Transcript Ionic Bonding

Chemistry
• Is the study of matter,
its properties and its
changes or
transformations
Matter
• Anything that has
mass and takes up
space
Physical Properties
• a change in the size or form of a substance which does
not change the chemical properties of the substance is
called a physical change
• a physical property is a characteristic of a substance
• these properties allow us to distinguish or tell the
difference between substances
• examples of physical properties - state, color, odor,
luster, texture, hardness, crystal form, mass, volume,
density, solubility, viscosity, malleability, ductility, melting
point, boiling point
Chemical Properties
• is a characteristic behavior that occurs when a
substance changes into a new substance\
• the change itself is called a chemical change
• the starting materials we call reactants - the
final materials we call products
• examples of chemical properties are: reactions
with water, reactions with acids, combustion
• example: iron + oxygen = rust
Introduction to Naming
• Today, most compounds are known
by their IUPAC names. IUPAC
stands for International Union of
Pure and Applied Chemistry. This
organization has determined a set of
rules to be used for naming
chemicals. Its purpose is to set
international guidelines so that all
scientists follow the same rules.
Classifying Compounds
•
There are three main types of
compounds that we will be dealing with:
1. Ionic – a combination of metals and non-metals
2. Molecular or covalent – a combinations of two
nonmetals
3. Inter-metallic or metallic – a combination of two
metals
Compound
Properties
Examples
Molecular
•Solid, liquid or gas at STP
•Relatively low melting and boiling
points
•Do not conduct electricity in aqueous
solutions (non electrolyte)
•May be soluble or insoluble in water
Sugar
Water
Propane
Metallic
•Ductile, malleable
•Good conductor of heat and electricity
(electrolyte)
•Shiny when freshly cut or polished
Brass
Steel
Ionic
•Crystalline solid at STP
•High melting and boiling point
•Usually soluble in water
•Conducts electricity in aqueous
solution
Copper
sulfate
Sodium
chloride
(cont’d)
• Molecules are combinations of two or
more atoms.
• Molecular element - if the atoms are all
the same
• For example oxygen gas is a molecule
composed of two atoms of oxygen. Since
there are two atoms the molecule is called
a diatomic molecule. (just remember
the gen's)
Diatomic molecular elements
As the heading suggest these are elements
composed of two (di) nonmetal atoms. We seen
these in the last lesson. (just remember the gen's)
oxygen
hydrogen
O2
H2
nitrogen
fluorine
the
chlorine
halogens
(group
bromine
17)
iodine
N2
F2
Cl2
Br2
I2
Mono-atomic elements
If we look at the name of the heading, mono - means
one, so these are the non-metals that that exist in
nature as individual atoms. Although these are not
compounds we have included them here because we
will reference them many times.
Noble Gases (group 18)
He
Helium
Ne
Neon
Ar
Argon
Kr
Krypton
Xe
Xenon
Rn
Radon
Polyatomic molecular elements
These are non-metal elements composed of many
(poly) atoms.
O3
S8
ozone
sulfur
P4
phosphorus
(red)
P10
Phosphorus
(white)
• Compound - a molecule that contains two or more
different types of atoms or ions.
• Water (H2O) is a compound because it contains both
Hydrogen and Oxygen, two different types of atoms.
• The formula for water (H2O) is a combination of symbols
and subscripts.
• H and O are the symbols for the two types of elements
(Hydrogen and Oxygen) found in water.
• The number 2 to the lower right hand corner of the
symbol for Hydrogen is called a subscript.
• a molecular formula indicates the total number of atoms in one
molecule
• an empirical formula is the simplest whole number ratio of atoms in
the compound.
• consider hydrogen peroxide (H2O2 ) as an example.
• - the molecular formula is H2O2.
• - the empirical formula is HO. (lowest ratio is 1:1)
• In some cases the molecular formula and empirical formula are the
same.
• For example, both the molecular formula and the empirical formula
for water is H2O. (the lowest whole number ratio is 2:1)
• It is important to recognize however that the empirical formula only
describes the ratio of one atom to another, and not the actual
number of atoms of each type in the compound.
