molecular bonds (putting elements together). molar mass each atom has an atomic mass molar mass is...
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Molecular Bonds
(Putting Elements Together)
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Molar Mass
• Each atom has an atomic mass
• Molar mass is the atomic mass of all the atoms in the molecule summed together
• For Example:
H2O = 2 x Atomic Mass of H +
1 x Atomic Mass of O
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Counting Atoms in a Molecule
In the example, NH3, the subscript 3 only applies to the hydrogen.– Therefore: there is 1 N and 3 H in ammonia
In the example, 3Ca3(PO4)2, the number of atoms changes due to the Coefficient in front of the molecule
The 3 is multiplied to the Ca, P and O
The subscript 2, multiplies the P and O
3Ca3(PO4)2
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3 Ca3 ( P O4 ) 2
This means that
there are 3 x 3 Ca,
3 x 2 P and 3 x (4 x 2) O
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Bonds. . .
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No, not that kind – bonds between atoms to form molecules
It all depends upon the atom’s valence (outer shell) electrons
These are the e- in the last Energy Level (n = 1 through 7)
Figure these out using the Periodic Chart and/or Lewis Dot Diagrams
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The Roman Numerals Tell You How Many Valence Electrons for the Primary or Representative Elements;
The Valence Electrons for the Transition Elements Vary
I II III IV V VI VII VIII
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Group I is monovalent; II is divalent; III is trivalent; IV is tetravalent; V is back to being trivalent (since three e- openings); VI is divalent; VII is monovalent and VIII has a complete octet, so these seldom react or bond
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Bond Types (In General):• Pure or Non-Polar Covalent
Χ difference = 0 to 0.5 on the Pauling EN ScaleThe pair of e- shared are done so equallyTwo nonmetals bonded together
• Polar CovalentA shared pair of e-, but not equallyχ difference = 0.5 to 1.6Molecule has Partial + and – Charges
• Ionic Bondsχ difference = 1.7 or higher to the maximum of 4.0Metal bonded with a nonmetal
• Metallic Bonds are similar to Ionic Bonds
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Metallic Bonds• Two or more metals mixed are called
alloys
• Two major formats– Interstitial and Substitutional
These bonds permit the
roaming of e- which creates
a sea of dissociated e-
Called the Electron Sea
Model
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Ionic Bonds
• These are the bonds between a metal and a nonmetal
• The metal Ion is positively charged and called a cation
• The nonmetal Ion is negatively charged and called an anion
• The bonded molecule should be neutrally charged when finished
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Knowing where the metals and nonmetals are on the table will make your life easier
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Let’s take a moment to discuss polyatomic ions. . .
• This is a molecule that acts as a cation or anion
• For example:NH4
+ ammonium N3- azide
ClO4- perchlorate CN- cyanide
HCO3- bicarbonate OH- hydroxide
CrzO7-2 chromate NO3
- nitrate
ClO3- chlorate C2H3O2
- acetate
• Don’t PANIC – I gave a list to you!
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In an Ionic Bond – one or more electrons are lost or gained by the atoms involvedThis allows the atoms to have a complete valence shell – following the octet rule
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In an Ionic Compound – balance the molecule using the criss-cross rule
Mg +2 + Cl-1
Mg Cl2 The one is understood.
This applies even if using a polyatomic ion
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NH4+ + O-2
(NH4)2O The parentheses are used to keep
the polyatomic together
Pb+4 + CO3-2
Pb2 (CO3)4 and this can be simplified by
reducing the subscripts to
Pb(CO3)2
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Naming Ionic Compounds is really simple:
1. Name the cation (metal) using its proper name; if it is a polyatomic, do the same2. Then, using the stem of the anion (nonmetal), simply add the suffix “ide”
Zinc + Chlorine = Zinc ChlorideIron + Oxygen = Iron OxideLithium + Cyanide = Lithium CyanideAmmonium + Fluorine = Ammonium FluorideCobalt + Phosphorous = Cobalt Phosphide
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Transition Metals present an issue for balancing and naming molecules since they can have varying oxidation states
For example:
Manganese can be a +2 or +3
Iron can be a +2 or +3
Lead can be a +2, or even a +4
Copper is a +1 or +2
Gold is usually a +1 or +3
And Hydrogen is a +1 or a -1!
