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((AA)) NNaammeess ffoorr tthhee bbiinnaarryy iioonniicc ccoommppoouunnddss formed from atoms of two elements are assigned by the following rule: Type I: The full name of the metallic element is given first , followed by a separate word containing the stem of the non-metallic element name and the suffix –ide. For example, NaF is called sodium fluoride. Type II : When naming compounds that contain metals with variable ionic charge behaviour, the charge on the metal ion must be incorporated in the name using Roman numerals. For example: FeCl2 is called
iron (II) chloride, and FeCl3 as iron (III) chloride. Likewise, CuOis named copper (II) oxide. This is called the Stock System of
naming ionic compounds. If it is uncertain about the charge on the metal ion in an ionic compound, use the charge on the nonmetal ion, which does not vary, to calculate it. For example, in CuO, we know the oxide ion has a –2 charge because oxygen is in Group VIA. This means that the copper ion must have a + 2 charge. All Group IA & Group IIA metals, and Ag (Grp. IB ), Zn, and Cd (Grp. IIB ), Al, Ga, & In (Grp. IIIB ) form only one type of ion, while transitional metals (Grp IIIA, Col 3 –Grp IIB, Col 12) form more than one.
Some common Cations and Anions
In the common names, “- ous” denotes lower valence, “- ic” denotes higher valence for multivalent ions. * In compounds of manganese, copper, and lead, the valence of the metal is often not revealed in the name of the compound. In such a
case, use the most common valence of that element.
Metal Cation IUPAC name Common name Nonmetal Anion Col 6 VIA Chromium
Cr
Cr
2
3
+
+
Chromium(II) Chromium(III)
Chromous Chromic
Col 1 IA Hydrogen
Hydride: H − a
relatively strong base
Col 7 VIIA Manganese
Mn
Mn
2
3
+
+ Manganese(II)*
Manganese(III) Manganous Manganic
Col 8 VIIIA Iron
Fe
Fe
2
3
+
+
Iron(II) Iron(III)
Ferrous Ferric
Col 9 VIIIA Cobalt
Co
Co
2
3
+
+
Cobalt(II) Cobalt(III)
Cobaltous Cobaltic
Col 10 VIIIA Nickel
Nickel (II) Nickel (III)
Col 11 IB Copper
Cu
Cu
+
+2
Copper(I) Copper(II) *
Cuprous Cupric
Col 11 IB Silver Ag
No more than 1 type of ion
Col 14 IVB Carbon
Carbide
C4− Col 11 IB Gold
Au
Au
+
+2 Gold(I)
Gold(II) Aurous Auric
Col 15 VB Nitrogen
Nitride
N3−
Col 12 IIB Zinc Zn
No more than 1 type of ion
Col 15 VB Phosphorus
phosphide
P3− Col 12 IIB Cadmium Cd
No more than 1 type of ion
Col 16 VIB Oxygen
Oxide
O2− Col 12 IIB Mercury
Hg Hg
Hg
22
2
+ +
+
( ) Mercury(I) Mercury(II)
Mercurous Mercuric
Col 16 VIB Sulfur
Sulfide
S2−
Col 13 IIIB Aluminum Al
No more than 1 type of ion
Col 16 VIB Selenium
Selenide
Se2− Col 13 IIIB Gallium Ga
No more than 1 type of ion
Col 17 VIIB Fluorine
Fluoride
F − Col 13 IIIB Indium In
No more than 1 type of ion
Col 17 VIIB Chlorine
Chloride
Cl− Col 14 IVB Tin
Sn
Sn
2
4
+
+
Tin(II) Tin(IV)
Stannous Stannic
Col 17 VIIB Bromine
Bromide
Br −
Col 14 IVB Lead
Pb
Pb
2
4
+
+ Lead(II) *
Lead(IV) Plumbous Plumbic
Col 17 VIIB Iodine
Iodide
I −
Col 15 VB Antimony
Sb
Sb
3
5
+
+
Antimony(III) Antimony(IV)
Stibnous Stibnic
2
((BB)) NNaammeess ffoorr tthhee bbiinnaarryy ccoovvaalleenntt ccoommppoouunnddss ((aallssoo ccaall lleedd mmoolleeccuullaarr ccoommppoouunnddss –– TTyyppee II II II )) are assigned by the following rule: The full name of the nonmetal with the lower electronegativity is given first, followed by a separate word containing the stem of the name of the more electronegative nonmetal and the suffix “-ide”. Greek numerical prefixes precede the names of both nonmetals.
Number 1 2 3 4 5 6 7 8 9 10 Greek Prefix mono- di- tri- tetra- penta- hexa- pepta- octa- ennea- deca-
Examples: (i) N O2 5 - dinitrogen pentoxide;
(ii) S Cl2 2 - disulfur dichloride;
(iii) CO2 - carbon dioxide (it is common to omit the initial prefix of mono-);
(iv) P S4 10- tetraphosphorus decasulfide.
� One exception to the use of Greek prefixes is when naming covalent compounds: When hydrogen is listed as the first element in the formula, no prefix is used, for examples: (i)H S2 - hydrogen sulfide,
and (ii)HCl - hydrogen chloride. � Another exception is the Peroxides:
Hydrogen oxide - , Hydrogen peroxide - ; Sodium oxide - , Sodium peroxide - ; Barium oxide – , Barium peroxide - ; Lead oxide – , Lead peroxide - .
� A third exception: Carbonyl chloride is an exception to the naming rule of –ide for binary compounds.
Some Common Covalent Molecules: No. Molecule Common Name Modern Name No Molecule Common Name Modern Name 1 H O2 Water Dihydrogen oxide 14 (i) NO Nitric oxide Nitrogen oxide
2 H O2 2 Hydrogen peroxide Dihydrogen dioxide 15 (ii) 2NO 4222 ONNO ↔ Nitrogen dioxide
3 SH 2 Sulfane; Hydrosulfide Hydrogen sulfide 16 (iii) ON2 Nitrous oxide Dinitrogen oxide
4 HBr Bromane Hydrogen bromide 17 (iv) N O2 3 Nitrous anhydride Dinitrogen trioxide
5 CH4 Methane Carbon tetrahydride 18 (v) 42ON 242 2NOON ↔ Dinitrogen tetroxide
6 NH3 Ammonia Nitrogen trihydride 19 (vi) N O2 5 Nitric anhydride Dinitrogen pentoxide
7 N H2 4 Hydrazine Dinitrogen tetrahydride 20 2SO Sulfurous anhydride Sulfur dioxide
8 3PH Phosphine Phosphorus hydride 21
104OP Phosphoric anhydride Tetraphosphorus decoxide
(Phosphorus pentoxide)
9 AsH3 Arsine Arsenic trihydride 22 2BeCl Beryllium chloride
10 XeF4 Xenon tetrafluoride 23 BrCl3 Bromine trichloride
11 SF4 Sulfur tetrafluoride 24 2COCl Phosgene Carbonyl chloride 12 S N4 4 Tetrasulfur tetranitride 25 BN Boron nitride
13 PBr5 Phosphorus
pentabromide
Note: Carbonyl chloride is an exception to the naming rule of –ide for binary compounds.
A polyatomic ion is a group of covalently bonded atoms that have acquired a charge through the loss or gain of electrons. Before this, the ions that are each formed from 1 atom are called monoatomic ions. Polyatomic ions are also called radicals. Numerous ionic compounds exist in which the positive or negative ion is polyatomic. Compounds that contain polyatomic ions offer an interesting combination of both ionic and covalent bonding; covalent bonding occurs within the polyatomic ion and ionic bonding occurs between it and the other ion. Polyatomic ions are very stable species that generally maintain their identity during chemical reactions. It would be a great advantage to memorize the following short list of some common polyatomic ions.
1. Most of the ions have a negative charge, which can vary from –1 to –3. Only two positive ions are listed in the table:
(i) NH4+ (Ammonium), and (ii)H O3
+ ( Hydronium).
2. Two of the polyatomic ions, OH− (hydroxide) and CN− (cyanide), have names ending in –ide. These names are the exceptions to the rule that the suffix –ide is reserved for use in naming binary compounds.
3. A number of –ate and –ite pairs of ions exist. The –ate ion will always have one more oxygen atom than the ite ion. Both ions carry the same charge.
4. A number of pairs of ions exist where one member of the pair differs from the other by having a hydrogen atom present, as in CO32−
(carbonate) and HCO3− (hydrogen carbonate); SO4
2− (sulfate) and HSO4
− (Hydrogen sulfate);
−34PO
(Phosphate) and HPO42− (
Hydrogen phosphate). In such pairs, the charge on the ion containing hydrogen is always one less than on the other ion.
3
Formulas and Names of Some Common Polyatomic Ions Key element
present Formula Name of ion Key element
present Formula Name of ion
1 Carbon CO2− Hypocarbonite 38 (Grp VIB) SO22−
Hyposulfite
2 (Grp IVB) CO22−
Carbonite 39 SO32−
Sulfite
3 CO32−
Carbonate 40 SO42−
Sulfate
4 CO42−
Percarbonate 41 SO52−
Persulfate
5 HCO− Hydrogen hypocarbonite 42 S O2 32−
Thiosulfate
6 HCO2−
Hydrogen carbonite 43 −252OS
Disulfite
7 HCO3−
Bicarbonate / hydrogen carbonate 44 −272OS Pyrosulfate
8 HCO4− Hydrogen percarbonate 45 HSO2
− Hydrogen hyposulfite
9 Hydrogen oxalate (bioxalate) 46 HSO3− Bisulfite or Hydrogen sulfite
10 C O2 42−
Oxalate (ethanedioate-an ester) 47 HSO4− Bisulfate or Hydrogen sulfate
11 Formate 48 HSO5− Hydrogen persulfate or Bisulfate
12 Acetate ( )CH COO3− 49 Fluorine FO− Hypofluorite
13 Stearate ( 50 (GrpVIIB) FO2− Fluorite
14 −485 NOHC Glutamate 51 FO3
− Fluorate
15 CN− Cyanide 52 FO4− Perfluorate
16 CNO− Cyanate 53 Chlorine ClO− Hypochlorite
17 SCN− Thiocyanate 54 (GrpVIIB) ClO2−
Chlorite
18 + or - thiocyanatoiron (III) 55 ClO3− Chlorate
19 −23SiO Silicate 56 ClO4
− Perchlorate
20 Nitrogen NO− Hyponitrite 57 Bromine BrO− Hypobromite
21 (Grp VB) NO2− Nitrite 58 (GrpVIIB) BrO2
− Bromite
22 NO3−
Nitrate 59 BrO3− Bromate
23 NO4−
Pernitrate 60 BrO4−
Perbromate
24 NH4+
Ammonium 61 Iodine IO− Hypoiodite
25 Phosphorus Hypophosphite 62 (GrpVIIB) IO2−
Iodite
26 (Grp VB) Phosphite 63 IO3− Iodate
27 −34PO Phosphate 64 IO4
− Periodate
28 −35PO Perphosphate 65 Metals +OH3 Hydronium
29 HPO22− Hydrogen hypophosphite 66 Hydrogen OH− Hydroxide
30 HPO32− Hydrogen phosphite 67 (Grp VIA) CrO4
2− Chromate
31 HPO42− Hydrogen phosphate 68 Cr O2 7
2− Dichromate
32 HPO52− Hydrogen perphosphate 69 (Grp VIIA) MnO4
2− Manganate
33 H PO2 2− Dihydrogen hypophosphite 70 MnO4
− Permanganate
34 H PO2 3− Dihydrogen phosphite 71 (Grp IIB) Zn OH( )4
2− Zincate
35 H PO2 4− Dihydrogen phosphate 72 (Grp IIIB) Al OH( )4
− Aluminate
36 H PO2 5−
Dihydrogen perphosphate 73 (Grp IVB) −23SnO
Stannate (IV)
37 Sulfur HS− Bisulfide or Hydrogen sulfide 74 (Grp VB) Arsenate
4
ions ions ions ions Ammonium, +OH3 Hydronium
Mercury (I)
FO−
Hypofluorite
FO2−
Fluorite
FO3−
Fluorate
FO4−
Perfluorate
4-d. Thiocyanato iron (III)
ClO− Hypochlorite
ClO2−
Chlorite
ClO3−
Chlorate
ClO4−
Perchlorate
BrO− Hypobromite
BrO2−
Bromite
BrO3−
Bromate
BrO4−
Perbromate
IO− Hypoiodite
IO2−
Iodite
IO3−
Iodate
IO4−
Periodate
NO− Hyponitrite
NO2−
Nitrite
NO3−
Nitrate
NO4−
Pernitrate
HCO− Hydrogen hypocarbonite
HCO2−
Hydrogen carbonite
HCO3−
Bicarbonate / hydrogen carbonate
HCO4−
Hydrogen percarbonate
−42OHC Hydrogen oxalate (bioxalate)
Hypocarbonite
Carbonite
Percarbonate | −42OC Oxalate
� Semiconductor material
HSO2−
Hydrogen hyposulfite
HSO3−
Bisulfite or Hydrogen sulfite
HSO4−
Bisulfate |* Permanganat
(per- naming exception )
HSO5−
Hydrogen persulfate or Bisulfate
SO22−
Hyposulfite
SO32−
Sulfite
SO42−
Sulfate |* Manganate
|* Chromate (yellow)
SO52−
Persulfate
------------------------------------------------------------------------------------------------------------------------------------------------------
Thiosulfate
Disulfite
Disulfate | Dichromate (Orange)
� (MnO4
−) permanganate(VII)
ion is a strong oxidizing agent � In aqueous solution, chromate
and dichromate anions exist in a
chemical equilibrium.
