Transcript Molecules, Ions and Their Compounds

Molecules, Ions and Their
Compounds
Chemistry 101
Chapter 3
Virginia State University
Dr. Victor Vilchiz
Summer 2008
Chemical Substances;
Formulas and Names
• Naming simple compounds
Chemical compounds are classified as
organic or inorganic.
Organic compounds are compounds
that contain carbon combined with other
elements, such as hydrogen, oxygen, and
nitrogen; they do not contain metals.
Inorganic compounds are compounds
composed of elements other than carbon
and usually contain at least one metal
atom.
Chemical Formulas; Molecular
Substances
• Organic compounds
An important class of molecular
substances that contain carbon is the
organic compounds.
Organic compounds make up the majority
of all known compounds.
The simplest organic compounds are
hydrocarbons, or compounds containing
only hydrogen and carbon.
Common examples include methane, CH4,
ethane, C2H6, and propane, C3H8.
Naming Covalent Compounds
A covalent compound as we said before is
formed by sharing electrons between 2
nonmetals or metalloids.
These compounds are usually molecular
and are named using a prefix system.
When naming these compounds name the
element further to the left (in the periodic
table) first, then the one on the right.
Naming Covalent Compounds
You name the first element using the exact
element name.
Name the second element by writing the root of
the element’s name and add the suffix “–ide.”
If there is more than one atom of any given
element, you add the Greek prefix denoting
how many atoms of that element are present.
Table lists the Greek prefixes used.
If only one atom of the second element is
present it gets the prefix “mono”
Naming Covalent Compounds
Here are some examples of prefix names for
binary molecular compounds.
PF5
phosphorus pentafluoride
SO2 sulfur dioxide
SF6
sulfur hexafluoride
N2O4 dinitrogen tetroxide
CO
carbon monoxide
Naming Acids
Acids are traditionally defined as compounds
that could donate an H+; however, they are
acids only in the presence of water. In other
words before they enter the liquid they are
covalent compounds and they are NOT acids.
There are two main types of acids:
Binary acids consist of a hydrogen ion and any
single anion. For example, HCl is hydrochloric
acid.
An oxoacid is an acid containing hydrogen,
oxygen, and another element. An example is a
HNO3, nitric acid. (see Figure 2.23)
Naming Acids
• Binary Acids
– Start with the prefix “Hydro” which represents the
Hydrogen, followed it with the root of the name of the
second element and append the ending –oic acid.
• Oxoacids
– Use the root of the “E” element if the ion taking part in
the acid had an ending in –ate to the root append the
ending –ic acid, if it ends on –ite then append the
ending –ous acid. If the ion had a prefix use the same
prefix.
Naming Acids
• Examples:
–
–
–
–
–
–
–
HCl(g) Hydrogen Chloride
HCl(aq) HydroChloric Acid
H2S(g) Dihydrogen Sulfide
H2S(aq) HydroSulfic Acid
H3PO4(aq) Phosphoric Acid
HClO4(aq) Perchloric Acid
HClO(aq) Hypochlorous Acid
Ionic Compounds Formulas
• How do we know how many atoms of each ion we need?
– A simple crossing of the charges can answer that question about
90% of the time.
• Example: Mg2+ and PO43-
Mg3(PO4)2
Check the charges… 3 x (+2) = +6
2 x (-3) = -6
– When they combined they cancel to yield a neutral
compound.
Ionic Compounds Formulas
• The crossing technique does not work if the
magnitude of the charges is the same
• Example: Mg2+ and CO32Mg2(CO3)2
This is incorrect since we want the lowest
ratio possible which is 1:1 to yield MgCO3
Ionic Compounds Properties
• Ionic compounds have properties completely
different from their component elements.
– Example: Table Salt (NaCl)
• Sodium (Na) in the presence of water reacts violently heating
up the water and producing hydrogen if the temperature of the
water is high enough the hydrogen can ignite explosively.
• Chloride (Cl) Green poisonous and corrosive gas. If inhaled
will destroy the nasal passages then dissolve in the stomach
producing high concentration of hydrochloric acid which will
destroy the stomach lining producing ulcers.
• Salt (NaCl) posses none of the properties mentioned above.
Ionic Structure
•Ions form a 3-D lattice.
•The coulombic (electrostatic)
attraction is so high that in
order to separate one ion from
the lattice requires a lot of
energy (DHlatt).
•The lattice energy depends
on charge and size of the ions.
