The Periodic Table and Compound Names and Formulas
The Periodic Table can be a very useful tool. The elements are arranged according to
increasing atomic number (number of protons). The elements are also grouped so that elements
with similar properties are in the same portion of the table.
The seven periods (rows) on the Periodic Table correspond to the number of energy
levels that are occupied by electrons for a neutral atom of the element.
The elements are assigned to one of the 18 groups or families (columns) on the Periodic
Table based on the number valence electrons (electrons in the outer most energy level – valence
shell) for a neutral atom of that element. The number of valence electrons for an atom plays a
major role in determining an element’s properties. Atoms gain, lose, or share electrons in order
to obtain a filled valence shell.
The elements are classified as belonging one of three groups: metals, nonmetals, or
metalloids (elements with some metallic and some nonmetallic properties). Please note that
other textbooks might list a slightly different group of elements as metalloids.
The nonmetals include: H2, He, C, N2, O2, F2, Ne, P4, S8, Cl2, Ar, Se, Br2, Kr, I2, Xe, and
Rn.
The metalloids include: B, Si, Ge, As4, Sb, Te, and At.
Due to the number of valence electrons in their neutral atoms, some elements are more
stable as covalently bonded (sharing valence electrons) molecules. There are seven diatomic
elements: H2, N2, O2, F2, Cl2, Br2, and I2. Other stable molecular forms are: P4, As4, and S8.
At room temperature (298 K) and a pressure of one – Cs, Fr, Hg, and Br2 are liquids; H2, He, N2,
O2, F2, Ne, Cl2, Ar, Kr, Xe, and Rn are gases; and the other elements are metals. Ga would melt
in your hand since its melting point is 302.9 K.
Elements in Group 1 (alkali metals) have one valence electron in the outermost occupied s–
sublevel and usually form ions with a 1+ charge (by donating its valence electron to another
atom) when combining with another element to form a compound.
Elements in Group 2 (alkaline earth metals) have two valence electrons in the outermost
occupied s–sublevel and usually form ions with a 2+ charge when combining with another
element to form a compound.
Element in Groups 3-12 (transition metals) have two valence electrons but have
electrons being added to the d–sublevel an “inner shell” and usually form ions with a + charge
when combining with another element to form a compound. These elements may have a variety
of + charges depending on what they are combining with when forming compounds. As a first
approximation, the following generalizations may be made regarding the most likely charge for
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an ion of an element in the following Groups: Group 3: 3+, Group 4: 4+, Group 5: 5+, Group
6: 6+, Group 7: 7+, Groups 8-10: 2+ or 3+, Group 11: 1+, and Group 12: 2+.
Please note: The transition (and the inner transition a.k.a. rare earth) elements and a few
of the Group 13-16 metals can form a variety of different charged ions depending on what they
are combining with … so we use Roman numerals when naming the ion in order to identify the
charge.
Elements number 57-70 (lanthanids) and 89-102 (actinids) on our printed Periodic
Tables are collectively known as the rare earth or inner transition elements. These elements
have two valence electrons but have electrons being added to the f–sublevel an “inner shell.”
These elements may form a variety of + charged ions depending on the compound that is formed.
A reasonable generalization would be that the most likely charge for an ion of these elements
would be 3+. Please note: The transition elements can form a variety of different charged ions
depending on what they are combining with … so we use Roman numerals when naming the
ion in order to identify the charge.
Elements in Group 13 (boron family) have three valance electrons, two in the outermost
occupied s–sublevel and one in the outermost occupied p–sublevel. These elements typically
form ions with a 3+ charge when forming a compound. Boron is classified as a metalloid and
the other Group 13 elements are metals.
Elements in Group 14 (carbon family) have four valance electrons, two in the outermost
occupied s–sublevel and two in the outermost occupied p–sublevel. Carbon is classified as a
nonmetal, silicon and germanium are classified as metalloids, and the other Group 14 elements
are metals. Carbon and, to a lesser extent, silicon typically forms molecular compounds
(covalent bonds where electrons are shared). The other Group 14 elements typically form ions
with a 4+ charge when forming a compound, though tin and lead can also form compounds
where their charge is 2+. Nonmetallic ions in this family would most like have a 4- charge.
Elements in Group 15 (nitrogen family) have five valance electrons, two in the
outermost occupied s–sublevel and three in the outermost occupied p–sublevel. Nitrogen and
phosphorous are nonmetals, arsenic and antimony are classified as metalloids, and bismuth is a
metal. Nitrogen, phosphorous, and arsenic typically forms molecular compounds (covalent
bonds where electrons are shared). Antimony and bismuth typically form ions with a 5+ charge
when forming a compound, though other charges are possible. Nonmetallic ions in this family
would most like have a 3- charge.
Elements in Group 16 (Chalcogen family) have six valance electrons, two in the
outermost occupied s–sublevel and four in the outermost occupied p–sublevel. Oxygen, sulfur,
and selenium nonmetals, tellurium is classified as a metalloid, and polonium is a metal.
Nonmetallic ions in this family would most like have a 2- charge.
Elements in Group 17 (halogens) have seven valance electrons, two in the outermost
occupied s–sublevel and five in the outermost occupied p–sublevel. Astatine is a metalloid and
the other elements are others are nonmetals. Halogens form ions with a charge of 1-.
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Elements in Group 18 (Noble Gases) have eight valance electrons (except helium which
has only two), two in the outermost occupied s–sublevel and six (helium has none) in the
outermost occupied p–sublevel. Since the noble gases already have a filled valence shell (two
valence electrons for helium and eight for the others), they are very stable and do not react
readily to form compounds. The lighter noble gases do not form compounds. KrF2, and various
compounds containing xenon with fluorine and/or oxygen have been synthesized.
Various elements combine by sharing electrons (molecular covalent bonds) or through
the transfer of electrons to form charged particles (ionic bonds) in order to more stable
substances know as compounds (compounds consist of two, or more, elements that have
combined to form a new substance).
Chemical formulas are used to represent the composition of a substance: element or
compound. There are several types of chemical formulas that each convey different types of
information about the substance. Molecular formulas show the actual number of atoms for
each element present in a molecule – the simplest unique unit of a covalent molecular compound.
An empirical formula shows the simplest ratio (representing a formula unit) of cations
(positive ions) to anions (negative ions) in an ionic compound.
Structural formulas show the arrangement of atoms within a molecule of the substance.
Naming Compounds:
1.
Binary nonmetallic compounds consist of only two elements (both nonmetals).
The name of the first element is used and the –ide suffix is added to the root-portion of
the second element’s name. Prefixes are used to indicate the number of atoms of each
element present. Typically, a prefix is not used for the first element if only one atom of
that element is present.
The prefixes are
mono- 1
di- 2
tri – 3
tetra- 4
penta – 5
hexa – 6
hepta – 7
octa – 8
nona – 9
deca – 10
sometimes the ending vowel for a prefix is omitted if the element name it is being added
to starts with a vowel
NO2 – nitrogen dioxide
N2O3 – dinitrogen trioxide
CO – carbon monoxide
CO2 – carbon dioxide
P4H10 – tetraphosphorous decahydride
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2.
Cations (positive ions)
a.
ions of metals that form only one charge have the same name as the metal
Li+ -- lithium
b.
Pb4+ -- lead (IV)
nonmetallic cation names end with –ium
NH4+ -- ammonium
3.
Al3+ -- aluminum
Roman numerals are used as part of the name for ions of metals that can
form ions with different charges to indicate the charge for that ion
Pb2+ -- lead (II)
c.
Ag+ -- silver
H3O+ -- hydronium
Anions (negative ions)
a.
monatomic (one-atom) anion names consist of the root-portion of the element
name with an –ide ending.
S2- -- sulfide
P3- -- phosphide
F- -- fluoride
Please note that a few polyatomic ions also have names ending with –ide
OH- -- hydroxide
CN- -- cyanide
O22- -- peroxide
Some anions consist of consist of multiple atoms of a single element, a prefix is
added to the anion name to indicate the number of atoms:
S22- -- disulfide
b.
I3- -- triiodide
Oxyanions -- polyatomic ions containing oxygen end with –ite or –ate
oxyanions whose name ends with –ate have more oxygen than does the –ite
PO33- -- phosphite
PO43- -- phosphate
Some oxyanions – especially those containing halogens – may have more than
two possible forms in which case the prefixes hypo- (used with the –ite suffix)
and per- (used with the –ate suffix).
ClO- -- hypochlorite
ClO2- -- chlorite
ClO3- -- chlorate
ClO4- -- perchlorate
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Some oxyanions are formed by the combination of two smaller oxyanions and the
lost of an oxygen atom:
c.
CrO42- -- chromate
Cr2O72- -- dichromate
SO32- -- sulfite
S2O52- -- disulfite
The names of anions consisting of hydrogen and an oxyanion add the word
hydrogen or dihydrogen to the oxyanion name
HSO3- -- hydrogen sulfite
d.
4.
