Transcript Nomenclature

Chapter 6 – Nomenclature
 Common names – were used before a
systematic method was developed.
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H2O = water
NH3 = ammonia
 Inorganic nomenclature
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Splits up into two main categories – Ionic and Molecular
(covalent) compounds.
Nomenclature
Compound
Ionic
Fixed
Charge
Molecular
Variable
Charge
Ionic Compounds
 Contains a cation (usually a _____) and an anion
(usually a _________).
 The metal can have a FIXED charge (or ________)
or a VARIABLE charge.
 Metals with fixed charges are:
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Main Group metals (group 1a, 2a, Al), but NOT Sn and Pb
 Metals with variable charges are:
 Transition metals, but NOT Zn(+2) and Ag(+1)
Ionic Compounds
 For a metal with a FIXED charge, the
rules are:
1. Name the metal first.
2. Name the non-metal second, but
change its suffix to –ide.
Learning Check
 Name the following ionic compounds whose metals
have a FIXED charge.
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Ex) CaCl2
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Ex) AlF3
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Ex) K2S
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Ex) Mg3N2
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Ex) ZnO
Ionic Compounds
 Converting a name to a formula.
 Rule: an ionic compound is always written
such that it has no net charge.
 Let’s see if that is true for the previous slide!
 Ex) Write the formula for the compound
formed between Li+1 and S-2.
 Ex) Write the formula for the compound
formed between Al+3 and O-2.
Learning Check
 Write the proper formula given the ionic
name.
 Ex) potassium bromide
 Ex) calcium iodide
 Ex) aluminum chloride
 Ex) lithium nitride
Ionic Compounds
 When the metal has more than one valence, then the rules
are changed.
 Ex) FeO and Fe2O3 are two common compounds formed
between iron and oxygen. We cannot name both as “iron
oxide” since they have a different ratio.
 Rules:
1. Determine the charge of the metal ion by deduction.
2. Name the metal followed by that charge in Roman
Numerals in parenthesis.
3. Name the non-metal, change suffix to –ide.
Ionic Compounds
 Roman numerals
 +1 = I
 +2 = II
 +3 = III
 +4 = IV
 +5 = V
 +6 = VI
Learning Check
 Name each of the following.
 Ex) Cu2S
 Ex) FeCl3
 Ex) CoO
 Ex) V2O3
Learning Check
 Write formulas for each of the names.
 Ex) manganese(II) chloride
 Ex) nickel(II) sulfide
 Ex) titanium(IV) oxide
 Ex) chromium(III) fluoride
Molecular Compounds
Consist of two non-metals or a metalloid and a
non-metal.
 Compounds share electrons, so charges do NOT
apply.
 Rules:
1. Name first element in formula as is.
2. Name second element in formula and change the
suffix to an –ide ending.
3. Add prefixes for each subscript – exception: mono
only used for second element.
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Molecular Compounds
 Prefixes:
 1 = mono
 2 = di
 3 = tri
 4 = tetra
 5 = penta
 6 = hexa
Learning Check
 Name each of the following molecular compounds.
 Ex) CO2
 Ex) N2O4
 Ex) NF3
 Ex) SF4
Learning Check
 Write formulas for:
 Ex) disulfur tetrachloride
 Ex) nitrogen monoxide
 Ex) diphosphorus pentoxide
 Ex) silicon tetrachloride
Polyatomic Ions
 An ion that contains two or
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more elements is called a
polyatomic ion.
These are very common!
Ex) CO3-2 = carbonate ion
Ex) SO4-2 = sulfate ion
Ex) PO4-3 = phosphate ion
Any compound containing a
polyatomic ion is IONIC.
Polyatomic Ions
 Be very careful…
 NO2 and NO2-1 are not the same thing!!!
 One and only one polyatomic ion is a cation – NH4+1.
 Polyatomic ions ALWAYS keep their names.
 Some compounds may have more than one of the
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same polyatomic ion in the formula.
Ex) MgSO4 = magnesium sulfate
Ex) magnesium sulfide = _______
Ex) LiNO3 = lithium nitrate
Ex) lithium nitride = _______
Learning Check
 Name the compounds with one polyatomic ion. Note
that the metals may have a FIXED or a VARIABLE
charge!
 Ex) Na3PO4
 Ex) FeSO4
 Ex) Zn(OH)2
 Ex) Mg(NO2)2
Learning Check
 From the name, write the correct formula.
 Ex) ammonium sulfide
 Ex) copper(II) hydroxide
 Ex) nickel(II) phosphate
 Ex) potassium chromate
Chapter 7 – Quantitative Composition
 Paper is not sold as a single sheet, rather it is sold as
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500 sheets called a _____.
Pencils are measured by the _____ (144 pencils).
Gasoline and other liquids are sold by the ______.
Atoms are too small to weigh on a scale!
Rather, chemists count atoms by the MOLE.
One mole of anything is 6.02 x 1023 “objects.”
602,000,000,000,000,000,000,000
Note: National debt = 11,900,000,000
The Mole
 The unit on a mole depends on the context.
