Download HE)HEPTA

Re-typed from The Ultimate Chemical Equations Handbook by Hague and Smith
SIMPLE INORGANIC
FORMULAS
AND NOMENCLATURE
Binary Molecules
I. A binary molecule is formed when two nonmetals or metalloids combine. Electrons are shared so the
bonding involved is known as CD VQ LJ:::N T
bonding.
2. Sometimes these compounds have generic or common names (water) and they also have systemic
names (dihydrogen monoxide). The common names must be memorized. The systemic name is more
complicated but it has the advantage that the formula of the compound can be deduced from the name.
3. Simple binary compounds consist of only a few atoms. Systemic naming of these compounds follow
the rules:
•
The elements, except for H, are written in order of increasing group number.
•
Prefixes are used to designate the number of each element present in the molecule.
4. The prefixes are:
rl0IJO
6
HE)<A
DI
7
HEPTA
3 II( \
8
otm
"-.ION 1-1
2
.~,..
.
4
TET12.A
9
5
Pi.::l'-iTA
10
5. Mono is
assumed.
*****
;Jt'l./l::7L. used in front of the first element. If there is only one atom, the mono is
Name the following binary molecules:
(a) C02
(b) N203
(c) P4010
(d)
DECIj
N20
cr·W_mAl
DIN
DIOX.I])~
IT/106EN
T!2IO'IJDl.:;-
It:"Tf(l::j PHoSPHoR.US
iJE(.o)(.1 D i2"
birJ rTQ.0680
SV\otJo '1..1j)t;"
Ionic Compounds
I. Ionic compounds are formed between metals and nonmetals. The electrostatic force of attraction
between the positive ion (cation) formed by the metal and the negative ion (anion) formed by the
nonmetal is what holds the compound together.
OVER
r
2. The cation is named first and written first when writing a formula while the anion is both named and
written second.
3. Cations can have one or more charges. These charges are known as oxidation numbers, or valences.
4. The transition metals and representative elements in Groups 13, 14, 15, and 16 have multiple oxidation
numbers.
5. Roman numerals, enclosed in parentheses, are used after the name of the cation to designate the
oxidation state of the cation - only if the cation has more than one positive oxidation state from which to
choose.
6. This system of nomenclature is known as the Stock system.
7. Here are some simple rules that should help in the determination of the oxidation numbers of cations
from the formulas of their compounds.
(a) The oxidation number of any element in its free state (uncombined with other elements) is zero. Fe
in a bar of iron is zero. 02 and N2 in the atmosphere are zero.
(b) The oxidation number of the alkali metals in a compound is always 1+.
(c) The oxidation number of the alkaline earth metals in a compound is always 2+.
(d) Fluorine in a compound is always assigned an oxidation number of 1-.
t.
(e) The oxidation number of oxygen is almost always 2- in a compound. Exceptions to this rule would
be peroxides, oi- where the oxidation number of each oxygen is 1-, and superoxides, Of where the
oxidation number of each oxygen is Yz-.
(0 In covalent compounds (with nonmetals) hydrogen is assigned an oxidation number of 1+ (examples
are HCI, H20, NH3, CH4).
(g) In metallic halides the halogen (F, CI, Br, I, and At) always has an oxidation number equal to 1-.
(h) Sulfide, selenide, telluride, and polonide are always 2- in binary salts.
(i) Nitrides, phosphides, and arsenides are always 3- in binary salts.
U) All other oxidation numbers are assigned so that the sum of the oxidation numbers of each element
equals the net charge on the molecule or polyatomic ion. In neutral compounds, the sum of the positive
and negative charges must equal zero.
*****
Determine the oxidation number of the underlined element:
+
(b) NaCI04
7
t
I~
+3
\
Re-typed from The Ultimate Chemical Equations Handbook by Hague and Smith
7. Free elements, no matter how complex the molecule, have an oxidation number (valence or charge)
equal to zero.
OJ.
Fe.
He
~~
<
T~
/1 1
L-td.
8. The following are diatomic or polyatomic elements in nature which must be committed to memory.
These elements exist as neutral molecules in nature:
(a)
i3r,J.
(g)
r,
(b)
J:'l
(h)
03
(c)
J\J.1
(i)
f
(d)
CJa
(j)
Sl?
