Am Cm Bk Cf
93
94
95
96
97
98
67
68
69
70
71
Es Fm Md No Lr
99
100
101
102
103
Types of Elements
The staircase line divides the elements into two
major categories: the metals and the nonmetals
The ratio of metals to nonmetals is about 4:1
The Metals
Metals are shiny, electrically conductive
elements.
They are also malleable (can be hammered into
shapes) and ductile (can be stretched into wire).
With the exception of mercury, they are all
solids at room temperature (25°C).
The Nonmetals
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Nonmetals are dull and are very poor
conductors or nonconductors of
electricity.
The solid nonmetals are brittle.
As a group, the nonmetals exhibit the
three states of matter at room
temperature.
eg, carbon is a solid, nitrogen is a gas,
and bromine is a liquid.
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The location of hydrogen in the
periodic table is unusual.
Hydrogen is a nonmetal, but in some
periodic tables it is located in the top
left hand corner of the periodic table
(i.e. on the metals side).
Due to the fact that hydrogen has
some metallic properties in addition
to nonmetallic properties.
2 groups: Alkali Metals and Halogens
(1 and 17).
The Metalloids



The metalloids are elements that
possess both metallic and
nonmetallic properties.
For example, silicon is shiny and
conducts electricity (like a metal) ,
but it is brittle (like a nonmetal).
Metalloids are also known as the
semimetals.
Representative Elements



The representative elements illustrate
the entire range of the properties of
the elements.
[Group 1, 2, 13-18].
Sometimes known as the group A
elements, they are organized into
chemical families based on their
specific chemical and physical
properties.
The Transition Elements

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The transition elements are all metals.
They are different from the representative
elements because of their electron
arrangements which in turn gives them
properties that are a little different from the
metallic representative elements.
Inner Transition Elements:
(Lanthanides, Actinides) Same as transition,
but removed from main table as a matter of
convenience in organizing table.


Elements in the periodic table are
organized based on shared
properties.
Metalloids
(staircase)
Molecular Substances


Exist as groups of atoms called
molecules
Molecules are substances composed
of nonmetallic elements
Nitrogen – N2
Methane – CH4
Mono-atomic molecular elements
Noble Gases (group VIIIA or 18)

He
Helium
Ne
Neon
Ar
Argon
Kr
Krypton
Xe
Xenon
Rn
Radon
You must memorize these
Do you know this one?
neon
Ne
Diatomic molecular Elements

eg
All have two identical atoms
O
iodine
=
I2
How about this one?
nitrogen
N2
Memorize the diatomic
molecular elements
Hydrogen
Oxygen
Fluorine
Bromine
Iodine
Nitrogen
Chlorine
H2
O2
F2
Br2
l2
N2
Cl2
Just remember the famous
chemist Dr. HOFBrINCl
remember the……
Molecular Elements
ozone
Sulfur
Phosphorus
(red)
Phosphorus
(white)
O3
S8
P4
P10
Molecular Compounds
Consist of two or more nonmetallic
elements
2 types of molecular compounds
1. Binary molecular compounds
2. Ternary molecular compounds

Binary vs ternary Molecular Compounds
Type of
Compound
Molecular
formula
name
Binary
H2O
water
Binary
H2O2
hydrogen peroxide
Binary
NH3
ammonia
Binary
CH4
methane
Ternary
CH3OH
methanol
Ternary
C2H5OH
ethanol
Ternary
C12H22O11
sucrose
Memorize the names and formulas of common
molecular substances as per the chemistry facts sheet
Binary Molecular Compounds
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Composed of 2 nonmetals
CO2 , CCl4 , BF3 are examples
Many are identified by common
names
ie., water = H2O ammonia = NH3
System for naming and writing
formulas established by I.U.P.A.C.
International Union for Pure and
Applied Chemistry

Requires a system of prefixes:
Number
prefix
1
mono
2
di
3
Tri
4
Tetra
5
Penta
6
hexa
7
Hepta
8
Octa
9
nona
10
deca
This table is on
your chemistry
facts page
RULES FOR NAMING BINARY
MOLECULAR COMPOUNDS