Solutions
• Aqueous solutions are those in which the
solvent is water. They form in at least
three ways depending on the nature of the
solute:
• molecular solvation
• dissociation
• ionization
Electrolytes and Non-electrolytes
• A solution that conducts electrical current is said
to be electrolytic and the solute is called an
electrolyte.
• The solute in a solution that does not conduct
electrical current is a non-electrolyte.
• Generally, dissociated ionic compounds are
electrolytes whereas dissolved molecular
compounds are non-electrolytes. The exceptions
to this rule are the molecular acids.
• Acids are defined theoretically as species that
ionize in water to produce hydrogen ions and
negative ions (anions).
Bohr’s Electron Energy Level
Theory
• Bohr was the first to attempt to describe
the way electrons are distributed in an
atom
• The type of bonding that occurs in a
substance is a function of the way the
electrons are distributed in an atom
Electron Energy Level Theory
• Electrons orbit the nucleus in specific
pathways called energy levels
• There is a fixed number of energy levels
• Each energy level is capable of holding a
certain number of electrons
• The location of an electron (with respect to
the nucleus) is an indication of the energy
it contains (closer more – further away
less)
• Electrons can move from one energy level
to the next by gaining or losing a specific
amount of energy called a quantum
• Energy levels are not equidistant from
each other
• The part of the energy level or region of
space in which we find electrons are also
called orbitals. These are not the fixed
pathways we normally think of when
discussing orbits, but rather they are a
region or area when the electron would be
located at some point in time.
Electron Energy Level Diagrams
•
These will demonstrate how electrons are
distributed in an atom
1. Determine the atomic number of the element
2. Draw a circle to represent the nucleus and
write the # of protons in the circle
3. Determine the period number of the element
(ie what row it is in). This is equal to the
number energy levels in the element
4. Each energy level can hold a specific number
of electrons. We always fill the inner level first.
5. For the first three levels it is 2 -8 – 8. We put
two hash marks - - for each level and put the
number of electrons in the level between the
marks.
6. The outermost or last energy level is called the
valence level.
7. The electrons in the valence level are called
valence electrons.
Example
•
Look up the atomic number of nitrogen and use it to
determine nitrogen's number of protons and electrons.
Since nitrogen is atomic number 7, it has 7 protons and 7
electrons.
•
Draw a circle to represent the nucleus and write in the
number of protons using the shorthand 7p+.
•
Locate the period number for nitrogen and use it to
determine the number of electron energy levels. The
period number (2) equals the number of energy levels.
Draw two broken lines above the circle to represent
nitrogen's two electron energy levels.
•Bohr theory, the first three energy levels hold a maximum of
two, eight, and eight electrons respectively. The energy levels
are filled from the inside out, so write "2e-" on the first line to
show the number of electrons in the first energy level.
•The outermost energy level of an atom is called the valence
level. In a nitrogen atom, the second energy level is the
valence level. Since a nitrogen atom has seven electrons, and
since two of them are in the first energy level, then the second
energy level holds five electrons. Write 5e- in the second
energy level. Note that the number of valence electrons in an
atom can be determined by looking at the last digit of its group
number. Nitrogen is in Group 15.
Important Point!!!
• Valence electrons are
the only ones thought
to be involved in
chemical reactions!
Your Turn
• Draw energy level diagrams for the
following elements:
• Lithium
• Calcium
• Neon
• Hydrogen
-- 2 –
-- 8 –
-- 1 --
-- 8 –
-- 8 –
-- 2 --
-- 2 --
-- 2 --
-- 1 --
3+
20+
10+
1+
Bohr Diagrams
Order of Filling Energy Levels
• Lower energy levels are filled before
higher energy level orbitals
• The electrons will fill the available orbitals
before pairing up.
Your Turn
Construct modified Bohr diagrams for the
first 20 elements
Why do certain atoms gain or lose
electrons?
• The loss or gain of electrons enables
atoms to achieve an octet of electrons (ie.
gain electrons) or expose a lower energy
octet of electrons (ie. lose electrons)
Nobel Gases and the Octet Rule
• Remember that the outer shell or energy level is called
the valence level and it contains the valence electrons
• It is commonly believed that the valence energy level
and the valence electrons are responsible for chemical
bonding
• A group or eight is called an octet
• The octet rule states that chemical stability is associated
with a group of eight valence electrons
• The noble gasses contain a complete outer shell of
electrons or they have 8 valence electrons
• The noble gasses are believed to be chemically stable
because their valence level have a complement of eight
electrons
Simple Ions
• Simple ions are atoms that no longer have equal numbers of
electrons and protons.