(e)
I--b
(k)
C &0
(f)
0)
(I)
C/o
y
Charges and the Periodic Table
1. The periodic table can be used to help determine charges on many ions. Cations come from metals
that lose electrons (
0 -y: I IJ tCl, TI 0 J-J
) in order to become .ISO 8-ECTl20 tJ' C
with a noble gas. Anions come from nonmetals that gain electrons (
lZ..Ebucno Iv
) in
order to become r50E1£ CIlLo,.) f c..
with a noble gas.
*****
Group 1
+1
Group 17
-I
Group 2
+~
Group 16
- .)
Group 3
+-3
Group IS
-3
Group 14
-~
2. Transition metals, representative metals with
typically have more than one oxidation state.
+
and ~
s«
+t..J
sublevels, and the inner transition metals
ei
3. Electrons for these metallic elements are lost in the following order:
S) J....
. Such
elements are ;...fa T isoelectronic with a noble gas when the outermost (valence) electrons are lost.
4. Inner transition elements are also known as the lanthanides, the actinides, the rare earth elements, and
the transuranium elements. These elements are rare and many exist for short periods of time.
5. Both inner transition elements and transition elements are known for their variable oxidation numbers.
The most common oxidation number for transition elements is f ~
OVER
r
6. The d: sublevel in transition elements is responsible for the various oxidation numbers that result.
Incomplete d sublevels are also responsible for the many colorful transition compounds that are known to
exist. Complete d sublevels in cations of silver and zinc result in white compounds.
SUMMARY OF CATIONS WITH VARIABLE OXIDATION NUMBERS
- STOCK SYSTEM
copper (I), Cu"; copper (II), Cu2+
mercury (I), Hgr"; mercury (II), Hg2+
1+,2+
gold (I), Au"; gold (Ill), Au3+
indium 0), In"; indium (III), £n3+
thallium (0, TI+; thallium ([II), T13+
chromium (II), Cr2+;chromium (III), cr3+
cobalt (Il), C02+; cobalt (III), C03+
iron (II), Fez+; iron (III), Fe3+
lead (II), Pb2+; lead (IV), Pb4+
platinum (II), peT; platinum (IV), pt+
tin (II), Sn2+;tin (IV), Sn4+
zirconium (II),
zirconium(lV), Zr4+
1+,3+
2+,3+
2+,4+
z>.
cerium (III), Ce3+; cerium (IV), Ce4+
3+,4+
antimony (1II),Sb3\ antimony (V), Sb5+
arsenic (III), As3+; arsenic (V), As5+
bismuth(III), Bi3+; bismuth (V), Bi5+
phosphorus (III), p3+;phosphorus (V), p5+
iridium (II), Ir2+;iridium (III), Ir3+;iridium (IV), Ir4+
titanium (II), Ti2+;titanium (HI), Ti3+; titanium (IV), Ti4+
manganese (II), Mnl+; manganese (III), Mn3+;manganese (IV), Mn4+
3+,5+
2+,3+,4+
2+,4+,5+
tungsten (II), W2+; tungsten (IV), W4+; tungsten (V), W5+
3+,4+,5+
uranium (III), U3+;uranium (IV), U4+; uranium (V), U5+
2+,3+,4+,
5+
vanadium (II), VZ+; vanadium (III), V3+;vanadium (IV), V4:;
vanadium (V), V5+
-
! '
Re-typed from The Ultimate Chemical Equations Handbook by Hague and Smith
Polyatomic Ions
I. Polyatomic ions are a group of atoms that behave as a single ion.
2. The bonding within a polyatomic ion is covalent, but because there is either an excess or a shortage of
electrons compared to the number of protons present, an ion results.
3. This short list of polyatomic ions must be MEMORIZED.
NH4+
ammonium
OR
hydroxide
NOf
nitrite
P043-
phosphate
N03-
nitrate
CIO-
hypochlorite
so>
sulfite
CIOf
chlorite
S042-
sulfate
C103•
chlorate
C032•
carbonate
CI04-
perchlorate
HC03-
bicarbonate
Cr042-
chromate
C2H)Of
acetate
Cr20l·
dichromate
CN-
cyanide
SCN-
thiocyanate
You also need to know the common Group ions such as (but not limited to) the ions in Groups 1 &
2, the halogens, 0, S, N, P, and "the triangle": Ag, Zn, Cd, AI, Ga, and In.
,.
r
\
OVER
© Adrian Dingle's Chemistry Pages 2004. All rights reserved.
These materials may NOT be copied or redistributed in any way, except for individual
class instruction.
