Download The Atomic Theory, and the Structure of Matter

Unit 1: Chemistry (P. 134-279)

Patterns and Compounds
– Periodic Table, Naming, Balancing Equations

Chemical Reactions
– Energy, 4 Types, Combustion

Acids and Bases
– Properties, pH, Reactions

Chemical Reactions in the Environment
– Factors affecting Rates, Chemicals and Us
Classification of Matter
MATTER
PURE SUBSTANCE
ELEMENTS
SUSPENSIONS
COMPOUND
MIXTURE
SOLUTION
EMULSIONS
MECHANICAL
MIXTURE
COLLODIAL
DISPERSION
Classification of Matter

PURE SUBSTANCE:A substance with a fixed composition
and constant properties
– ELEMENT: A substance that cannot be broken down into simpler substances
by chemical means. Atoms are the simplest particles that cannot be broken
down by chemical means
– COMPOUND: A substance that is made up of two or more different atoms
(molecules). These substances can be broken down only by chemical means.

MIXTURE: A mixture consists of two or more kinds of
matter, each keeping its own characteristic properties.
– SOLUTIONS: A mixture that is homogeneous. If the solution is a liquid or
gas, it is transparent.
– MECHANICAL MIXTURE: A heterogeneous mixture with parts that are
visibly distinguishable
Metals Metalloids Nonmetals
Metals vs. Non Metals
•Shiny
•Dull
•Ductile
•Brittle
•Malleable
•Does Not Conduct Electricity
•Conducts Electricity
•Does Not Conduct Heat
•Conducts Heat
Elements and the Periodic Table:
A Review

Every Element has a unique
– Name
– Symbol
– Atomic Mass Number (A)

Represents the number of protons + number of neutrons
– Atomic Number (Z)

Represents the number of protons and the number of electrons in a
neutral atom
– The number of neutrons can be calculated by subtracting
A-Z
Atomic Number Atomic Mass
Atomic Number
Atomic Mass
Number of electrons,
protons
2
4
He
Number of Protons
and Neutrons
Number of Neutrons = Atomic Mass – Atomic Number
Example
19
K
39.098
Name
Potassium
Atomic Mass
40
Atomic Number
19
Electrons &
Protons
Neutrons
19
20
Grouping of Elements

Elements are subdivided into:
– “groups” or "families" (vertical columns)
– and “periods” (horizontal rows)





Metals elements are on the left
Non-metal elements on the right
separated by a dark "staircase line".
Elements bordering this division line exhibit some
properties of both metals and non-metals and are
called metalloids.
Copy table 5.1 on page 140 into your notes.
Alkali
Metals
Group Names
Alkaline
Earth
Metals
Group 3
Carbon Group
Nitrogen
Group
Oxygen Group
Halogens
Noble Gases
Bohr-Rutherford Diagrams

The following information is required:
– 1. Number of Electrons:

Is the same as the number of protons in a neutral atom. The
electrons are organized into shells in the following order.
–
–
–
–
up to 2 electrons in the first shell
up to 8 electrons in the second shell
up to 18 electrons in the third shell
up to 32 electrons in the fourth shell
– 2. Number of Protons

Is the same number as the atomic number
– 3. Number of Neutrons

Can be determined by subtracting the atomic mass from the
atomic number
Bohr Rutherford Diagram
Nucleus with
protons and
Neutrons
18p
36n
Electron Orbits
(shells) with a
2,8,8, pattern
Predicting Chemical Reactivity




Elements with 8 electrons in their outer energy level
appear to have a special significance. Elements with this
arrangement do not react easily and are considered stable.
All noble gases (Neon, Krypton, Xenon, Radon) have 8
electrons in their outer energy level and are very nonreactive elements (Helium is a special gas that is very
stable with 2 electrons in its first level).
All elements want to be stable and therefore want to gain
or lose electrons in order to achieve a stable 8
configuration (Stable Octet).
Whenever an atom gains or loses electrons they become
negative or positive and they are called ions.
Two main factors determine chemical
activity(reactivity)

1) The number of electrons in the outer energy
level
– i) Elements with 1-3 electrons on outer level lose
electrons (become positive)
– ii) Elements with 5-7 electrons in outer level gain
electrons (become negative)
– iii) Elements with 4 electrons in outer level are special
(tend to become positive)

2) The number of energy levels
– As the number of energy levels increase, the attraction
between those electrons in the outermost energy level
and the positive nucleus decrease
Ions: To gain or lose an
Electron