• As a result of this imbalance they now have a positive or negative
charge
• Cation
• A cation is a positve (+) ion
• Metal atoms lose electrons to become positive thus, metals form
cations.
• Example: Na+, Mg 2+
• Anion
• An anion is A Negative ION (A N ion)
• Non-metal atoms gain electrons to become more negative thus,
non-metals form anions:
• Example: F- , O 2-
Energy level diagrams for IONS
Atoms with 1, 2, or 3 valence electrons
lose them to form 1+, 2+ or 3+ ions
respectively.
•naming metallic ions - the full name of
the atom is followed by the word ion.
Mg2+ is the magnesium ion
Group 1 (1+) (lose 1e)
Group 2 (2+) (lose 2e)
Group 13 (3+) (lose 3e)
Atoms with 5, 6, or 7 valence electrons
gain electrons to form 3-, 2-, or 1- ions,
respectively.
Group 15 (3-) (gain 3 e)
•naming non-metallic ions - the name of Group 16 (2-) (gain 2 e)
the atom is shortened and the suffix Group 17 (1-) (gain 1 e)
ide is added.
O2- is oxide
Atoms with valency of 4 generally do
not form ions.
•These atoms do not gain or lose
electrons. They become stable by
sharing electrons - (recall molecular
compounds - covalent bonding).
Group 14
(do not form ions)
Your Turn
Construct modified Bohr diagrams for the
ions formed by the first 20 elements and
name each ion formed
Complete the simple ions sheet
Your Turn
• Complete sheet u3 s2 l3 Simple Ions
Chemical Bonding
• A chemical bond will form between two atoms if
the attractive forces between two atoms is
greater than the repulsive forces
• Attractive forces are electrical forces that hold
the atoms, ions, or molecules together
• Bonds are formed through the valence electrons
in the atom
• Valence electrons are usually transferred from
the outer shell of one atom to the outer shell of
another atom or shared among the outer shell of
combining atoms
Ionic vs Molecular Compounds
• Ionic Compounds
– Involve the transfer of electrons resulting in ionic
bonding
– Made up of two oppositely charged ions (metal and
nonmetal)
– Exist in the form or an ionic crystal lattice
• Binary Molecular Compounds
– Involve the sharing of electrons resulting in covalent
bonding
– Composed of two nonmetals
– Exist as individual molecules
Molecular Compounds
• Always form between two or more nonmetals
• In order to get full outer energy levels the
atoms that make up molecular compounds
share their valence electrons (bonding electrons)
• If they dissolve in water, their solutions do
not conduct electricity
• Molecular compounds are non electrolytes
Molecular Compounds (cont’d)
• Individual parts of a molecular compound
are called molecules.
• The molecular formula tell us how many
atoms of which elements are being
shared.
• Example: H2O tells us that in every
molecule of water there are two hydrogen
atoms and one oxygen atom
Naming Molecular Compounds
(cont’d)
• When we want to come up with
(derive/make) a formula for a molecular
compound there are two routes to follow
as well:
• 1. It has a trivial name and its formula
must have been memorized, like water
• 2. We use the prefixes given to tell us how
many of each element, like silicon dioxide
Naming Molecular Compounds
• Two basic Ways
1. the compound has a trivial (common)
name that has been used for so long, it’s
too late to change it. We must memorize
trivial names and their formulas.
