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IONIC BONDING
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CHAPTER OPENER
Final Grade Review on Smartboard
Seating Change
Final Book Accountability
Tell them about syllabus change
Go over new lunches for classes
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Why does hard water bubble
less than distilled water?
Minerals decrease the
effectiveness of the detergent.
Atoms and Ions
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Atoms are electrically neutral.
Because there is the same number of
protons and electrons.
Ions are atoms, or groups of atoms, with
a charge (positive or negative)
ionic compounds consist of a combination of
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cations and anions
• the sum of the charges on the cation(s) and anion(s) in each
formula unit must equal zero
The ionic compound NaCl
They have different numbers of protons and
electrons.
Only electrons can move, and ions are
made by gaining or losing electrons.
2.6
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An Anion is…
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A Cation is…
• A negative ion.
• Has gained electrons.
• Nonmetals can gain electrons.
A positive
ion.
by losing electrons.
More protons than electrons.
Metals can lose electrons
Formed
• Charge is written as a superscript on the
right.
F1-
Has gained one electron (-ide
is new ending = fluoride)
O2-
Gained two electrons (oxide)
K1+
Ca2+
Has lost one electron (no
name change for positive ions)
Has lost two electrons
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Common Multiple 2
F1-
Number Line
-4 -3 -2 -1 0 +1 +2 +3 +4
Ca2+
Penny Example
Breathing Example
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Common Multiple 2
Ca2+
2(F1-)
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WHAT DETERMINES A CATION
OR AN ANION
• THE PATH OF LEAST RESISTANCE TO
BEING NOBLE (AKA LESS ENERGY)
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•Why do elements form compounds in the first place?
Bonding lowers the potential energy between positive and
negative particles. BONDING BALANCES THINGS
• There are two major types of bonding between
elements:
1. ionic bonds —result from a transfer of
electrons from one species (usually a metal) to
another (usually a nonmetal or polyatomic ion).
• We will use NaCl (salt) as our main
example of an ionic reaction
2. covalent bonds —result from the sharing of
electrons by two or more atoms (usually
nonmetals).
– Seawater is 3% NaCl
IONIC BONDS CLIP
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Number Line
-4 -3 -2 -1 0 +1 +2 +3 +4
Teeter Totter Example
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Na
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Octet Rule (Rule of 8)
• When 2 or more atoms combine, they
tend to get a complete, outermost shell
with 8 electrons.
• Atoms try to fill valence shell by gaining,
losing, or sharing electrons during
reactions.
- - - - - - - - - - - - - - - -- - - - - - - - - - - - - - - -
Cl
In a reaction, sodium loses an electron, has
unequal numbers of electrons and protons.
In a reaction, chlorine gains an electron, so it now
now has one more electron than proton.
Because it now has one more proton than electron,
it has a charge of 1+.
It therefore has a charge of 1-
If electrons are lost from an atom, positive ions are
formed.
It takes Ionization Energy To Make This Happen
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If electrons are gained by an atom, negative ions
are formed
A negative ions are called anions.
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CRYSTAL LATTICE
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FORMATION OF NaCl CLIP
Ionic Bonding
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• The oppositely charged ions
• attracted to each other by electrostatic forces
• Joined by an ionic bond (ions)3-D structure
called a crystal (cations surrounded by anions –v v)
Each Na+ and Cl– ion has 6 of the oppositely
charged ions clustered around it .
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Naming Overview
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Ionic Bonds Easily Ionize in Water
• Cation gets named first (keeps name)
– If it is a transition metal it needs roman numeral
• Anion gets named second
Shortcut to Hydrolosis of Salt.lnk
– Add “ide” to the end of the anion name
– Polyatomic always keep their name.
Ions Increase Electrical Conductivity
HYDROLOSIS OF WATER
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Lattice Energy
Properties of Ionic Compounds
• ionic compound, such as rock salt (NaCl):
• Lattice Energy is how much required to
separate the ionic bonds.
1. hard (doesn’t dent)
2. rigid (doesn’t bend)
3. brittle (cracks without deforming)
• Water gives in energy (endothermic) and then
reaction gives off energy (exothermic)
• Overall reaction is exothermic
ASSNT DIS IONIC BONDING
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Why are ionic solids brittle?