2 CrO42−
+ 2 H+ Cr2O7
2− + H2O
�Note: The disulfite ion is a dimer
of the bisulfite ion (HSO3−):
2 HSO3− (aq) S2O5
2− (aq)
+ H2O (l) H PO2 2
− Dihydrogen hypophosphite
H PO2 3− Dihydrogen phosphite
H PO2 4− Dihydroge phnosphate
H PO2 5− Dihydrogen perphosphate
HPO22− Hydrogen hypophosphite
HPO32− Hydrogen phosphite
HPO42−
Hydrogen phosphate
HPO52−
Hydrogen perphosphate
Hypophosphite
Phosphite
Phosphate|* Arsenate Perphosphate
---------------------------------------- Tripolyphosphate
Note:
Special naming of some
compounds � (for G11/12)
1-a. Hydroxide, 1-b.
Aluminate
----------------------------------------------------------------------------------------------------------------------------------------------------------------------------------------------------------------------------------------------------
2-a. Cyanide 2-b. Cyanate 2-c. Thiocyanate ----------------------------------------------------------------------------------------------------------------------------------------------------------------------------------------------------------------------------------------------------
3. Bisulfide /Hydrogen sulphide --------------------------------------------------------- 5-a. Formate 5-b.
Acetate ( )CH COO3
−
5-c. Stearate ( -------------------------------------------------------------------------------------------------------------------------------------------------------------------
--
6.
Glutamate ( - MSG)
.................................................................... 1-c. Zincate -------------------------------------------------------------------------------------------------------------------------------------------------------------------------------------------------------------------------------
= = = = = = = = = = = = = = = = = = = = = = = = = = = = = = = = = = = = = = = = = = = = 2-d. Thiocyanato iron (III) -------------------------------------------------------------------------------------------------------------------------------------------------------------------------------------------------------------------------------
4. Peroxide
................................................................................................................................................................................................................................................................................................................
� Note the exception: 2+ ion
5
Nomenclature - There are two "rules" that can be used for learning the nomenclature of polyatomic ions. First, when the prefix bi- is
added to a name, a hydrogen is added to the ion's formula and its negative charge is decreased by 1, the latter being a consequence
of the hydrogen ion carrying a +1 charge. An alternate to the bi- prefix is to use the word hydrogen in its place: the anion derived
from H+ + CO3
2−, HCO3
− can be called either bicarbonate or hydrogen carbonate.
Note that many of the common polyatomic anions are conjugate bases of acids derived from the oxides of non-metallic elements.
For example the sulfate anion, SO42−
, is derived from H2SO4 which can be regarded as SO3 + H2O.
The second rule looks at the number of oxygens in an ion. Consider the chlorine oxoanion family:
oxidation state −1 +1 +3 +5 +7
anion name chloride hypochlorite chlorite chlorate perchlorate
formula Cl− ClO
− ClO2
− ClO3
− ClO4
−
structure
First, think of the -ate ion as being the "base" name, in which case the addition of a per- prefix adds an oxygen. Changing the -ate
suffix to -ite will reduce the oxygens by one, and keeping the suffix -ite and adding the prefix hypo- reduces the number of oxygens
by two. In all situations, the charge is not affected. The naming pattern follows within many different oxyanion series based on a
standard root for that particular series. The -ite has one less oxygen than the -ate, but different -ate anions might have different
numbers of oxygen atoms.
These rules will not work with all polyatomic ions, but they do work with the most common ones (chlorate, nitrate, carbonate,
sulfate, phosphate,).
AAnniioonn __________ iiddee ((ee..gg.. cchhlloorriiddee,, )) __________ iittee ((ee..gg.. cchhlloorriittee,, )) __________ aattee ((ee..gg.. cchhlloorraattee,, ))
CCoovvaalleenntt bboonndd wwiitthh ccaattiioonn ooff HH ee..gg.. HHyyddrrooggeenn CChhlloorriiddee ee..gg.. HHyyddrrooggeenn CChhlloorriittee ee..gg.. HHyyddrrooggeenn CChhlloorraattee
CCoovvaalleenntt bboonndd wwiitthh ccaattiioonn ooff HH
iinn aaqq.. ssoolluuttiioonn
HHyyddrroo________ iicc aacciidd
ee..gg.. HHyyddrroocchhlloorriicc AAcciidd((HHCCll))
__________ oouuss aacciidd
ee..gg.. CChhlloorroouuss AAcciidd ((
__________ iicc aacciidd
ee..gg.. CChhlloorriicc AAcciidd (( ))
IIoonniicc bboonndd wwiitthh ccaattiioonn ooff mmeettaall ee..gg.. SSooddiiuumm CChhlloorriiddee
(( aa ssaalltt)) ee..gg.. SSooddiiuumm CChhlloorriittee
(( aa ssaalltt))
ee..gg.. SSooddiiuumm CChhlloorraattee
(( ,, aa ssaalltt))
IUPAC nomenclature of inorganic chemistry - The IUPAC nomenclature of inorganic chemistry is a systematic method of naming
inorganic chemical compounds, as recommended by the International Union of Pure and Applied Chemistry (IUPAC). The rules are
commonly known as "The Red Book" There is also an IUPAC nomenclature of organic chemistry.
System - The names "caffeine" and "3,7-dihydro-1,3,7-trimethyl-1H-purine-2,6-dione" both refer to the same chemical. The
systematic name encodes the structure and composition of the caffeine molecule in some detail, and provides an unambiguous
reference to this compound, whereas the name "caffeine" just names it. These advantages make the systematic name far superior to
the common name when absolute clarity and precision are required. However, for the sake of brevity, even professional chemists
will use the non-systematic name almost all the time, because caffeine is a well-known common chemical with a unique structure.
Similarly, H2O is most often simply called water in English, though other chemical names do exist.
1. Single atom anions are named with an -ide suffix: for example, H− is hydride.
2. Compounds with a positive ion (cation), the name of the compound is simply the cation's name (usually the same as the
element's), followed by the anion. For example, NaCl is sodium chloride, and CaF2 is calcium fluoride.
3. Cations able to take on more than one positive charge are labeled with Roman numerals in parentheses. For example, Cu+ is
copper(I), Cu2+
is copper(II). An older, deprecated notation is to append -ous or -ic to the root of the Latin name to name
ions with a lesser or greater charge. Under this naming convention, Cu+ is cuprous and Cu
2+ is cupric. For naming metal
complexes see the page on complex (chemistry).
4. Oxyanions (polyatomic anions containing oxygen) are named with -ite or -ate, for a lesser or greater quantity of oxygen. For
example, NO2− is nitrite, while NO3
− is nitrate. If four oxyanions are possible, the prefixes hypo- and per- are used:
hypochlorite is ClO−, perchlorate is ClO4
−,
5. The prefix bi- is a deprecated way of indicating the presence of a single hydrogen ion, as in "sodium bicarbonate" (NaHCO3).
The modern method specifically names the hydrogen atom. Thus, NaHCO3 would be pronounced sodium hydrogen
carbonate.
Positively charged ions are called cations and negatively charged ions are called anions. The cation is always named first. Cations can be
metals or polyatomic ions. Therefore the name of the metal or positive polyatomic ion is followed by the name of the non-metal or negative
6
polyatomic ion. The cation retains its element or polyatomic name, so does the polyatomic anion, whereas for a single non-metal anion the
ending is changed to -ide. Example: sodium chloride, potassium oxide, ammonium nitrate, or calcium carbonate.
When the metal has more than one possible ionic charge or oxidation number the name becomes ambiguous. In these cases the
oxidation number (the same as the charge) of the metal ion is represented by a Roman numeral in parentheses immediately
following the metal ion name. For example in uranium (VI) fluoride the oxidation number of uranium is 6. Another example is the
iron oxides. FeO is iron(II) oxide and Fe2O3 is iron(III) oxide.
An older system used prefixes and suffixes to indicate the oxidation number, according to the following scheme:
Oxidation state Cations and acids Anions
Lowest hypo- -ous hypo- -ite
-ous -ite
-ic -ate
per- -ic per- -ate
Highest hyper- -ic hyper- -ate
Thus the four oxyacids of chlorine are called hypochlorous acid (HOCl), chlorous acid (HOClO), chloric acid (HOClO2) and perchloric
acid (HOClO3), and their respective conjugate bases are the hypochlorite, chlorite, chlorate and perchlorate ions. This system has
partially fallen out of use, but survives in the common names of many chemical compounds: the modern literature contains few
references to "ferric chloride" (instead calling it "iron(III) chloride"), but names like "potassium permanganate" (instead of
"potassium manganate(VII)") and "sulfuric acid" abound.
Traditional Naming
(1) Naming simple ionic compounds [see also Names for Ionic Compounds on page 9 and Names for Bases on page 11] - An ionic
compound is named by its cation followed by its anion. See polyatomic ions for a list of possible ions. For cations that take on
multiple charges, the charge is written using Roman numerals in parentheses immediately following the element name) For example,
Cu(NO3)2 is copper(II) nitrate, because the charge of two nitrate ions (NO3-1
) is 2 × −1 = −2, and since the net charge of the ionic
compound must be zero, the Cu ion has a 2+ charge. This compound is therefore copper(II) nitrate.