Trivial names, formulas, and
Special Situations
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The following names and formulas must be memorized
Water H20
Ammonia NH3
Glucose C6H12O6
Sucrose C12H22O11
Methane CH4
Propane C3H8
Octane C8H18
Methanol CH3OH
Ethanol C2H5OH
Hydrogen Peroxide H2O2
O3 ozone
Naming Molecular Compounds
(cont’d)
2. We use a system of prefixes to
signify how many of each type of atom is
being shared to make up the molecules
1 = mono
6 = hexa
2 = di
7 = hepta
3 = tri
8 = octa
4 = tetra
9 = nona
5 = penta
10 = deca
Naming Molecular Compounds
(cont’d)
• First element in the compound is the
regular element name
• The last element in the molecule must
have its name changed to an “ide” ending
• The only time we are allowed to leave off a
prefix is if there is only one atom of the
first element in the molecule
Naming Molecular Compounds
(cont’d)
• Example:
• C02 - carbon dioxide
• CCl4 – carbon tetrachloride
• P2O5 – diphosphorus pentaoxide
Your Turn
• Complete sheet u3 s2 l2 Molecular
Compounds
Ionic Bonding
• An ionic bond is a force of attraction
between a metallic ion and a non-metallic
ion
• It is an electron transfer between two
elements or polyatomic ions
• This type of bond created a crystal with a
definite, repeated pattern
• Consider that chemical reactions between
metals and non-metals occur because metals
lose their loosely held electrons to highly
electronegative non-metals.
• This results in ions of the opposite charge
• These oppositely charged ions them mutually
attract each other to form ionic compounds.
• This attraction between oppositely charged ions
is called an ionic bond.
Chlorine atom
Sodium atom
Writing formulas for ionic
compounds
• Write the symbols and charges for the ions that
make up your ionic compound
• Remember
The positive ion (cation) is always written first, the
negative ion (anion) is last
Almost always your ionic compound will start with a
metal. (The only time it doesn’t is when it starts with
the word ammonia
Metal ions has the same name as the metal elements
Non-mental ions always end in “ide”
Writing Formulas for Ionic
Compounds (Cont’d)
• Short cut: criss-cross the charge and put
in lowest terms
Put the two on the charge of calcium as the
number of chlorines you need, and the one on
the charge of chlorine as the number of
calcium's you’ll need.
Writing IUPAC names from
formulas
• Since every ionic compound is made up of
some cation in combination with some
anion all you do is name the cation
(positive ion) and then name the anion
(the one with the negative charge).
• Don’t use prefixes
NaCl – sodium chloride
Li3N – lithium nitride
Writing formulas for ionic
compounds (cont’d)
• Determine the lowest whole number ration of
ions to give an overall net charge of zero. In
other words, cancel each other’s charges.
Example: calcium chloride
Calcium is the cation, its symbol and charge is
Ca2+
Chloride is the anion, its symbol and charge is
ClIt would take two chlorides to cancel the charge
of one calcium, so the formula is CaCl2
Your Turn
• Complete sheet u3 s2 l4 Molecular and
Binary Ionic
Writing Formulas for Ionic
Compounds (Cont’d)
–.
– Multi-valent metal ions is the stock system for
naming. The stock system says that roman
number in brackets after the metal name is
the number of the positive charge.
– Example: Iron (III) oxide is made up of Fe3+
ions in combination with oxide ions
Writing IUPAC Names From
Formulas (Cont’d)
Some ions require special treatment:
Multi – valent ions – some metals, the ones in B groups – they can
form more that one positive ion
Example
Fe can be Fe2+ or Fe3+
We have to say which ion was used to make the compound by using
the Roman numeral for the number on the charge in brackets after
the name of the element.
Cu2S is copper (II) sulfide
We know this because it took two copper ions to balance the charge on
the sulfide ion which is always S2• PbO2 is lead (IV) oxide. Lead can be a 4+ ion or a 2+ ion. In this
case it must be 4+ because it took two O2- to cancel out the charges
Example
• Compound: Iron(III) oxide
• Ions present: Fe3+
O2-
• Molecular formula:
Fe2O3
• Note: # of oxygen's is same as charge on
iron and # of irons is same as the charge
on oxygen
Writing ionic formulas from names
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First write down the name of the metal
Then write down the name of the non-metal
Next – determine the charge on the non-metal
To find the overall negative charge multiple the
charge by the number of atoms present
• To determine the charge on the metal – divide
the total negative charge by the number of
atoms present in the molecule
Example
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4.
5.
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7.