Strong
repulsive
forces
IONIC LATTICE STRUCTURE
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• Ionic lattices are extremely difficult
structures to break apart. As a result, all ionic
substances are solids with high melting
points.
• Potassium iodide, magnesium chloride and
calcium oxide. All of these ionic compounds have
melting points over 500 °C.
• There is one method of breaking up the lattice dissolve the ionic compound in water. Water
has the ability (polarity) to separate the ions
from the lattice and allow them to move freely as
a solution.
(it poles them apart. JOD)
IONIC BONDS
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STRONG in crystal
BROKEN
in H2O
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Ionic Bonds Are the Reaction of a
Metal and a Nonmetal
How Do We Tell The Difference?
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IONIC BONDS DO NOT
INCLUDE METALLOIDS
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The elements in the middle, metalloids, exhibit
some metal-like behavior but also form
covalent-like bonds. Hydrogen tends to behave
more like a nonmetal in bonding except at high
pressure at which it actually assumes a phase
with the properties of a liquid metal. Since we
don't deal with such high pressures normally,
hydrogen is often just considered a nonmetal.
Ionic bonding:
• NaCl
• MgO
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•
•
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Transfer of electrons
Network of cations and anions
Metals (left) with non-metals (right)
Atoms achieve pseudo-Noble configurations
Strong interactions due to large electrostatic
forces (+/-)
• Lattice energy: quantitative description of
binding energy in ionic compounds
• They (NaCl, MgO, ect.) are not molecules
but lattice Structures
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Ionic compounds:
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Sodium Chloride
Sodium gives away electron to become more
stable.
Chlorine has 1 electron short of a stable noble gas
structure, so it takes it.
• CaCl
• K2O
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Magnesium Oxide
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Calcium Chloride
•MgO ionic bonding is stronger than NaCl
•Due to 2+ ions attracting 2- ions.
Greater charge = greater attraction
•Two chlorines to use up 2 electrons in the calcium.
•The formula of calcium chloride is therefore CaCl2.
•Ionic Bonds (Lattice Energy) Similar to NaCl
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Potassium Oxide
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MP/BP is proportional to the lattice bond strength
Higher bond strength = Higher Temperature
•Two Potassium donate 2 electrons to one oxygen.
•The formula of calcium chloride is therefore K2O.
•Ionic Bonds (Lattice Energy) Similar to NaCl
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Predicting Ionic Charges
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Group 1A:
1A: Lose 1 electron to form 1+ ions
H1+ Li1+
Na1+ K1+ Rb1+
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Predicting Ionic Charges
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Group 2A:
2A: Loses 2 electrons to form 2+ ions
Predicting Ionic Charges
B3+
Al3+
Ga3+
Be 2+ Mg2+ Ca2+ Sr2+ Ba2+
Predicting Ionic Charges
Neither! Group 4A
elements rarely form
ions.
Oxide
S2- Sulfide
Se 2- Selenide
Group 6A:
6A: Gains 2
electrons to form
2- ions
Nitride
Phosphide
As3- Arsenide
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Group 3A:
3A: Loses 3
electrons to form
3+ ions
Predicting Ionic Charges
N3 P3-
Group 4A:
4A: Lose 4
electrons or gain
4 electrons?
Predicting Ionic Charges
O2-
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Group 5A:
5A: Gains 3
electrons to form
3- ions
Predicting Ionic Charges
F1- Fluoride
Cl1- Chloride
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Group 7A:
7A: Gains
Br1- Bromide 1 electron to form
I1- Iodide
1- ions
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Predicting Ionic Charges
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Formula of Ionic Compounds
2 x +3 = +6
Group 8A:
8A: Stable
noble gases do not
form ions!
Al3+
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3 x -2 = -6
Al2O3
O2-
ALUMINUM OXIDE
1 x +2 = +2
2 x -1 = -2
CaBr2
Ca2+
Br-
CALCIUM BROMIDE
+1 x 2 = +2
Na+
1 x -2 = -2
Na2S
S2-
SODIUM SULFIDE
Write Formulas for these:
• Aluminum Oxide
Al3+
O2-
Al2O3
• Cesium Sulfide
Cs+
S2-
Cs2S
• Cobalt (II) Oxide
Co2+
O2-
CoO
• Barium Chloride
Ba 2+
Cl-
BaCl2
• Boron Sulfide
B3+
S2-
B2S3
• Lithium Selenide
Li+
Se2-
Li2Se
Predicting Ionic Charges
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Predicting Ionic Charges
2.6
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Group B elements:
elements: Many transition elements
have more than one possible oxidation state.
