Chemistry
Sections 5.1, 5.2, 5.5, 5.6, 5.7, 5.8,
5.9, 5.11, 5.12, 5.13, 6.1, 6.5, 6.6, 6.7,
6.19, 7.1, 7.2, 7.3, 7.5,
An introduction to chemistry
• Chemistry can be defined as the study of
chemicals and their reactions.
• Chemicals may be described by their physical
characteristics or their chemical characteristics;
– Physical characteristics include things like colour, state
at room temp., smell, boiling or melting points.
– Chemical characteristics mean how a chemical reacts
with other chemicals. A chemical change occurs when
a substance changes to a new substance.
Mixtures
• Most chemicals exist in nature as mixtures,
made up of 2 or more substances.
• These mixtures may be either homogeneous
or heterogeneous.
– Homogeneous mixtures are those in which the
components are not distinguishable, is completely
uniform. Ex coffee or chocolate ice cream
– Heterogeneous mixtures are those in which the
components are distinguishable. Ex rocky road ice
cream, stew
• Homogeneous mixtures-
• Heterogeneous mixtures-
Pure Substances
• Are not as common as mixtures, consist of
elements or compounds
• Elements are the simplest form of matter that
can exist under natural conditions. Ex. Hydrogen,
carbon, sodium
• Compounds are pure substances that contain two
or more different elements in fixed proportions.
• Compounds are usually identified with a chemical
formula, a combination of letters and numbers to
tell you what type and how many of each
element is present. Ex. H2O
HO
2
Matter
Matter = any material
substance with Mass
& Volume
Matter
comes in 3 phases
Solid
Gas
Liquid
Solid
Definite Shape
Definite Volume
Liquid
Indefinite Shape –
takes the shape of
the container
Definite Volume
Gas
Indefinite Shape –
takes the shape of
the container
Indefinite Volume –
can expand and be
compressed
Elements one of the 100+ pure substances
that make up everything in the universe
Examples of Elements
C = Carbon
Na = Sodium
O = Oxygen
Ca = Calcium
H = Hydrogen
K = Potassium
N = Nitrogen
I = Iodine
Cl = Chlorine
S = Sulfur
P = Phosphorus
Atom
the smallest particle making up elements
WHMIS
• Stands for Workplace Hazardous Materials Information
System
• Is a system to inform those using or exposed to
chemicals the hazards they may encounter.
• Every chemical used in the school (cleaners included!)
comes with a MSDS (Material Safety Data Sheet) that
describes hazards associated with the chemical,
disposal procedures etc.
• Health Canada
• Complete questions 1, 4, 10 – 12 on page 175
Section 5.5
ELEMENTS AND THE PERIODIC
TABLE
The periodic table
• Organizes elements according to their atomic
structure, physical and chemical properties.
• The columns (up and down) are known as
groups and the rows (across) are known as
periods
• Chemical families are groups of elements that
have similar properties
• We can use the organization of the elements
in the periodic table to predict their reactivity
(how well an element will react)
The Periodic Table
• Interactive Periodic Table
• Periodic Table: Groups and Trends
Elements
An element is a substance made up of only 1 type of atom.
There are about 112 different elements that make up the
periodic table of the elements.
On the periodic table each atom type has its information. For
example…
Atomic no.
Symbol
Name
Mass no.
Periodic Table
Atomic no.
Any atom can be identified by the
atomic no., the symbol or by the
name. For instance...
Name
Symbol
Name
Symbol
Atomic No.
Hydrogen
H
1
Iron
Fe
26
Magnesium
Mg
12
The information from the table can also be shown as:
11
5
B
Questions pg 184- 186
1. Using table 1 on page 185, compare metals to non
metals.
2. Where can metals and nonmetals be found on the
periodic table.
3. Describe the four chemical families of the periodic
table.
4. Fill in the following table about sub atomic particles
Particle
Proton
Electron
Neutron
Location
Charge
Symbol
What it means
The Atomic Number:
11
= number of protons
5
= number of electrons (as an
atom has the same of each)
B
The Mass Number:
= number of protons + neutrons - why are electrons not included in
the mass no?
So for Boron…
Protons =
Electrons =
Neutrons =
What about Phosphorus?
