Transcript Chemical Bonding Notes

Standards:
 C.7.A name ionic compounds containing main group transition metals,
covalent compounds, acids, and bases, using International Union of
Pure and Applied Chemistry (IUPAC) nomenclature rules
 C.7.B write the chemical formulas of common polyatomic ions, ionic
compounds containing main group or transition metals, covalent
compounds, acids
 C.7.C construct electron dot formulas to illustrate ionic and covalent
bonds
 C.7.D describe the nature of metallic bonding and apply the theory to
explain metallic properties such as thermal and electrical
conductivity, malleability, and ductility
 C.7.E predict molecular structure for molecules with linear, trigonal
planar, or tetrahedral electron pair geometries using Valence Shell
Electron Pair Repulsion (VSEPR) theory
Naming Chemical Compounds… slide 4
Naming Acids and Bases… slide 11
Writing Chemical Formulas… slide 16
Chemical Bonding… slide 22
Molecular Geometry, VSEPR Theory… slide 31
C.7.A name ionic compounds containing main
group transition metals, covalent compounds, using
International Union of Pure and Applied Chemistry
(IUPAC) nomenclature rules.
Types of Compounds
 There are three main types of compounds when
working on Naming Compounds.
 Metal Binary Compounds (Ionic) – Contain a metal
and a non-metal. They form an ionic bond.
 Non-Metal Binary Compounds (Molecular)–
Contain two non-metals. They form a covalent bond.
 Ternary Compounds (Polyatomic)– Contain
polyatomic ions. The formula will have three or more
elements in it.
Metal Binary Compounds
(Ionic)
 Name the first element. (This will always be the
metal.)
 Replace the ending on the second element with an
“ide” ending. ( This element will be the non-metal)
 Example:
 NaCl
sodium and chlorine becomes sodium
chloride
 MgS
magnesium and sulfur becomes magnesium
sulfide
Naming Compounds with a
Transition metal
 When some atoms can have more than one possible
charge, you name the charge on the atom.
 The following elements must have a roman numeral:
Cr-Cu, Au, Hg, Sn, & Pb
 Copper +1 and +2
Iron +2 and +3
Cu +1 is copper (I)
Fe +2 is iron (II)
Cu +2 is copper (II)
Fe +3 is iron (III)
CuCl is copper (I) chloride FeCl2 is iron (II) chloride
CuCl2 is copper (II) chloride FeCl3 is iron (III) chloride
Non-Metal Binary Compounds
(Molecular)
 Name the first element
 Replace the ending on the second element with “ide”
 Use Prefixes to indicate the number of atoms in the
formula.
 *Exception: A prefix is not required when the first
element only has 1 atom.
 Ex:
 CO2 carbon and oxygen is carbon dioxide
 N2O
nitrogen and oxygen is dinitrogen monoxide
Pre-fixes
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1 atom = mono
2 atoms = di
3 atoms = tri
4 atoms = tetra
5 atoms = penta
6 atoms = hexa
7 atoms = hepta
8 atoms = octa
9 atoms = nona
10 atoms = deca
Ternary Compounds (Polyatomic)
 Name the first part of the compound. Element or
polyatomic ion.
 Name the second part of the compound. Element or
polyatomic ion.
 Example:
 MgSO4
NH4OH
magnesium sulfate
ammonium hydroxide
 K3PO4
potassium phosphate
C.7.A name, acids using International Union of
Pure and Applied Chemistry (IUPAC)
nomenclature rules
Naming Acids without Oxygen
 Acids without Oxygen are named with the prefix
“hydro” and end in “ic”
 Examples:
 HCl
hydrochloric acid
 HF
hydrofluoric acid
 HBr
hydrobromic acid
Naming Acids with Oxygen
 Acids with oxygen have several forms.
 The “ic” or regular ending for an acid comes from the
polyatomic ion with the “ate” ending. This gives the
regular count for the oxygen for this type of acid.
 Example:
 H2SO4
 SO4 is sulfate so this acid is called sulfuric acid
 Once you know the “ic” ending you count the number
of oxygen in the other forms to find the name for the
acid. (REMEMBER: The regular “ic” form comes from
the polyatomic ion that ends with “ate”)
 Two less oxygen hypo ________ ous acid
 One less oxygen
________ ous acid
 Regular “ic” form
________ ic acid
 One more oxygen per ________ ic acid
 The other names for the acids will come from the
count based from the “regular acid name”
 H2SO4 -ate ending so it is sulfuric acid
 H2SO3 -ite ending so it is sulfurous acid
 H2SO2 two less oxygen will have the prefix hypo and
the –ous ending. hyposulfurous acid.
 H2SO5 one more oxygen will have a prefix per and the
regular -ic ending. persulfuric acid
C.7.B write the chemical formulas of common
polyatomic ions, ionic compounds containing main
group or transition metals, covalent compounds, acids
Writing Formulas:
Ionic Compounds
 Write chemical symbol for each part of the compound.
 Write the charge (oxidation #) for the element.
Do the charges add together and equal zero?
 Yes, Stop this is the formula. The number of electrons
given away is the same as what is being taken by the
second atom.