Lattice Energy
• Since the lattice energy is an electrostatic interaction the
more separated the charges are the weaker the interaction
is.
– Bigger ions have lower lattice energies
• The higher the charge of the ions the stronger they will
attract ions of the opposite charge.
DH latt 
q1q2
r12
• When size and charge point to opposite trends the charge
will outweigh the size.
– From smallest atom to biggest atom there is only 1.7x factor. From
a +1 to +2 that is already a 2x factor.
Properties of Ionic Substances
• Dues to the charged interaction a blow to a
crystal leads to the possibility of splitting
the crystal since we will force like charged
particles to interact.
• Ionic compounds have high melting/boiling
points since in order to move the ions from
their respective spots it will require
breaking the lattice.
Ionic Solutions
• However, if we do melt an ionic compound
it will be able to conduct current.
• When ionic compounds are placed in a
solvent the produced solution conducts
electricity. The higher the number of ions
the higher the conductivity.
• More when we cover chapter 4.
Naming Hydrates
A hydrate is a compound that contains
water molecules weakly bound in its
crystals.
Hydrates are named from the anhydrous
(dry) compound, followed by the word
“hydrate” with a Greek prefix to indicate
the number of water molecules per formula
unit of the compound.
For example, CuSO4. 5H2O is known as
copper(II)sulfate pentahydrate. (see Figure
2.24)
Determining Chemical
Formulas
• Determining both empirical and
molecular formulas of a compound
from the percent composition.
The percent composition of a compound leads
directly to its empirical formula.
An empirical formula (or simplest formula)
for a compound is the formula of the substance
written with the smallest integer (whole
number) subscripts.
Determining Chemical
Formulas
• The percent composition of a compound is the
mass percentage of each element in the compound.
We define the mass percentage of “A” as the
parts of “A” per hundred parts of the total, by
mass. That is,
mass of " A" in whole
mass % " A" 
 100%
mass of the whole
Mass Percentages from
Formulas
• Let’s calculate the percent composition of
butane, C4H10.
First, we need the molecular mass of C4H10.
4 carbons @ 12.0 amu/atom  48.0 amu
10 hydrogens @ 1.00 amu/atom  10.0 amu
1 molecule of C4 H10  58.0 amu
Now, we can calculate the percents.
amu C
% C  5848.0
.0 amu total  100%  82.8%C
amu H
% H  5810.0
.0 amu total  100%  17.2%H
Determining Chemical
Formulas
• Determining the empirical formula
from the percent composition.
Benzoic acid is a white, crystalline powder used as a
food preservative. The compound contains 68.8% C,
5.0% H, and 26.2% O by mass. What is its empirical
formula?
In other words, give the smallest whole-number ratio
of the subscripts in the formula
Cx HyOz
Determining Chemical
Formulas
• Determining the empirical formula
from the percent composition.
For the purposes of this calculation and making
calculations simpler, we will assume we have 100.0
grams of sample benzoic acid.
Then the percentage of each element equals the mass
of each element in the sample.
Since x, y, and z in our formula represent mole-mole
ratios, we must first convert these masses to moles.
Determining Chemical
Formulas
Determining the empirical formula from
the percent composition.
Our 100.0 grams of benzoic acid would contain:
1 mol C
68.8 g C 
 5.73( 3) mol C
12.0 g
1 mol H
5.0 g C 
 5.0 mol H
1.0 g
1 mol O
26.2 g O 
 1.63( 7 )mol O
16.0 g
This isn’t quite a whole number ratio, but if we divide each number by
the smallest of the three, a better ratio might emerge.
Determining Chemical
Formulas
Determining the empirical formula from
the percent composition.
Our 100.0 grams of benzoic acid would contain:
5.73 mol C  1.63(7)  3.50
5.0 mol H  1.63(7)  3.0
1.63(7) mol O  1.63(7)  1.00
now it’s not too difficult to see that the smallest whole number ratio
is 7:6:2. The empirical formula is C7H6O2 .
Determining Chemical
Formulas
• Determining the “true” molecular
formula from the empirical formula.
An empirical formula gives only the smallest wholenumber ratio of atoms in a formula.
The “true” molecular formula could be a multiple
of the empirical formula (since both would have the
same percent composition).
To determine the “true” molecular formula, we must
know the “true” molecular weight of the
compound.
Determining Chemical
Formulas
• Determining the “true” molecular
formula from the empirical formula.
For example, suppose the empirical formula of a
compound is CH2O and its “true” molecular weight
is 60.0 g/mol.
The molar weight of the empirical formula (the
“empirical weight”) is only 30.0 g/mol.
This would imply that the “true” molecular formula
is actually the empirical formula doubled 2(CH2O) or
C2H4O2
Molecular and structural formulas
and molecular models.
Return to Lecture
A model of a portion of a Sodium Chloride crystal.
Return to Lecture
Common Ions of the transition metals
Return to Lecture
List of Polyatomic Ions
Return to Lecture
Greek Prefixes for Covalent
Compounds Nomenclature
Return to Lecture
Making and Acid
Return to Lecture
Molecular model of nitric acid.
Return to Lecture
Figure 2.24: Copper (II) sulfate.
Photo courtesy of James Scherer.
Return to Slide 44
Naming Flow Chart
Return to Lecture
Naming Flow Chart II
Return to Lecture
Naming Acids Flow Chart
Return to Lecture
Quantities of
Reactants and
Products
Chapter 4
Dr. Victor Vilchiz
What is a mole?
• A mole is a unit of measurement used to
specified amounts of chemical substances.
– It is not a unit of mass.
– It is similar to “a dozen”
• A dozen eggs is not the same as a dozen cars but
they are still both a dozen.
What is a mole?
• A mole is defined as the number of
atoms of carbon in 12 g of Carbon-12.
(examples)
• 1mole=6.022x1023 atoms and can be
applied to any moiety
– 6.022x1023 is also known as Avogadro’s
Number (NA)
• 1mol of Carbon=12g Carbon = 6.022x1023 C atoms
• 1mol of water= 18g H2O =6.022x1023 water molecules
Why the mole?
• The mole helps determine amounts of
substances and allows for conversion
between species.
– CaCO3(s) + 2HCl(aq)  CaCl2(aq) + H2CO3(aq)
• From this we cannot say 1g of CaCO3 will react with 2 grams of HCl;
however, we can say 1mol of CaCO3 reacts with 2 moles of HCl.
– 1 mole of CaCO3 is not the same as 1 mole of HCl mass wise,
but both have 6.022x1023 molecules.
Molar Mass Examples
• Molar Mass of Ca(C2H3O2)2, Calcium
Acetate.
2x(40.1)+4(12.0)+6(1.01)+4(16.0)=198.3g/mol
• Molar Mass of Ethylene Glycol, C2H4O2.
2x(12.0)+4(1.01)+2(16.0)=60.0g/mol
• Molar Mass of Ammonium Oxalate,
(NH4)2C2O4.
2x(14.0)+8x(1.01)+2(12.0)+4(16.0)=124.1g/mol
Stoichiometry: Quantitative
Relations in Chemical
Reactions
• Stoichiometry is the calculation
of the quantities of reactants and
products involved in a chemical
reaction.
It is based on the balanced chemical equation
and on the relationship between mass and
moles.
Such calculations are fundamental to most
quantitative work in chemistry.
Chemical Reactions:
Equations
• Writing chemical equations
A chemical equation is the symbolic
representation of a chemical reaction in
terms of chemical formulas.
For example, the burning of sodium and
chlorine to produce sodium chloride is
written
2Na  Cl 2  2NaCl
The reactants are starting substances in a
chemical reaction. The arrow means
“yields.” The formulas on the right side of
the arrow represent the products.
Chemical Reactions:
Equations
• Writing chemical equations
In many cases, it is useful to indicate the
states of the substances in the equation.
When you use these labels, the previous
equation becomes
2Na(s )  Cl 2 (g )  2NaCl(s )
s=solid, l=liquid, g=gas, aq=aqueous
Molar Interpretation of a
Chemical Equation
• A balanced chemical equation:
2H2 +1O2  2H2O
can be interpreted to read 2 moles of
Hydrogen react with one mole of oxygen to
produce 2 moles of water.
– In the balanced equation the 2, 1, and 2 are
known as the stoichiometric coefficients.
• At the molecular level they refer to the number of molecules
reacting.
Molar Interpretation of a
Chemical Equation
Because moles can be converted to mass,
you can also give a mass interpretation of a
chemical equation.
2H2 +1O2  2H2O
2(2.02g)H2 react with 1(32.0g) O2 to yield
2(18.0g)H2O
4.04g H2 react with 32.0g O2 to yield 36.0g H2O