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Nomenclature,
Chemical Equations
and Reactions
Nomenclature is a system of
naming chemical
compounds.
Background terms
Oxidation: gain or loss of electron from neutral atom
Oxidation state (valence): "charge" on atom
General rules for determination of oxidation states:Group A elements1. positive oxidation state is equal to column
number2. negative oxidation state is equal to
column number minus 8 (beginning with column IVA)
3. if an atom is in an even numbered column, it will have an even number oxidation state; similarly with odd numbered columns
4. maximum change is going to be "8" (magic number)
Group B elements1. one possible positive oxidation state is the column number except for column VIIIB
2. another possible positive oxidation state is "2+" except for IIIB which will only have "3+"
More rules:always
IA +1IIA +2Al, Sc +3Zn, Cd +2
Metallic endingsnew system old system iron (III) Fe+3 ferric iron (II) Fe+2 ferrous-ic suffix usually corresponds to highest oxidation state element usually takes
Example:Cr (group VIB) maximum + charge=
group number 6+ 3+ chromic (usually exists this way)
2+ chromous Sntin (IV) or stannic SnF4 stannic fluoride
tin (II) or stannous SnF2 stannous fluoride
BINARY COMPOUNDS are compounds made up of two elements
metal + non-metalThe metal always takes its given name; a commonsuffix for a metal is -ium
non-metal suffixes are usually -on, -gen, and -ine The metal takes its given name, the non-metal suffixchanges to -ide ex. NaCl sodium chloride KF potassium fluoride AlN aluminum nitride
Cross-over process1. write the symbols and charges for the ions
next to each other, always writing the cation first
2. cross over the charges by using the absolute value of each ion’s charge as the subscript for the other ion
3. check the subscripts and divide them by their largest common factor to give the smallest possible whole-number ratio of ions.
The purpose of this process is to ensure that the compound is electrically neutral.
Binary two non-metalsHere we need the use of various prefixes to tell ushow many of each kind of atom we have present
mono- 1 hexa- 6di- 2 hepta- 7tri- 3 octa- 8tetra- 4 nona- 9penta- 5 deca- 10
When we name a binary compound composed of two
non-metals we state how many atoms of each kind we have by using the above prefixes. If we only have one of the first element listed, we do not need to state that by using the prefix mono-. However, we do need to state any other quantity of the elements.
Examples:ex. NO mononitrogen monooxide N2O dinitrogen monoxide
NO2 nitrogen dioxide
N2O3 dinitrogen trioxide
N2O4 dinitrogen tetraoxide
When we have similar vowels together, such as two O's or an A and an O we cancel out the first vowelto make for more sensible spelling and pronunciation.
Ternary compoundsCompounds made of 3 or more elements, one of which is
usually oxygen.
metal + radical (polyatomic ion) (takes its (2 or more elements joined
name) together with a residual charge; when written as the second component of a
compound it usually has a negative charge)
The radical name is based on the amount of oxygen it
contains and is named according to this definition even if it doesn't contain any oxygen.
RADICAL NAMING CHARTEXPLANATION EXAMPL
ESUFFIX FORMAT
NAME
1 more O than normal
ClO4- per-root-
ate perchlorate
normal amount of O
ClO3- -ate chlorate
1 less O than normal
ClO2- -ite chlorite
2 less O than normal
ClO- hypo-root-ite
hypochlorite
More “radical” rulesIn group VIIA: Cl, Br, and I follow the precedingrules; F does not
In group VIA: S and those elements below it followthe above rules (4 oxygens is the "normal" amount)
In group VA: P and As follow the above rules (4 oxygens is "normal“ for P and As) [3 oxygens is the
“normal“ amount for N]
An important point to remember is that the oxidation state of the radical does not change as the amount of oxygen changes.
Examples that do not follow rules: CN- cyanide CNO- cyanate SCN- thiocyanate
(thio- tells us there is sulfur present)
Examples using polyatomic ions: NaClO3 sodium chlorate
MgSO4 magnesium sulfate
AlPO3 aluminum phosphite
ACID NAMINGbinary acidsTo recognize a binary acid, it must be made up oftwo elements (one of which is hydrogen) andmust have the word aqueous, abbreviated,following the formula. The name must includethe prefix hydro- and the word acid, i.e.