5
5
5.811
Protons =
Electrons =
Neutrons =
15
15
16
Electron configuration
• Electrons travel in orbits or orbitals around
the nucleus. The atomic number on the
periodic table tells you how many electrons
each element has.
• Because atoms are electrically neutral, the
number of electrons equals the number of
protons.
ELECTRON ARRANGEMENT
Electrons are very fast moving. They are arranged in shells around the nucleus. The
first shell fits…
2e
The second fits…
8e
The third fits…
8e
So the electron shell for 12Mg would be…
2, 8, 2
Interactive periodic table
Ionic Bonding
• Na
2,8,1
+ Cl
2,8,7
IONIC FORMULAE
So Mg2+ will be attracted to Cl-.
Because Mg is 2+ and Cl is only 1-, Mg will attract 2 Cl’s.
The compound formed will be MgCl2.
The subscript shows that
there are 2 Cl’s for each Mg.
If the starting ions were Cu2+ and S2-, the 2 ions have the same charge. So each Cu will only
attract 1 S.
The compound formed will be CuS.
There are never any
charges on the final
product - they balance out
• The mass number tells you the mass of the
element and when rounded to the nearest
whole number can be used to determine
the number of neutrons inside the nucleus.
• Mass number – atomic number = # of
neutrons.
6e
2e
• Ex. Oxygen
8p+
Atomic # = 8
8nº
Mass # = 16
Ions
• Elements are most stable where their outer
electron shell or orbit is full.
• Elements whose orbit are almost full lose or
gain electrons, and become ions to achieve
stability
• Elements that gain electrons (and therefore a
negative charge) form anions.
• Elements that lose electrons (and therefore
have a positive charge) form cations.
Anions
• Are formed when non-metals gain electrons.
• What was once a neutral atom becomes a
negatively charged ion.
• The value of the charge is equal to the
number of electrons gained.
Cations
• Are formed when metals lose electrons
• What was once a neutral atom becomes a
positively charged ion.
• The value of the charge is equal to the
number of electrons lost
Naming Ions
• Cations are named by simply stating the
element from which it forms followed by the
word “ion”
– Ex. Sodium ion
• Anions are named by stating the elements
from which it forms and replacing the ending
with “ide”
– Ex. Chloride
Compounds: Ionic Bonding
How do atoms become stable ions?
Ionic bonding animation
Types of Ions…
• Anions…Number of
• Cations…number of
electrons is greater than
electrons is less than
the number of protons
the number of protons.
•
Negative charge
•
Positive charge
Determining ion…general guidelines
• Metals form cations
• Non-metals form anions
Writing formulas…
•
•
•
•
•
•
•
Five step rule…
1. Write the symbol.
2.Write the charges.
3. Cross over the charges from top to bottom.
4. Remove the charge.
5. Simplify the numbers.
Formulas...
Predicting Ionic Charges
Group 1:
H+
Lose 1 electron to form 1+ ions
Li+
Na+
K+
Predicting Ionic Charges
Group 2:
Be2+
Loses 2 electrons to form 2+ ions
Mg2+
Ca2+
Sr2+
Ba2+
Predicting Ionic Charges
B3+
Al3+
Ga3+
Group 3: Loses 3 electrons
to form 3+ ions
Predicting Ionic Charges
Neither! Group 13 elements
rarely form ions.
Group 4: Lose or gain
4 electrons?
Predicting Ionic Charges
N3P3As3-
Nitride
Phosphide
Arsenide
Group 5: Gain 3 electrons
to form 3- ions
Predicting Ionic Charges
O2S2Se2-
Oxide
Sulfide
Selenide
Group 6: gain 2 electrons to
Form 2- ions
Predicting Ionic Charges
F1-
Fluoride
Br1-
Cl1-
Chloride
I1-
Bromide
Iodide
Group 7: gain 1 electron
to form 1- ion
Work for today...
Chapter 5.5
• 1. Describe the alkali metals.
• 2. How are the alkali metals different from the
alkali earth metals?
• 3. Describe the noble gases.
• 4. Describe the halogens.
Chapter 5.6
Do #’s 1, 2, 3, 4
Chapter 5.8
Do #’s 1,2,3,4,5,6
Transition Metals
• Transition Elements include
those elements in the B
families.
• These are the metals you
are probably most familiar:
copper, tin, zinc, iron, nickel,
gold, and silver.