RuO2
Name of metal → ruthenium
Name of non – metal → oxygen
Charge of oxide ion → 2Number of oxygens present → 2
Total negative charge → 2 x 2 = 4
Number of metal atoms present → 1
Charge of metal → 4 divided by the number of
metal atoms (1) = 4
8. Name of compound → ruthenium (IV) oxide
Example 2
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1.
2.
3.
4.
5.
6.
7.
Nb2O5
Name of metal → niobium
Name of non – metal → oxygen
Charge of oxide ion → 2Number of oxygens present → 5
Total negative charge → 2 x 5 = 10
Number of metal atoms present → 2
Charge of metal → 10 divided by the number
of metal atoms (2) = 5
8. Name of compound → niobium (V) oxide
Your Turn
Complete sheet u3 s2 l4
Nomenclature Involving Multivalent Ions
Polyatomic Ions
• Are composed of several elements (more
than one capital letter) bonded together
with an overall charge.
• The charge is almost always negative
• You should recognize these special ions
Writing IUPAC Names From
Formulas (Cont’d)
• We have to recognize the polyatomic ions.
• Once we have identified the cation, and the rest
of the formula still has more than one capital
letter in it , there must be a polyatomic ion in it.
The polyatomic ions all have special names.
• Example
• Na2CO3 is called sodium carbonate. Na is
positive part (cation) so that means that the CO3
part has a special name called carbonate
Writing IUPAC Names From
Formulas (Cont’d)
• Al(OH)3 is aluminum hydroxide –
hydroxide us the name given to the
polyatomic ion made up of O and H joined
together as OH-
Example 1
• Compound: Sodium carbonate
• Ions present: Na+
CO32-
• Molecular formula:
Na2C03
• Note: # of sodium's is same as charge on
carbonate ion and # of carbonate ions is
same as the charge on sodium
Example 2
• Compound: Ammonium sulfide
• Ions present: NH4+
S2-
• Molecular formula:
(NH4)2S
• Note: # of ammoniums is same as charge
on sulfate and # of sulfates is same as the
charge on ammonium
Your Turn
Complete Sheet u3 s2 l5
Polyatomic Ions (two sheets)
Special Ionic Compounds
• Some ionic compounds have a number of
water molecules hanging around with
them.
• These special compounds are called
hydrates.
• A hydrate or formula has two parts – the
first part is the ionic compound, the
second part is the water. We separate the
two parts with a giant asterisk or dot.
Special Ionic Compounds (Cont’d)
• CuSO4 * 5H2O is named in two steps
1. First name the ionic compound in the
regular way (Copper (II) Sulfate)
2. Next use your molecular prefixes along with
the word hydrate to tell how many water
molecules there are. (pentahydrate)
Deriving Formulas
1. First come up with the formula for the ionic
compound the usual way by balancing the
charges on the ion.
2. Check the prefix to find out how many
water molecules are with the ionic compound
Deriving Formulas
Example:
magnesium sulfate heptahydrate
• Mg2+ with SO42- is
MgSO4
• heptahydrate means 7 water
• Full formula: MgSO4 * 7H2O
Your Turn
Complete sheet u3 s2 l5 Hydrates
Nomenclature of Acids
• Acids are hydrogen compounds, that
release hydrogen ions (H+) when they are
dissolved in water. The formula of an acid
is always followed by (aq), which means –
”in aqueous solution” – in other words
dissolved in water
Properties of Acids:
• There are several key properties that distinguish acids
from other classes of compounds.
• acids turn blue litmus paper red
• acids react with metals like zinc, magnesium, and iron to
produce hydrogen gas. (Hydrogen gas is flammable.)
• acids neutralize bases
• acids have low pH (pH < 7)
• acids that are edible, such as those found in foods, taste
sour (e.g. vinegar, and citric acid found in oranges,
lemons, and other citrus fruits, sour candy.)
Naming Acids
• Naming Acids.
1. Look at the formula and name the acid as a hydrogen
compound. It will be named hydrogen ___________.
The blank represents some ion.
2. The names of hydrogen compounds can only have
three possible ending:
hydrogen ______ ide.
hydrogen ______ ate.
Hydrogen ______ ite.
