Transcript Chapter 7: Naming Compounds

Chapter 7: Naming
Compounds
Naming Ionic
Compounds
• Review: Chemical formula
indicates the name and number
of atoms in a formula
• Ex: C8H18
• Al(SO4)3
• Monoatomic ions-ions with just one
atom
• Ex: Li+
• Naming Monoatomic Ions:
• Cations are simply the element’s
name
• Anions-drop the ending of the
element’s name and add –ide
• Ex: F- Fluoride
You Try It!
• Name the following monoatomic
anions:
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Carbon
Nitrogen
Oxygen
chlorine
Binary Ionic compounds
• Compounds composed of two
different compounds
• The bond must be between a
negative and a positive ion and
the charges must balance to a
zero
Writing Binary
Compound Formulas
• Writing the formula:
• Write each symbol with its
charge
• Write the cation first and then
cross the charges
• Al+3 O-2
• Al2O3
You Try It!
• Zinc and iodine
• Zinc and sulfur
• Potassium and iodine
• Sodium and sulfur
Naming Binary Ionic
Compounds
• Write the name of the cation
first and then the anion
(remember that the ending to
the anion is changed)
• Ex: Al2O3
• Aluminum Oxide
You Try It!
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AgCl
ZnO
CaBr2
SrF2
BaO
CaCl2
Stock System of
Nomenclature
• Some cationic elements can have
more than one correct charge so a
Roman Numeral in parentheses is
used after the symbol to show the
correct charge
• Ex: Iron (II) Fe+2
• Iron (III) Fe+3
Naming an ionic Compound
using the Stock System
• CuCl2
• Copper (II) Chloride
You Try It!
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Write the formula and name the
following compounds Using the
Stock System:
Cu+2 and BrFe+2 and O-2
Pb+2 and ClHg+2 and S-2
Sn+2 and FFe+3 and O-2
Polyatomic Ions
• Pg. 210 Table 7-2
• A group of atoms with a
common charge
• Oxyanions-polyatomic ions that
contain oxygen
• Ex: OH-
• Some have the same two elements
but are different
• Most common ion is given the ending
–ate
• One less oxygen is given the ending
–ite
• One less oxygen than the –ite is
given the prefix hypo• One more oxygen than the –ate is
given the prefix per-
Examples
• Hypochlorite ClO-
• Chlorite ClO2• Chlorate ClO3• Perchlorate ClO4-
Naming Binary
Molecular Compounds
• There are two methods to
naming binary molecules
• The older is the prefix system
• The newer is the Stock System
Prefix System
• Table 7-3
• Rules for naming:
• The less electronegative element is
first. Prefix given only if it has more
than one atom
• The second word is begun with a
prefix for the number of atoms,
followed by the root word of the
element, and ending with an –ide
(indicating it has only 2 elements)
Examples
• P4O10
• Tetraphophorus decoxide
• SO3
• Sulfur trioxide
You Try It!
• Write the name of the
compound using prefixes:
• ICl3
• PBr5
• N2O3
• Write the formulas for the
compounds using the prefix
system:
• Carbon tetraiodide
• Phosphorus trichloride
• Nitrogen monoxide
Acids
• Binary acids-acids that consist
of two elements, usually
hydrogen as the cation and
another element, dissolved in
water.
• Oxyacids-acids that contain
hydrogen, oxygen, and a third
element
• Look at some examples on pg.
214 and Table 7-5
How acids are named
Compound
-ide
Acid Name
hydro- -ic acid
-ite
-ous acid
-ate
-ic acid
Examples
• HClO3 Hydrogen Chlorate
Acid: chloric acid
• HClO2 Hydrogen chlorite
Acid: Chlorous Acid
• HCl
Hydrogen chloride
Acid: hydrochloric acid
Salt
• An ionic compound composed of
a cation and the anion from an
acid
Acid Salts
• Ionic compounds that still
contain an acidic hydrogen,
such as NaHSO4
• Sodium hydrogen sulfate
Hydrocarbons
• Contain only carbon and
hydrogen and are the simplest
organic compounds
• Alkanes are the simplest
hydrocarbons and are only
carbon-carbon single bonds.
