Bonding and Chemical Formulas
Unit IVA -
.2 Chemistry
October 2006- Revised Nov. 2007
Bonding
· There are two major types of bonding:
· Ionic and covalent:
· Covalent bonds occur when two atoms SHARE
valence electrons
· Ionic bonds are due to a very strong
electrostatic attraction between two ions (atoms
who just “EXCHANGED” electrons)
· In other words, one atom stole electrons from
the other
Ions
· Let’s start by deciding what an ion is
· An ion is an atom which has either
gained or lost an electron
· When an atom gains an electron, the ion
is negatively charged
· Why?
· When an atom loses an electron, the ion
is positively charged
· Why?
Ions
· A positively charged ion is referred to as
a cation
· A negatively charged ion is called an
anion
· Metals tend to be cations because they
will frequently lose electrons
· Nonmetals tend to be anions because
they will take on more electrons
Electronegativity
· The difference in electronegativities
between two atoms, determines whether
an ionic or covalent bond will form.
· So, what is electronegativity?
Electronegativity is defined as :
an atom’s pull on its electrons
As a rule of thumb, electronegativity
increases as you go
and to the
Electronegativity
· Look at your table of electronegativities
· Metals tend to have very LOW electronegativity
values and NonMetals tend to have quite HIGH
values
· The difference between any two atoms will decide
whether an ionic or a covalent bond will form
· If the diff is >2, then an ionic bond will form
· If the diff is <1.69, then a covalent bond will form
· (between 1.7-2, if M-NM then ionic and if 2 NM
then covalent)
Ionic vs Covalent
· What type of bond will form between each of the
following?
H and Cl?
· C and O?
· Na and Br?
· Be and F?
· What trend do you see?
· Typically nonmetals form ionic bonds with metals
and they form covalent bonds with other nonmetals
Valence Electrons
· As you recall, the electrons in the outermost
shell or energy level are called valence
electrons
· These valence electrons are typically the only
ones involved with bonding
· Let’s use a technique called electron dot
structures to represent the valence electrons
of an atom (also called Lewis Dot structures)
Valence Electrons
· Each dot represents an electron
· First, determine how many valence electrons an
atom has
· Where will you find this information?
· Then write the chemical symbol for that atom
· Place the correct number of dots around each of
the 4 sides of the atomic symbol
· Be sure to only put one dot per side until all
sides are “1/2 filled”, then you can start doubling
up
· Why?
Electron Dot Structures
•Write the electron dot structure for each of the following:
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Na
C
O
Cl
Kr
Mg
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Li
B
P
S
Ne
Ca
•Notice the patterns that will occur within families
Electron Configurations for
Ions
· We already know how to write the electron
configuration for an atom, let’s apply it to
ions
· We do it the exact same way except we take
into account the electrons that have been
added or the missing electrons
· The most stable form of the ion will be that
which shares the electron configuration with
a noble gas
Electron Configurations for
Ions
· Noble gases all have the same ending to their
electron configuration, s2 p6, giving them 8
electrons in their outermost energy level
· This is where the octet rule comes from
· We know that noble gases are the most stable
elements, scientists gathered that the reason for
this is because they have 8 electrons in their
valence shell
Electron Configurations for
Ions
· What is the noble gas configuration for the
following ions?
· Ca+2
· O 2· Li+1
· Al+3
· H-1
Electron Configurations for
Ions
· What is the noble gas configuration for the
following ions?
· Ca+2
[Ar]
· O 2[Ne]
· Li+1
[He]
· Al+3
[Ne]
· H-1
[He]
Ionic Bonding
· Ionic bonding involves the TRANSFER
of electrons or stealing of electrons as we
mentioned before
· Do you recall how we determine which
element is going to become a cation and
which will become an anion?