Naming Acids (cont’d)
• The acid name depends on which ion
ending you have in the formula:
– Hydrogen __ ide becomes hydro ___ ic acid
– Hydrogen __ ate becomes ___ ic acid
– Hydrogen __ ite becomes ___ ous acid
Naming Acids (cont’d)
• HF(aq) would be called hydrogen flouride
so it would be called hydrofluoric acid
• HClO4(aq) would be hydrogen perchlorate
so it would be called perchloric acid
• NNO2(aq) would be called hydrogen nitrite,
so it would be called nitrous acid
Writing Formulas for Acids
1. Look at the name and work backwards to
change it to a hydrogen compound.
Hydro __ ic acid becomes hydrogen ___ ide.
______ ic acid becomes hydrogen ____ ate.
______ous acid becomes hydrogen __ ite.
2. Make your formula by balancing the charges.
The hydrogen in acids is always 1+ and the
charge on the other ions can be determined
from the periodic table or the polyatomic chart.
Writing Formulas for Acids
(Cont’d)
• Example
Hydrochloric acid must be hydrogen chloride
Hydrogen is 1+ and chlorine is 1- so the formula is
HCl(aq)
Nitric acid must be hydrogen nitrate. If hydrogen is 1+
and nitrate is 1- so nitric acid is HNO3 (aq)
Chlorous acid must be hydrogen chlorite. If hydrogen is
1+ and chlorite is 1- then the formula is HClO2 (aq)
Special Situations and
Exceptions to the Rules
• Any acid names that contain the root word “sulf”
or “phosph”, add an extra syllable to make them
sound better:
H2SO4 (aq) should be called sulfic acid according
to our rules but we call it sulfuric acid,l with an
extra “ur” syllable
H3PO4 (aq) should be called phosphic acid but we
call it phosphoric acid with an extra “or” in it
Special Situations (cont’d)
• The other exception is when you put
together a formula with certain polyatomic
ions that end in ____COO. For these
acids the hydrogen gets put at the end of
the formula rather than at the beginning. It
is still balanced with the negative charge,
but is put at the end of the formula.
Special Situations (cont’d)
• Example:
Acetic acid – from hydrogen acetate
Hydrogen is 1+
Acetate ion is 1To balance the charges you need one of each,
but the hydrogen has to go at the end
CH3COOH (aq)
Your Turn
Complete sheet u3 s2 l7 Acids (2 pages)
Complets sheet u3 s2 l2 – l4 (2 pages)
Bases
Bases are substances that behave in
opposition to acids. In this course, we will
restrict our discussion of bases to one
particular type - ionic compounds that
contain the hydroxide ion (OH-).
• example: Sodium hydroxide - NaOH(aq)
Properties of Bases
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turn red litmus blue
neutralize acids
have high pH (> 7)
form slippery solutions
tend to have a bitter taste
The pH scale is a measure of the concentration
of hydrogen ions (H+) in aqueous solution, and
can be used to compare the strength of acid or
base. The pH scale ranges from pH = 1 (strong
acid), to pH = 14 (strong base). The value pH =
7 is neutral (neither acid nor base)
Neutralization Reactions and the
Formation of Salts
• Salts are formed as a result of the
reaction between an acid and a base.
• The reaction between an acid and a base
results in the formation of salt and water.
The type of salt formed would be
determined by the type of acid and base
that react in the neutralization reaction.
• ACID + BASE → SALT + WATER
Chemical Reactions
• A chemical reaction is the result of
chemical species colliding.
• If the collision is successful bonds will be
broken and new bonds formed.
• This will result in a new chemical species
being formed.
• Energy is either required (endothermic)
or released (exothermic) during these
reactions.
Evidence of a Chemical
Reaction
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Energy change
Color change
Gas evolved
Precipitate formed (a solid)
Odor change (never smell directly)
Every reaction must obey the law of
conservation of mass – matter cannot be
created nor destroyed only transformed
5 Basic Reaction Types
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Formation
Decomposition
Single replacement
Double replacement
Combustion
Formation
• A formation reaction involves creating a compound from
the elements it contains
• General formula:
Element + element  compound
A + B  AB
Ex.