Hydro -root-ic acid
ex. HCl(aq) hydrochloric acid H2S(aq) hydrosulfuric acid
Ternary acidsORIGINAL RADICAL SUFFIX
EXAMPLE ACID SUFFIX
NAME
per-root-ate
HClO4(aq) Per-root-ic Perchloric acid
-ate HClO3(aq) -ic Chloric acid
-ite HClO2(aq) -ous Chlorous acid
Hypo-root-ite
HClO (aq) Hypo-root-ous
Hypochlorous acid
BasesIn general, the presence of the hydroxide ion andthe term aqueous are sufficient to denote a base.
ex. NaOH(aq) sodium hydroxide (present in lye, Drano, liquid plumber)
Bi(OH)3(aq) bismuth hydroxide
(present in pepto bismol)
Mg(OH) 2(aq) + Al(OH) 3(aq) magnesium hydroxide and aluminum hydroxide
(present in mylanta)
Crossing over processEx: magnesium phosphate We know that Mg has a +2 charge, and PO4 has a -3
charge. These two numbers do not add up to zero. Thus, we find a least common denominator andfind out what we must multiply each number by toget this result. Out LCD is 6, thus we multiply +2 by3 and -3 by 2. This results in +6 and -6 cancellingout to zero.
Mg3(PO4) 2
lead (II) nitrite = Pb(NO2) 2
A mole represents 6.02 x 1023 particles of a substance
Type of matter Type of particle
Elements atoms
ionic compounds formula units
molecular compounds
molecules
Amadeo Avogadro di Quarengo (Italian 1776-1856) is credited with discovering this number; therefore it is called Avogadro's number
Examples:1 mole NaCl = 6.02 x 1023 formula units1 mole N2O4 = 6.02 x 1023 molecules1 mole Fe = 6.02 x 1023 atoms1 mole Ca2+ = 6.02 x 1023 ions
(A mole of molecules contains more than a mole of atoms)Calculation example: H2O1 mole molecules = 6.02 x 1023 molecules
1 molecule = 3 atoms6.02 x 1023
molecules x 3 atoms = 1.80 x 1024
1 mol H2O molecule atoms/mole H2O
The mass of a mole varies with the substance 1 mole C = 6.02 x 1023 atoms of C = 12 g C
Formula mass: the sum of the average atomic masses of all atoms represented in its formula (in a.m.u.)CO2 (12.01 amu + 2(16.00 amu) = 44.01 amu
Molar mass: the mass of 1 mole of any element or compound (in grams/mole) [numerically equal to formula mass]For an element the molar mass is the atomic mass found on the periodic table.
H2O 2(1) + 16 = 18 g/molCO2 12 + 2(16) = 44 g/mol
Percent composition: % by mass of each element in a compoundFor H2O
%H = 2 g x 100 = 11.11% 18 g
%O = 16 g x 100 = 88.89% 18 gFor CO2
%C = 12 g x 100 = 27.27% 44 g
%O = 32 g x 100 = 73.73% 44 g
Empirical formula/simplest formula gives the lowest whole-number ratio of the elements in a compound (or the lowest whole-number ratio of moles of atoms in a compound) ex. Empirical formulas Molecular formulas
CH C2H2 acetyleneC6H6 benzene
CH2O CH2O
formaldehydeC2H4O2 acetic acidC6H12O6 glucose
Determining the empirical/simplest formulaUsing the % composition, determine the mass of each of the elements in 100 g of that compoundConvert the masses in grams to molesDetermine the simplest whole number mole ratio between elements and set that equal to the atoms ratio in the simplest formula
Example: What is the empirical formula of acompound containing 25.9% N and 74.1 % O?
25.9 g N x 1 mole N = 1.85 mole N ÷ 1.85 = 1 x 2 = 2 14 g N
74.1 g O x 1 mole O = 4.63 mole O÷1.85 = 2.5 x 2
= 5 16 g O
So, the simplest formula is N2O5
Alternate methodIf you are not provided with the percent
composition of the elements in the compound, you may follow this method:
1. Divide the grams provided of each element by its molar mass (atomic mass)2. Divide all mole values by the smallest value to obtain the smallest whole number ratio possible.3. Set those values equal to the subscripts in the chemical formula.
To calculate the molecular formula from simplest formula you need the molecular weight of the compound and the simplest formula.