Positively Charged: Cations
– When a neutral atom gives up one or more electrons,
the positively charged ion that results is called a
Cation.
– For example:

Negatively Charged: Anions
– When a neutral atom gains one or more electrons, the
negatively charged ion that results is called an Anion.
– For example:
Electron Dot Diagrams

A Bohr-Rutherford diagram represents an atom
and all its electrons.
 A simpler way to represent atoms and ions of
atom is with electron dot diagrams
 Electron Dot Diagrams show only the outer
energy level (valence shell) of an atom. Only these
electrons are represented because they are
responsible for an atom’s chemical properties. For
example:
Lewis Dot / Electron Dot
diagrams
C
N
Chemical Bonds: Forming
Compounds




Most substances on earth do not exist as elements, they are composed
of two or more different elements joined together to make compounds.
When two atoms collide, valence electrons on each atom interact. A
chemical bond forms between them if the new arrangement of their
valence electrons have less energy than their previous arrangement.
For many atoms that new arrangement of their electrons will be that of
their closest noble gas.
Atoms may acquire a valence shell like that of its closest noble gas in
one of three ways:
– 1. An atom may give up electrons and forma ion
– 2. An atom may gain electrons and form an ion
– 3. An atom may share electrons
Ionic Compounds
Substances held together by ionic bonds are called
Ionic compounds e.g. NaCl, KCl. Ionic Bonds
occur because of the attraction of cations and anion
for each other. Electrons are transferred between the
atoms during bond formation.
Properties include:
•
High melting point (i.e. strong bonds)
• Conduct electricity when dissolved in water or molten
• Form crystal lattice structures
• Soluble in water
Molecular Compounds



Substances that are composed of molecules are called
molecular compounds. Many non-metals form compounds
with other non-metals. When this occurs there is no
transfer of electrons between the two atoms instead they
share electrons forming a covalent bond.
Although bond between atoms are strong, bonds between
molecules are weak. eg. Moth crystals, nitrogen gas etc.
Properties Include:






Low melting and boiling points
Often have an odour
Don’t conduct heat
Don’t conduct electricity (non-electrolytes)
Diatomic molecules (e.g. O2, F2 etc.) are also the
result of covalent bonds.
Chemical Naming and Formulas

Binary Ionic
 Transition Metals
– Stock versus Classical

Polyatomic Ions
 Binary Molecular
General Rules

The Metal is always written first
 The nonmetal suffix in a compound is either
“ide” or “ate”
 Every compound must be electrically
neutral
– All Positive charges must equal Negative
charges
Binary Compounds: Formula to Name

Composed of two Elements
 One metal and one nonmetal
 Write the name of the metal first unchanged
 Write the name of the nonmetal second
 Change the ending to an “ide”
– LiCl  Lithium Chloride
– MgI2  Magnesium iodide
Binary Compounds: Name to Formula

Write the symbol for each element with the metal
written first
 Find the ionic charge for each element
 Cross the number value of the charge and place it
as the subscript of the other element
 Reduce the values to lowest ratio

Magnesium Oxide  Mg2+
O2Mg2O2
MgO
Transition Metals: Groups 3-12+
Name to formula

Almost all are able to form more than one
cation
 When writing the formula the charge of the
metal cation will be indicated by roman
numerals after the metal
 Lead (III) chloride  PbCl3
 Iron (II) oxide  FeO
Transition Metals: Groups 3-12+
Formula to Name