Write the name of the first element of
the formula in full.
Shorten the name of the second element
and add the “ide” ending.
Use prefixes to indicate the number of
atoms of each element in the molecular
formula.
The prefix mono on the first name is
optional.
Sample
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
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
Write the IUPAC name for CCl4
The first element is C.
Its full name is carbon.
The second element is chlorine.
Its name is shortened to “chlor”, and the
suffix “ide” is added to give chloride.
The prefix mono (1) is added to carbon,
and the prefix tetra (4) is added to
chloride to give the name:
monocarbon tetrachloride.
The prefix mono can be omitted from the
first element name to give:
carbon tetrachloride.
Sample 2

Name N2O4

Nitrogen oxide

Dinitrogen tetraoxide (tetroxide)
U do

Name:

B2H6

Diboron hexahydride
Chemical Bonding




Molecular compounds like B2H6
are held together by bonds.
A chemical bond is the force of
attraction between atoms.
In molecular compounds the bond is
the force of attraction occurs
between nonmetallic elements
This type of bond is called a
covalent bond
Atomic Theory


Atom – the smallest particle of an
element that retains the properties
of that element.
Atoms are thought to be composed
of negatively charged particles called
electrons and a dense central
region called the nucleus


Within the nucleus are found
positively charged particles called
protons and neutral particles known
as neutrons
Electrons are believed to exist a
specific distances from the nucleus
called energy levels
Electron Energy Level Diagrams
Representative Elements
Show electron arrangements within
an atom’s energy levels.
We can predict them using the following:
1. Atomic number
2. Period number
3. Group number
4. Electrons per energy level

Electron Energy Level Diagrams
Representative Elements
Atomic number
 Atomic # represents the number of
protons in the nucleus of an atom.
 the # protons = # electrons.
 Example: Carbon has atomic # 6
which means C has 6 protons and 6
electrons)
nucleus
6 electrons
6+
Electron Energy Level Diagrams
Representative Elements



Period Number
tells us how many
energy levels contain
electrons
Eg Carbon is in the
second row of the
periodic table thus it
is in period 2 and has
electrons in 2 energy
levels.
2nd
1st
6+
Electron Energy Level Diagrams
Group # (family number)
 The group # tells us about electrons in
the outer energy level of an atom
(valence electrons)
Electron Energy Level Diagrams
Representative Elements


We will deal with the
representative elements
only
(for groups 13 and above
valence electrons = the
last digit):
13, 14, 15, 16, 17, 18
C has 4 valence electrons
since it is group 14
4e-
2nd
1st
6+
Electron Energy Level Diagrams
Representative Elements
Period
1
#of
Energy
elements in level
the period
2
1
max #
of e2
2
8
2
8
3
8
3
8
4
18
4
18
Electron Energy Level Diagrams
Representative Elements
4e2e6+
2nd
1st
Electron Energy Level Diagrams
Representative Elements

Example: Draw an electron energy level
diagram for an atom of aluminum
Group
13
3e-
Period 3
8e2e-
Atomic
number
13
13+
Fill in the
nonvalence
electrons with the
maximums per
energy level
Energy level diagrams - ions


Recall the noble gases are unreactive
All noble gases have filled valence
energy levels




atoms react by changing the number
of electrons to try and get the same
structure of the nearest noble gas
In other words, atoms either gain or
lose electrons to become stable
Metals lose electrons to have the
same electron arrangement as the
nearest noble gas
Nonmetals gain electrons to have an
electron arrangement of the nearest
noble gas.
Ions



Metals lose electrons to form positive
ions called cations
Nonmetals gain electrons to form
negative ions called anions
Both ions have noble gas stability
Ionic compounds

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

Composed of oppositely charged ions
Held together by ionic bonds
3 categories of ionic compounds :
1)
Binary ionic compounds
• simple ions (only single charges)
• multivalent ions (more than one
charge)