Note the use of Roman
Iron (II) = Fe 2+
numerals to show charges
Iron (III) = Fe 3+
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Some of the postpost-transition elements also
have more than one possible oxidation state.
Tin (II) = Sn2+ Lead (II) = Pb 2+ Bismuth (II) =Bi2+
Tin (IV)=Sn4+ Lead (IV)= Pb 4+ Bismuth (V)=Bi5+
Predicting Ionic Charges
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Group B elements:
elements:Some transition elements
have only one possible oxidation state, such
as these three that are always:
Silver = Ag1+
Zinc = Zn2+ Cadmium = Cd2+
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Exceptions YOU HAVE TO
MEMORIZE:
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• Some of the transition metals have
only one ionic charge:
–Do not use roman numerals for
these :
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Naming Ionic Compounds
Binary Ionic Compounds
Simplest Ionic compounds: Metal + Non-metal
• Silver is always 1+ (Ag1+)
• Cadmium and Zinc are always 2+
(Cd2+ and Zn 2+)
Polyatomic Ionic Compounds
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Formulas and Names of Some Polyatomic Ions
NH4+
ammonium
CO3 2– carbonate
H3O+
hydronium
HCO 3– hydrogen carbonate (bicarbonate)
OH–
hydroxide
OCN– cyanate
cyanide
SCN– thiocyanate
CN–
O22peroxide
S2O3 2– thiosulfate
N3azide
CrO4 2– chromate
NO 2– nitrite
Cr2O7 2– dichromate
NO 3– nitrate
SO4 2– sulfate
ClO–
hypochlorite
SO32– sulfite
HSO4 – hydrogen sulfate (bisulfate)
ClO2 – chlorite
ClO3– chlorate
PO4 3– phosphate
ClO 4– perchlorate
HPO4 2– monohydrogen phosphate
MnO4 – permanganate H2PO4– dihydrogen phosphate
C2H3O 2– acetate (OAc-) HSO3 – hydrogen sulfite (bisulfite)
C2O4 2– oxalate
Writing formulas for binary ionic compounds.
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Classifying Compounds
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Naming Ionic Compounds
Ionic compounds are named simply by naming the ions present.
Is there a metal or a polyatomic ion present?
There are, however, two complicating factors:
I. Some metals form more than one ion.
II. Identifying polyatomic ions
If the answer is yes, use the system for naming ionic
compounds.
If the answer is no, use the system for naming covalent
compounds.
I. Metals that form more than one ion, such as iron, add a
Roman numeral to the name to indicate the charge:
Fe 2+ is called iron (II) and Fe 3+ is called iron (III)
Assume a Roman numeral is required for any metal except
1. metals in groups IA and IIA on the periodic table
2. aluminum, cadmium, silver, and zinc
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Naming cations
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Practice by naming these:
• We will use the Stock system.
• Cation - if the charge is always the
same (like in the Group 1&2 metals)
just write name of the metal.
• Transition metals can have more
than one type of charge.
•
•
•
•
•
•
•
– Indicate their charge with roman
numerals in parenthesis after the
name of the metal
Write Formulas for these:
• Potassium ion
• Magnesium ion
• Copper (II) ion
• Chromium (IV) ion
• Barium ion
• Mercury (II) ion
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Mg 2+
Cu 2+
Cr 4+
Ba 2+
Hg 2+
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• Cl • N 3• Br • O 2-
Chloride Ion
Nitride Ion
Bromide Ion
Oxide Ion
Na+
Cd2+
Ag+
Fe3+
Fe2+
Pb2+
Zn 2+
Sodium Ion
Cadmium Ion
Silver Ion
Iron (III) Ion
Iron (II) Ion
Lead (II) Ion
Zinc Ion
Naming Anions
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• Anions are always the same
charge
• Change the monatomic
element ending to – ide
• F- a Fluorine atom
becomes a Fluoride ion.
K+
Practice by naming these:
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Write symbols for these:
• Sulfide ion
• Iodide ion
• Phosphide ion
• Selenide ion
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S2IP3Se2-
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Write Formulas for these:
• Potassium Bromide
K+
Br-
KBr
• Magnesium Nitride
Mg2+
N 3-
Mg3N2
• Copper (II) Chloride
Cu 2+
Cl -
CuCl2
• Chromium (IV) Oxide
Cr 4+
CrO2
• Barium Sulfide
Ba2+
O2S2-
• Mercury(II) Fluoride
Hg2+
F-
HgF2
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Naming Binary ionic compounds.
BaS
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Formula
Reasoning
Cl has a 1- charge, and there are 2 of them for a total
of 2-, so the Fe must be 2+
iron (II)
chloride
Fe2O3
O has a 2- charge, and there are 3 of them for a total
of 6-, so the Fe must have a total charge of 6+ split
equally between the two iron atoms, so each must
have a 3+ charge
iron (III) oxide
PbS2
S has a 2- charge, and there are 2 of them for a total
of 4-, so the Pb must be 4+
lead (IV)
sulfide
Cu3N
N has a 3- charge, so the Cu must have a total
charge of 3+ split equally between the 3 copper
atoms, so each must have a 1+ charge
MONOATOMIC NAMING RACE
Name
FeCl2
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1) copper (II) chloride
2) sodium hydroxide
3) lithium oxide
4) cobalt (III) chloride
5) aluminum sulfide
6) francium bromide
7) iron (III) phosphide
8) vanadium (V) nitride
9) calcium iodide
10) manganese (III) fluoride
copper (I)
nitride
CuCl2
NaOH
Li2O
CoCl3
Al2S3
FrBr
FeP
CaI2
MnF3
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QUESTIONS ON BINARY IONIC
BONDING
Writing formulas for polyatomic ionic
compounds.
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Polyatomic ions are…
• Groups of atoms that stay together and
have an overall charge, and one name.
• Usually end in –ate or -ite
• Nitrite: NO21• Permanganate: MnO41• Hydroxide: OH1- and Cyanide: CN1-?
-orhypo
(One of the few positive
polyatomic ions)
If the polyatomic ion begins with H, then combine the
word hydrogen with the other polyatomic ion present:
H1+ + CO32- →
HCO31hydrogen + carbonate → hydrogen carbonate ion
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-ate ions
suffix
3A
per-
• Phosphite: PO33• Ammonium: NH41+
• Carbonate: CO32• Chromate: CrO42• Dichromate: Cr2O72-
• Nitrate: NO31-
Nomenclature of
Polyatomic Oxy Anions
• Phosphate: PO43-
• Sulfate: SO42• Sulfite: SO32-
• Acetate: C2H3O21-
prefix
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-ate
+ Root +
identifies the
element other
than O (or H)
4A
CO32carbonate
-or-
-ite
5A
6A
7A
SO42-
ClO31-
sulfate
chlorate
NO31nitrate
PO43phosphate
contains 1 fewer
O atoms than does
the -ate ion
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-ite ions
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Prefixes
3A
4A
CO22-
5A
NO21-
carbonite
nitrite
PO33phosphite
6A
7A
SO32- ClO21sulfite
chlorite
per: contains 1 more O atom than
the - ate ion
hypo: contains 1 fewer O atoms
than the - ite ion
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Prefixes
per chlorate
ClO41-
Writing Ionic Compound
Formulas
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Barium nitrate
ClO3
1-
chlorate
ClO1-
Ba2+ ( NO3-)
chlorite
ClO21-
Now balanced.
Not balanced!
hypo chlorite
= Ba(NO3)2
Writing Ionic Compound
Formulas
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Writing Ionic Compound
Formulas
Aluminum phosphate
Calcium sulfate formula?
Ca2+
(SO4
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Al3+ (PO4)3-
)2-
They ARE balanced!
Now balanced.
= AlPO4
= Ca (SO4)
Stopped here for CHEM 1
Naming Ionic Compounds w/ Polyatomic Ions
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A few of the more common polyatomic ions
polyatomic cation polyatomic anion
MgSO4
1. magnesium sulfate
Formula
Name
C2H3O21-
acetate
or polyatomic cation anion root + ide
2. ammonium chloride
NH4Cl
CO32-
carbonate
HCO31-
bicarbonate
NH41+
ammonium
NO31-
nitrate
OH1-
hydroxide
PO43-
phosphate
SO42-
sulfate
or cation polyatomic anion
3. ammonium sulfate
(NH4)2SO4
C2H3O21CO32HCO31NH41+
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acetate
carbonate
bicarbonate
ammonium
NO31-
nitrate
OH1-
hydroxide
PO43-
phosphate
SO42-
sulfate
* Groups I & II, Al, Zn, Cd,
and Ag need no Roman
numeral.