• They are good conductors
of heat and electricity.
Transition Elements
• Transition elements have properties similar to
one another and to other metals, but their
properties do not fit in with those of any other
family.
• Many transition metals combine chemically
with oxygen to form compounds called oxides.
Predicting Ionic Charges
Many Transition metals have more than one possible ionic charge
Iron(II) = Fe2+
Iron(III) = Fe3+
Predicting Ionic Charges
Some transition elements have only one possible charge
Zinc = Zn2+
Silver = Ag+
Transition metals…
• Ionic compounds with transition elements...
Polyatomic Ions
Writing Ionic Compound Formulas
Example: Barium nitrate
1. Write the formulas for the cation and anion, including
CHARGES!
2. Check to see if charges are balanced.
2+
Ba
3. Balance charges , if necessary, using subscripts. Use
parentheses if you need more than one of a polyatomic
ion.
( NO3-)
Not balanced!
2
Writing Ionic Compound Formulas
Example: Ammonium sulfate
1. Write the formulas for the cation and anion, including
CHARGES!
2. Check to see if charges are balanced.
( NH)4+
3. Balance charges , if necessary, using subscripts. Use
parentheses if you need more than one of a polyatomic
ion.
SO42-
2
Not balanced!
Writing Ionic Compound Formulas
Example: Iron(III) chloride
1. Write the formulas for the cation and anion, including
CHARGES!
2. Check to see if charges are balanced.
3. Balance charges , if necessary, using subscripts. Use
parentheses if you need more than one of a polyatomic
ion.
Fe3+ ClNot balanced!
3
Writing Ionic Compound Formulas
Example: Aluminum sulfide
1. Write the formulas for the cation and anion, including
CHARGES!
2. Check to see if charges are balanced.
3+
Al
2
3. Balance charges , if necessary, using subscripts. Use
parentheses if you need more than one of a polyatomic
ion.
2S
Not balanced!
3
Writing Ionic Compound Formulas
Example: Magnesium carbonate
1. Write the formulas for the cation and anion, including
CHARGES!
2. Check to see if charges are balanced.
Mg2+ CO32They are balanced!
Writing Ionic Compound Formulas
Example: Zinc hydroxide
1. Write the formulas for the cation and anion, including
CHARGES!
2. Check to see if charges are balanced.
2+
Zn
3. Balance charges , if necessary, using subscripts. Use
parentheses if you need more than one of a polyatomic
ion.
(OH-) 2
Not balanced!
Writing Ionic Compound Formulas
Example: Aluminum phosphate
1. Write the formulas for the cation and anion, including
CHARGES!
2. Check to see if charges are balanced.
3+
Al
3PO4
They ARE balanced!
Questions…
• Please do questions 1,3,4,6 and 7 on page
198.
Naming Ionic Compounds
(continued)
Metals with multiple oxidation states
• -
• -
some metals form more than one cation
use Roman numeral in name
• PbCl2
• Pb2+ is cation
• PbCl2 = lead(II) chloride
• Roman numeral is equal to the charge of the cation
Complex Ions (Polyatomic Ions)
• Mg 2+, I-, Li +, S2– are all called simple ions or monatomic ions
• Complex, or polyatomic ions, are tightly bound groups of ions that
behave as a unit and carry a charge. Example : sulfate ion. A sulfate
ion is composed of 1 sulfur atom and 4 oxygen atoms. These 5
atoms together form a unit with a charge.
– SO4 2-
• Recognizing complex ions is a key in naming chemical compounds
and writing chemical formulas. Polyatomic Ion Rap...
Polyatomic Ions
•
•
•
•
•
•
•
•
•
•
•
Ammonium……………...
Nitrate……………………
Permanganate…………. .
Chlorate…………………
Hydroxide……………….
Cyanide………………….
Sulfate…………………...
Carbonate……………….
Chromate………………..
Acetate…………………..
Phosphate……………….











NH4+
NO3MnO4ClO3OHCNSO4 2 CO32CrO42C2H3O2PO43-
• cobalt (III) carbonate
– Co2(CO3)3
• beryllium nitrate
– Be(NO3)2
– Polyatomic tutorial...