Revised June 2005
AP Common Ions
CAT10NS (+ve)
Name
Symbol
ANIONS (-ve)
Alternative"
Name
Symbol
Alternatlve~
Bf
Aluminum
At'+
Bromide
Ammonium
NH:
As3+
Bromate (I)
BrO'
(Hypobromite)
Bromate (III)
BrOz'
As~
Bromate (V)
BrO,'
(Bromite)
(Bromate)
Barium
Ba2+
Bromate (V")
BrO•.
(Perbromate)
Bismuth (III)
8i3+
Carbonate
C0,2'
Arsenic (III)
Arsenic (V)
Bismuth (V)
Bi~
Chlorate (I)
Cadmium
Calcium
c<f+
Ca2+
" Chlomte (III)
CIO'
CI02'
(Hypochlortte)
(Chlorite)
: Chlorate (V)
CIO;
(Chlorate)
Chromium (II)
Cr·
,Chlorate (VII)
CIO"
(Perchlorate)
Chromium (III)
Cr+
. Chloride
Cobalt (II)
C02+
Cobalt (III)
Co'·
Cu+
Chromate
, Cyanide
Copper (I)
Hydrogen
Cu2+
H+
Hydronium
H,O+
Iron (II)
Fe2+
Iron (III)
Fe'+
ptf+
Copper (II)
Lead (II)
Lead (IV)
cr
crO.2,
CN"
CrzOr2,
(Cuprous)
Dichromate
(Cupric)
Dihydrogen Phosphate
H2POi
Ethanoate
CzH,Oi
F"
Ruoride
(Ferrous)
(Ferric)
Hydride
(Plumbous)
Hydrogen Oxalate
Hydrogen Carbonate
H'
He03'
(Acetate)
(Bicarbonate)
/:
HCzO"
HPO/'
(Binoxalate)
(
,,Hydrogen Sulfate
HSOi
(Bisulfate )
..Hydrogen Sulfide
HS'
HSO)'
(Bisulfide)
(Bisulfite)
Hydrogen Phosphate
\'
Pb"
L(
Mg1+
Mn~+
Mn.•.•
(Plumbic)
Mercury (I)
H!hz+
(Mercurous)
Iodate (I)
10'
(Hypoiodile)
Mercury (II)
Hg2•
(Mercuric)
Iodate (III)
102'
(Iodite)
Nickel (II)
Ni2+
(Iodate)
K'"
Iodate (V)
Iodate (VII)
10;
Potassium
10.
(Periodate)
lithium
Magnesium
Manganese (II)
Manganese (IV)
Silver
Na+
Strontium
Sr+
Snz+
Tin (IV)
Zinc
Hydroxide
Ag+
Sodium
,Tin (II)
Hydrogen Sulrrte
Sn'"
Znz+
Iodide
Manganate (V")
OH'
r
MnO •.
, Nitrote
NO,'
(Stannous)
Nitride
(Stannic)
NiDite
N""
NOz'
Oxalate
Oxide
Peroxide
Phosphate
'~
CzO/'
(Permanganate)
(Ethandioate)
ry-
0/'
PO.3-
Phosphide
p""
Phosphite
PO/'
Sulfate
SO.z,
Sulfide
Sulfite
S1'
5032-
Thiosulfate
S20)2-
Thiocyanate
SCN"
• In the case of the cations, the alternative names are genemlly redundant in modem chemistry, but the anions sometimes
use the older, alternate names, For example, the oxyhatogen ions (bromate, chlorate, iodate stc.) are usually referred to
by the altemate names, but HSOJ' is much more likely 10 be called Hydrogen Sulfrte rather than Bisufite,
,---,-------
~-----
Re-typed from The Ultimate Chemical Equations Handbook by Hague and Smith
Ternary
Nomenclature:
Acids and salts Containing
Halogens and/or Oxygen
1. The halogens, with their variable oxidation numbers, allow for a great variety of compounds.
!3/1~£"
2. A good way to learn ternary nomenclature is to start with a certain
HOI'\<lE
polyatomic ion. This is polyatomic ion ending with the suffix
- C~ te
Number of Oxygen Atoms
(compared to home base)
Polyatomic
Ion Name
Plus one oxygen atom
Cl04'
perchlorate ion
Home base
CI03'
ch lorate ion
Minus one oxygen atom
CI02'
chlorite ion
Minus two oxygen atoms
ClO'
hypochlorite ion
No oxygen atoms
CI'
chloride ion
3, Water solutions of binary hydrides form acids. The root derived from the hydride is given the prefix
li'/ Dr2..0 and the suffix
- Ie and the name ends with the word
ACt D
4. The binary hydride HCl is known as H'I Dt2.0GeJ CI-ILo.e.IOlE'
when aqueous it is known as 14 'f D1200-l W.e.IL !~
C ID
(hydrogen monochloride) gas, but
5. Many common acids contain only oxygen, hydrogen, and a nonmetallic ion or polyatomic ion. These
acids are called _-=O...:..Y--...:'1..:.,i-\,!.,:C:..;I.::::.b..=S=-_
6, If the name of the polyatomic ion ends in
the word acid.