 No, Cross the absolute value of the charge to the
opposite element as a subscript. Multiply the new
subscript by the charge and see if the new values will
add together and equal zero. If yes, Stop you have the
formula
 potassium bromide
K +1
Br -1
+1 + -1 = 0 Yes
Formula
KBr
 magnesium chloride
Mg +2
Mg 1
Cl -1 +2 + -1 = +1 No
Cl 2 Mg (1 x +2)= +2 Cl (2 x -1)= -2
Yes MgCl2
Transition Elements
 Same rules as normal ionic compounds. The charge for
the transition metal will come from the name of the
compound.
 iron (III) chloride
 Fe +3 Cl -1 +3 + -1 = +2 No
 Fe1
Cl 3 Fe (1 x +3) +3 Cl (3 x -1) -3
Yes
FeCl3
Molecular Compounds
 Use the prefix to determine the subscript of each
element in the formula.
 NO PREFIX on the first element indicates a subscript
of 1
 Write the correct formula using the correct symbol
and subscript for each element.
 Ex: carbon dixoide
CO2
Polyatomic Ions
 The rules for polyatomic ions will be the same as ionic
compounds.
 *Polyatomic ions must be placed in parenthesis if the
subscript is larger than 1 when criss-crossing.
 magnesium sulfate
 Mg +2
SO4 -2
MgSO4
 iron (III) phosphate
 Fe +3
PO4 -3
FePO4
 sodium hydroxide
 Na +1
OH -1
NaOH
 calcium hydroxide
 Ca +2
OH -1
aluminum phosphate
Ca(OH)2
Acids without Oxygen
 Write the symbol and charge (oxidation #) of each element.
 If the charges do not add up to zero, criss-cross the oxidation #.
 Ex: hydrosulfuric acid
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H+1 S-2
= H2S
Acids with Oxygen
 Write the symbol and the charge for the polyatomic ion (oxyanion).
 If the charges do not add up to zero, criss-cross the oxidation numbers.
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1 more oxygen
MEMORIZED(-ate)
1 less oxygen (-ite)
2 less oxygen
H2SO3
sulfurous acid
per ____________ic
____________ic
____________ous
hypo___________ous
C.7.C construct electron dot formulas to
illustrate ionic and covalent bonds
 There are three main types of chemical bonding.
ionic, covalent, and metallic.
 Ionic bonding occurs when there is a transfer of
electrons.
 Covalent bonding occurs when atoms share electrons.
 Metallic bonding consist of the attraction of free
floating valance electrons for positively charged metal
ions.
 Electro negativities are used to determine what type of
bond is formed when atoms come together in a
chemical reaction.
 To find the type of bond find the difference in the
electronegativities.
 If the difference is greater than 1.67 an ionic bond is
formed.
 If the difference is less than 1.67 a covalent bond is
formed.
 All atoms want to obtain eight electrons in the valence
energy level. To do so they will give, take, or share
electrons.
Rules for Ionic Bonds
 The element with the fewest atoms goes in the center.
 The other atoms go around the central atom.
 Show the transfer of the electrons with a positive for
the atom that lost the electrons and a negative for the
atoms that gain the electrons.
 NaCl sodium chloride
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sodium: (1.01)
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Na: 1s22s22p63s1
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chlorine: (2.83)
Cl: 1s22s22p63s23p5
Sodium transfers the 3s1 to chlorine to complete
the 3p energy level.
The electronegativity difference is 1.72
An ionic bond is formed.
Rules for showing Covalent Bonds
 The element with the fewest atoms goes in the center.
 The other elements go around the central atom.
 A bonding pair can only form where there is an
unpaired electron.
 Shared pairs or bonding pairs are shown with a dash.
One dash equals two electrons.
 AsI3
arsenic triiodide
arsenic (2.20)
iodine (2.21)
As: 1s22s22p63s23p64s23d104p3
I: 1s22s22p63s23p64s23d104p65s24d105p64d105p5
The electronegativity difference is .01
A covalent bond is formed. The atoms share the
electrons.
C.7.E predict molecular structure for molecules with linear,
trigonal planar, or tetrahedral electron pair geometries
using Valence Shell Electron Pair Repulsion (VSEPR) theory
Molecular Geometry
 The shape that a covalently bonded substance will take
is referred to as its Molecular Geometry.
 The shape is determined by the central atom, and the
number of shared and unshared electron pairs around
the atom.
 Electron pairs around the central atom will spread out
as far as possible to minimize the repulsive forces.
 This gives bond angles depending on the shape.
Total number Number of
of electron
shared pairs
pairs.
2
2
Number of
unshared
pairs
0
Shape
Linear
Bond Angle
180 0
Total number Number of
of electron
shared pairs
pairs.
3
3
Number of
unshared
pairs
0
Shape
Trigonal
Planar
Bond Angle
120 0
Total number Number of
of electron
shared pairs
pairs.
4
4
Number of
unshared
pairs
0
Shape
Tetrahedral
Bond Angle
109.5 0
Total number Number of
of electron
shared pairs
pairs.
4
3
Number of
unshared
pairs
1
Shape
Trigonal
Pyramidal
Bond Angle
107.3 0
Total number Number of
of electron
shared pairs
pairs.
4
2
Number of
unshared
pairs
2
Shape
Bent
Bond Angle
104.5 0
Linear
Tetrahedral
Trigonal Planar
Trigonal Pyramidal
Bent