– Please do #’s 1,2,3,4,6,7 on page 189 Chapter 5.9
Molecular Compounds
Section 5.11
Properties of Molecular Compounds
• Composed of 2 or more non-metals
• Form covalent bonds in which the electrons
are shared (not lost or gained… friendlier!)
• Many are gases at room temperature, they do
not conduct electricity and most are not
soluble in water.
• Ionic and covalent bonds
Naming Molecular Compounds
• Uses Prefixes
–
–
–
–
1 — mono
4 — tetra
7 — hepta
10 — deca
2 — di 3 — tri
5 — penta
6 — hexa
8 — octa
9 — nona
• Example: CCl4 — carbon tetrachloride
– Try: 1. P2O5
2. N2O
3. ICl3
1. Diphosphorous pentaoxide
2. Dinitrogen monoxide
3. Iodine trichloride
Name These
•
•
•
•
•
•
•
N2O
NO2
Cl2O7
CBr4
CO2
BaCl2
H2O
•
•
•
•
•
•
•
Dinitrogen monoxide
Nitrogen dioxide
Dichlorine heptoxide
Carbon tetrabromide
Carbon dioxide
Barium dichloride
Dihydrogen monoxide
Write Formulas for These
•
•
•
•
•
•
•
Diphosphorous pentoxide
Tetraiodine monoxide
Sulfur hexaflouride
Nitrogen trioxide
Carbon tetrahydride
Phosphorous trifluoride
Aluminum chloride
•
•
•
•
•
•
•
P2O5
I4O
SF6
NO3
CH4
PF3
AlCl3
Covalent Bonding
• To illustrate how the bonding occurs between
two non-metals, Lewis Dot Structure is used.
• Only the valence electrons are used and each
valence electron is represented by a dot
• Write the symbol for the element, then draw dots
around the symbol to represent the number of
valence electrons.
• Electrons are placed one on each side going around
the symbol.
It is only the electrons in the outermost orbits that can form
bonds.
Diatomic molecules are molecules that have
only 2 atoms of the same element.
• They prefer to share electrons in covalent bonds than
to exist on their own.
• Element Chemical Symbol Formula and State
• hydrogen
H
H 2 (g)
• oxygen
O
O 2 (g)
(g) = gas
• nitrogen
N
N 2 (g)
(l) = liquid
• fluorine
F
F 2 (g)
(s) = solid
• chlorine
Cl
Cl 2 (g)
• bromine
Br
Br 2 (g)
• iodine
I
I 2 (g)
• “I have no bright or clever friends”
Writing Chemical Formulas for
Molecular Compounds
• There is no need to balance the charges (there
are no ions remember!)
• Simply look at the prefix used in the name to
determine how many atoms of each element
is present.
• Ex. Carbon tetrachloride C – 1, Cl – 4
CCl4
• Ex. Pentaphosphorous Trisulfide , P – 5, S – 3
P5S3
• When using the prefixes to illustrate the
number of atoms in the compound, the rules
are clear;
– NEVER use the prefix mono on the first element
– All compounds , whether ionic or molecular, end
in “ide”
– If by adding the prefix you create a double vowel,
drop the first for ease of pronunciation.
– Please do the following questions on page
204…1,2,3,4,5,6
Acids and Bases…
Acids…
• Acids are sour-tasting, water soluble substances found
in many common products.
• They are very reactive and good conductors of
electricity.
• All acids contain hydrogen atoms in combined form
and when dissolved in water they release H+.
Examples of common acids…
 Vinegar (acetic acid)
Salad dressing
 Citric acid
oranges, lemons
 Acetylsalicylic acid (ASA)
Aspirin
 Sulfuric acid
car batteries
 Carbonic acid
carbonated drinks
Acids
• A dilute acid has lots of water and a small
amount of acid
• A concentrated acid has lots of acid and not
much water so must be handled carefully
• A strong acid releases lots of H+
• A weak acid releases fewer H+
Bases…
• Bases are bitter tasting, water soluble and feel
slippery.
They release hydroxide ions (OH-) when dissolved
in water and are good conductors of electricity.
