Transcript Formula writing

Formula Writing and
Nomenclature
What is an ion?
charged particle.
 An ion is a ______________.
+
 It may be a ____
or ___charge.
 Lose electrons  cation (+)
 Gain electrons  anion (-)
Why do atoms gain or lose
electrons?
stable
 To become more ______.
other atoms
 Electrons come from ___________.
loses
 One ion ______
e- and the other
gains
bond
_____eto form a _____.
Do ions follow rules? YES!
The Rule of 8
 Octet Rule: _____________
 Tendency of valence electrons to
stable/full valence shell
rearrange to form a ________________
.
Happy Ion
 THE MAGIC NUMBER=________
Examples:
 Na
F
+1
-1
2-8 -1
2-78
Opposites attract! NaF
Do ions follow rules? YES!
The Rule of 2
 Duet Rule:___________
 For atoms so small their valence shell is
first
the ______
energy level which can only
two electrons
hold ___________.
Examples:
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H can gain ___ e- to form _____.
1
H
H can lose ____
e- to form_____.
1
Li
Li loses _____
e- to form _____.
2
Be
Be loses ____
e- to form _____.
3
B
B loses _____
e- to form _____.
H-1
1
+1
+1
+2
+3
Writing Formulas
 All compounds have a charge of zero.
 When writing formulas, all ions have to
add up to zero.
IUPAC
 IUPAC- International Union of Pure and
Applied Chemists (created this naming
system)
Forming Binary Ionic
Compounds
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Binary = two elements
Ionic = one metal and one nonmetal
Transfer of electrons
Not called molecules!
Writing Binary Ionics
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Write each ion (metal first)
Crisscross the charges
Drop the + and –
Write numbers as subscripts
Binary Ionic Compounds
 Example:
magnesium chloride
+2
Mg
-1
Cl
MgCl2
Why does this work?
MgCl2 means…
+2
Mg
-1
Cl
-1
Cl
Binary Ionic Compounds
 Example:
barium oxide
+2
Ba
-2
O
BaO
Ba2O2
Practice
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Sodium bromide NaBr
Calcium fluoride CaF2
Magnesium oxide MgO
Lithium oxide Li2O
Aluminum oxide Al2O3
Magnesium fluoride MgF2
Potassium iodide KI
Aluminum sulfide Al2S3
Forming Ternary Ionics
 Contain 3 or more elements
 Combination of a metal and a polyatomic
ion (Table E)
Writing Ternary Ionics
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Write each ion (positive first)
Crisscross the charges
Drop the + and –
Write numbers as subscripts
Keep polyatomic ions in parentheses if
more than 1
 Never change a polyatomic ion!!!!!
Ternary Ionic Compounds
 Example:
ammonium chloride
NH4
+1
-1
Cl
NH4Cl
Ternary Ionic Compounds
 Example:
lithium carbonate
+1
Li
CO3
-2
Li2CO3
Ternary Ionic Compounds
 Example:
calcium hydroxide
+2
Ca
-1
OH
Ca(OH)2
Why does this work?
Ca(OH)2 means…
+2
Ca
-1
OH
-1
OH
Practice
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Magnesium hydroxide Mg(OH)
Potassium sulfate K SO
Sodium phosphate Na PO
Calcium nitrate Ca(NO )
2
2
4
3
3 2
4
Stock System for Ionics
 Some metals can have more than one
oxidation state (i.e. transition metals)
 Use roman numerals
Examples:
Copper (I) chloride
Cu+1
Cl-1 CuCl
Copper (II) chloride Cu+2 Cl-1
CuCl2
Practice
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Mercury (II) oxide HgO
Vanadium (V) bromide
Copper (I) oxide Cu O
Tin (IV) bromide SnBr
2
4
VBr5
Forming Molecular
Compounds
 Composed of two non-metals
 Electrons are shared so no ions are
formed (covalent bonding).
 Called molecules
 Prefix system- tells you how many atoms
of each element
Prefixes
prefix
MonoDiTriTetraPentaHexaHeptaOctaNonaDeca-
meaning
1
2
3
4
5
6
7
8
9
10
Molecular Compounds
 Examples
Carbon monoxide
Carbon dioxide
CO
CO2
Phosphorous trichloride
PCl3
Phosphorous pentachloride
Carbon tetrachloride
CCl4
PCl5
Practice
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Phosphorous trichloride
Dichlorine monoxide Cl O
Sulfur tetrafluoride SF
Dinitrogen trioxide N O
Iodine monochloride ICl
2
4
2
3
PCl3
Naming Ionic Compounds
Ionic Compounds (formula units)- Metal and non-metal
1.
2.
Write cation name first (use roman
numerals if more than one oxidation state).
Write the first syllable of the anion and add
–ide. OR just name the polyatomic ion.
Examples: Ionics
 LiBr lithium bromide
 Na2SO4 sodium sulfate
 CuCl2 copper (II) chloride
Naming Molecular Cmpds
Molecular Compounds (molecules)- two non-metals
1.
2.
Use prefix system on first element (except
Mono).
Use prefix system on the second element
(including mono) and add –ide ending.
Examples: Molecular
 N 2O 5
 CO
 PCl3
Dinitrogen pentoxide
Carbon monoxide
Phosporous trichloride
Empirical Formulas vs.
Molecular Formulas
 Compounds exist with a definite ratio of
atoms (ex: water has 2 H per 1 O)
 Empirical formula: lowest whole number
ratio
 Molecular formula: actual formula (can be
empirical also).
Examples
 C2H8 is a molecular formula (can be
reduced).
 CH4 is an empirical formula (can’t be
reduced).