Unit 5 Nomenclature.
Balancing Chemical Reaction
Equations
Molecules
• Many compounds exist as molecules.
• Examples
• In a molecule formula, the cation is always written first, on the left side,
followed by the anion. In our example, phosphorus is the cation end and
sulfur is the anion end.
Common Names
• H2O = water, steam, ice
• NH3 = ammonia
• CH4 = methane
• NaCl = table salt
• C12H22O11 = table sugar
Classifying Compounds
Naming Starts with Classifying Compounds
• Binary Compounds = only 2 elements
• Compounds containing polyatomic ions
• Acids = formula often starts with H (HCl, H2S)
• Names of chemical compounds have two parts: the cation name part
and the anion name part.
• The cation part is the name of the element in the cation position of
the molecule, on the left side; in our example, it is Na, sodium.
• The anion has its name changed from its original element name by
adding the ending -ide. In our example, the original anion element is
chlorine, but its name in the compound is chloride.
• So, NaCl name is sodium chloride.
Type I Binary Compounds Examples
Practice – Type I Binary Compounds
• 11) Cs2S
• 12) KCl
• 13) Sr3P2
• 14) BaI2
• 15) NaF
• 16) CaBr2
• 17) BeO
• 18) SrS
• 19) BF3
• 20) AlP
Practice
• 11) Cs2S
• 12) KCl
• 13) Sr3P2
• 14) BaI2
• 15) NaF
• 16) CaBr2
• 17) BeO
• 18) SrS
• 19) BF3
• 20) AlP
Answers
Type II Binary Compounds
• Type II binary compounds are made of:
•
cations: transition (or transitional) metals and metaloids from
'under the staircase' in the periodic table, except group 3A (transition
metals are called 'transition' because they form cations with variable
charges)
•
anions: non-metals
• Type II compounds are named with the same method as compounds
in the first category, except the charge of the metal ion, the cation) is
specified by a Roman numeral in parentheses after the name of the
metal.
Determining the Charge of the Cation
• The charge of the metal ion is determined from the formula of the
compound and the charge of the anion.
• For example, consider binary ionic compounds of iron and chlorine.
Iron typically exhibits a charge of either 2+ or 3+, and the two
corresponding compound formulas are FeCl2 and FeCl3. The simplest
name, “iron chloride,” will, in this case, be ambiguous, as it does not
distinguish between these two compounds.
• Type II binary compounds are also called ionic compounds.
Determining the Charge of the Cation
• Determine the charge on the anion
•
Au2S3 - the anion is S, since it is in Group 6A, its charge is -2
• Determine the total negative charge
•
since there are 3 S in the formula, the total negative charge is -6
• Determine the total positive charge
•
since the total negative charge is -6, the total positive charge is
+6
• Divide by the number of cations
•
since there are 2 Au in the formula & the total positive charge is +6, each
Au has a +3 charge
Type II Binary Compounds Examples
Practice – Type II Binary Compounds
Practice – Type II Binary Compounds
Type III Binary Compounds
• Type III binary compounds are made of:
• cations: non-metals, for example, carbon C, nitrogen N, sulfur S,
fluorine F
• anions: non-metals, for example, carbon C, nitrogen N, oxygen O,
hydrogen H, phosphorus P
• Type III binary compounds are often referred to as molecular
compounds or covalent compounds.
Naming Type III Binary Compounds
• To name type III binary compounds,
• 1. Identify the elements in the molecule from its formula.
• 2. Begin the name with the element name of the first element.
• 3. If there is more than one atom of this element in the molecular formula,
use a numerical prefix to indicate the number of atoms, as listed in the
table in the next slide. Do not use the prefix mono- if there is only one
atom of the first element.
• 4. Name the second element by using three pieces: a) a numerical prefix
indicating the number of atoms of the second element, plus b) the stem of
the element name (e.g., ox for oxygen, chlor for chlorine, etc.), plus c) the
suffix -ide. Combine the two words, leaving a space between them.
Prefixes
Practice – Type III Binary Compounds
• 11) P2O5
• 12) S2Cl2
• 13) ICl2
• 14) SO2
• 15) P4O10
• 16) UF6
• 17) OF2
• 18) ClO2
• 19) SiO2
• 20) BF3
Answers – Type III Binary Compounds
• 11) P2O5
• 12) S2Cl2
• 13) ICl2
• 14) SO2
• 15) P4O10
• 16) UF6
• 17) OF2
• 18) ClO2
• 19) SiO2
• 20) BF3
Naming Polyatomic Compounds
Polyatomic Compounds
• Polyatomic compounds contain three or more elements. Polyatomic
compounds are also often called ionic compounds. Polyatomic
compounds can be Type I or Type II.
• Type I polyatomic compounds are compounds made of:
• cations: metals from groups 1A, 2A, and 3A in the periodic table
•
anions: polyatomic anions from the chart
Selected Polyatomic Ions Chart
Extended Polyatomic Ions Chart
• https://www.templateroller.com/template/93693/commonpolyatomic-ions-chart.html
Naming Type I Polyatomic Compounds
• Let's name this compound: Ca(NO3)2
• Names of chemical compounds have two parts: the cation name part
and the anion name part.