The Roman numerals in fact show the oxidation number, but in simple ionic compounds (i.e., not metal complexes) this will always
equal the ionic charge on the metal.
(2) Naming hydrates [see also page 9]- Hydrates are ionic compounds that have absorbed water. They are named as the ionic
compound followed by a numerical prefix and -hydrate.
(3) Naming molecular compounds [see also Names for Acids on page 10] - Inorganic molecular compounds are named with a prefix
(see list above) before each element. The more electronegative element is written last and with an -ide suffix. For example, H2O
(water) can be called dihydrogen monoxide. Organic molecules do not follow this rule. In addition, the prefix mono- is not used with
the first element; for example, SO2 is sulfur dioxide, not "monosulfur dioxide". Sometimes prefixes are shortened when the ending
vowel of the prefix "conflicts" with a starting vowel in the compound. This makes the name easier to pronounce; for example, CO is
"carbon monoxide" (as opposed to "monooxide").
(3-a) Naming acids [see also Names for Acids on page 10] - Acids are named by the anion they form when dissolved in water. If an
acid forms an anion ending in ide, then its name is formed by adding the prefix hydro to the anion's name and replacing the ide with
ic. Finally the word acid is appended. For example, hydrochloric acid forms a chloride anion. With sulfur, however, the whole word is
kept instead of the root: i.e.: hydrosulfuric acid. Secondly, anions with an -ate suffix are formed when acids with an -ic suffix are dissolved,
e.g. chloric acid (HClO3) dissociates into chlorate anions to form salts such as sodium chlorate (NaClO3); anions with an -ite suffix are formed
when acids with an -ous suffix are dissolved in water, e.g. chlorous acid (HClO2) disassociates into chlorite anions to form salts such as
sodium chlorite (NaClO2).
In summary there are basically four choices of traditional naming systems for chemicals:
1. Stock System (cations))
(for compounds with
multivalent metals)
2. Prefix System:cations+anions
(the best for covalent
compounds)
3. Classical System (-ous/-ic)
(for compounds with multivalent
metals with only 2 valences)
4. Standard System (-ide)
(for binary compounds
with a univalent metal)
• English Names
• Roman Numerals tell
the valence of the
elements (mostly
metal) preceding
them, e.g.:
Gold (II) iodide;
Iron (III) oxide.
• English Names
• Greek prefixes &
suffixes e.g.:
Carbon dioxide;
Diarsenic pentoxide;
Copper monchloride
• Latin/Greek Names
• Ending (ous/ic) for
low/high valence, e.g.
Cuprous/cupric;
Stannous/stannic
Extension: hypo _ ous =>hypo _ ite
_ ous => _ ite
_ ic => _ rate
Per _ ic => per _ rate
• English Names
• -ide endings, e.g.:
Sodium oxide;
Calcium hydride;
Aluminum
carbide
7
IUPAC (International Union of Pure and Applied Chemistry) naming system adopts a combination of these system, but mostly the Stock and Prefix Systems.
8
((CC)) NNaammeess ffoorr iioonniicc ccoommppoouunnddss ccoonnttaaiinniinngg ppoollyyaattoommiicc iioonnss (Type IV-Tertiary Compounds) are derived in a similar manner as
naming for the binary ionic compounds: If the polyatomic ion is positive, its name is substituted for that of the metal. If the polyatomic ion is negative, its name is substituted for the “nonmetal stem plus –ide” in the binary compound.
Examples: (i) Ca PO3 4 2( ) is named as calcium phosphate; (ii) ( )NH CO4 2 3 is named as ammonium carbonate.
Two rules are to be followed in writing formulas containing polyatomic ions: (i) When more than one polyatomic ion of a given kind is required in a formula, the polyatomic ion is enclosed in parenthesis and a
subscript placed outside the parenthesis is used to indicate the number of polyatomic ions needed. Example: Mg NO( )3 2
(ii) To preserve the identity of polyatomic ions, the same elemental symbol may be used more than once in a formula. Example: NH CN4 instead of N H C2 4 for ammonium cyanide.
--------------------------------------------------------------------------------------------------------------------------------------------------------------------------- Ionic compound that absorb water into their solid structures form substances called hydrates. Hydrates typically have properties different
from anhydrous substances, which are water-free substances. For example, anhydrous copper (II) sulfate is nearly colourless. When it
absorbs water, it becomes bright blue. When it becomes fully hydrated, 5 water molecules are present for every copper ion:
CuSO H O4 25• . Its name is called copper (II) sulfate pentahydrate. MgSO H O4 27• is called magnesium sulphate heptahydrate.
Some exceptions to naming rules: (i) the prefix mono- is usually not written with the first word of a compound’s name: CO2 - carbon
dioxide, not monocarbon dioxide. (ii) Prefixes are sometimes shortened to make a name easier to say: CO - carbon monoxide, not carbon
monooxide; 42ON - dinitrogen tetroxide, not dinitrogen tetraoxide. (iii) Common names other than formal names are used for some
compounds: O2 - diatomic oxygen, molecular oxygen, or simply oxygen, rather than dioxygen; O3 -ozone; NH3 -ammonia; N H2 4 -
hydrazine; H O2 - water rather than dihydrogen monoxide; H O2 2 -hydrogen peroxide; PH3 -phosphine; AsH3 -arsine; CH4 -
methane; C H2 4 -ethane.
Classification of Compounds - Among the most important compounds are the acids (a molecular compound with covalent bond), and bases (an ionic compound). In a neutralization reaction, an acid reacts with a base to form an ionic compound called salt, which is any ionic
compound that does not contain OH− .
Acids & Bases – Arrhenius definition:
• An acid is a substance that releases hydrogen ions ( H + ) in aqueous solution.
• A base is a substance that releases hydroxide ions ( OH− ) in aqueous solution (this restricts bases to compounds containing (OH− only).
Acids & Bases – Bronsted-Lowry definition:
• An acid is any substance that can donate H + ions ⇒Lewis Theory: accept electron pairs from other atoms, ions, or molecules.
• A base is any substance that can accept H + ions. ⇒ Lewis Theory: donate electron pairs to other atoms, ions, or molecules.
Limitations of Arrhenius Definition: (1) It restricts acids and bases to water solutions while similar reactions occur in the gas phase and in solvents other than water. (2) It oversimplifies what happens when acid dissolves in water. In an acid such as HCl the hydrogen atom forms covalent bond with
the chlorine atom. The Arrhenius definition does not explain why this bond breaks to produce H + ion. (3) The Arrhenius definition does not include certain compounds that have the characteristics of a base such as 3NH .
The Bronsted-Lowry definition expands the Arrhenius definition in two important ways: (1) it defines acids and bases independently of how
they behave in water. (2) It focuses solely on H + ions (OH− ions are not part of the definition of a base). (3) It allows substance to be defined as acid or base in terms of chemical reaction. AAcciiddss - According to Bronsted-Lowry, an acid is defined as a molecular substance that dissolves in water to produce hydrogen ions
( )H + . When acids are in pure state (not in solution), they are covalent compounds; that is they do not contain H + ions. These ions are
formed through a chemical reaction when the acid is mixed with water. The reason is that H + is a proton, which is strongly attached to the electrons of surrounding water molecules forming the hydronium ions, H O3
+ . The chemical reaction between water and the acid molecules
therefore results in the breaking of the covalent bond in the acid molecules and setting free the H + ion. In this way, although acids are molecular compounds, in water they behave somewhat like ionic compounds. Acids, therefore are exceptions to the rule that molecular substances are non-electrolytes. A strong acid is one that ionizes essentially completely in aqueous solution. In a molecule of an acid, a
hydrogen atom that can be donated is called acidic hydrogen. The cation is always ( )H + . The anion depends on the particular acid. For
example: HCl yields H + cations and Cl− anions; HNO3yields H + cations and NO3− anions. An acid that can donate only one H + ion per
molecule is called a monoprotic acid. , two H + ions a diprotic acid; and three H + ions a triprotic acid. The concepts of electronegativity and the bond polarity can used to determine the relative strengths of acids based on their chemical formulas.
9
Acid Strength - The strength of the acid is inversely proportional to the bond strength between hydrogen and the rest of the molecule. Most acids are weak; only a few are strong.
Most acids fall into 3 categories based on their composition and structure:
(1) Binary Acids – Molecules have a hydrogen atom (but no oxygen) and another element usually from Group VIA or Group VIIA. Strongest binary acids: hydrochloric acid (HCl), HBr, HI. Weak binary acids: HF, H S2 , H Se2 .
The strength of binary acids increases 1. from left to right in a period (e.g. PH H S HCl3 2< < corresponding to increasing electronegativity), and 2. top to bottom in a group in the periodic table (e.g. HF HCl HBr< < Electronegativity cannot explain this trend; instead,
the size of the anion- decreasing bond strength is important).
(2) Oxy Acids or Oxoacids– These acids contain hydrogen, oxygen, and one other element. Sulfuric acid (H SO2 4 ) is called the
“King of Chemicals”, also called battery acid because it is used in the lead storage batteries of automobiles. Others: nitric acid ( HNO3) and phosphoric acid (H PO3 4), H SO2 4 , and HClO4 . Most other oxy acids are weak acids. The acidic hydrogens of oxy
acids are attached to oxygen atoms. Oxoacids in which the number of oxygen atoms exceeds, by two or more, the number of protons that can be dissociated: for example: , .
The strength of oxoacids with the same number of oxygens increases 1. From left to right in a period (as the number of oxygen atoms per hydrogen atom also increases), and 2. Bottom to top in a group in the periodic table (contrary to binary acids and is due to electronegativity).
Trends in the relative strength of oxoacids are explained by effects of electronegativity and bond polarity on the ease of donating a proton. (increasing electronegativity weakens the O-H bond, e.g.: telluric acid < selenic acid < sulfuric acid)
For a given nonmetal, the acid strength increases with increasing number of oxygens e.g.: hypochlorous acid < chlorous acid < chloric acid < perchloric acid.
A correct understanding of the effects of oxygen involves two considerations: a. Double-bonded oxygen atoms (=O) attract electrons more strongly than do oxygen atoms bonded to hydrogen atoms (-O-H),
and the effects of double-bonded oxygen atoms are additive. b. Double-bonded oxygen atoms provide a pi electronic structure to delocalize and stabilize the electron that remains behind
when H + ionizes. The more double-bonded oxygens present, the more stabilization of the anion results.
(3) Carboxylic Acids – These are organic acids, which include the carbon atom. Most of them are weak. Their name comes from a group of atoms called a carboxyl group( )−COOH , which donates the hydrogen atom to form a relatively stable anion called a carboxylate ion. One common carboxylic acid is acetic acid (HC H O2 3 2), which is found in vinegar.