2) Polyatomic ions (complex ions)
3) Hydrates
Binary ionic compounds

Binary ionic compounds are
composed of a metal ion (+) and
non-metal ion (-).
Naming binary ionic compounds:
1.
2.
Name the cation (+) by writing the
full name of the metal.
Name the anion (-) by shortening
the name of the element and add
the -ide ending.
Binary ionic compounds
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
NaCl
sodium and chlorine
sodium chloride
CaF2 calcium and fluorine
calcium fluoride
K2O potassium and oxygen
potassium oxide
IMPORTANT: Do Not use prefixes - they are for
molecular compounds (two non-metals)
Rules for Writing Binary Ionic
Formulas:



Write down the symbols of the ions
involved.
Determine the lowest whole
number ratio of ions that will give a
net charge of zero.
Write the formula removing all
charges.
Write a chemical formula for a compound that
Sample
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


contains Calcium ions and Bromide ions.
Write down the symbols of the ions
involved
Calcium is group IIA, Ca2+
Bromide is group VIIA, Br –
Determine the lowest whole number
ratio of ions that will give a net charge
of zero. Use the crossover method
Ca2+ Br –
Ca Br2
Binary Ionic Compounds
The Stock System





Ions of a certain elements can have more
than one possible charge.
Such elements are called multivalent
species.
Example 1: tin forms two common ions:
Sn2+ and Sn4+
The Stock System is used to name ions
like these
Sn2+ is called tin (II) and Sn4+ is called tin
(IV)
Stock System


Example too! Cu+ is copper (I)
Cu2+ is copper (II)
The periodic table lists the ions that have stock
names.
Ion
Cu+
Stock name
Copper (I)
Cu2+
Copper (II)
Fe2+
Iron (II)
Fe3+
Iron (III)
Sn2+
Tin (II)
Sn4+
Tin (IV)
Pb2+
Lead (II)
Pb4+
Lead (IV)
Stock System

1.
2.
3.
4.
Egg Sample: Write the chemical
formula for iron(II) chloride.
Write the symbols of the ions
involved
Iron (II) (the roman numeral tells us
it has a 2+ charge) Fe 2+
Chloride (has a 1- charge) Cl –
Determine the lowest whole
number ratio of ions that will give a
net charge of zero.
Use the criss cross method:
Fe 2+ Cl –
FeCl2
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

Write formulas for:
Titanium (IV) fluoride
Titanium (II) fluoride
Nickel (II) oxide
Lead (IV) sulfide
Naming Ionic Compounds Polyatomic ions



polyatomic ion (complex ion) - is a
group of atoms that are covalently bonded
which then gain or lose electrons to
become stable
Example:
The ammonium ion, NH4+, consists of
one nitrogen atom and four hydrogen
atoms which as a group have lost one
electron.
Table of Complex Ions
Writing Chemical Formulas for
Compounds with Polyatomic Ions:


write the cation symbol first and the
anion symbol last.
balance the charges by providing the
appropriate numerical subscript for
each ion.

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Write the a chemical formula for each
compound:
magnesium chlorate
Mg2+ ClO3put brackets around the complex ion
Mg2+ (ClO3) –
criss cross the charges
Mg (ClO3)2
iron(III) sulfate
Fe2(SO4)3
Naming Ionic Compounds
Ionic Hydrates


An ionic hydrate is a compound that
decomposes upon heating to release water
Water is part of its crystalline structure.
sample
CuSO4●5H2O is
copper(II)sulfate pentahydrate
Each ionic hydrate has two parts
to its name:
A number of
molecules of
water
Ionic salt
CoCl2
●
2 H2O
A
separator

cobalt (II)
chloride
dihydrate
U do



zinc sulfate heptahydrate
potassium sulfate decahydrate
nickel (II) nitrate tetrahydrate
Acids – hydrogen compounds