Li (s) + O2(g)  LiO2 (g)
K (s) + S8  K2S (s)
F2 (g) + Al (s)  AlF3 (s)
Formation (cont’d)
• To correctly predict the product write the metal and nonmetal ions together and balance the charges (ionic
compounds)
• NO subscripts from the reactant side are brought to the
product side
Ex.
Na (s) + P4 (s) 
Na3P(s)
Ba (s) + O2 (g) 
BaO(s)
Cl2 (g) + K (s) 
KCl(s)
Formation (cont’d)
• Sometimes the metal that reacts form a multivalent ion in the ionic compound.
• To determine which ion should be used, use the
ions that is on “top” on the periodic table
Ex.
V (s) + O2 (g)  V2O5 (s)
Mn (s) + F2 (g)
 MnF2 (s)
Decomposition
• Involves particles of the same compound colliding
and forming the elements found in the compound
Compound → element 1 + element 2 …….
ABC → A + B + C
Example
C6H12O6 → C (s) + H2 (g) + O2 (g)
VBr5 (s) → V (s) + Br2 (l)
Decomposition (cont’d)
To correctly predict results:
Write each element that you identify in the reactant
compound as an element.
Do not bring any subscripts with the element.
Write each element as it would normally exist.
It does not matter what order you write the elements.
Decomposition (cont’d)
• U2O 4 (s) → U (s) + O2 (g)
• AlP (s) → Al (s) + P4 (s)
Decomposition (cont’d)
• If the compound contains polyatomic ions, follow the same
procedure.
• Al2(SO4)3 (s) → Al (s) + S8 (s) + O2 (g)
• Examples
C4H10 (g) →
(NH4)2Cr2O7 (s)
→
Single Replacement
• Is a reaction that is a result of an element colliding
with a compound
• The element will “switch” or replace the
appropriate element in the compound
• Two possibilities
a) A + BC → B + AC (A is a metal)
b) A + BC → C + BA ( A is a non-metal)
Single Replacement (cont’d)
• Example
• Al (s) + BaCl2 (s) → Ba (s) + AlCl3 (s)
• P4 (s) + BaCl2 (s) → Cl2 (g) + Ba3P2 (s)
Single Replacement (cont’d)
• The ionic bonds in the compound were broken
• The element then bonds with the oppositely charged ion
from the ionic compound
• The other ion forms an element
• The new ionic compound must be correctly written (criss –
cross method)
• No subscripts are moved across the equation
• In does not matter the order in which the chemical
species are written
Single Replacement (cont’d)
•
•
•
•
K (s) + AlP (s) → (metal with metal)
NaF (s) + Al (s) →
(metal with metal)
O2 (g) + LiBr (s) →
(nonmetal with nonmetal)
AgI (s) + S8 (s) →
(nonmetal with nonmetal)
Single Replacement (cont’d)
Example:
• If the element is a metal you must check to see if
it is multi-valent and use the most abundant
charge (the top one)
U (s) + LiF (s) → Li (s) + UF6 (s)
Ba (s) + MnCl4 (s) → Mn (s) + BaCl2 (s)
• If opposite, the metal to be replaced in the
compound is multi-valent, it does not matter
Single Replacement (cont’d)
If the compound contains a polyatomic ion, the
polyatomic ion will not change since it remains
part of the compound.
Example
Ca (s) + NaOH (s) → Na (s) + Ca(OH)2 (s)
Al (s) + CuSO4 (s) → Cu (s) + Al2(SO4)3 (s)
S8 (s) + NH4F (s) → F2 (g) + (NH4)2S (s)
Double Replacement
• Is the result of two compounds colliding
• Some/all bonds are broken and new bonds are
formed
General Form
compound1 + compound 2 → compound 3 + compound 4
AB + CD → AD + CB
note: A and C were metals
B and D were nonmetals
It does not matter which order the reactants or
products are written in
Double Replacement (cont’d)
• This type of reaction is very similar to a
single replacement except there is a
double switch.