Electronegative atoms such as fluorine, chlorine, bromine, iodine, oxygen, and sulfur on nearby carbon atoms will withdraw electron density from the O-H bond and increase the strength of the organic acid. The effects of these electronegative atoms are approximately additive:
H O H H O H H O H Cl O HH
H O
H
Cl O
Cl
Cl O
Cl
Cl O
C C C C C C C C− − − − < − − − − < − − − − < − − − −|
| ||
|
| ||
|
| ||
|
| ||
acetic acid chloroacetic acid dichloroacetic acid trichloroacetic acid weakest �----------------------------------------------------------------------� strongest
(4) Polyprotic Acids - Polyprotic acids are able to donate more than one proton per acid molecule, in contrast to monoprotic acids
that only donate one proton per molecule. All polyprotic acids are weak acids except sulfuric acid, which is unique in that its first proton dissociates completely but the second proton does not.
((DD)) NNaammeess ffoorr AAcciiddss && TThheeiirr SSaallttss (1) Binary Acids: The names of all binary acids start with the prefix hydro- and end with the suffix –ic, and the word acid is added. Salts of
these acids contain the negative ion of the nonmetal and always end in –ide. For example, when gaseous HCl (hydrogen chloride) is dissolved in water, it forms hydrochloric acid. Similarly, hydrogen cyanide (HCN) and dihydrogen sulphide ( ) dissolved in water are called hydrocyanic acid and hydrosulfuric acid, respectively.
In (a), the atom X is electropositive, and so extra electron density (blue areas) accumulates on the OH group. The OX − bond then breaks easily, making the compound a base. In (b), X is electronegative, and so electron density is drawn from the H atom to the OX − bond. Now it is the HO − bond that breaks easily, and the compound is an acid.
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(2.i) Acids with polyatomic anions ending in –ate: the –ate ending of the anion is changed to –ic and the word acid is added. Their salts end in –ate.
(2.ii) Acids with polyatomic anions ending in –ite: the –ite ending of the anion is changed to –ous and the word acid is added. Their salts end in –ite.
(2.iii) The halogens can occur in as many as four different oxoacids. The oxoacid with the most oxygens has the prefix per-, and the one with the least has the prefix hypo-. For example: HClO hypochlorous acid, HClO2 chlorous acid, HClO3 chloric acid,
HClO4 hyperchloric acid.
(3) Organic Acids: In the systemic name for organic acid, the suffix –oic and the word acid are added to the name of the molecule. (4) Polyprotic Acids and their Salts –Polyprotic acids are acids that can supply more than one hydrogen ion. Polyprotic acids may
be partially neutralized to give ions that still contain hydrogen. For example, H SO NaOH H O Na SO2 4 2 2 42 2+ → + , but
H SO2 4 can be partially neutralized to give the HSO4− ion: H SO NaOH H O NaHSO2 4 2 4+ → + These ions form salts called
acid salts. In naming them we specify the number of hydrogens that can still be neutralized if the salt were to be treated with
additional base: NaHSO4 sodium hydrogen sulfate, NaH PO2 4 sodium dihydrogen phosphate. For acid salts of diprotic acids,
the prefix bi- is still often used: NaHCO3 sodium bicarbonate. The prefix bi- does not mean “two”; it means that there is acidic
hydrogen in the compound. With polyprotic acids, the product of neutralization depends on the proportions of acid and base used in the reaction. For example: (i) H SO NaOH H O NaHSO2 4 2 42+ → + , (ii) H SO NaOH H O Na SO2 4 2 2 42 2+ → +
The salt NaHSO4 , which can be isolated as crystals by evaporating the reaction mixture, is referred to as an acid salt because it
contains the anionHSO4− , which is still capable of giving outH + . In fact, solution ofNaHSO4 , are acidic.
BBaasseess - According to Bronsted-Lowry a base is a compound that contains an unshared pair of electrons which allows it to accept
an H + ion from water formingH O3+ . For example, a molecule of ammonia (NH3 ) has an unshared pair of electrons on the nitrogen
atom. With this unshared pair, an ammonia molecule attracts and bonds withH+ . ((EE)) NNaammeess ffoorr BBaasseess (1) A hydroxide base is designated by using the name of the metal, with a roman numeral if necessary, and then the word hydroxide. (2) Nitrogen bases related to ammonia are amines. Replacing a hydrogen on ammonia with a methyl, −CH3 , group produces
methylamine. If two methyl groups replace two hydrogen atoms, the compound is called dimethylamine. The chloride salt of ammonia is called ammonium chloride. The name of the chloride salt of methylamine is methlammonium chloride, or alternatively, methylamine hydrochloride. Chloride salts of nitrogen bases with common names, such as hydrazine, usually use the hydrochloride ending, as in hydrazine hydrochloride.
Bases and their Strength: (1) Anions – Because of their negative charges and available electrons pairs, many anions function as bases. For examples,
−H (hydride ion), −2NH (amide ion),O OH2− − and are relatively strong bases. The strong acids have corresponding weak
conjugate bases. NaOH and Ca OH( )2 are strong bases. However, from a strict interpretation of the Bronsted-Lowry definition,
the bases are not the compounds themselves but the OH− ions that they yield in water.
(2) Metal Hydroxides - all metal hydroxides are strong bases; that is, the hydroxide ion dissociates completely when compound is dissolved in water. However, most metal hydroxides are also very slightly soluble. Only the hydroxides of Group IA metals, and barium, have appreciable solubility; strontium and calcium hydroxides are moderately soluble. Soluble hydroxides may cause severe skin burns. Group IIA elements form the strong hydroxides but these hydroxides are only slightly soluble in water. Insoluble hydroxides are much less harmful; e.g. Mg OH( )2 can safely be swallowed as an antacid to neutralize excess stomach acid.
Insoluble hydroxides readily dissolve in acids because the neutralization reaction forms water as one of the product and water is such a weak electrolyte that reactions in which it can be formed tend to be driven to completion even when one of the reactant is insoluble.
(3) Nitrogen Bases - These compounds related to ammonia (NH3 ) and are called amines. All amines contains a nitrogen atom that has
an unshared pair of electrons, which makes amines weakly basic. The relative strengths of the weak bases may be evaluated based on the electronegativities of the organic functional groups that replace the hydrogen atoms of ammonia. Electronegatve substituents such as chlorine increase the strength of organic acids; the reverse is true of organic bases: e.g. chloromethylamine is a weaker base than methylamine. Many drugs, such as the caffeine and cola drinks, are amines. The decomposition of dead organisms can produce a variety of foul-smelling amines, such as putrescine (H N CH NH2 2 4 2( ) ) and cadaverine (H N CH NH2 2 5 2( ) )
The nitrogen atom on pyridine features a basic lone pair of electrons. Because this lone pair is not delocalized into the aromatic pi-system, pyridine is basic with chemical properties similar to tertiary amines. The pKa of the conjugate acid is 5.21. Pyridine is
11
protonated by reaction with acids and forms a positively charged aromatic polyatomic ion called pyridinium. The weak bases are covalent compounds that are sometimes called molecular bases because they exist in acqueous solution primarily as nonionized molecules.
Metal hydroxides are examples of strong bases (strong electrolytes) but only slightly soluble in water. Ammonia is an example of a weak base (weak electrolyte) even though it is quite soluble in water.
Acids ����
(1) Binary Acids (H)
(2) Oxy Acids (H, O)
(3) Carboxylic Acids (Organic-COOH)
Bases ����
(1) Anion Bases
( OH− )
(2) Amine Bases ( NH3 : 1 H replaced
with a methyl −CH3)
1 HI hydroiodic acid (S)
HClO4 perchloric acid (S)
HCOOH methanoic acid (formic acid)
1 LiOH lithium hydroxide (S)
NH3
ammonia (W) 2 HBr
hydrobromic acid (S) HClO3 chloric acid (S)
CH COOH3 ( HC H O2 3 2)
ethanoic acid (acetic acid in vinegar) (W)
2 NaOH sodium hydroxide (S)
H NNH2 2
hydrazine (W)
3 HCl hydrochloric acid (S)
H SO2 4 sulfuric acid (S-d)
CH CH COOH3 2
propanoic acid
3 KOH potassium hydroxide (S)
CH NH3 2
Methylamine 4 HF
hydrofluoric acid (W) HNO3
nitric acid (S-m)
CH CH CH COOH3 2 2 butanoic acid (butyric acid)
4 RbOH Rubidium hydroxide (S)
H N CH NH2 2 4 2( )
putrescine 5 H S2
hydrosulfuric acid (W)
HClO hypochlorous acid (W)
CH CH CH CH COOH3 2 2 2 pentanoic acid (valeric acid)
5 CsOH cesium hydroxide (S)
H N CH NH2 2 5 2( )
cadaverine
6 H PO2 4 hypophosphoric acid (W)
C H COOH6 5 benzoic acid ( aromatic acid)
6 Ca OH( )2 calcium hydroxide (S)
7
H PO3 4 phosphoric acid (W-t)
oxalic acid (W-relatively S)
7 Sr OH( )2 strontium hydroxide (S)
8 HNO2
nitrous acid (W)
ClCH2CO2H. chloroacetic acid
8 Ba OH( )2
barium hydroxide (S)
9 33BOH boric acid (W-t)
HCO3−
hydrogen carbonate ion (W)
9 Mg OH( )2 magnesium hydroxide (S)
10 H SO2 3 sulfurous acid (W)
32COH carbonic acid (W-d)
10 PO43−
phosphate ion (W)
11 HCN hydrocyanic acid (W)
11
carbonate ion (W)
12 CaO calcium oxide (W)
Note that the formulas for all organic acids may be written in two forms, one with an initial hydrogen and the other with the organic functional group for acids, -COOH. The acetic acid in the table is an example. (S) – denotes strong; (W) – denotes weak.
Lewis Definitions of Acids and Bases: To completely generalize acid-base theory and to account for the formation of complex ions, G. N. Lewis proposed that acids are substances that accept electron pairs from other atoms, ions, or molecules, and bases are substances that donate electron pairs, in forming chemical bonds. The reaction between boron trichloride , 3BCl , an electron-deficient compound, and ammonia, 3NH ,
which has a non-bonding pair of electrons, is an acid-base reaction according to the Lewis definitions. The ammonia is an electron-pair donor and is a base, while the boron trichloride is the electron-air acceptor and is an acid. The bond that forms is a coordinate covalent bond, in which both electrons are supplied by a lone pair on the Lewis base.