Alkanes (intro to organic
chemistry)
• Straight chain alkanes are
where the carbons are all linked
in a straight chain (no
branching)
• CnH2n+2
Examples
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CH4 Methane
C2H6 Ethane
C3H8 Propane
C4H10 Butane
C5H12 Pentane
C6H14 Hexane
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C7H16 Heptane
C8H18 Octane
C9H20 Nonane
C10H22 Decane
Oxidation NumbersRules
• A pure element has an oxidation
number of zero
• The more electronegative
element is assigned the
negative charge. The least
electronegative element is
assigned the number of its
positive charge
• Fluorine is -1 always
• Oxygen is almost always -2
• Hydrogen has +1 in compounds with
elements more electronegative than
it. It has -1 with metals
• The sum of all oxidation numbers
must equal zero
• The oxidation number for polyatomic
ions is equal to the sum of its
charges
Assigning Oxidation
Numbers
• UF6
Oxidation number for F is -1
(times 6 because there are 6
atoms= -6)
So, to equal 0, the oxidation
number of U must be +6
• H2SO4
• Most electronegative? Oxygen, so it
has -2 (times 4 for a total of -8)
• Hydrogen is least electronegative so
it is +1(times 2 for a total of +2)
• If we subtract +2 from -8 we get -6
so Sulfur must have +6
You Try It!
• Assign oxidation numbers:
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HCl
CF4
HNO3
P4O10
HF
H2O
H2CO3
PI3
CI4
ClO2IO3-
• Table 7-6 gives oxidation numbers
for some nonmetals that can have
more than one.
• These are also used in the Stock
System learned earlier (roman
numeral)
• SO2 SO3
• Sulfur (IV) oxide and Sulfur (VI)
oxide
Formula Masses
• Sum of the average atomic masses
of all atoms in the formula
• H2O
• 2H X 1.01 = 2.02
• 1 O X 16.00= 16.00
• 16.00 + 2.02 = 18.02 amu for H2O
You Try It!
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Find the formula masses:
KClO3
H2SO4
Ca(NO3)2
PO4-3
MgCl2
Molar Mass
• Grams/mole for a molecule
• Numerically equal to the
formula mass
• Ex: H2O is equal to 18.02 g/mol
Conversion Factors
• What is the mass in grams of
2.5 mol of Oxygen gas (O2)?
• Amount in moles X molar mass
= mass in grams
• 2.5 mol X 32.00g = 80 g
Sample Problem 7-9
• a. Grams to moles
• B. Moles to molecules
• C. Moles to moles to grams
You Try It!
• Practice Problems pg. 226
Percent Composition
• Percentage of one element in a
compound
• Mass of element in a compound X 100 %
Mass of a sample of compound
• Find the percent composition of
copper(I) sulfide, Cu2S
• Mass of copper = 127.1 g
• Mass of Sulfur = 32.07 g
• Molar Mass = 159.2 g
• % Cu (127.1g/159.2g) X 100= 79.84%
• % S (32.07 g/159.2 g) X 100= 20.14%
You Try It!
• Find the percent compositions
of the following:
• PbCl2
• Ba(NO3)2
Empirical Formulas
• Consists of symbols for the elements
combined in a compound, with
subscripts showing the smallest
whole-number mole ratio of the
different atoms in the compound
• Ex: B2H6
• Empirical formula BH3
How to Calculate
Empirical Formula
• BH3
• 78.1% B
• 21.9% H
• 78.1 g B X (1mol B/10.81 g B)
=7.22 mol B
• 21.9 g H X (1 mol H/1.01 g H) =
21.7 mol H
• These numbers give us numbers of
moles but not smallest whole
number ratios so divide each by the
smallest number:
• 7.22 mol B/7.22 = 1 B
• 21.7 mol H/7.22 = 3.01 H
• So, the formula is BH3
You Try It!
• Pg. 231 Practice Problems 1-3
Molecular Formula
• Actual formula of a molecule
instead of the smallest wholenumber formula
• X(empirical formula)=molecular
formula
• Diborane’s formula mass is
27.67 amu and the empirical
formula mass is 13.84 amu
• X= molecular formula mass
empirical formula mass
X = 27.67 amu = 2.00
13.84 amu
2.0(BH3) = B2H6
You Try It!
• Pg. 233 1-2 Practice Problems