Finding the charge on the metal can be done two
ways
Reverse Cross-Over Method
1.
–
–
The subscript of the nonmetal becomes the charge of
the metal
Sometimes the charge is misleading
Charge Balancing
2.
–
Charge = Subcript of the nonmetal multiplied by the
charge of the nonmetal divided by the subscript of the
metal
Chemical Equations and Reactions
A chemical equation is a description of a chemical
reaction using chemical symbols, not words
Steps:
1) The reactants are written first
2) The products are written second
3) The state for each atom is indicated
(g) gas, (s) solid, (l) liquid, (aq) aqueous
4) The reactants and products are separated by an
"arrow" (  )
e.g. Word Equation
Hydrogen gas plus chlorine gas produces
hydrogen chlorine gas
e.g. Chemical Equation
H (g) + Cl (g)  HCl(g)
Balanced and Unbalanced Chemical
Equations
The Law of Conservation of Mass states:
Matter cannot be created or destroyed; it can
only be changed from one form to another.
Therefore, the number of atoms in the reactants
must equal the number of atoms in the products
An unbalanced or skeleton equation does not
follow the Law of Conservation of Mass. The
number of atoms on the left side (reactants)
does not equal the atoms on the right side
(products)
e.g.
H2(g) + Cl2(g) 
4 atoms (2 H, 2 Cl)
HCl(g)
2 atoms(1 H, 1 Cl)
A balanced chemical equation follows the Law
of Conservation of Mass. The number of atoms
on the left side (reactants) equals the atoms on
the right side (products)
e.g.
1H2(g) + 1Cl2(g) 
4 atoms (2 H, 2 Cl)
2HCl(g)
4 atoms(2 H, 2 Cl)
Writing Balanced Chemical Equations
1. Write the chemical formula for each reactant and product
followed by the state of each: solid (s); liquid (l); gas (g);
aqueous(aq)
2. Adjust the numbers of molecules until there are the same
number of atoms of each type on both sides of the equation.
This balances the mass of both the reactants and products.
3. Usually, balancing is easiest when hydrogen and oxygen atoms
are left until the end
NOTE:
Do not change the subscript in a formula to balance an
equation. Changing these numbers changes the
molecular structure of the molecule.
Energy Changes and Chemical
Reactions
Chemical reactions, physical changes of state and
dissolving processes often involve energy changes.
Exothermic Processes:
Processes that release energy (e.g. heat and
light) and increase the temperature of the
surroundings.
Endothermic Processes:
Processes that absorb energy and decrease the
temperature of the surroundings.
Factors Affecting Chemical Reaction
Rate
The Rate of Reaction is defined as:
The time it takes for a given product to form, or for
given amounts of reactant to react.
Reaction rate is determined by:
i. Measuring how fast reactants are used up.
ii. Measuring how fast the products are formed.
Factors affecting Reaction Rate
1. Concentration and Reaction Rate
 Concentration (amount of substance in a given volume)  Rate
2. Surface Area and Reaction Rate
 Surface Area (area exposed)  Rate
4. Catalysts and Reaction Rates
A Catalyst is defined as:
A substance that speeds up the rate of a chemical
reaction without being used up in the reaction.
Catalyst lower the energy required to break the
bonds that hold substances together. Examples
include: enzymes (biological catalysts),
platinum, rhodium and palladium (chemical
catalyst used in catalytic converters)
Types of Chemical Reactions
There are four basic patterns that most chemical
reactions follow:
1) Synthesis Reactions
This type of reaction fits the general pattern:
A + B  AB
e.g.
N2(g) + 3H2(g)  2NH3(g)
CaO(s) + H2O(l)  Ca(OH)2
A synthesis reaction involves the formation of a
new compound from simpler elements or
compounds
Combustion reactions (involving the reaction
with O2) are examples of Synthesis Reactions
2) Decomposition Reactions
These type of reactions are opposite to direct
combinations. They fit the general pattern:
AB  A + B
e.g.
CuCO3(s)  CuO(s) + CO2(g)
2KClO(s)
 2KCl(s) + 3O2(g)
A decomposition reaction involves the breaking
down of a compound into simpler compounds or
elements
3) Single Displacement Reactions
A single displacement or substitution reaction
fits the general pattern of:
A + BC  AC + B
This type of reaction involves a change in
partners. One element displaces or knocks off
another element in a compound..
e.g. Zn(s) + 2HCl(aq)  ZnCl2(aq) + H2(g)
3C(s) + Fe2O3(s)  3CO(g) + 2Fe(s)
4) Double Displacement Reactions
A double displacement reaction fits the following
general pattern:
AB + CD  AD + CB
This type of reaction involves a change of both
partners. The cation (positive element or polyatomic
ion) of one compound changes place with the cation
of the second compound.
e.g.
Na2S(aq) + ZnCl2 (aq)  ZnS(s) + 2NaCl(aq)
AgNO3(aq) + KBr(aq)  AgBr(s) + KNO3(aq)
SF4(s) + 2H2O(l) 