Octet-deficient compounds involving elements of Group III such as boron and aluminum are often strong Lewis acids, because Group III atoms can achieve octet configurations by forming coordinate covalent bonds. Atoms and ions from Group V through VII have the necessary lone pairs to act as Lewis bases. Compounds of the main-group elements from the later periods can also act as Lewis acids through valence expansion. In such reaction, the central atom accepts a share in additional lone pairs beyond the 8-electrons needed to satisfy the octet rule. For example is a Lewis acid that accepts electrons from chloride ion lone pairs. After the reaction, such tin atom is surrounded by 12 rather than 8 valence electrons
The Lewis Theory is used mainly to explain the formation of substances called complexes., For example, silver chloride (AgCl) is an insoluble salt,
but it does dissolve in ammonia solutions in the following reaction: −+ +→+ ClNHAgNHAgCl s 2)(2 33)(
. The same reaction occurs, without the
−Cl ions, when silver ions in solution react with ammonia. This reaction is a Complexation Reaction. The silver ion starts as a silver atom with the
12
electron configuration [Kr]10145 ds . In forming the +Ag ion, it loses the 15s electron, resulting in [Kr] 104d electron configuration. Therefore it
has empty 5s and 5p orbitals that may accept pairs of electrons. In fact, silver ions accept only two pairs of electrons from two ammonia molecules. In a similar fashion, all metal ions have available orbitals that may accept electron pairs. Copper accepts four, the rest of the metal ions tend to accept six electron pairs in complexes. Metal ions are generally Lewis acids (Groups 1 & 2 metal cations do not act as acid except ). The Lewis acid-
base theory can also explain why small, highly positive ions such as form complex ions in water: . The water
molecules each possess nonbonding pairs of electrons and so act as Lewis bases, and the ion possesses empty 3s, 3p and 3d orbitals that may accommodate electron pairs. The electron configuration of the ion can be represented a : [Ne] . The complex
is formed when the ion, acting as a Lewis acid, bonds with six water molecules ( by relatively weak electrostatic forces of attraction), each acting as a Lewis base. The high-charge density of the ion increases the polarity of the –OH bonds in the molecules. That is, the molecules in are more likely to transfer a proton in the solvent ( ) than are molecules in pure water, and experiments show that only one of these six will. Thus, behaves as a weak monoprotic acid in aqueous solution in accordance with the equation shown in page 19. Comparison of Arrhenius, Bronsted-Lowry, and Lewis Definitions: The neutralization reaction between HCl and NaOH : HCl + NaOH NaClOH +→ 2 shows the progressive generality in these
definitions. By the Arrhenius definition, HCl is the acid and the NaOH is the base. By the Bronsted-Lowry definition, +OH 3is the
acid and −OH is the base. According to Lewis, +H is the acid and −OH is the base, since the proton accepts the lone pair donated by −OH in the reaction )(2)()( laqaq OHOHH →+ −+
Bronsted-Lowry acids and bases are acids and bases in the Lewis model, but the reverse is not always true.
Properties of Acids and Bases Properties Acids Bases 1. Taste Sour when dissolved in water Bitter taste when dissolved in water 2. Touch Sharp sting when in contact with injured skin Fell smooth, soothing, and slippery 3. Reactions with Metals
React vigorously with many metals, including magnesium, zinc, iron, and aluminum, to produce hydrogen gas
Do not react with most metals
4. Electrical Conductivity
Solutions of acids conduct electricity quite well; are examples of electrolytes because they form ions
Solutions of bases conduct electricity quite well; are examples of electrolytes because they form ions
5. Indicators Litmus paper turns from blue to red Litmus paper turns from red to blue 6. Neutralization Acid and base when mixed together produce salt Acid and base when mixed together produce salt
The H + ion is simply a proton; it is strongly attracted to the electrons of surrounding water molecules. Essentially, this interaction
forms an hydronium ion ( H O3+ ): H H O H O+ ++ →2 3 .
From the Bronsted-Lowry definition , whenever one compound in a reaction acts as an acid – or donates an H + ion – another compound acts as a
base – or receives an H + . For example, in the reaction:HCl g H O l H O aq Cl aq( ) ( ) ( ) ( )+ → ++ −2 3 , water acts as the base by accepting an H + ion
to become an H O3+ ion. In the reaction:NH g H O l NH aq OH aq3 2 4( ) ( ) ( ) ( )+ → ++ − , ammonia ( NH3 ) acts as the base, and water acts as the acid
by donating an H + ion to become an OH − ion . Water being able to acts an acid and as a base is described as amphoteric (amphiprotic). All
polyatomic ions whose chemical formula begins with H (e.g., ,)(3−
aqHCO −)(4 aqHSO ) are amphoteric.
When an acid loses an +H ion, it becomes its conjugate base, and when a base gains an +H ion, it becomes its conjugate acid:
HA aq H O l H O aq A aqacid base conjugateacid conjugatebase( ) ( ) ( ) ( )+ → ++ −
2 3 .
For example, in above reaction, Cl− is the conjugate base of HCl, and NH4+ is the conjugate acid of NH3 . The stronger the acid,
the weaker its conjugate base, and vise versa.
13
Let’s consider a generic weak acid HA, where H is a hydrogen atom and A is the rest of the acid molecule. Because HA is a weak
acid, it very partially ionizes in water: )()(3)(2)( aqaqlaq AOHOHHA −− +⇔+ . The equilibrium constant is . However, for
dilute solutions, the concentration of OH2 is essentially constant. So, the equation can be re-written as .
aK is called the acid ionization constant. [In some textbook the equilibrium constant is written as .
The acid ionization constant, Ka , is a measure of the strength of an acid. It is defined as KH O A
HAa =+ −
3. The greater theKa , the
stronger the acid.
The base dissociation constant, Kb , is a measure of the strength of a base. It is defined as , where H is the
hydrogen atom and B is the generic base in the following reaction: B aq H O l HB aq OH aq( ) ( ) ( ) ( )+ ⇔ ++ −2 . The greater the Kb , the
stronger the base. For strong bases like NaOH their dissociation in water is essentially complete. On the other hand, weak Bronsted bases
are usually molecules or ions that react with water to remove a proton from water and generate a hydroxide ion. For example:
and . In these cases, is or
Note: Polyprotic acids do not donate all of their protons simultaneously when they react. They always ionize in a stepwise fashion. In general, for a
polyprotic acid, ...221 aaa KKK >>> . Usually the first ionization constant is much larger than the subsequent Ka values, so only the first
ionization step is used in determining the pH of the solution.
14
Self – ionization of water: Water acts as both an acid and a base in the same reaction: H O H O H O OHl l aq aq2 2 3( ) ( ) ( ) ( )+ ← +→ + − .
As indicated by the long and short arrows, the self – ionization proceeds only minimally to the product side. In pure water at 25oC ,
both H O3+ ions and OH − ions are found at concentrations of 10 10 7. × − M , which is very small.
The equilibrium constant for the self-ionization of water can be written as follows: .
In pure water the molar concentration of water is calculated as follows: 1000 .
So that .
The product of H O3+ and OH− concentrations is always constant: K H O OHw = = × × = ×+ − − − −
37 7 1410 10 10 10 10 10( . )( . ) . . We
often write H O3+ as H + , so K H OHw = = ×+ − −10 10 14. . Kw is called the ion-product constant for water. It is important to
recognize that in any aqueous solution at 25oC , no matter what it contains, the product of H + and OH− must always equal to
Kw = × −10 10 14. so that if H + goes up, the OH− must go down.
In summary: 1. A neutral solution, where H OH+ −= ; pH = 7
2. A acidic solution, where H OH+ −> ; pH < 7
3. A basic solution, where H OH+ −< ; pH > 7
Because the H + in an aqueous solution is typically quite small, using a scale called the pH scale provides a convenient way to represent
solution’s acidity. The pH is defined as pH=-log H + . Similarly, pOH is defined as pOH=-log OH− . So, pH+ pOH = 14.00
---------------------------------------------------------------------------------------------------------------------------------------------------------------------------
for some metal ions at SATP Metal Ion
The aqueous metal ion is a hydrated complex ion (e.g. ). Aqueous ions of transition metals are usually written in a simplified form, without showing the number of water molecules present in the actual hydrated complex ion as shown in the table.
Groups 1 & 2 metal cations do not act as acid except .
15
Liquids
Atoms of molecules of gases and liquids are held together by covalent bonds. Liquids and solids are collectively referred to as the condensed states of matter because substances in these states have substantially higher densities. Energy requirement for the changes of state – the bonding forces that hold the atoms of a molecule together are called intramolecular forces. The forces that occur among molecules that cause them to aggregate to form a solid or a liquid are called intermolecular forces. The physical properties, such as viscosity and surface tension, of liquids are determined mainly by the nature and strength of the intermolecular forces. Only intermolecular forces are involved in changes of state. It takes energy to overcome the intermolecular forces in order for the changes of state (solid to liquid, and liquid to gas) to take place. The energy to melt 1 mol of a substance is called the molar heat of fusion. The energy to change 1 mol of liquid to its vapor is called the molar heat of vaporization.
Solutions A solution is a homogeneous mixture of two or more substances in a single physical state.
Type of Solutions Solute Solvent Example 1. Gaseous Solution Gas Gas Air (oxygen and nitrogen) 2. Liquid Solution Gas Liquid Seltzer (carbon dioxide in water) 3. Liquid Solution Liquid Liquid Antifreeze (ethylene glycol in water) 4. Liquid Solution Solid Liquid Ocean water (salt in water) 5. Aqueous Solution Liquid / Solid Water Salt (electrolyte); Sugar (non-electrolyte) 6. Solid Solution Gas Solid Charcoal filter (poisonous gases in carbon) 7. Solid Solution Liquid Solid Dental filling (mercury in silver) 8. Solid Solution (alloy) Solid Solid Sterling silver (copper in silver), an alloy
Some pairs of liquids can mix in any amount. These are said to be miscible in all proportions.
1. Solvent – the substance present in the largest amount; Solute – the other substance.
2. Aqueous solutions – solutions with water as the solvent.
3. Mass Percent = mass of solute
mass of solution×100%=
grams of solute
grams of solute + grams of solvent×100%
4. Concentration of a solution is the amount of solute in a given amount of solvent or solution. The most common measurement of concentration are the following:
Molarity (M) = moles of solute
liters of solution= mol
L
How a Solution Forms? Intermolecular forces also operate between solute and solvent particles in a solution as in pure substances. Sodium chloride dissolves in water because
the water molecules have a sufficient attraction for theNa+ and Cl− ions – enough to overcome the attraction of these two ions for one another in the crystal. Water molecules orient themselves on the surface of the NaCl crystal so that they can separate, dissociate, the ions and pull them into
solution. Once separated from their crystal, theNa+ and Cl− ions are surrounded by water molecules. The interaction between solute and solvent particles is called solvation. The interaction is called hydration when the solvent is water. The water molecules are also separated from one another to make room for the solute particles, so that the solute and solvent particles are intermingled. So the formation of solution of NaCl in water involves the breaking of attractions among solute particles, the breaking of attractions among solvent particles (endothermic process - energy absorbing), and the formation of attractions between solute and solvent particles (exothermic process – energy releasing). Whether energy, in the form of heat, is absorbed or given off in the overall process depends on the balance between the processes. So, the forming of sodium hydroxide solution is exothermic while that of ammonium nitrate (NH NO4 3 ) is endothermic. Instant cold pack used to reduce swelling cause by an injury is equivalent
to the melting a great deal of ice because when the pack is hit the breaking of attractions in forming this solution absorbs more energy than is released. The heat is absorbed from outside the pack thereby cooling the injured area. In an opposite process, a supersaturated solution ofNa S O2 2 3is used to
make instant heat pack. When the pack is squeezed, a crystal ofNa S O2 2 3is released from a small compartment in the pack. The crystal causes
excess solute to come out of solution. The process during which the crystal pulls particles out of solution is exothermic. The heat pack can be recycled by placing it in boiling water to re-dissolve theNa S O2 2 3.
Solubility - The solubility is the amount of a solute that will dissolve in a specific solvent under given conditions, which is the amount of solute required to form a saturated solution, and are usually expressed in grams of solute per 100 grams of solvent at a specified temperature and pressure.