Ex. If the simplest formula for acetic acid is CH2O and the molecular mass is 60, what is the molecular formula? CH2O Simplest formula mass = 30 g/mole
30 X = 60 X = 2
so multiply all subscripts in the empirical formula by X to get C2H4O2
A chemical equation allows us to describe in a concise manner, on paper, a chemical reaction that has taken place. It represents, with symbols and formulas, the identities and relative molecular or molar amounts of the reactants and products in a chemical reaction.Evidence for a chemical reaction:Energy release as heat or light. Color changeEvolution of gas (bubbles and/or odor)
Appearance of a solid (precipitate)
Parts of a chemical equation: The items on the left of the arrow are
called reactants; (arrow) means "yields; the right products
States of matter are described: s (↓), l, g (↑), aq
Other symbols used in chemical equations are on p. 266, Table 2
Coefficients represent relative numbers of particles that take part in the reaction
#atoms are conserved mass is conserved
Equations must be balanced; the number of atoms on both sides of the arrow must be equal. Basics for balancing:1. Start from left to right.2. Balance polyatomic ions as a single
entity whenever possible.3. Balance the H's and O's last, and 4. NEVER manipulate subscripts in a
formula.5. Try and keep the coefficients to the
smallest whole numbers possible (fractional coefficients are acceptable)
Types of chemical reactions
Combination/synthesis reactions: 2 or more substances react to form a single substance.Reactants are often elements and/or simple compounds, often H2OProducts are compoundsDifficult to guess product of nonmetals (must be told)Often liberates energyGeneral format: A + X AX
Examples: 8Ca (s) + S8 (s) 8CaS (s)
2Mg (s) + O2 (g) 2MgO (s) 2Fe (s) + O2 (g) 2FeO (s)
Na2O (s) + H2O (l) 2NaOH (aq)
Decomposition/analysis reactions: one compound broken down into two or more simpler productsProducts are often elements and/or compounds in any combination; difficult to predictBinary compounds break down into their elementsEnergy is required for the reaction to take place General format: AX A + X
Decomposition/analysis reactions Examples:a. metallic carbonates metallic oxides + CO2 (g)
CaCO3 (s) CaO (s) + CO2 (g) (NH4) 2CO3(s)2NH3(g)+H2O(g) +CO2(g)
b. metallic hydroxides metallic oxides + H2O
Ca (OH) 2 (s) CaO (s) + H2O (g)NaOH and KOH are exceptions
c. metallic chlorates metallic chlorides + O2
2KClO3 (s) 2KCl (s) + 3O2 (g)d. some acids nonmetallic oxides + H2O H2CO3 (aq) H2O(l) + CO2 (g)
H2SO3 (aq) H2O(l) + SO2 (g)e. some oxides decompose upon heating
2HgO (s) 2Hg (l) + O2 (g) 2PbO2 (s) 2PbO (s) + O2 (g)
f. electric current - electrolysis2H2O(l) 2H2 (g) + O2 (g)2NaCl (s) 2Na (s) + Cl2 (g) 2HI (g) H2 (g) + I2 (g)
Single replacement reactions: atoms of one element replace atoms of a 2nd similar element in a compounddetermined by relative reactivities of the 2 metals (Activity Series is a list of elements arranged according to the ease with which the elements undergo certain chemical reactions.)Halogens are nonmetals that are replaced (activity decreases going down Group VIIA) General format:
A + BX AX + B Y + BX BY + X
Displacement of Hydrogen in water or acid by a metal
Double replacement reactions: exchange of positive ions between 2 compoundsGenerally are reactions between ionic compounds in aqueous solutions.
General format: AX + BY AY + BX
At least one statement below is usually true of one of the productsOnly slightly soluble and ppt. from solutionis a gas and bubbles out of solutionis a molecular compound such as H2O
Examples:2NaOH (aq) + H2SO4 (aq) Na2SO4 (aq) + 2H2O (l) (neutralization rxn)
AgNO3 (aq) + NaCl (aq) AgCl (s) + NaNO3 (aq) (precipitation rxn)
ionic reactions:Ag+ (aq) + NO3
- (aq) + Na+ (aq) + Cl-(aq) AgCl(s) + Na+ (aq) + NO3
-(aq)net ionic: Ag+ (aq) + Cl-(aq) AgCl(s) Net ionic reactions DO NOT include spectator ions!
Combustion reactions: oxygen reacts with another substance often producing energy in the form of heat and/or lightCommonly involve hydrocarbonsComplete combustion produces water and carbon dioxideIncomplete combustion produces CO and C in addition to CO2 and H2O due to decreased amounts of O2 Examples:
CH4 (g) + 2O2 (g) CO2 (g) + 2H2O (g) 2C6H6 (l) + 15O2 (g) 12CO2 (g) + 6H2O(g)
Sub-category – Redox reactions
Oxidation-reduction reactions, typically called “redox rxns” involve two simultaneous processes
Oxidation – loss of electronsReduction – gain of electrons
Mnemonic device to remember this: OILRIG – oxidation is loss, reduction is gain
Identifying Redox rxns
0 +2 -2 0 +1 -2
H2 (g) + CuO (s) Cu (s) + H2O (l)
Hydrogen is oxidized; copper is reduced
DisproportionationThis occurs when one substance is both oxidized and reduced during the same chemical reaction.
+1 -1 +1 -2 0
2H2O2 (aq) 2H2O (l) + O2 (g)Oxygen is both oxidized and reduced in this reaction.