SO2(g) + 4HF(aq)
Carbon Chemistry
Organic Chemistry: The study of carbon
containing compounds and their properties e.g.
hydrocarbons
When hydrocarbons (contain carbon and hydrogen)
are burned in enough oxygen complete combustion
occurs.
Hydrocarbon + oxygen gas  carbon dioxide + water + E
(good supply)
If hydrocarbons are burned in a poor supply of
oxygen, incomplete combustion occurs.
Hydrocarbon + oxygen gas  carbon dioxide + water + E
(poor supply)
+ carbon monoxide + residue
Classification of Substances by Their
Behaviour
The process of grouping substances
according to common properties is called
classification.
Previously we have classified substances
according to:
i) State (e.g. solid, liquid or gas)
ii) Composition (e.g. pure substances, mixtures etc.)
Matter can also be classified by
chemical behaviour.
Acids and Bases
Acids:
An acid is a compound that dissolves in water to
produce hydrogen ions (H +) in solution. e.g. HCl
Bases:
A base is a compound that dissolves in water to
produce hydroxide ions in solution (OH -) e.g. NaOH
Copy Table 7.3 “Acids and Bases: A Summary”
found on page 230 in your text.
Preparation of Common Acids
A common way to prepare an acid is to react a nonmetal oxide
with water. An oxide is an element combined with only
oxygen e.g.
sulphur trioxide + water  sulphuric acid
carbon dioxide + water  carbonic acid
Some common acids in the laboratory include:
i)
ii)
iii)
iv)
sulfuric acid ( H2S04 )
nitric acid (HNO3)
hydrochloric acid (HCl)
acetic acid, (CH3COOH)
Other common acids include:
Preparation of Common Bases
A common way to prepare a base is to react a metal oxide
with water. e.g.
sodium oxide + water  sodium hydroxide
calcium oxide + water  calcium hydroxide
Some common bases in the laboratory include:
i) Sodium hydroxide (NaOH)
ii) Calcium hydroxide (Ca(OH)2)
iii) Potassium hydroxide (KOH)
iv) Magnesium hydroxide (Mg(OH)2)
Indicators
An indicator is a chemical that changes colour as
the concentration of H+ (aq) and OH- (aq)
changes. e.g.
i) Litmus:
• blue litmus turns red in acid
• red litmus turns blue in base
ii) Phenolphthalein
• turns pink in base
Indicators can be made from flowers, fruits,
vegetables, leaves (e.g red cabbage, tea etc.)
Synthetic Indicators are more easy to use than
natural indicators because they:
• last longer than natural indicators
• can be produced in large quantities
The pH Scale
The pH scale describes the "strength of the
hydrogen ion (H+)".
The scale is numbered from 0 to 14
• acids have a pH less than 7
[H+] > [OH-]
• bases have a pH more than 7
[H+] < [OH-]
• neutral substances have a pH of 7 [H+]= [OH-]
The change in 1 pH unit represents a tenfold
increase in the concentration of hydrogen ions in
solution. e.g.
A pH of 2 is 10 x's stronger than a pH of 3
A pH of 2 is
stronger than a pH of 5
A pH of 2 is _
stronger than a pH of 7
The Strength Of Acids And Bases
The strength of an acid or base is dependant on two
factors:
1. Concentration
The concentration of an acid or base is the amount
of the pure substance dissolved in 1 L of water.
2. Ionization
When acids and bases are dissolved in water, they
ionize (break apart into charged particles). The term
“Percent Ionization” refers to the number of
molecules that will ionize for every 100 molecules
that dissolve. e.g. HCl + H2O  H3O+ + ClSolutions that form ions in water are called
electrolytes. Electrolytes conduct electricity.
The Strength of Acids
Strong acids: ionize completely in water e.g H2SO4
Weak acids: ionize partially in water e.g. CH3COOH
The Strength of Bases
Strong Bases: ionize completely in water e.g NaOH
Weak Bases: ionize partially in water e.g. NH3
Neutralization
Neutralization occurs when hydroxide ions
(base) and hydrogen ions (acid) are mixed to
make water. The general word equation is:
Acid + Base 
Water
+
Salt
e.g
hydrochloric
acid
(HCl) (aq)
+
sodium
hydroxide

+ ( NaOH)(aq) 
water
+ sodium chloride
( H2O)(l) +
( NaCl) (aq)
After neutralization, the solution no longer has
a high concentration of either ion.
Soaps and Detergents
What makes up soap ?
1. fatty acid (lipid)
2. strong base (NaOH)
The word equation is:
fat + base  soap + glycerol
Soap curds cling as scum to whatever it comes
into contact with, and does not rinse away easily.
This problem led to the development of synthetic
detergents called syndets. Advantages include:
1.
2.
3.
4.
5.
good at removing dirt
more soluble in water
prevented dirt from collecting back onto clothes
did not form a curd
mild to hands and fine fabrics
How soap cleans
A soap or detergent molecule consists of
two ends:
1. Hydrophillic (water loving)
The end with the sodium ion is attracted to water and
becomes soluble
2. Hydrophobic (water hating)
Hydrocarbon end is attracted to insoluble dirt (grease)
on clothes etc.
For example: