Transcript Modern Chemistry

Chemical Formula & Naming
Chapter 7
CCl4
MgCl2
• Guess the name of each of the above compounds
based on the formulas written.
• What kind of information can you discern from
the formulas?
• Guess which of the compounds represented is
molecular(covalent) and which is ionic.
 Chemical formulas form the basis of the language
of chemistry and reveal much information about
the substances they represent.
• A chemical formula indicates the relative
number of atoms of each kind in a chemical
compound.
example: octane — C8H18
Carbons = 8
Hydrogens = 18
• example: aluminum sulfate — Al2(SO4)3
• Parentheses surround the polyatomic ion
to identify it as a unit. The subscript 3 refers to
the unit.
• Note also that there is no subscript for sulfur:
when there is no subscript next to an atom,
the subscript is understood to be 1.
Monatomic Ions
• Many main-group elements can lose or gain
electrons to form ions.
• Ions formed from a single atom are known as
monatomic ions.
– example: To gain a noble-gas electron configuration,
nitrogen gains three electrons to form N3– ions.
• Some main-group elements tend to form
covalent bonds instead of forming ions.
– examples: carbon and silicon
Naming Monatomic Ions
• Monatomic cations are identified by the element’s name.
– examples:
• K+ is called the potassium cation
• Mg2+ is called the magnesium cation
• For monatomic anions, the ending of the element’s name is
dropped, and the ending -ide is added to the root name.
– examples:
• F– is called the fluoride anion
• N3– is called the nitride anion
Binary compounds
• Compounds composed of two elements are
binary compounds.
• In a binary ionic compound:
positive charges = negative charges
• The formula for a binary ionic compound can
be written cation 1st , anion 2nd .
– example: magnesium bromide
Ions combined: Mg2+, Br–, Br–
Chemical formula: MgBr2
Example
Aluminum Chloride
+3
Al Cl
-1
AlCl3
Example
Sodium Sulfide
+1
Na S
-2
Na2S
Example
Magnesium Oxide
+2
Mg O
-2
MgO
Sample Problem
Write the formulas for the binary ionic
compounds formed between the following
elements:
a. iodine and zinc
b. zinc and sulfur
Polyatomic Ions
• Many common polyatomic ions are
oxyanions—polyatomic ions that contain
oxygen.
• Examples:
– PO4, SO4, NO2
Polyatomic Ions (cont.)
• Some elements can combine with oxygen to
form more than one type of oxyanion.
• example: nitrogen
NO3NO2nitrate
nitrite
• The name of the ion with the greater number
of oxygen atoms ends in -ate. The smaller
number of oxygen atoms ends in -ite.
Example
Magnesium Sulfate
+2
-2
-ate and -ite
endings are
polyatomics
Mg SO4 MgSO4
Example
Aluminum Nitrite
+3
-1
Al NO2 Al(NO2)3
Example
Calcium Hydroxide
+2
-1
Ca OH
Ca(OH)2
Example
Iron (III) Oxide
+3
-2
Fe O
Fe2O3
Sample Problem
Write the formula for tin(IV) sulfate.
Example
Ba3N2
barium nitride
Example
Na3PO4
sodium phosphate
Example
CuO
Which
copper is it?
-2 X 1 atom = -2
Copper (II)
( ) Oxide
Copper (I) – Cu+1
Copper (II) – Cu+2
Example
FeCl3
Which
iron is it?
-1 X 3 atom = -3
Iron (III) Chloride
Iron (II) – Fe+2
Iron (III) – Fe+3
Example
Cu3PO4
Which
copper is it?
-3 X 1 atom = -3
+3/3 Cu atoms = +1
Copper (I
( ) Phosphate
Copper (I) – Cu+1
Copper (II) – Cu+2
Naming Binary Molecular Compounds
• Molecular compounds are composed of
individual covalently bonded units.
• Naming molecular compounds is based on the
use of prefixes.
– examples: CCl4 — carbon tetrachloride
CO — carbon monoxide
CO2 — carbon dioxide
Prefixes
1 – mono
2 – di
3 – tri
4 – tetra
5 – penta
6 – hexa
7 – hepta
8 – octa
9 – nona
10 - deca
Naming (cont.)
Rule – Use prefixes to tell how many of each
atom; 1st atom name from periodic table, 2nd
element gets –ide ending
Exception – If only 1 of first atom, NO mono-
Prefixes
Example
Dichlorine heptabromide
Cl2Br7
Example
Sulfur trioxide
SO3
Note: no “mono” when only
one of first element
Sample Problem
a. Give the name for As2O5.
b. Write the formula for oxygen difluoride.
Example
F3P4
trifluorine tetraphosphide
Example
CO2
carbon dioxide
Example
CO
carbon monoxide
Warm-UP
•
•
•
•
•
Potassium carbonate
Dinitrogen pentachloride
CBr4
AuP
Mercury(I) selenide
Acids
• Can be either:
– Binary acids are acids that consist of two
elements, usually hydrogen and a halogen.
– Oxyacids are acids that contain hydrogen, oxygen,
and a third element (usually a nonmetal).
Naming Acids – Create the rules
• Examples of binary acids:
– Hydrochloric acid
– Hydrofluoric acid
HCl
HF
• Examples of oxyacids:
– phosphoric acid
– nitric acid
– sulfuric acid
H3PO4
HNO3
H2SO4
Acid Chart
Ion ending
Acid name is…
____-ide
hydro-___-ic acid
____-ate
_____-ic acid
____-ite
_____-ous acid
Example
HBr
bromide = hydrobromic acid
Ion ending
____-ide
____-ate
____-ite
Acid name is…
hydro-___-ic acid
_____-ic acid
_____-ous acid
Example
H3N
nitride = hydronitric acid
Ion ending
____-ide
____-ate
____-ite
Acid name is…
hydro-___-ic acid
_____-ic acid
_____-ous acid
Example
HNO3
nitrate = nitric acid
Ion ending
____-ide
____-ate
____-ite
Acid name is…
hydro-___-ic acid
_____-ic acid
_____-ous acid
Example
Hydrosulfuric acid
+1
H
-2
S
Ion ending
____-ide
____-ate
____-ite
= H 2S
Acid name is…
hydro-___-ic acid
_____-ic acid
_____-ous acid
Example
H2SO3
sulfite = sulfurous acid
Ion ending
____-ide
____-ate
____-ite
Acid name is…
hydro-___-ic acid
_____-ic acid
_____-ous acid
All Mixed Up
1. Determine compound type
A. Ionic – starts with metal or NH4
B. Covalent – starts with nonmetal
C. Acid – starts with hydrogen
2. Use proper rules to write formula/name
Section 7-3
• The chemical formula for water is H2O.
• How many atoms of hydrogen and oxygen are
there in one water molecule?
• How might you calculate the mass of a water
molecule, given the atomic masses of
hydrogen and oxygen?
• Chemical formulas allow chemists to calculate
a number of characteristic values for a
compound:
1. formula mass
2. molar mass
3. percentage composition
• The formula mass of any molecule, formula
unit, or ion is the sum of the average atomic
masses of all atoms represented in its formula.
– example: formula mass of water, H2O
average atomic mass of H: 1.01 amu
average atomic mass of O: 16.00 amu
average mass of H2O molecule: 18.02 amu
Sample Problem:
Find the formula mass of potassium chlorate,
KClO3
formula mass of KClO3 = 122.55 amu
• A compound’s molar mass is numerically
equal to its formula mass.
• Ex.) the molar mass of pure calcium, Ca, is 40.08
g/mol because one mole of calcium atoms has a
mass of 40.08 g.
• Ex.) molar mass of H2O molecule: 18.02 g/mol
Sample Problem G
What is the molar mass of barium nitrate,
Ba(NO3)2?
molar mass of Ba(NO3)2 = 261.35 g/mol
Molar Mass as a conversion factor
Sample Problem
What is the mass in grams of 2.50 mol of oxygen
gas (O2)?
• moles O2  grams O2
• amount of O2 (mol)  molar mass of O2 (g/mol) = mass of O2 (g)
Sample Problem
Ibuprofen, C13H18O2, is the active ingredient in many
nonprescription pain relievers. Its molar mass is
206.31 g/mol.
a. If the tablets in a bottle contain a total of 33 g of
ibuprofen, how many moles of ibuprofen are in
the bottle?
b. How many molecules of ibuprofen are in the
bottle?
c. What is the total mass in grams of carbon in 33 g
of ibuprofen?
Percent Composition
• To find the mass percentage of an element in
a compound, the following equation can be
used.
• Mass of element/mass of total sample x 100 =
percent comp.
Sample Problem
Find the percentage composition of copper(I)
sulfide, Cu2S.
More problems
A. From data: What is the percent composition
of a compound made from 222.6 g Na and
77.4 g O?
B. From formula: Find the percent composition
of sodium sulfate.
Empirical Formula
• The simplest ratio of atoms in a compound
• Need percent composition to find it
Empirical formula, cont.
Example: Find the empirical formula of a
compound that is 79.9% C and 20.1 % H.
1. Assume 100 g
2. Convert to moles
3. Divide by smallest number of moles
*If all are close to whole numbers, stop
*If NOT, multiply all to make them all whole
4. Write the formula
Empirical formula, cont.
Ex: Find the empirical formula of a compound
that is 17.6% Na, 39.7% Cr, and 42.7% O.
1. Assume 100 g
2. Convert to moles
3. Divide by smallest number of moles
*If all are close to whole numbers, stop
*If NOT, multiply all to make them all whole
4. Write the formula
• Example: diborane
• The percentage composition is 78.1% B and
21.9% H. What is the empirical formula?
• Sample Problem:
• Quantitative analysis shows that a compound
contains 32.38% sodium, 22.65% sulfur, and
44.99% oxygen. Find the empirical formula of
this compound.
Molecular Formulas
• NOT the simplest ratio
• Is a MULTIPLE of the empirical formula
• Ex.
C2H4  CH2
Molecular
formula
Empirical
formula
Molecular formulas, cont.
Ex. A compound that is 58.8% C, 9.8% H, and
31.4% O, has a molecular/formula mass of 204
g/mol. Find its formula mass.
1. Find the empirical formula
2. Find mass of the empirical formula
3. Divide molecular mass by empirical mass to
find multiplier
4. Write the molecular formula
Chapter Review
• Pg. 251
• #’s 2-8, 10-12, 14, 23, 30-32,
35-38
Mole review
1 mole = 6.02E23 atoms
1 mole = 6.02E23 molecules
1 mole = 6.02E23 formula units
1 mole = ____g
1. How many formula units are in 4.3 grams silver
nitrate?
2. How many moles is 3.43 X 1024 molecules
carbon dioxide?
3. How many grams is 2.65 X 1023 atoms of
aluminum?