H2PO3- -- dihydrogen phosphite
Some polyatomic ions are formed when one of the oxygen atoms in the
polyatomic ion is replaced with a sulfur atom. The prefix thio- is added to the
oxyanion’s name
SO42- -- sulfate
S2O32- -- thiosulfate
CNO- -- cyanate
SCN- -- thiocyanate
Ionic Compounds
The names for ionic compounds consist of the cation name followed by the anion name
NaI -- sodium iodide
LiNO3 – lithium nitrate
Fe2O3 – iron (III) oxide
Cr2(SO4)3 – chromium (III) sulfate
For the transition and rare earth metals, we can determine their charge from the
compound formula using the following guidelines: In a compound,
(a)
(b)
(c)
hydrogen typically has a 1+ charge
oxygen (oxide) typically has a 2- charge
the charge for other monatomic anions can be estimated based on its position on
the Periodic Table – i.e., N would form N3- (nitride) since Group 15 elements
have
(d)
(e)
(f)
five valence electrons and gain three to get an octet
polyatomic ion charges are used as shown
the sum of the individual ion charges for a compound is zero
the sum of the individual ion charges for a polyatomic ion equals the net charge
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Examples:
Fe2O3 : Let “x” represent the charge for each Fe and each oxide’s charge is 2-. Since the
formula contains two iron atoms and three oxygen atoms, the equation would be:
2x + 3(2-) = 0
thus: 2x – 6 = 0
so x = 3+
therefore each Fe has a 3+ charge and the compound name would be: iron (III) oxide
Pb(SO4)2: Let “x” represent the charge for Pb and each sulfate’s charge is 2-. Since the
formula contains lead atom and two sulfate ions, the equation would be:
x + 2(2-) = 0
thus: x – 4 = 0
so x = 4+
therefore the Pb has a 4+ charge and the compound name would be: lead (IV) sulfate
Sn3(PO3)2: Let “x” represent the charge for each of the three tin atoms. Each phosphite
ion has a 3- charge. So the equation would be:
3x + 2(3-) = 0
5.
thus 3x – 6 = 0
so x = 2+
Bases (Arrhenius) are named using the name of the cation’s name followed by the work
hydroxide.
NaOH – sodium hydroxide
Ba(OH)2 – barium hydroxide
NH4OH – ammonium hydroxide
6.
Inorganic Acid Names are formed from the compound names as follows:
compound
acid
hydrogen _______ide
hydro____ic acid
hydrogen hypo_____ite
hypo______ous acid
hydrogen _____ite
_______ous acid
hydrogen _____ate
________ic acid
hydrogen per_____ate
per______ic acid
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Examples:
compound name and formula
acid name and formula
hydrogen iodide – HI(s)
hydroiodic acid – HI(aq)
hydrogen hypoidodite – HIO(s)
hypoiodous acid – HIO(aq)
hydrogen idodite – HIO2(s)
iodous acid – HIO2(aq)
hydrogen idodate – HIO3(s)
iodic acid – HIO3(aq)
hydrogen peridodate – HIO4(s)
periodic acid – HIO4(aq)
hydrogen sulfide – H2S(g)
hydrosulfuric acid – H2S(aq)
Polymers are very large molecules that consist of long chains of repeated monomer units.
Teflon contains long chains of –CF2CF2- monomer units.
Organic compounds consist of carbon bonded to one or more other elements. Since carbon can
form four equal bond sites, numerous types of organic compounds are possible. The simplest
type is the hydrocarbons – a special class of organic compounds that contain only carbon and
hydrogen.
Functional groups are reactive groups of atoms that undergo predictable reactions. Organic
compounds may contain one, or more, functional groups: alcohols, carboxyl, ethers, ester,
amides, amines, etc.
Hydrates contain water molecules that are weakly bound to the compound’s crystal structure.
Copper (II) sulfate pentahydrate, CuSO4•5 H2O, has five water molecules “trapped” within each
formula unit. When heated, the water “escapes” leaving behind anhydrous copper (II) sulfate.
A chemical equation uses chemical symbols to represent a chemical reaction. The starting
materials (reactants or reagents) react to produce new substances (products). An arrow ()
points in the direction of the products. Phases are indicated as follows: gas (g), liquid (l), solid
(s), and aqueous solution (aq). The presence of a catalyst is shown by writing the catalyst’s
formula above the arrow. If heat must be added to start the reaction, the symbol Δ is written
above the arrow. If the reaction requires a minimum temperature, that temperature is written
above the arrow.
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Chemical equations must be “balanced” to satisfy the law of conservation of mass. Coefficients
are used to indicate the relative number of molecules of each reactant and product. As a general
“rule of thumb” I would recommend that you start by attempting to balance:
(1)
any “heavy” metals that are present,
(2)
any “light” metals that are present,
(3)
any polyatomic ions that are present,
(4)
any non-metals (except oxygen and hydrogen) that are present,
(5)
finish with any oxygen and/or hydrogen that is present
With practice, you will be able to identify patterns that will aid you in the balancing process.
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