-
['L
7. If the name of the polyatomic ion ends in - /
the word acid.
8, The mnemonic aid is:
*****
(cfe-
ic.
+e , the suffix
te. , the suffix
('fe -
Name the following compounds:
(a) HI04 (aq)
(b) NaBr04
Sdb \v yt,
PERIOble.
I~(
11)
eEQ8~Oi!~\(~TE
OVER
OU.5
1/
-
f
c..-
is substituted followed by
- OUSis substituted followed by
*****
Write the formulas for the following compounds:
(a) calcium hypochlorite
Cl\. ( OD)::I
(c) cyanic acid
Ii (NO
(b) hydrotelluric acid
H.)Ie
O!L
iioe iJ
(d) chlorous acid
HC~O,;!
BALANCING EQUATIONS HINTS
SYNTHESIS & DECOMPOSITION
Chemists write balanced equations to illustrate what is happening during a chemical reaction. Bonds are
broken, atoms are rearranged, and new bonds are formed. Every chemical reaction supports the Law of
conservation of Matter. This means that in every reaction, the number of atoms of each type of element
contained within the reactants must be the same as the number of atoms of each type of element contained
within the products.
Balancing equations is a process which assures that equations are written properly to support the Law of
Conservation of Matter; however, balancing cannot be done until each reactant and product formula is
written correctly. It is important to properly write the seven elements that are diatomic in their elemental
form and also use subscripts and parentheses appropriately when considering the oxidation number of
ions. All compounds must be made neutral before beginning to balance the atoms.
Balancing is accomplished by adding coefficients that multiply the number of atoms represented by the
formula. For example, a coefficient of2 in front of oxygen (e.g., 2 02) means that 4 oxygen atoms are
represented. Unlike algebra, in chemistry a coefficient does not need to be outside parentheses or
brackets to b distributed. A coefficient applies to the complete substance; however, it no longer applies
when a plus sign (+) or arrow (--+) is encountered. For example:
3 (NH4)2C03 shows 6 nitrogen, 24 hydrogen, 3 carbon, and 9 oxygen atoms.
3 MgCb + NaBr indicates 3 magnesium, 6 chlorine, 1 sodium, and 1 bromine atom.
PREREQUISITE KNOWLEDGE
Before you do anything, you must know and understand the following areas of nomenclature and formula
writing:
I. Ionic compounds
2. Covalent compounds
3. Acids and bases
4. Complex ions (coordination chemistry)
5. Organic nomenclature
Re-typed from The Ultimate Chemical Equations Handbook by Hague and Smith
TIPS FOR BALANCING
EQUATIONS
I. Ensure each molecular formula is written correctly and each compound is neutral.
2. Mentally count or tally how many of each type of atom is present on each side of the equation.
3. Begin by balancing elements that are only found in one substance on each side.
4. Balance oxygen and hydrogen LAST - they usually balance out at the end or perhaps only the number
of water molecules needs to be adjusted.
5. If there is an odd number of an element on one side and an even number on the other, the odd number
will need to be evened out - so use a coefficient of 2 for that substance.
6. Ifthere are polyatomic ions that remain together as a unit during the reaction, count the polyatomic ion
as a unit.
7. When tallying, be sure to adjust the count for each and every element that an added coefficient affects.
8. Combustion reactions that don't seem to balance will often come out better if a coefficient of 2 is used
for the hydrocarbon.
SYNTHESIS REACTIONS
Synthesis reactions occur when two or more reactants combine to form a single product. There are
several types of synthesis reactions.
I.
A metal combines with a nonmetal to form a binary salt.
Example: A piece of lithium metal is dropped into a container of nitrogen gas.
2. Metallic oxides and water form bases (metallic hydroxides)
Example: Solid magnesium oxide is added to water.
MgO + 2 HOH
-+
Mg(OH)2
3. Nonmetallic oxides and water form acids. The nonmetal retains its oxidation number.
Example: Dinitrogen pentoxide is bubbled into water.