• The cation part is the name of the element in the cation position of
the molecule, on the left side; in our example, it is Ca, calcium.
• The anion part will be the name of the anion NH3, nitrate. So, the
name of our compound is calcium nitrate.
Practice – Naming Type I Polyatomic
Compounds
• 1) AlPO4
• 2) KNO2
• 3) NaHCO3
• 4) CaCO3
• 5) Mg(OH)2
Answers
• 1) AlPO4
• 2) KNO2
• 3) NaHCO3
• 4) CaCO3
• 5) Mg(OH)2
Naming Type II Polyatomic Compounds
• named with the same method as type II binary compounds
• the charge of the metal ion, the cation, specified by a Roman numeral in
parentheses after the name of the metal.
• The charge of the metal ion is determined from the formula of the compound
and the charge of the anion.
• For example, consider Fe(OH)2 and Fe(OH)3 . From the polyatomic ions' chart,
we know that OH charge is -1. In the first compound, there are two hydroxide
ions with -1 charge; therefore, the total negative charge is -2. There is only one
atom with a positive charge, Fe, and the entire positive charge balancing -2
(which will be +2) has to be on the Fe ion. So, Fe is +2 in the first compound; its
name is iron (II) hydroxide.
• Similarly, the second compound's name is iron (III) hydroxide. There is also an
older system of naming type II compounds; please, go back to the type II binary
compounds to review.
Practice – Naming Type II Polyatomic
Compounds
• 11) Sn(NO3)2
• 12) FePO4
• 13) Cu2SO4
• 14) Ni(C2H3O2)2
• 15) HgCO3
Answers
• 11) Sn(NO3)2
• 12) FePO4
• 13) Cu2SO4
• 14) Ni(C2H3O2)2
• 15) HgCO3
11) tin(II) nitrate [stannous nitrate]
12) iron(III) phosphate [ferric phosphate]
13) copper(I) sulfate [cuprous sulfate]
14) nickel(II) acetate [nickelous acetate]
15) mercury(II) carbonate [mercuric carbonate]
Writing Formulas from Names
•
•
For Type III compounds, use the prefixes to determine the
subscripts
For Type I, Type II, polyatomic Compounds and Acids
• Determine the ions present
• Determine the charges on the cation and anion
• Balance the charges to get the subscripts
Examples
Example 1: Write the formula from the following name: sodium
bromide
• Step #1 - Write down the symbol and charge of the first word. Result
= Na+
• Step #2 - Write down the symbol and charge of the second word.
Result = Br¯
• Step #3 - Use the minimum number of cations and anions needed to
make the sum of all charges in the formula equal zero. In this case,
only one Na+ and one Br¯ are required.
• The resulting formula is NaBr.
Examples
Example 6: Write the name of the following formula: aluminum oxide
• Step #1 - Write down the symbol and charge of the first word. Result = Al3+
• Step #2 - Write down the symbol and charge of the second word. Result =
O2¯
• Step #3 - Use the minimum number of cations and anions needed to make
the sum of all charges in the formula equal zero. In this case, two Al3+ are
required and three O2¯
• This is the only possible way to get the positive and negative charges equal and keep
the numbers to a minimum. Note that the positive charge is a +6 and the negative
charge is a -6.
• Also, keep in mind that you cannot change the charges to make a formula correct.
• The resulting formula is Al2O3.
• Warning: beware of the temptation to write the above formula as Al3O2.
Examples
Example #1 - write the formula for copper(II) chlorate
• Step #1 - the first word tells you the symbol of the cation. In this case
it is Cu.
• Step #2 - the Roman numeral WILL tell you the charge on the cation.
In this case it is a positive two.
• Step #3 - the polyatomic formula and charge comes from the second
name. In this case, chlorate means ClO3¯.
• Step #4 - remembering the rule that a formula must have zero total
charge, you write the formula Cu(ClO3)2.
Practice
• Write the correct formula for:
• 1) magnesium oxide
• 2) lithium bromide
• 3) calcium nitride
• 4) aluminum sulfide
• 5) potassium iodide
• 6) strontium chloride
• 7) sodium sulfide
• 8) radium bromide
• 9) magnesium sulfide
• 10) aluminum nitride
Answers
• Write the correct formula for:
• 1) magnesium oxide
• 2) lithium bromide
• 3) calcium nitride
• 4) aluminum sulfide
• 5) potassium iodide
• 6) strontium chloride
• 7) sodium sulfide
• 8) radium bromide
• 9) magnesium sulfide
• 10) aluminum nitride
Practice - Worksheet
Naming Acids
• Contain H+ cation and anion
• Binary acids have H+ cation and a nonmetal anion
• Oxyacids have H+ cation and a polyatomic anion
• Acid examples:
Naming Binary Acids
• The word “hydrogen” is changed to the prefix hydro• The other nonmetallic element name is modified by adding the suffix
–ic
• The word “acid” is added as a second word
• For example, when the gas HCl (hydrogen chloride) is dissolved
• in water, the solution is called hydrochloric acid.
• Other examples of this nomenclature are shown in the figure.