Factors affecting solubility: 1) Nature of solute and solvent: Like Dissolve Like Rule for a solid in a liquid:
Solute Polar solvent Non-polar solvent polar soluble Insoluble nonpolar insoluble Soluble Ionic (similar to polar) soluble Insoluble
16
Some Polar and Nonpolar Substances
Polar Non-polar
Water ( OH 2 ) In general, greases, petroleum oils, vegetable oils, waxes, tars, gasoline Alcohols Methyl alcohol ( OHCH3 )
Ethyl alcohol ( OHHC 52 )
Isopropyl alcohol ( OHHC 73 )
Hexane ( 146HC )
Acetone ( OHC 63 ) Heptane ( 177HC )
Acetic acid ( 232 OHHC ) Octane ( 188HC )
Formic acid ( 2HCHO ) Carbon tetrachloride ( 4CCl ) Chloroform ( 3CHCl )
Some compounds contain both polar and nonpolar components, yet exhibit properties more like one component than the other. Cholesterol (C H O27 26 ) is such a compound. It is considered nonpolar, because it is insoluble in water and soluble in
non-polar solvent such as fat tissue.
2) Temperature – Solutions of gases in liquids are greatly affected by changes in temperature. As the temperature increases, the kinetic energy of the solute gas becomes greater. The gas particles acquire more of a tendency to escape from the solvent. Thus as the temperature increases, the solubility of a gas in a liquid decreases. The effect of temperature changes on the solubility of solids in liquids is very different from that for gases. Generally, the solubility of a solid solute increases as the temperature increases. The fact that solubility of a solid solute changes with temperature is the key to preparing supersaturated solutions. To prepare a supersaturated solution, the solution must be heated and then excess solute added. If the solution is then cooled slowly, the extra solute will stay in the solution. Shaking or disturbing a supersaturated solution or adding a tiny crystal of the solid solute can destroy the super-saturation and cause the excess solid solute to crystallize, leaving a saturated solution.
3) Pressure (for gases) – While the solubility of solids and liquids is not appreciably affected by pressure, the solubility of a gas in a liquid is strongly influenced by pressure. According to Henry’s Law , the solubility of a gas was proportional to the partial pressure of the gas above the liquid.
Factors affecting rate of a solid solute dissolving in a solution: 1) Surface area – grinding the solute into smaller particles thereby increasing the surface increases the rate of dissolving. 2) Stirring – similar effect as grinding solute into small particles. 3) Temperature – raising the temperature of a solvent increases the rate at which a solute dissolves, because as temperature
increases, solvent particles move faster. As solvent particles move faster, more particles come into contact with the solute.
Colligative Properties: Some physical properties of liquid solutions differ from those of the pure solvent. A property that depends on the concentration of solute particles but is independent of their nature is called colligative property.
1) Vapor pressure reduction – The extent to which a nonvolatile solute lowers the vapor pressure is proportional to its concentration.
2) Boiling point elevation – The difference is directly proportional to the number of solute particles per mole of solvent particles, that is, it is proportional to the molality of the solute: mKT bb =∆ , where bK is the molal boiling point elevation
constant. The value of bK depends on the solvent. This property can be used to determine molar mass:
solvent kg
solute mol=∆=b
bK
Tm → solvent kgsolute mol ×= m →
solute of massmolar
solute mass=solute mol , →
solute mol
solute mass=massmolar .
(Note: Any of the four colligative properties can be used to determine the molar mass of an unknown substance in this fashion.)
3) Freezing point depression – Like boiling point elevation, the decrease in the freezing point ( fT∆ )is directly proportional to
the to the molality of the solute: mKT ff =∆ .
4) Osmotic pressure – The process that allows a net flow of solvent molecules from the less concentrated solution to the more concentrated solution is called osmosis. The pressure required to prevent osmosis is known as the osmotic pressure (π ) of the solution. When two solutions with identical osmotic pressure are separated by a semi-permeable membrane, there is no osmosis, and the solutions are said to be isotonic. Fluids administered intravenously to people needing replacement of body fluids must be isotonic with body fluids because the membranes of red blood cells are semi-permeable. If one solution has a lower osmotic pressure than another, it is said to be hypotonic, conversely, the other solution is called hypertonic.
17
Many chemical reactions are carried out in aqueous solutions because particles of the solute are dispersed and can intermingle freely. When solutions of the solutes react, it is the ions themselves that participate in the reaction; that is why these reactions are called ionic reactions. Neutralization reaction of an acid and a base involves the reaction of the hydronium ions and the hydroxide ions in water.
In order for a solution to conduct an electric current, ions must be present. Pure acetic acid is made of molecules. It contains no ions. If we pass an
electric current into acetic acid, nothing will happen because no ions are present. An important change takes place, however, when acetic acid is
added to water. Water molecules have the ability to tear acetic acid molecules apart, breaking them down into hydrogen ions and acetate ions. Now
that ions are present, the water solution of acetic acid can conduct an electric current. This process is known as ionization because ions are produced
from a substance (acetic acid in this case) that did not contain them originally. Percent ionization values vary with the concentration of the acid. In
general, the more dilute a weak acid solution, the greater the percent ionization, and vice versa.
A similar story about the conductivity and non-conductivity of sodium chloride could be told. If the two ends of a battery are attached to a large
crystal of sodium chloride, no electric current will flow. One might guess that this result indicates that no ions are present in sodium chloride.
However, that is not the case. Indeed, a crystal of sodium chloride is made up entirely of ions, positively charged sodium ions and negatively charged
chloride ions. The problem is, however, that these ions are held together very tightly by electrical forces. Sodium ions are bound tightly to chloride
ions, and vice versa.
The situation changes, however, when sodium chloride is added to water. Water molecules are able to tear apart sodium ions and chloride ions in
much the same way they tear apart acetic acid molecules. Once the sodium ions and chloride ions are no longer bound tightly to each other, they are
free to roam through the salt/water solution. The name given to this change is dissociation. The term means that ions already existed in the sodium
chloride crystal before it was put into water. Water did not create the ions, it only set them free. It is this difference between creating ions and setting
them free that distinguishes ionization from dissociation. ================================================================================================ The word anhydride means “without water”, and the acidic and basic anhydrides are compounds that, when added to water, become common acids and bases. Acid anhydrides are the oxides of nonmetals; basic anhydrides are the oxides of metals. In general, when metal oxides react with water they form bases, and when nonmetal oxides react with water they form acids.
For example: ( )
( ) ( )
i Na O H O Na OH
ii CaO H O Ca OH
sodium hydroxide
calcium hydroxide2 2 2
2 2
2+ →+ →
. Calcium oxide is an important ingredient in cement. When water is
added to the cement the reaction above is one of many that occurs.
Typical examples of the formation of an acid are the reactions of CO2 and SO2 : acid sulfurous )(
acid carbonic )(
3222
3222
SOHOHSOii
COHOHCOi
→+→+
Carbonic acid and sulfurous acids are too unstable to be isolated as pure compounds, but water solutions of them are quite common. Atmospheric CO2 dissolved in groundwater exists partly as CO2 molecules and partly as carbonic acid. It slowly dissolves limestone and is
responsible for the large limestone caves found in various locations. Pollution of the air by SO2 makes rain slightly acidic. This has caused a
great deal of damage to marble statues in many parts of the world (marble is a form of limestone 3CaCO ).
18
Oxides of metals in the middle groups of the periodic table (III through V) lie on the border between ionic and covalent behaviour and are frequently amphoteric. An example is aluminum oxide ( ), which dissolves to only a limited extent in water but much more readily in either acids or bases: Acting as a base: Acting as an acid: Some oxides do not show behavior as either acid or base.
The oxides of the chemical elements in their highest oxidation state are predictable and the chemical formula can be derived from the number of valence electrons for that element. Even the chemical formula of O4, tetraoxygen, is predictable as a group 16 element. One exception is copper for which the highest oxidation state oxide is copper(II) oxide and not copper(I) oxide. Another exception is fluoride that does not exist as expected as F2O7 but as OF2. Since F is more electronegative than O, OF2 does not represent an oxide of fluorine, but instead represents a fluoride of oxygen. Phosphorus pentoxide, the third exception is not properly represented by the chemical formula P2O5 but by P4O10.
Insolubility in water The oxide ion, O2−, is the conjugate base of the hydroxide ion, OH−, and is encountered in ionic solid such as calcium oxide. O2− is unstable in aqueous solution − its affinity for H+ is so great (pKb ~ −22) that it abstracts a proton from a solvent H2O molecule: O2− + H2O → 2 OH−
Nomenclature In the older literature and continuing in industry, oxides are named by contracting the element name with "a." Hence alumina, magnesia, and chromia are respectively names for Al2O3, MgO, Cr2O3.
Two other types of oxide are peroxide, O22−, and superoxide, O2
−. In such species, oxygen is assigned higher oxidation states than oxide. Other oxygen ions ozonide, O3
− and dioxygenyl, O2+.
19
(a) Acids are often used to dissolve insoluble hydroxides or oxides of metals (base anhydrides). For instance, lanthanum ions serve to suppress interferences in atomic absorption spectroscopy. Dissolution of lanthanum oxide, 32OLa , is achieved by reacting it with
either HCl or 3HNO :
(i) OHLaClOLaHCl 2332 326 +→+ ; (ii) OHNOLaOLaHNO 233323 3)(26 +→+
The chloride and nitrate salts are soluble in water.
(b) Acid anhydrides such as 23 & COSO react with bases as in the following reactions:
(i) )(2)(42)(3)(2 laqgaq OHSONaSONaOH +→+ ; (ii) )(2)(3)(2)(2)( lsgaq OHCaCOCOOHCa +→+
(c) Acid anhydrides react with water to form an acid solution: (d) Acid and basic anhydrides may react with each other without any water present. Lime can react with sulfur trioxide in this trioxide
in this reaction: )(4)()(3 ssg CaSOCaOSO →+
======================================================================================= Salts - These are ionic compounds that contain any negative ion except hydroxide ion (that would be a base) and any positive ion except hydrogen ion (that would be an acid) or oxide ion (that would be base or acid). All common soluble salts are dissociated into ions in solution. Even if a salt is only slightly soluble, the small amount that does dissolve completely dissociates. Thus the term “weak” and “strong” are not applicable to salts.
Salts are formed by a chemical reaction between:
• A base and an acid, e.g. NH3 + HCl → NH4Cl
• A metal and an acid, e.g. Mg + H2SO4 → MgSO4 + H2
• A base and an acid anhydride, e.g. 2 NaOH + Cl2O → 2 NaClO + H2O
• An acid and a basic anhydride, e.g. 2 HNO3 + Na2O → 2 NaNO3 + H2O
• Salts can also form if solutions of different salts are mixed, their ions recombine, and the new salt is insoluble and precipitates (see: solubility equilibrium ), for example:
Pb(NO3)2(aq) + Na2SO4(aq) → PbSO4(s) + NaNO3(aq)
Salts are strong electrolytes (Solutions of salts containing highly charged ions, like magnesium sulphate, , are better conductors of electricity than equally concentrated solutions of salts with less highly charged ions, like potassium bromide). When they dissolve in water, they dissociate into their component cations and anions. In many cases, these ions are weak Bronsted-Lowry acids or bases. The reactions of
ions from salts to form +OH3 or −OH ions are called salt hydrolysis reactions. Hydrolysis does not occur with all ions, only with those
that are conjugate acids of weak bases or conjugate bases of weak acids. Chlorine ion is the conjugate base of the strong acid HCl and consequently is ineffective as a base, unlike and . Its interaction with water would therefore scarecely change the concentration. For this reason, a solution of NaCl is neutral, while one of NaF is slightly basic. It is possible to predict whether a salt
hydrolysis reaction produces an acidic solution (containing −OH3 ions) or a basic solution ( −OH ions). One simple way is to consider the acid
and base from which the salt is formed. There are seven (mainly four) possibilities:
(i) Salts of strong acids and strong bases – solutions are neutral. E.g. : neutral acid strong base strong
)(2)()()( laqaqaq OHNaClHClNaOH +→+ The solution is neutral.