4. Metallic oxides and nonmetallic oxides form salts.
Example: solid calcium oxide is added to sulfur trioxide.
CaO + S03 ~ CaS04
OVER
DECOMPOSITION
REACTIONS
Decomposition reactions occur when a single reactant is broken down into two or more products.
Ie
>
1.
Metallic carbonates decompose into metallic oxides and carbon dioxide.
Example: A sample of magnesium carbonate is heated.
2.
Metallic chlorates decompose into metallic chlorides and oxygen.
Example: A sample of magnesium chlorate is heated.
3. Ammonium carbonate decomposes into ammonia, water and carbon dioxide.
Example: A sample of ammonium carbonate is heated.
4.
Sulfurous acid decomposes into sulfur dioxide and water.
Example: A sample of sulfurous acid is heated.
5. Carbonic acid decomposes into carbon dioxide and water.
Example: A sample of carbonic acid is heated.
6. A binary compound may break down into two elements.
Example: Molten sodium chloride is electrolyzed.
2 NaCI
7.
-+
2 Na + Ch
Hydrogen peroxide decomposes into water and oxygen.
8. Ammonium hydroxide decomposes into ammonia and water.
"
from The Ultimate Chemical Equations Handbook by Hague and Smith
Re-typed
NOMENCLA TURE PRACTICE
I. Name each of the following
compounds.
(a) CaF2
(h) PFs
Pt:NTt\ FU)
,:)H O:Sr l-\0LU.s
F LUOe.1 Dt=
CA Lt lIJ M
ce. I DE
(i) (NH4hS03
Si..'LFlll:
1~i\,nION'viH
U)
(c) NaH
WI Dfll j) l:
:30 b I Vi'Vl
(m)
GeLD
(d) HIO (aq)
H'IpcIO
AU2(C204)3
(k) KSCN
DoLlS
{JOTAS S \ v r11
I~LI.t)
(e) Hg2Ch
'/l-11 OC'f A A.1 A n=
(I) Pb(C2H302)4
n18Z.. C 1'Q.'( (I)
C H L02. I O<i-
LE'AD
(t) Ba02
(IV)
ACETATE
(m) LiH
i=t-e.o Y... I .D~
rM)'fLl U WI
U T\-l h..H1I\
(g) Mn(HC03)3
rY\ i\NGArJESE
(n) H2Cr04 (aq)
( ill.."
H'I DI(OGe-J
2. Write formulas for the following
(a) vanadium
(V) oxide
CAf2&iJ.RTE
substances.
(g) tin (IV) chromate
'Vd. O.s
Sn~ (Cr OLl)'f
(b) zinc hydroxide
(h) nitrous acid
l'1,(oH)J
/f,JO(J
(c) silver chromate
Aj ,:J
H '11)/2..\ DE
(i) magnesium
c-o;
(d) tungsten
(V) thiosulfate
U)
calcium carbide
~()
(k) mercury (II) acetate
C e (A103)-3
(t) hydroiodic
acid
phosphate
Il1!J /-I1tJ C;
~J~(S.;l 03)S
(e) cerium (III) nitrate
hydrogen
J/j (C,; 1-/3 Q;J)
(I) perbromic
(-1-1
acid
!--I erO 'I
OVER
J
BALANCING SYNTHESIS AND DECOMPOSITION
EQUATIONS
Predict and balance the following synthesis and decomposition reactions. Use abbreviations to indicate
the phase of reactants and products where possible [(aq) (s) (I) (g)]
I. A sample of calcium carbonate is heated.
2. Sulfur dioxide gas is bubbled through water.
3. Solid potassium oxide is added to a container ofcarbon dioxide gas.
4. Liquid hydrogen peroxide is warmed.
5. Solid lithium oxide is added to water.
"
L .,
D
"\
(5)
/ti 0
to
-'7
6. Molten aluminum chloride is electrolyzed.
-----
-------------
Re-typed from The Ultimate Chemical Equations Handbook by Hague and Smith
7. A pea-sized piece of sodium is added to a container of iodine vapor .
.;J J4.(s)+-
I,J-
Cj) -7
,J Aft:-...-l.
(5)
8. A sample of carbonic acid is heated.
9. A sample of potassium chlorate is heated.
')
..v
L.Ji)
AQ
1\ L..,( .3 (S)
-~
d.
k'tf
(5)
-I-
.J 002
~J)
..
10. Solid magnesium oxide is added to sulfur trioxide gas.
OVER
--~----
.. ---