Ionic Bonding
· Each of you will be given a card with an
elemental symbol on it
· For that symbol, determine the number of
valence electrons (this is how many dots
you would put around the symbol in your
Lewis dot structure)
· Then, come up and grab $1 bill for each
valence electron you have
Ionic Bonding
· Using the list of compounds, find your partner
and EXCHANGE electrons accordingly
· Once you have exchanged electrons, you are
transformed into your respective ions
(cation/anion)
· You are electrostatically attracted to each other
so… BOND (stand tightly shoulder to shoulder
until I come around and check your compound)
Properties of Ionic
Compounds
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A solid ionic compound is called a salt
All salts share 5 characteristics:
1. Made of crystals
2. Conduct electricity
3. Have high melting and boiling points
4. Are hard
5. Are brittle
1. Made of Crystals
· Attraction between opposite charged ions
is SO great that there is more than one
bond
· A tightly packed cluster of repeating
units forms the crystal structure
· Ex. NaCl is the formula
· unit for table salt
· crystals
2. Conducts Electricity
· Electricity needs charged particles that
are free to move (in solution)
· In salt water, the particles spread out and
can carry electricity from ion to ion
throughout the solution
· Electrolytes - ions in solution which
carry an electric current
3. High BP and MP
· Because of the strong attractions between the
oppositely charged ions, it takes a lot of
energy to “break up” the particles
· BP - boiling point is the temperature at which
you have a phase change from a liquid to a
gas
· MP - melting point is the temperature at
which you have a phase change from a solid
to a liquid
4. Hard and 5. Brittle
· Salts are hard due to the strong attraction of
opposite charges and the layering of crystals
· Salts are brittle, or break up to make a powder
· Layers usually line up so that - and + alternate, but
added energy or pressure can cause + to be next to
+ and - to be next to -.
· Does it like it…NO so it breaks apart into powder
Hydrates
· Some salts can hold water molecules
between their bonds
· These are called Hydrated Salts
· Possible Uses: drying agents or moisture
indicators
· Ex. CuSO4 . 5 H2O
· In this hydrate, for every salt unit of Copper
sulfate, 5 molecules of water are trapped
Percent Composition of Hydrates
· What % of CuSO4 . 5 H2O is water?
· First let’s find the formula mass of copper
sulfate by itself
· 63.5 + 32.1 + 64.0 = 159.6 g/mol
· What about the water that is trapped?
· 1.0 (10) + 16.0 (5) = 90.0 g/mol
· Total molar mass: 249.6 g/mol
· % mass of water = 90.0 / 249.6 x 100 = 36%
Try another hydrate problem:
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Calculate the percent water in NiCl2. 6H2O ?
NiCl2 = 129.7 g/mol
6 H2O = 108.0 g/mol
Total mass: 237.7 g/mol
108.0 / 237.7 x 100 = 45%
Percent Composition
· The law of definite proportions refers to the
chemical make-up of ONE compound
· Within that compound, the proportion or
ratio of one element to another will remain
the same no matter how much of the
compound is present
· The law of multiple proportions compares
the compositions of two different
compounds which contain the same
elements
Percent Composition
· The Law of definite proportions explains
why we can have a “formula unit” for a
compound
· This is the simplified version of the
elemental ratios within the compound
· For example, when joining Mg and Cl
what is the smallest whole number ratio that
can be used to join these two together
· Make sure that the charges balance out
Percent Composition - (refresher)
· If I have an ionic compound of MgCl2,
what percentage of the whole mass does Mg
make up?
· First, find the atomic mass of Mg
· Then, find the atomic mass of Cl and double
it because there are two atoms of Cl
· What is the total mass?
· Divide the mass of the Mg by the total mass
and multiply by 100. This is the percent of
the whole that Mg makes up.
Percent Composition
· Calculate the percent composition of
water in the following hydrate:
· (NH4)2SO4 . 5 H2O
Mole Ratios and Hydrate Predictions
· For our hydrated salt lab, you took the mass
of the hydrated salt before heating
· You then heated it up for 10-15 minutes as
instructed and took the mass of the
anhydrous salt
· Anhydrous salt – salt without the water
· To determine the amount of water that was
in your original hydrated salt sample,
subtract the mass of anhydrous salt from the
mass of hydrated salt
· Now you can do two things:
· 1. determine the percent of the salt
sample that was water
· Do this by taking the mass of water
and dividing by the total mass of the
hydrated salt (before heating)
· Multiply by 100 and this is your
experimental % water.
· 2. The second thing you can calculate is the # of
waters trapped in the salt per formula mass unit.