The anion of a strong acid is a very weak base which has no tendency to capture protons from water molecules and therefore doesn’t hydrolyze; the cation of a strong base (generally ions of metals in Group IA) is a very weak acid and doesn’t hydrolyze either.
(ii) Salts of strong acids and weak bases – solutions are acidic. e.g.: . The solution is acidic because
+4NH ion is a Bronsted-Lowry acid and donates +H ions to water: )(3)(3)(2)(4 aqaqlaq OHNHOHNH ++ +⇔+ . The Chloride ion, being the
conjugate base of a strong acid (HCl ) has virtually no affinity for +H ions. It is merely a spectator ion.
(iii) Salts of weak acids and strong bases – solutions are basic. e.g.: basicslightly acidk weabase strong
2 2 )(2)(32)(32)( laqaqaq OHCONaCOHNaOH +→+. The solution is basic
because the −23CO ion is a weak Bronsted-Lowry base and accepts +H ions from water: −−− +⇔+ )()(3)(2)(
23 aqaqlaq OHHCOOHCO
(iv) Salts of weak acids and weak bases – solutions can be acidic, basic, or neutral, depending on the relative strengths of the acids and bases from which the salt is formed. In this case, both the cations and anions react with water. Prediction is a little bit harder. (Concentration of weak acid or a weak base in water is difficult to measure directly, but can be calculated using a procedure called acid-base titration.)
((vv)) ((aa)) AAcciidd SSaall ttss aass ddeessccrriibbeedd iinn tthhee ffooll lloowwiinngg sseeccttiioonn aanndd ootthheerrss ssuucchh aass ClNH 3 aanndd BrHN 52 [[ +
)(44 aqNH iiss tthhee ccaattiioonn ((aacciidd)) ooff wweeaakk
bbaassee 3NH aanndd +
)(52 aqHN iiss tthhee ccaattiioonn ((aacciidd)) ooff wweeaakk bbaassee 42HN ]]
→→ HHyyddrroollyyssiiss rreeaaccttiioonn:: 11.. )(3)(32)(44 aqaqaq NHOHOHNH +⇔+ ++ 22..
)(42)(32)(52 4 aqaqaq HNOHOHHN +⇔+ ++
20
(b) All polyatomic ions whose chemical formula begins with H (e.g., ,)(3−
aqHCO −)(4 aqHSO ) are amphoteric. They are also called
amphiprotic because they can either donate or accept a hydrogen ion. Examples:
(1) −+ +→
)(3)()(3 aqHCONaNaHCO aqs (Dissociation)
−+− +⇔+ 2)(33)(2 )()(3 aql COOHOHHCO
aqs (acid hydrolysis:)
)(32)()(2)(3 aqaql COHOHOHHCOs
+⇔+ −−
(base hydrolysis)
The solution is tested basic.
(2) −+ +→ )(4)()(4 aqaqs HSONaNaHSO (Dissociation)
−+− +⇔+ 2)(43)(2)(4 )( aqlaq SOOHOHHSO
aq (acid hydrolysis)
)(42)()(2)(4 aqaqlaq SOHOHOHHSO +⇔+ −−
(base hydrolysis
The solution is tested acidic.
(3) −+ +→
)(32)()(32 aqBOHNaBONaH aqs (Dissociation)
−+− +⇔+ 2)(33)(22 )()(3 aql HBOOHOHBOH
aqs (acid hydrolysis)
)(33)()(22 )(3 aqaql BOHOHOHBOHs
+⇔+ −−
(base hydrolysis) The solution is tested basic.
Use Ka and Kb to decide whether the solution to acidic or basic.
(vi) Group 1 & 2 metal ions (except for ) do not produce acidic solutions, but that highly charged small ions do form acidic
solution. Ions such as have large (positive) charge densities - a large amount of charge in a small volume. These cations produce hydronium ions indirectly by a slightly different reaction than the ammonium ion example above. When a
highly charged metal ion such as dissolves in water, it becomes hydrated with six water molecules (water hydration)
according to the following equation: ++ ⇔+ 3
)(62)(23
)( )(6 aqlaq OHAlOHAl . This means that six water molecules bond to the ion with relatively weak electrostatic forces of attraction. The high-charge density of the ion increases the polarity of the –OH bonds
in the OH2molecules. That is, the OH2
molecules in the +3)(62 )( aqOHAl are more likely to transfer a proton to the solvent (
)(2 lOH )
than are )(2 lOH molecules in pure water, and experiments show that only one OH2
of the six will transfer the proton:
+++ +⇔+ 2)(52)(3)(2
3)(62 )()()( aqaqlaq OHOHAlOHOHOHAl
(***see pages 12, 13, 19)
Note that no cation with low charge density acts as an acid in this way. This includes all of the singly charged ions of Groups 1 and 2 metals except Be.
(vii) Metal oxides (basic anhydrides) react with water to produce basic solutions; nonmetal oxides (acid anhydrides) react with
water to produce acidic solution. (a) Most metal oxides have low solubility in water, but the solid state oxide ions are converted completely into aqueous hydroxide ions by
reacting with water to form a basic solution: in which
(b) ----------------------------------------------------------------------------------------------------------------------------------
TThhee aacciiddiicc pprrooppeerrttiieess ooff nnoonnmmeettaall ooxxiiddeess aarree rreessppoonnssiibbllee ffoorr mmaannyy nnaattuurraall pprroocceesssseess ssuucchh aass tthhee wweeaatthheerriinngg ooff mmiinneerraallss,, tthhee aabbssoorrppttiioonn ooff nnuuttrriieennttss bbyy tthhee rroooottss ooff ppllaannttss,, aanndd tthhee cchheemmiissttrryy ooff ttooootthh ddeeccaayy.. ((aacciiddii ff iiccaattiioonn ooff rraaiinn oorr ggrroouunndd wwaatteerr))
(dissolution of in acidic solution) AAbboovvee rreeaaccttiioonnss aarree rreessppoonnssiibbllee ffoorr tthhee ffoorrmmaattiioonn ooff ccaavveess aanndd ssttaallaaccttii tteess ((cceeii ll iinngg)) aanndd ssttaallaaggmmii tteess ((ggrroouunndd)) iinn tthhee ccaavveess.. NNoottee:: LLiikkee wwaatteerr,, hhyyddrrooggeenn ccaarrbboonnaattee iioonn iiss aammpphhootteerr iicc aanndd ccaann aacctt aass aacciidd oorr bbaassee::
((ii )) iinn tthhiiss rreeaaccttiioonn::
((ii ii )) iinn tthhiiss rreeaaccttiioonn::
21
Song of Solubility Rules (sung to the tune of 99 Bottles of Beer on the Wall)
Alkali metals and ammonium salts,
Whatever they may be, All acetates, chlorate, perchlorate, and nitrates
except beryllium acetate Show solubility.
Most every halide’s soluble At least we’ve always read Save silver, mercury one
And halides of lead.
Every single sulfate Is soluble, ‘tis said
‘Cept barium, strontium, silver, mercury one And calcium and lead.
Sulfides, sulphite, and Hydroxides
Don’t dissolve at all But barium, strontium and calcium
Are slightly soluble.
The chromate, phosphates and the carbonates Aren’t soluble you know
Or else our rocks and statues Would melt away like snow.
Refrain:
Alkali metals and ammonium salts,
Whatever they may be, All acetates, chlorate, perchlorate, and nitrates
Show solubility.
22
Predicting Precipitation: When there are more dissolved ions than the solvent can hold, a supersaturated solution is formed, which is an unstable, non-equilibrium state. Some ions will have to precipitate out of the solution. Thus a precipitate will form in a supersaturated solution. To find out if a solution is supersaturated, the reaction quotient, Q, which in this case is the ion product, is calculated. The actual concentrations are measured at a given point and inserted into the formula for Ksp , the valued thus calculated is
called the ion product (Q) and is compared to the Ksp for the substance. If Q is greater than Ksp , that indicates the concentration of
ions is too high and the solution is supersaturated. To attain equilibrium, ions must precipitate from the solution.
I) Precipitation Reaction (Double Displacement)–Driven by Formation of Solids: Use the Solubility Rules to predict precipitates:
Soluble Ionic Compounds (Salts): ���� ���� Exceptions (Insoluble): sodium, potassium (alkali metals), ammonium salts
NO3
− , ClO3− ,ClO4
− , −COOCH3 nitrate, chlorate, & perchlorate, acetate salts Those containing:
Cl Br I− − −, , chloride, bromide, iodide salts Those containing:
SO42− sulfate salts Those containing:
Exceptions (Soluble): ���� ����Insoluble Ionic Compounds (Salts & Hydroxide):
Those containing also alkali metals or ammonium ions: sulfide, sulfite, carbonate, phosphate, & chromate salts
Those containing: Grp I: ; Grp II: hydroxide compounds
e. g. 1) K CrO Ba NO BaCrO KNOaq aq s aq2 4 3 2 4 3( ) ( ) ( ) ( )( )+ → + , BaCrO s4( ) is a yellow solid precipitated. This reaction is
also called a Double-Displacement Reaction. 2) KCl AgNO AgCl KNOaq aq s aq( ) ( ) ( ) ( )+ → +3 3 , AgCl s( ) is a white solid precipitated.
3) MgSO BaCl MgCl BaSO4 2 2 4+ → +
A reaction in which two solutions are mixed and a precipitate forms is called a precipitation reaction. In this kind of reactions, the solubility rules can be used to predict whether a product will be a solid or will remain as dissolved ions in solution. Whether or not a precipitate actually forms in a double-displacement reaction depends upon the concentrations of the dissolved ions after the solutions are mixed. If the final solution is too dilute, no precipitate can form and the double – displacement does not occur. Only if the ion product exceeds the solubility product will a precipitate form, which will last until the ion concentrations decrease to the equilibrium level.
An equation that shows all soluble ionic substances as ions is called a complete ionic equation. The following is an example of a
complete ionic equation: Cu NO Na OH Cu OH NO Naaq aq aq aq s aq aq+ − + − − ++ + + → + +( ) ( ) ( ) ( ) ( ) ( ) ( )( )2 2 2 2 23 2 3 . Ions such as NO aq3( )
− and
Na aq+
( ) that do not take part in a chemical reaction and are found in solution both before and after the reaction are called spectator ions. If these spectator ions are removed, the equation is called the net ionic equation.
The common – ion effect is a shift in equilibrium that occurs because the concentration of an ion that is part of the equilibrium is changed. In other words, the presence of a common ion lowers the solubility of a sparingly soluble substance. Examples: (i) CaSO Ca SOs aq aq4
242
( ) ( ) ( )⇔ ++ − . Calcium sulfate is sparingly soluble in water. Suppose Na SO s2 4( ) is now added, then Na SO s2 4( ) will
dissolve into Na+ and SO42− ions, thus increasing the concentration of the common ion SO4
2− . In accordance with Le Chatellier’s
principle, this solution will return to equilibrium by precipitating additional CaSO s4( ) : CaSO Ca SOs aq aq42
42
( ) ( ) ( )← ++ − .