•
•
•
•
•
Examples of bases…
Sodium hydroxide
Potassium hydroxide
Aluminum hydroxide
Sodium bicarbonate
drain cleaner
soap, cosmetics
antacids
baking soda
• In our home we
often use bases to
clean things… Bleach
and toothpaste
• Some things are not
acids or bases: we
say that they are
neutral…eg. water
Recognizing acids and bases from their chemical formulas…
• 1. Acids are easy! They begin with hydrogen
H2SO4 – sulfuric acid, or H2CO3 – carbonic acid.
• 2. Bases are more difficult. They usually
contain OH but not always, ex. NaOH. An
exception would be NaHCO3 (baking soda) is a
base because it reacts with water to produce
Oh-.
• Questions…
1. What is the most important acid in the chemical industry?
2. What is it used for?
3. #3, 4 on page 295
Chapter 6 - Understanding Chemical
Reactions
• A word equation is one way of representing a
chemical reaction. It tells you what reacts and what
is produced.
• Word equations are written like this:
reactants
products
DO NOT COPY:
When hot steel wool (iron) is put into a bottle of oxygen,
there is a spectacular reaction and iron (III) oxide is
produced. The word equation would be:
iron +
oxygen
iron (III) oxide
• Write the word equation for the following
example:
When zinc is added to hydrochloric acid,
hydrogen and zinc chloride are produced.
Zinc + hydrochloric acid
hydrogen + zinc chloride
• Please do questions 2 and 3 on page 219
6.5 Balancing Chemical Equations
• A skeleton equation represents all chemicals
by their formulas.
Word Equation
methane + oxygen
water + carbon dioxide
Skeleton Equation
CH4
+
O2
H2O
+
CO2
6.3 Conserving Mass
• The Law of Conservation of Mass states that
in a chemical reaction, the total mass of the
reactants is always equal to the total mass of
the products.
How to balance Equations:
• 1. Determine the correct formulas and write the
skeleton equation:
Fe +
O2
Fe2O3
• 2. Count the number of atoms of each element
in the reactants and products. (Polyatomic ions
appearing unchanged on each side are counted
as a single unit).
Type of Atom
Reactants Products
Fe
1
2
O
2
3
• 3. Balance the elements one at a time by
using coefficients. The coefficient is a whole
number that appears in front of the formula.
When no coefficient is written, it is assumed
to be 1.
•
4 Fe +
3 O2
2 Fe2O3
• Always check to be sure that the equation is
balanced.
• Type of Atom
Reactants
Products
• Fe
4
4
• O
6
6
5.
Make sure all coefficients are in the lowest
possible ratio. (Reduce if possible)
Balancing Chemical Equations
Section 6.5
• Because we cannot change the chemical
formulas of compounds in the reaction, we
need to use coefficients to balance the
number of atoms.
• Coefficients are numbers placed in front of the
compound and apply to all elements in the
compound (unlike subscripts which only apply
to that element).
Equation types
• We began by writing word equations
– Iron + Oxygen
Iron II oxide
• Writing chemical formulas based on the word
equations is known as skeleton equations.
– Fe + O2
FeO
• Because of the Law of Conservation of Mass,
we now need to write a balanced equation.
Steps to balancing equations
1. Count the number of atoms of each type in
the reactants and products.
Fe + O2
FeO
Type of Atom
Reactants
Products
Fe
1
1
O
2
1
2. Multiply each of the formulas by the
appropriate coefficient to balance the
number of atoms. Re-write the equation.
2Fe + O2
2 FeO
Tips
• Look for larger molecules (ie polyatomic) or
complex molecules and balance them first,
especially if they appear on both sides.
• If you have an odd number on one side and an
even number on the other, fix the odd side
first.
– Al + O2
Al2 O3 , fix products first
• Leave diatomic molecules and elements that
appear more than once on the same side to
the end.
6.7 Types of Reactions
• There are five main categories of chemical
reactions:
1. Combustion
2. Synthesis
3. Decomposition
4. Single Displacement
5. Double Displacement
• Combustion –the rapid reaction of a
substance with O2 to produce compounds
called oxides (often call this process burning).
• The fuel can be a variety of things but it is
often a hydrocarbon (ex. gasoline) The
formula for combustion of a hydrocarbon is
C4 H10 + O2
CO2 + H2 O + energy
(C4H10 is Butane)
The products of a combustion reaction are always carbon
dioxide and water. (C4H10 is butane)
Skeleton...