· Do this by first converting your number of grams
of anhydrous salt to moles (use the molar mass of
the salt)
· Then convert your mass of water to moles (using
the molar mass of water)
· Now divide both numbers (of moles) by
whichever is smaller to get a ratio.
· Clue: the ratio will by 1 to ___ (rounded to the
nearest whole number).
· If the ratio is very close to a HALF number, then
double both. Ex. 1 : 2.5 gets doubled to 2:5 ratio
Hydrate Prediction Example
· In the lab, you measure out 5.25g of
barium chloride (BaCl2 )
· After heating, the mass is 4.50g.
· Calculate the amound of water that was in
the hydrated salt.
· 5.25g – 4.50g =
· 0.75 g
· Calculate the % of water in the sample:
· 0.75g / 5.25 g = 14.3% water
· Now, let’s determine the # of water molecules
trapped in the hydrated salt per formula unit.
· Convert the grams of water to moles.
· 0.75g / 18.01 g/mol = 0.0416 moles H2O
· Convert the grams of anhydrous salt to moles.
· 4.50g / 208.24g/mol = 0.0216 moles BaCl2
· Divide each by the lesser of the two
· 0.0416/0.0216 = 1.93 (round to 2)
· 0.0216/0.0216 = 1
· Write your whole number ratio.
· 1 BaCl2 : 2 H2O
· Finally, write your formula for the hydrated salt
· BaCl2 . 2 H2O
· To determine how close you were (or to calculate %
error): calculate your theoretical percent water in the
above formula as usual
· 2(18.01) / ( 208.24 + 2(18.01) )
· Mass of water / total mass =
· 14.75%
· Compare with your original 14.30% experimental
percent water: 14.75 – 14.30 = 0.45
· % error is this difference divided by the theoretical:
· 0.45 / 14.75 =6.6% error (not too bad for a first
year chemistry lab student)
Polyatomic Ions
· As you saw when we discussed and calculated
oxidation numbers, you sometimes will see a
“cluster” of atoms that have a combined overall
charge
· These are called poly (many) atomic (atoms)
ions (with a charge)
· Ex. MnO4 -1 the permanganate ion
· I have provided you with a list of polyatomic
ions that you’re responsible for memorizing
Naming Ionic Compounds
1. When naming ionic compounds, begin with
the cation.
This can either be a metal from the periodic
table OR it could be a polyatomic cation
· What are some cations?
· Family I, II, or III, transition metals, or
ammonium
Ionic Compounds
· 1.When writing the name, start with the cation
· 2.Follow it by the anion (for now we’ll start with
polyatomic ions that have a known charge)
· Choose a polyatomic anion that would balance with
Cu+2
· Let’s use sulfate for this example
· Write the formula for this compound
· Now write the name of it
· CuSO4
Using Roman Numerals
· When the cation can vary in its oxidation
number/ charge, we must use a Roman
numeral to indicate its charge
· Ex. Iron can vary in its ox #
· Write the formula for iron bonded to sulfate
IF the iron has a +3 charge
· Fe2(SO4)3
Using Roman Numerals
Fe2(SO4)3
When we name this compound, we must use a
Roman numeral to indicate iron’s charge so
the name of this compound is
Iron (III) sulfate because the iron has a +3
charge
Do NOT confuse this with the subscript,
which indicates only how many atoms of
iron we have
Ionic Compounds
· Do we have to use the Roman numeral for
family I and II elements?
· No, they have a set charge that is understood
· Practice writing the names of the following:
· ZnSO4
CaCO3
Fe2(SO4)3
·
Zinc sulfate
Calcium Carbonate Iron (III) Sulfate
Other Anions
· Polyatomic ions do not all end in ide
· but other ions do
· When naming, you still place the name
of the cation first
· Followed by the anion (ending in –ide)
· What would you call MgO ?
· Magnesium oxide
Ionic Compounds (cont.)
· Name each of the following:
· NaCl
KI
· Sodium chloride
Potassium iodide
FrBr
Francium bromide
· What happens when you have multiple anions
with a cation? For example: MgCl2 ?
· This is still called Magnesium chloride
because in order to form this ionic compound,
there HAS to be 2 chlorines