(ii) AgCl Ag Cls aq aq( ) ( ) ( )⇔ ++ − . If sodium chloride is added, the precipitation of AgCl will occur: AgCl Ag Cls aq aq( ) ( ) ( )← ++ − .
The Effects of pH on Solubility Some solids are only weakly soluble in water but dissolve readily in acidic solutions. Copper and nickel sulfides from ores, for example, can be brought into solution with strong acids. Solubility of Hydroxides – one direct effect of pH on solubility occurs with the metal hydroxides.For example,
and its solubility product expression is
As the solution is made more acidic, the concentration of hydroxide ion decreases, causing an increase in the concentration of ion. Thus Zinc hydroxide is more soluble in acidic solution than in pure water.
Solubility of Salts of Bases (see p. 13) – the solubility of salts in which the anion is not the hydroxide ion but a different weak or strong base are also affected by pH. For example,
and
As the solution is made more acidic, some of the fluoride ion reacts with hydronium ion through .
23
Because this is the reverse of the acid ionization of HF, its equilibrium constant is the reciprocal of or
As acid is added, the concentration of the fluoride ion is reduced, so the calcium ion concentration must increase to maintain the solubility product equilibrium for As a result, the solubility of fluoride salts increases in acidic solution. The same apply to other ionic substances in which the anion is a weak or a strong base. By contrast, the solubility of a salt such as AgCl is only very slightly affected by a decrease in pH. The reason is that HCl is a very strong acid, and so is ineffective as a base. The reaction occurs to a negligible extent in acidic solution.
Solubility of Complex Ions – Many transition-metal ions form coordination complexes in solution or in the solid state; these consist of a metal ion surrounded by a group of anions or neutral molecules called ligands. The interaction involves the sharing by the metal ion of a lone pair on each ligand molecule, giving a partially covalent bond with that ligand. Such complexes frequently have strikingly deep colours.
When exposed to gaseous ammonia, greenish white crystals of copper sulfate ( give a deep blue crystalline solid with the chemical formula . The anions in the solid are still sulfate ions ( ), but the cations are now complex ions of the central ion with 4 ammonia molecules, . The ammonia molecules coordinate to the copper ion through their lone-pair electrons acting as Lewis bases toward the metal ion, the Lewis acid. When the solid is dissolved in water, the deep blue color remains. This is evidence that the complex persists in water, because when ordinary (without ammonia ligands) is dissolved in water, a much paler blue color results.
.
When silver ions are dissolved in an aqueous ammonia solution, doubly coordinated silver-ammonia complexes form in two stepwise reactions:
(i)
(ii)
If these two equations are added (and their corresponding equilibrium laws are multiplied), the result is
Where is the formation constant of the full complex ion . The larger the formation constant , the more stable the corresponding complex ion, for ions with the same number of ligands.
Sample Calculation: Suppose 0.100 mol of is dissolved in 1.00 L of a 1.00 M solution of . Calculate the concentration of the and ions present at equilibrium.
Solution: Assume that most of the is present as (this will be checked later).
Then and after each silver ion has become complexed with two ammonia molecules. The two stages of the dissociation of the ion are the reverse reaction of the complexation:
(ii i)
(i ii)
If mol /L of dissociates at equilibrium according to the first equation,
(ii i)
I 0.100 0 0.80
C - x + x + x
E 0.100 – x x x
We can then calculate the concentration of the free ions from the equilibrium law for the second step of the dissociation of the complex ions:
24
So,
It is clear that the original assumption was correct, and most of t he silver present is tied up in the complex. Note: the above calculation is similar to the calculation of polyprotic acid equilibriua in the example on page 23 of III: Thermodynamic & Equilbrium. The only difference is that in complex ion equilbria it is conventional to work with formation constants, which are the inverse of the dissociation constants used in acid-base equilibria. The formation of coordination complexes can have a large effect on the solubility of a compound in water. Silver bromide is only very weakly soluble in water, But addition of thiosulfate ion ( ) to the solution allows the complex ion to form: This greatly increases the solubility of the silver bromide. The formation of this complex ion is an important step in the development of photographic images; thiosulfate ion is a component of the fixer that brings silver bromide into solution from the unexposed portion of the film.
25
Another interesting effect of complex ions on solubility is illustrated by the addition of iodide ion to a solution containing mercury (II) ion. After a moderate amount of iodide ion has been added, an orange precipitate forms through the reaction:
With further addition of iodide ion, however, the orange solid redissolves because complex ions form:
(i) (ii)
In the same way, silver chloride will dissolve in a concentrated solution of sodium chloride by forming soluble complex ions.
Complex ion formation affects solubility in the opposite direction from the common-ion effect.
Acidity and Amphoterism of Complex Ions (see pages 12, 13, 19)
When dissolved in water, many metal ions increase the acidity of the solution. The iron (III) ion is an example: each dissolved
ion is strongly solvated by six water molecules, leading to a complex ion . This complex ion can act as a Bronsted-Lowry
acid, donating hydrogen ions to the solvent, water.
Example 11.13: Calculate the pH of a solution that is 0.100 M in .
Solution: The iron (III) is present as , which reacts as weak acid:
with
If mol/L of reacts, then (neglecting the ionization of water itself)
;
,
and so the pH is 1.62. Solutions of iron (III) salts are strongly acidic.
When strong acid is added, the equilibrium is driven back to the left, and the color fades.
Those metal ions that form strong complexes with hydroxide ion have low pH whreas those that do not form such complexes give neutral solutions. Different cations behave differently as water ligands are replaced by hydroxide ions in an increasingly basic solution. A particularly interesting example is . It forms a series of hydroxo complex ions:
In the Bronsted-Lowry theory, a polyprotic acid, , donates hydrogen ions in succession to make all the product ions. The
second product, , is amphoteric; it can react as either acid or base. It is only slightly soluble in pure water (its
). If enough acid is added to solid , the ligands are removed, forming the soluble ion; if enough
base is added, ligands attach to form the soluble (zincate) ion. Thus, is soluble in strongly acidic or
strongly basic solutions, but is only slightly soluble at intermediate pH values. This ampghoterism can be used to separate from other cations that do not share this property. For example, adds a maximum of two ions to form a sparingly
26
soluble hydroxide. If a mixture of and ions is made sufficiently basic, precipitates as and zinc remains in solution as . In the same way, precipitates while dissolves in strong basic solution as
27
Summary of Types of Compounds and Their Names Type of Compounds Cation (+) Anion (-)
Polyatomic Ion (+)
Polyatomic Ion (-) Formula
IUPAC name (Stock System)
Classical (Common) Name
Type I - Binary Ionic Compounds (Singular
Valent: salts)
metal (Groups I & II): (name)
non-metal: (root)-ide
NaCl KI
CaS CsBr Ca P3 2
Sodium chloride potassium iodide calcium sulphide cesium bromide
Calcium phosphide
salt
Type II - Binary Ionic Compounds (Multivalent)
(salts)
metal (transition):
(name) (roman numeral)
non-metal : (root)-ide
CuCl
2CuCl
copper (I) chloride copper (II) chloride
cuprous chloride cupric chloride
Type I&II (A)- Binary Ionic Compounds
(basic anhydrides)
metal (Grp. I & II, and
transition): (name; roman
numeral )
non-metal: "O"
oxide
CaO MgO FeO
32OFe
calcium oxide magnesium oxide
iron (II) oxide iron (III) oxide
quicklime
ferrous oxide ferric oxide
Type III (A) - Binary
Molecular Compounds
(acid anhydrides)
non-metal: prefix-(name)
non-metal: "O"
prefix-oxide
ON2
NO
32ON
52ON
dinitrogen oxide nitrogen oxide
dinitrogen trioxide dinitrogen pentoxide
nitrous oxide nitric oxide
nitrous anhydride nitric anhydride
Type III (B) - Binary
Molecular Compounds
(bases or acids)
non-metal: prefix-(name)
non-metal: prefix-(root)-
ide
3NH
42HN
6SF
PCl3
5PCl
nitrogen trihydride
dinitrogen tetrahydride
sulphur hexafluoride
phosphorus trichloride
phosphorus pentchloride
Ammonia (base)
Hydrazine (base)
Type III (C) - Binary
Molecular Compounds (gaseous & aqu acids)
non-metal: (i).hydrogen in gaseous
state: (ii).hydrogen
in aqu solution: (aq. hydrogen)
non-metal: prefix-(root)-ide
HF HCl HBr HI
SH2
(aq.) gaseous acids: (aq.) hydrogen fluoride (aq.) hydrogen chloride (aq.) hydrogen bromide
(aq.) hydrogen idoide (aq.) hydrogen sulfide
aqueous acids: hydrofluoric acid hydrochloric acid hydrobromic acid
hydroiodic acid hydrosulfuric acid
Type IV (A)- Tertiary
Compounds (Bases)
metal: aqueous (name)
Hydroxide (OH-)
NaOH 2)(OHCa
sodium hydroxide calcium hydroxide
caustic soda
Type IV (B)- Tertiary
Compounds (salts)
metal: (name)
oxyanions: hypo-(root)-ite
(root)-ite (root)-ate
per-(root)-ate
CaCO3 .calcium carbonate limestone
Type IV(C) - Tertiary
Compounds (salts
hydrates)
metal: (name)
oxyanions: hypo-(root)-ite
(root)-ite (root)-ate
per-(root)-ate
+ water: prefix-
hydrate:OHCuSO 24 5•
copper (II) sulphate penthydrate
bluestone
Type IV (D)- Tertiary
Compounds (acid salts)
metal + hydrogen:
(name) prefix-hydrogen
oxyanions: hypo-(root)-ite
(root)-ite (root)-ate
per-(root)-ate 4
3
42
MgHPO
NaHCO
POKH
potassium Dihydrogen phosphate
sodium hydrogen carbonate
magnesium monohydrogen phosphate
Type IV (E)- Tertiary
Compounds (acids)
non-metal: (i).hydrogen
in gaseous state: (ii).hydrogen in aqu. solution: (aq. hydrogen)
oxyanions: hypo-(root)-ite
(root)-ite (root)-ate
per-(root)-ate
e. g.
52
4
3
2
SOH
HSO
HSO
HSO
.
aq. hydrogen hypo-(root)-ite aq. hydrogen (root)-ite
aq. hydrogen (root)-ate aq. hydrogen per-(root)-ate
hypo-(root)-ous acid (root)-ous acid
(root)-ic acid per-(root)-ic acid
Type IV (F) - Tertiary
Compounds (acids)
non-metal: hydrogen
Non-Oxy: cyanide(CN-)
HCN
aq. hydrogen cyanide
hydrocyanic acid
.
Type IV (G)- Tertiary
Compounds
other non-metal:
oxyanions: hypo-(root)-ite
(root)-ite (root)-ate
per-(root)-ate
Type IV (H) - Tertiary
ammonium oxyanions: hypo-(root)-ite
(root)-ite
28
Compounds (root)-ate per-(root)-ate
44ClONH
ammonium perchlorate
solid rocket fuel
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