C4 H10 + O2
CO2
+ H2 O + energy
8CO2
+ 10H2 O + energy
Balanced ...
2C4 H10 +13O2
Synthesis Reactions
• Involves the combination of smaller atoms or
compounds into larger compounds. (also
known as combination reactions).
• They have the following general formula:
•
A+B
AB
• If both reactants are elements then the
reaction MUST be synthesis.
•
Example: 2H2 +
O2
2H2O
Examples
• 2Na + Cl2 => 2NaCl
• 2Al + 3Br2 => 2AlBr3
• Synthesis reactions sometimes involve joining
two compounds into a larger one.
hydrogen chloride + ammonia
HCl
+
NH3
ammonium chloride
NH4Cl
Decomposition Reactions
• Involves the splitting of a large compound
into smaller molecules or elements.
• They have the following general formula:
•
AB
A + B
• If there is only 1 reactant then the reaction
MUST be decomposition.
• Example: 2H2O
2H2 + O2
What types of Reactions are these?
1. H2CO3
CO2+ H2O
2. 2Fe + O2
2FeO
3. C10H8 + 12 O2
10 CO2 + 4 H2O
Answers:
1. Decomposition
2. Synthesis
3. Combustion
Please do the following…
• Page 235…#’s 1,2,3,4
Single Displacement
• This is when one element trades places with
another element in a compound. These
reactions come in the general form of:
A + BC ---> AC + B
• Example: Fe + CuSO4 => FeSO4 + Cu
• The reactants MUST be an element and a
compound
• Single displacement can involve metals:
Na + KCl
K + NaCl
• Single displacement can involve nonmetals:
F2
+ 2LiCl
2 LiF + Cl2
• Remember - If the single element is a
nonmetal it will replace the nonmetal.
• If the single element is a metal it will replace
the metal.
Double Displacement
• Involves two elements replacing one another.
• The reactants must be compounds (usually
happens in solution).
• The positive ions stay in the same position (A
and C) and the negative ions change partners
(B and D). The general formula is:
AB + CD
AD + CB
• NaOH + FeCl3
NaOH + FeCl3
Fe(OH)3 +
NaCl
• Pb(NO3)2 + 2 KI
Pb(NO3)2 + 2 KI
PbI2 + 2 KNO3
List what type the following reactions
are:
•
•
•
•
•
•
1)
2)
3)
4)
5)
6)
NaOH + KNO3 --> NaNO3 + KOH
CH4 + 2 O2 --> CO2 + 2 H2O
2 Fe + 6 NaBr --> 2 FeBr3 + 6 Na
CaSO4 + Mg(OH)2 --> Ca(OH)2 + MgSO4
Pb + O2 --> PbO2
Na2CO3 --> Na2O + CO2
• 1) double displacement
2) combustion
3) single displacement
4) double displacement
5) synthesis
6) decomposition
Please do the following…
• Page 241…#’s 1,2,3
• Chemical reactions and balancing equations...
Rates of reaction
Objectives
• To understand that a chemical reaction
involves collisions between particles
• To be able to describe the four factors which
will affect the rate of a chemical reaction.
How do we make the reaction go faster?
• There are four things that we can change
to make the reaction go faster.
They are:
• Temperature
• Surface area
• Concentration
• Using a catalyst
Temperature
• When we increase the temperature
we give the particles energy
• This makes them move faster
• This means they collide with other
particles more often
• So the reaction goes faster.
Surface area
• If we make the pieces of
the reactants smaller
we increase the number
of particles on the
surface which can react.
• This makes the reaction
faster.
The particles
on the surface
can react
When cut into
smaller pieces
the particles on
the inside can
react
Concentration
• If we make one reactant
more concentrated (like making a
drink of orange squash more concentrated)
• There are more
particles in the same
volume to react
• So the reaction goes
faster.
There are less red
particles in the same
volume so there is
less chance of a
collision
There are more red
particles in the same
volume so there is
more chance of a
collision so the
reaction goes faster
Using a catalyst
• A catalyst is a chemical which is added to
a reaction.
• It makes the reaction go faster.
• The catalyst does not get used up in the
reaction.
• It gives the reaction the energy to get
started
Click here ...to complete exercise 2
Click here ...to complete exercise 3
Rates of reaction