 Ex) One mole of Carbon = 6.02 x 1023 atoms of Carbon
 Ex) One mole of SO2 = 6.02 x 1023 molecules of SO2
 Ex) One mole of CaCl2 = 6.02 x 1023 formula units of CaCl2
Molar Mass
 The atomic weights on the periodic chart were
originally interpreted on the atomic mass unit scale
(amu).
 With the mole concept, this is now interpreted as
grams.
 A molar mass is the amount of mass, in grams, that
one mole of a substance would weigh. Where one
mole is 6.02 x 1023 particles.
Learning Check
 What is the molar mass of:
 Ex) Ar =
 Ex) CH4 =
 Ex) Na2SO4 =
 Ex) Al(NO3)3 =
Using a Molar Mass
 A molar mass can be used to convert between moles
and grams.
 Ex) 5.00g of Ar = ? moles
 Ex) 0.250 moles of CH4 = ? grams
 Ex) 2.85g of Na2SO4 = ? moles
 Ex) 4.52 moles of Al(NO3)3 = ? grams
Using Avogadro’s Number
 Ex) How many molecules are present in 0.0155
moles of H2O?
 Ex) How many moles are present in 4.25 x 1024
molecules of NF3?
Grams to Molecules
 We can combine the two problems into a single
problem.
 Ex) How many molecules are present in a 2.20g
sample of CO2?
Molecules to Grams
 How many grams are present in a sample of PF3
containing 6.25 x 1021 molecules?
Grams to Atoms
 An important distinction – atoms and molecules are
NOT the same thing.
 One molecule of CH4 contains one atom of C and four
atoms of H.
 Ex) How many H atoms are present in 0.400 grams
of CH4?
Percent Composition
 One calculation that involves using your molar mass
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is called the percent composition.
Step 1: Calculate the molar mass of the compound.
Step 2: % A = total mass of A  molar mass x 100
Note that the sum of all the mass percents = 100
Ex) Calculate the mass percents for Na2CO3.
Empirical Formulas
 An empirical formula is one that has the lowest whole
number coefficients for each element.
 Empirical formula from mass percentages:
1. Assume a 100.g sample.
2. Convert grams to moles by dividing each gram
amount by the element’s atomic weight.
3. Divide each mole amount in step 2 by the smallest
number of moles.
4. If necessary, multiply to get rid of decimal
equivalent of a fraction.
Empirical Formulas
 0.500 = ½ ; therefore, multiply all by 2.
 0.250 = ¼; therefore, multiply all by 4.
 0.333 = 1/3 or 0.667 = 2/3; multiply all by 3.
 Ex) A compound is found to be 56.58% K, 8.68% C,
and 34.73% O by mass. What is the empirical
formula?
Empirical Formulas
 Ex) A compound is found to be 43.7% P and 56.3%
O by mass. What is the empirical formula?
 Ex) A compound is analyzed and found to be 17.5%
Na, 39.7% Cr, and 42.8% O. What is its empirical
formula?
Chapter 8: Chemical Equations
 A Chemical Equation is a symbolic representation of
a chemical reaction.
 Involves a rearrangement of atoms or ions into new
combinations.
Chemical Equations
Chemical Equations
 A Chemical Equation is always written so that the
total numbers of atoms on each side of the equation
are equal.
 To do this, coefficients are added in front of each
substance.
 Must use lowest whole number coefficients!
2 Al(s) + Fe2O3(s)  2 Fe(l) + Al2O3(s)
Coefficients (note: a “1” is not written!)
coefficients
Chemical Equations
Balancing an Equation
 __Al + __Cl2  __AlCl3
 Reaction is NOT balanced!
 Note: can NOT alter any subscripts!
 Make a tally sheet.
 Always start with elements that occur only once on
each side of the equation.
Balancing an Equation
 Ex) __Mg + __O2  __MgO
 Ex) __KClO3  __KCl + __O2
 Ex) __H3PO3  __H3PO4 + __PH3
 Ex) __AgNO3 + __H2S  __Ag2S + __HNO3
Types of Reactions
 Combination Reaction
+ O2(g)  2MgO(s)
 CaO(s) + CO2(g)  CaCO3(s)
 2Mg(s)
Types of Reactions
 Decomposition Reactions
 (NH4)2Cr2O7(s)  Cr2O3(s) + 4H2O(g) + N2(g)
Types of Reactions
Types of Reactions
 Single Replacement
 2Al(s) + Fe2O3(s)  Al2O3(s) + 2Fe(s)
Types of Reactions
 Double Replacement
 Pb(NO3)2(aq) + K2CrO4(aq)  PbCrO4(s) + 2KNO3(aq)
Types of Reactions
Types of Reactions
 Combustion – a rapid reaction with Oxygen (O2)
usually involving hydrocarbons (C, H compounds)
 Products are always carbon dioxide and water.
 Balancing these: C, H, then O last.
 Odd, Even oxygen in an equation.
Types of Reactions
 Balance the following combustion reactions.
 Ex) __C3H8 + __O2  __CO2 + __H2O
 Ex) __C5H10 + __O2  __CO2 + __H2O