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.