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Na2SO4
sodium sulfate
Fe(NO3)2
iron (II) nitrate
AlCl3
aluminum chloride
PbI4
lead (IV) iodide
(NH4)3PO4
ammonium phosphate
Mg3N2
magnesium nitride
AgC2H3O2
silver acetate
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POLYATOMIC NAMING RACE
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POLYATOMIC NAMING RACE
scandium hydroxide
POLYATOMIC NAMING RACE
1)
2)
3)
4)
5)
6)
7)
8)
9)
10)
sodium nitrite
scandium (III) hydroxide
vanadium (III) sulfate
ammonium fluoride
calcium carbonate
nickel (III) phosphate
lithium sulfite
zinc phosphate
strontium acetate
copper (I) oxide
Be(NO3)2
Ni2(SO3)3
K NO2
Zn3(PO4)2
Al2(SO4)3
FrOH
Fe(PO3)
(Bi)3(PO4)5
Ca(MnO4)2
1) beryllium nitrate
2) nickel (III) sulfite
3) potassium nitrite
4) zinc phosphate
5) aluminum sulfate
6) francium hydroxide
7) iron (III) phosphite
8) bismuth (V) phosphate
9) calcium permanganate
10) manganese (III) hypochlorite
Na (NO2)
V2(SO4)3
NH4F
CaCO3
NiPO4
Li2SO3
Zn3(PO4)3
Sr(C2H3O2)2
Cu2O
1.
Ag3PO4
silver phosphate
2.
YClO3
Yttrium (I) chlorate
3.
SnS2
tin (IV) sulfide
Ti(CN)4
titanium (IV) cyanide
5.
KMnO4
potassium permanganate
6.
Pb3N2
7.
CoCO3
lead (II) nitride
cobalt (II) carbonate
8.
CdSO3
cadmium sulfite
9.
Cu(NO2)2
copper (II) nitrite
10.
Fe(HCO3)2
iron (II) bicarbonate
4.
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CuC2H2O2
Copper (I) Acetate
Cobalt (II) Cyanide
Co(CN)3
NaHCO3
Iron (III) Oxide
Fe2O3
Ammonium Nitrate
CaCO3
LiH2PO4
Sodium Hydrogen
Carbonate (Bicarbonate)
Calcium Carbonate
Lithium Dihydrogen
phosphate
Iron (II) Phosphite
Fe3(PO3)2
(NH4)3PO4
ammonium phosphate
Potassium Acetate
KC2H2O2
Mg3N2
magnesium nitride
Zinc Phosphate
Zn3(PO4)2
AgC2H3O2
silver acetate
silver acetate
CaCO3
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AgC2H3O2
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Naming Acids
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• What the book doesn’t teach, yet is
essential for future success in Chem:
•
acid: Any compound with Hydrogen as
a cation: HX (or HA if you like)
• H+ = hydrogen ion, X = any negative
ion
• Naming is determined by the original
name of X (anion)
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If X- ends in -ide --- acid name
becomes Hydro-X-ic acid
HCl Hydrochloric HBr Hydrobromic HF HydroFlouric
Acid
Acid
Acid
11)H3PO4
Phosphoric Acid
12)HClO2
Chlorous Acid
12)HNO2
Nitrous Acid
13)HC2H3O2
Acetic Acid
15)H3PO3
Phosphorous Acid
16)H2CrO4
17)HCO3
Chromic Acid
Carbonic Acid
18) H2S
Hydrosulfic Acid
19) HNO3
Nitric Acid
20) HI
Hydroiotic Acid
Acid
Acid
3. if X- ends in -ite --- acid name becomes
X-ous acid
HNO2Nitrous H2SO3 Sulfurous HCl02 Chlorous
Acid
POLYATOMIC NAMING RACE
Acid
2. if X- ends in -ate --- acid name
becomes X-ic acid ( No Hydro-)
HNO3 Nitric H2SO4 Sulfuric
HClO3 Chloric
Acid
Acid
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