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Transition Metals
• To determine the correct Roman Numeral to place after the metal:
Roman Numeral = - (Charge # anion)(#anions)
(# cations)
This is needed because, for example,
iron chloride can be either FeCl2 or FeCl3;
or iron (II) chloride or iron (III) chloride
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Therefore – Ionic Bonds are:
Metal + Nonmetal• + ion - ion• cation anion• monatomic monatomic or
(except NH4+) polyatomic
• left of steps right of steps
Reactions are Exothermic
Form Crystal Lattice Structures
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Covalent Compounds
• These can be monatomic or polyatomic compounds
• It is a bond between two nonmetals• They share a pair of electrons• They can be subgrouped into polar or
nonpolar• If a binary compound (2 atoms) – use
the same naming rules as in Ionic Compounds
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• If it has more than two atoms – need to use the prefixes
Number Prefix Number Prefix
1 Mono 7 Hepta
2 Di 8 Octa
3 Tri 9 Nona
4 Tetra 10 Deca
5 Penta 11 Undeca
6 Hexa 12 Dodeca
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Naming Covalent Compounds
Process:
1. Prefix Indicating # + full name of first
nonmetal
2. Prefix Indicating # + root name of second nonmetal + suffix “ide”
3. Watch for polyatomics and use their proper names
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For Example:• P4S10 becomes Tetraphosphorous Decasulfide
• P2O5 Becomes Diphosphorous Pentaoxide
• SF6 becomes Sulfur Hexafluoride
• SiBr4 becomes Silicon Tetrabromide
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Covalent Bonds can be Polar or Nonpolar
A nonpolar has no discernable
negative or positively charged sides
(EN difference is 0)
A polar covalent bond means one
side is negative and the other positive
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Electronegativity Percent Ionic BondDifference Character Type• 0.2 1 % Non-polar• 0.4 4 Covalent• 0.5• --------------------------------------------------------------------------• 0.6 9• 0.8 15• 1.0 22 Polar• 1.2 30 Covalent• 1.4 39• --------------------------------------------------------------------------• 1.6 47 Ionic if metal/nonmetal• 1.8 55 Polar Cov. if non/nonmetal• 2.0 63• --------------------------------------------------------------------------• 2.2 70• 2.4 76 Pure Ionic• 2.6 82• 2.8 86• 3.0 89• 3.2 92
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• Some elements are able to form more than one oxyanion (polyatomic ions that contain oxygen), each containing a different number of oxygen atoms.
• For example, chlorine can combine with oxygen in four ways to form four different oxyanions: ClO4
-, ClO3-, ClO2
-, and ClO- (Note that in a family of oxyanions, the charge remains the same; only the number of oxygen atoms varies.)
• The most common of the chlorine oxyanions is chlorate, ClO3
-. In fact, you will generally find that the most common of an element’s oxyanions has a name with the form (root)ate.
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• The anion with one more oxygen atom than the (root)ate anion is named by putting per- at the beginning of the root and -ate at the end. For example, ClO4
- is perchlorate.
• The anion with one fewer oxygen atom than the (root)ate anion is named with -ite on the end of the root. ClO2
- is chlorite. • The anion with two less oxygen atoms
than the (root)ate anion is named by putting hypo- at the beginning of the root and -ite at the end. ClO- is hypochlorite.
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Oxyanion Example
• ClO- Hypochlorite
• ClO2-Chlorite
• ClO3-Chlorate
• ClO4-Perchlorate
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• Some compounds have common names as well as their scientific names – you should learn these and others!
– NO nitrogen monoxide nitric oxide
– H2O dihydrogen monoxide water
– NH3 nitrogen trihydride ammonia
– CH4 carbon tetrahydride methane
– C4H10 tetracarbon decahydride butane
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Some atoms are Diatomic – KNOW THESE!
H2 N2 O2 F2 Cl2 Br2 and I2
and P is usually found as P4
while Sulfur is found as S8
Other elements will bond beyond the octet rule – like PCl5, and the noble gas Xe bonds with F in XeF6, XeF2, XeF4, and XeO4 – and this is due to a thing called “hypervalence” or “expanded octet”
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Molecular Geometry
• The 3-Dimensional Shapes of Molecules depend upon the valence e-’s of the atoms involved
• Valence Bond Theory and VSEPR Model both use the same shapes– Basically – they focus on covalent bonds with
the shared bonding pairs of electrons (BP)– The assumption is made that the molecule will
adopt a geometry to minimize the repulsion between e-’s
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• The General Shapes:
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Basic Geometry Bond Angles
• Linear 180o
• Trigonal Planar 120o
• Tetrahedral 109.5o
• Trigonal Bipyramidal 90o and
120o
• Octahedral 90o
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Molecular Orbital TheoryMOT uses atomic orbitals (AO), e- λ’s and e-
density regions to examine bonds
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This is the end of Part I
• Next:– Van der Waals and London Dispersion Forces– Polarity– Intermolecular Forces– Lewis Dot Diagrams with Covalent Bonds– Determining Molecular Structure– Resonance Structures