All are hydrogen compounds dissolved in
water
Acids can be simply defined as
substances that release hydrogen ions
(H+) in water
Substances dissolved in water are
denoted by a subscript (aq) written after
their formula
Eg. HCl(aq)
Acids turn blue litmus red
Rules for Naming Acids:

name the hydrogen compound as if it were an
ionic compound.
•




(all of these compounds should end in - ide, -ate,
or -ite.)
depending on the ending convert the ionic
name to the acid name.
Ionic name
acid name
hydrogen _______ide → hydro _____ic acid
hydrogen _______ate → _________ ic acid
hydrogen _______ite → ________ous acid
Rule #1: hydrogen…ide
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
If the aqueous hydrogen compound
begins with “hydrogen” and ends in
“ide”, then:
Replace hydrogen with hydro and ..
replace the “ide” ending of the anion
with ic acid
Example: HCl (aq)
Hydrogen chloride → hydrochloric
acid
Rule #2: hydrogen…..ate :



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
If the aqueous hydrogen compound
begins in “hydrogen” and ends in
“ate” then:
drop the name hydrogen (do not
replace it)
replace the “ate” ending of the anion
with -ic acid
Example: HClO 3 (aq)
hydrogen chlorate → chloric acid
Rule #3: hydrogen…ite





If the aqueous hydrogen compound
begins in “hydrogen” and ends in
“ite” then:
drop the name “hydrogen” (do not
replace it)
replace the “ite” ending with -ous
acid
Example: HNO2 (aq)
hydrogen nitrite → nitrous acid
Bases


Bases are substances that behave in
opposition to acids.
ionic compounds that contain the
hydroxide ion (OH-).
Sample

Sodium hydroxide – NaOH
(aq)
Properties of Bases:





turn red litmus blue
neutralize acids
have high pH (> 7)
form slippery solutions
tend to have a bitter taste
pH
scale
Chemical Change

communicated in sentence form or
as chemical equations.
Chemical equations have four parts:
1 chemical formulas
2 subscripts for states of matter
 (s) solid
 (l) liquid
 (g) gas
 (aq) aqueous - dissolved in water
3 numerical coefficients
 indicates how many atoms/molecules are
involved
4 reaction symbols
 the "+" sign on the reactants (left) side is read as
"reacts with"
 the arrow ( → ) is read as "to produce"
 the "+" sign on the products (right) side is read
as "along with".
Chemical Equations

Two molecules of diesel fuel react
with 49 molecules of oxygen to
produce 32 molecules of carbon
dioxide and 34 molecules of water.
Evidence for Chemical Change:


Chemical changes involve changes
in make up - new substances are
formed with new properties
Physical changes involve changes
in state without a change in make
up.
Evidence of Chemical Change

1.
2.
3.
4.
4 indicators of chemical reaction:
energy change
colour change
precipitate formation
gas formation
1.
2.
3.
4.
5.
a rock warmed by the sun all day
loses its heat at night
milk goes sour when left out of the
fridge
bubbles form in a glass of cold
water as it warms
bubbles and steam rise out of a
kettle of boiling water
paint dries on a hot day.
The Law of Conservation of Mass



In a chemical reaction the mass of
the reactants before a chemical
reaction equals the mass of the
products after the reaction is
complete.
Antoine Lavoisier
placed mercury(II) oxide powder (a red
powder) in a test tube, sealed it, and then
weighed it carefully



heated it and observed that the red powder
gradually changed into a grey liquid
reweighed the sealed tube after the reaction
was complete and observed that its mass
had not changed
opened the tube and noticed a rapid release
of a gas which was later learned to be
oxygen. The grey liquid was mercury metal.
Balancing Chemical Equations




All representations of a chemical
change must reflect the law of
conservation of mass
Chemical equations must obey the law
of conservation of mass
The number of atoms of each element
must be the same on both sides of the
equation
For this reason all chemical equations
we write must be balanced.
Balancing Chemical Equations
sample









Mercury(II)oxide forms mercury and oxygen
Reactants
Products
HgO(s)
→ Hg(l) + O2(g)
1 Hg
1 Hg
1O
2 O equation is unbalanced
2HgO(s)
→ 2 Hg(l) + O2(g)
2 Hg
2 Hg
2O
2 O equation is balanced
2HgO(s)
→ 2 Hg(l) + O2(g)