• Simply replace the metal in one compound
with the metal in the second compound
• Since each product is an ionic compound,
balance the charges to find the correct
products
Double Replacement (cont’d)
Examples:
BaF2 (aq) + AlP → Ba3P2 (s) + AlF3 (aq)
Na2O (s) + Mg(NO3)2 (aq) → NaNO3 (aq) + MgO (s)
NH4I (aq) + Al2(SO4)3 (aq) → (NH4)2SO4 (aq) + AlI3 (s)
Double Replacement (cont’d)
UF6 (s) + KMnO4 (aq) →
U(MnO4)6 (aq) + KF (aq)
LiCN (aq) + BaBr2 (aq) →
LiBr (aq) + Ba(CN)2 (aq)
Ca(NO3)2 (aq) + FeSO4 (aq) →
CaSO4 (s) + Fe(NO3)2(aq)
Double Replacement (cont’d)
• A special case of double replacement reactions
involves the addition of a base with an acid
• This is called a neutralization reaction
• A “salt” and water are produced
Example
HCl(aq) + LiOH(aq) → LiCl(aq) + H2O (l)
Al(OH)3 (aq) + H2SO4 (aq) → Al2(SO4)3 (aq) + H2O (l)
Combustion
• Involves the collision of oxygen with a
hydrocarbon
• Two types:
A) complete and B) incomplete
Complete Combustion
• The reaction has the same products no matter what
the reactants
• Products always – carbon dioxide and water
Example:
CH4 (g) + O2 (g) → CO2 (g) + H20 (g)
C4H10 (g) + O2 (g) → CO2 (g) + H20 (g)
C6H12O2 (g) + O2 (g) → CO2 (g) + H20 (g)
Incomplete Combustion
• Same as complete combustion, but not all of the
reactant is burnt
• This results in two additional products carbon
and carbon monoxide
Example:
CH4 (g) + O2 (g) → CO2 (g) + H20 (g ) + C (s) + CO (g)
C4H10 (g) + O2 (g) → CO2 (g) + H20 (g) + C (s) + CO(g)
C6H12O2 (g) + O2 (g) → CO2 (g) + H20 (g) + C (s) + CO (g)
note: you will have to be told if the reaction is incomplete to know the
products
Combustion (cont’d)
• Examples:
CH4 (g) + O2 (g) →
C20H42 (g) + O2 (g) →)
C12H22O11 (g) + O2 (g) →
Balancing Chemical Equations
• The law of conservation of mass states that matter is
conserved during a chemical reaction
• Therefore we are required to “balance” chemical
equation, we need to have the same amount of
products and reactants
• To balance chemical equations we use molar coefficients to ensure that we have the same number
and type of atom on each side of the equation
• Most molar co-efficients are simply numbers written in
front of each reagent and product that, when
multiplied, result in equal numbers of atoms on both
sides of the equation
Balancing Chemical Equations
1. Start with the largest subscript present
(not the ones found within the polyatomic
ions)
2. Balance the polyatomic ions as a group
3. Balance in a logical progression
4. Leave the elements until last
5. If it seems hard to balance, you may
have made a mistake in the formulas of
your chemical species
Balancing Chemical Equations
Example
F2 (g) + Al (s) → AlF3 (s)
1.
2.
3.
4.
5.
6.
7.
8.
9.
The largest subscript is 3
Compare the number of F’s in the products vs the reactants (3
vs 2)
These numbers are not divisible, therefore, find the LCM
LCM = 6
Determine the number of times that AlF3 goes into the LCM
Place the number 2 in front of AlF3
Now compare the number of F’s on both sides of the
equation. F2 would require the number 3 to be placed in front
of it to make the number of F’s equal on both sides (6)
Next balance the number of Al’ s. Since there are two on the
product side, multiple the one on the reactant side by 2.
The balanced equation would be
3 F2 (g) +2 Al (s) → 2 AlF3 (s)
Complete the following equations
•
•
•
•
•
•
•
Na + Cl2 →
K + O2 →
H 2 + O2 →
H2 + Cl2 →
N 2 + H2 →
N2H4 + O2 →
CH4 + O2 →
Complete the following equations
•
•
•
•
•
•
•
•
•
Ca + HBr →
Al + O2 →
KNO3 + HBr →
Ba + H3PO4 →
CaCl2 + AL2(SO4)3 →
C3H8 +O2 →
Mg + HNO3 →
AgNO3 + NaCl →
K3PO4 + CoCl3 →