Naming Oxyacids
• Omit “hydrogen”
• Start with the root name of the anion
• Replace –ate with –ic, or –ite with –ous
• Add “acid”
• For example, consider H2CO3 (which you might be tempted to call
“hydrogen carbonate”). To name this correctly, “hydrogen” is omitted; the
–ate of carbonate is replace with –ic; and acid is added—so its name is
carbonic acid.
• Other examples are given in the table. There are some exceptions to the
general naming method (e.g., H2SO4 is called sulfuric acid, not sulfic acid,
and H2SO3 is sulfurous, not sulfous, acid).
Practice
Answers
Balancing Chemical Reaction
Equations
Chemical Reactions - Basics
• A chemical reaction expresses a chemical change. For example, one chemical property of
hydrogen is that it will react with oxygen to make water. We can write that as follows:
• hydrogen reacts with oxygen to make water
• We can represent this chemical change
• hydrogen + oxygen → water
• where the + sign means that the two substances interact chemically with each other and
the → symbol implies that a chemical reaction takes place.
• Substances can also be represented by chemical formulas.
• we can rewrite our chemical change as
• H2 + O2 → H2O
• This is an example of a chemical equation, which is a concise way of representing a
chemical reaction. The initial substances are called reactants, and the final substances
are called products.
Examples of Chemical Reactions
Evidence of Chemical Reactions
• A chemical change occurs when new substances are made.
• Look for visual clues (permanent).
• color change, precipitate formation, gas bubbles, flames, heat release,
cooling, light
• Look for other clues.
• new odor, permanent new state
Chemical Equations
• Shorthand way of describing a reaction
• Provides information about the reaction
• Formulas of reactants and products
• States of reactants and products
• Relative numbers of reactant and product molecules that are required
• Can be used to determine weights of reactants used and of products that can
be made
Practice – Information about Equations
Answers
Conservation of Mass
• Matter cannot be created or destroyed.
• In a chemical reaction, all the atoms present at the beginning are still
present at the end.
• Therefore, the total mass cannot change.
• Therefore, the total mass of the reactants will be the same as the
total mass of the products.
Combustion of Methane
• Methane gas burns to produce carbon dioxide gas and liquid water.
• Whenever something burns it combines with O2(g).
CH4(g) + O2(g)  CO2(g) + H2O(l)
• To show the reaction obeys the Law of Conservation of Mass, it must
be balanced.
CH4(g) + 2 O2(g)  CO2(g) + 2 H2O(l)
Writing Equations
• Use proper formulas for each reactant and product.
• Proper equation should be balanced.
• obey Law of Conservation of Mass
• all elements on reactants side also on product side
• equal numbers of atoms of each element on reactant side as on product side
• Balanced equation shows the relationship between the relative
numbers of molecules of reactants and products.
• can be used to determine mass relationships
• symbols used after chemical formula to indicate state
• (g) = gas; (l) = liquid; (s) = solid
• (aq) = aqueous, dissolved in water
Balancing by Inspection
Count atoms of each element
• polyatomic ions may be counted as one “element” if it does not change in
the reaction
Al + FeSO4 Al2(SO4)3 + Fe
1 SO4 3
• if an element appears in more than one compound on the same side, count
each separately and add
CO + O2  CO2
1 + 2 O 2
Balancing by Inspection
Pick an element to balance
3. Find Least Common Multiple and factors needed to make both
sides equal
4. Use factors as coefficients in equation
• if already a coefficient then multiply by new factor
5. Recount and Repeat until balanced
6. Examples and practice:
https://www.chemteam.info/Equations/Balance-Equationproblem-list.html
Suggested Order of Balancing
• There are various approaches
• One of the approaches is the MINOH method (Me know chemistry, said Tarzan as he climbed the stoichiome-tree.)
• M - metals. Balance metals such as Fe or Na first.
I - ions. Looks for polyatomic ions (such as PO4¯3 or SO4¯2 that cross from reactant to product unchanged.
Balance them as a group.
N - non-metals. Look for Cl or S, these are common ones.
O - oxygen, and then H - hydrogen.
• Often, balancing H and O will involve water on one side or the other.
• look for elements which occur in only one place on each side of the arrow. These should be balanced before
examining elements that are spread over several compounds.
• Often, either H or O will be spread out over several compounds. This is the one to leave to the last.
• cannot change a subscript to balance the equation,
• nor can you add in new compounds.
• Cannot leave equations as coefficients
• Coefficients must be the lowest numbers possible
Balancing Example
Types of Chemical Reactions
• Precipitation reactions – when a solid forms
• Acid-base reactions – also called neutralization reaction, when acid and
base react
• Oxidation-reduction reactions – involves electron transfer; changes in
charge happen
• Synthesis reaction – two or three reactants produce a more complex
product
• Decomposition reactions – one complex reactant breaks down into simpler
product
• Combustion reactions – also called ‘burning’; reactions involving oxygen
Other Types of Reactions
• Double displacement reactions – also called double replacement, two
groups of ions exchange places
• Single displacement – also called single replacement, one group or
one ion switches places as a result of the reaction
Examples
More
Examples
Practice
Answers