Sorts of substance bonds

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Hydrogen fluoride is a sample of an atom that has bond extremity. ... Reverberation Structures: More than one Lewis Structure can be drawn for an atom. ...

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Slide 1

Sorts of compound bonds Bond : Force that holds gatherings of at least two molecules together and makes the iotas work as a unit. Case: H-O-H Bond Energy : Energy required to break a bond. Ionic Bond : Attractions between oppositely charged particles. Case: Na + Cl -

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Types of concoction bonds Ionic Compound : A compound coming about because of a positive particle (normally a metal) joining with a negative particle (as a rule a non-metal). Case: M + X -  MX Covalent Bond : Electrons are shared by cores. Illustration: H-H Polar Covalent Bond : Unequal sharing of electrons by cores. Illustration: H-F Hydrogen fluoride is a case of a particle that has bond extremity.

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Lewis structures Lewis Structure : Representation of a particle that shows how the valence electrons are orchestrated among the iotas in the atom. Holding includes the valence electrons of iotas. Case: Na ● H-H

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Lewis structures of components Dots around essential image Symbolize valence electrons Thus, one must know valence electron setup

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Lewis Structures of particles Single Bond : Two molecules sharing one electron combine. Illustration: H 2 Double Bond : Two particles sharing two sets of electrons. Illustration: O 2 Triple Bond : Two iotas sharing three sets of electrons. Case: N 2 Resonance Structures : More than one Lewis Structure can be drawn for a particle. Illustration: O 3

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Rules for Lewis structures of particles Write out valence electrons for every iota Connect solitary electrons in light of the fact that solitary electrons are destabilizing Become two shared electrons Called a "bond" Check to check whether octet administer is fulfilled Recall electron arrangement looking like respectable gas as it were, there must be 8 electrons (fortified or non-reinforced) around particle Non-fortified electron-match Called "solitary combine"

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Let's do a few cases on the board H 2 Duet lead F 2 Octet govern O 2 N 2

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Lewis structures Example Write the Lewis Structure for the accompanying atoms: H 2 O CCl 4 Where does the carbon go & why? PH 3 H 2 Se C 2 H 6

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Lewis structures proceeded with CO 2 C 2 H 4 C 2 H 2 SiO 2

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Polyatomic particles If positive charge on particle Take away electron from focal species If negative charge on particle Add electron to focal species Example: H 3 O +

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Your turn NH 4 + ClO - OH -

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Resonance structures When structures can be composed in more than one way O 3 Actual atom is "in the middle of" Resonance half and half Another case HCO 3 - What might its reverberation cross breed resemble?

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Practice NO 2 - NO 3 -

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Formal Charge figured on iota in light of Lewis structure Yields best Lewis structure of contenders FC = VE - [LE + ½(BE)] Rules: Sum of all FC's must equivalent to charge on species, if any Smaller or zero FC's on iotas superior to anything substantial FC's Negative FC ought to be on most electronegative species

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Examples HBr FC on H = 1-[0 + ½ (2)] = 0 FC on Br = 7 – [6+ ½ (2)] = 0 Net whole of FC's = charge on particle = 0 OH - FC on O = 6 – [6 + ½(2)] = - 1 FC on H = 1 – [0 + ½(2)] = 0 Net total of FC's = charge on particle = - 1

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Practice H 2 O 2 H 3 O +

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Aberrant mixes Odd-electron species NO 2

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Aberrant mixes Incomplete octet BH 3

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Aberrant mixes Expanded octet Some focal molecules can surpass an octet Third period and higher components can do this E.g., Al, Si, P, S, Cl, As, Br, Xe, and so forth d-orbitals can oblige additional electrons

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Examples AsI 5 XeF 2

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Practice SCl 6 XeF 4

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Aberrant mixes Write this out: SO 4 2-Can we decrease the formal charges? Assuming this is the case, how? We can likewise locate the normal FC Let's investigate

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Aberrant intensifies Its OK to grow the octet for those particles that can take it with a specific end goal to lower FC's

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Practice SO 3 2-PO 3-SO 2 SO 3 H 2 SeO 4

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Electroneutrality rule Electrons appropriated so that charges on iotas are nearest to zero If "- " charge present, ought to be on most electronegative iota (in this way, "+" charge ought to be on slightest electronegative iota) Good to decide which reverberation structure is best Example: OCN -

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Electronegativity : The relative capacity of a molecule in an atom to draw in shared electrons to itself. Case: Fluorine has the most noteworthy electronegativity. Comparable electronegativities between components give non-polar covalent bonds (0.0-0.4) Different electronegativities between components give polar covalent bonds (0.5-1.9) If the distinction between the electronegativities of two components is around 2.0 or more prominent, the bond is ionic

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Electronegativity Example For each of the accompanying sets of bonds, pick the bond that will be more polar. Al-P versus Al-N C-O versus C-S

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Dipole minute Dipole Moment A particle that has a focal point of positive charge and a focal point of negative energize Will line on electric field In Debye units 1 D = 3.34 x 10 - 30 C  m

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Examples F 2 CO 2 H 2 O NH 3 BF 3 CCl 4

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Molecular extremity Net-dipole minute prompts to sub-atomic extremity Thus the accompanying two that have net-dipole minutes are polar: H 2 O NH 3

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Molecular structure Molecular Structure : or geometric structure alludes to the three-dimensional course of action of the iotas in a particle. Bond Angle : The edge shaped between two bonds in a particle.

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Molecular structure: VSEPR The VSEPR Model : The valence shell electron combine aversion model is valuable for anticipating the sub-atomic structures of particles framed from nonmetals. The structure around a given particle is dictated by minimizing aversions between electron sets. The holding and nonbonding electron sets (solitary sets) around a given iota are situated as far separated as could be expected under the circumstances.

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Molecular Structure: VSEPR Steps for Predicting Molecular Structure Using the VSEPR Model 1. Draw the Lewis structure for the particle. 2. Tally the electron matches and orchestrate them in the way that minimizes aversion (that is, put the solitary combines as far separated as could be expected under the circumstances). 3. Decide the places of the molecules from the way the electron sets are shared. 4. Decide the name of the sub-atomic structure from the places of the iotas.

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Example Br2 CO2 CF4 PF3

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Your turn NH 4 + XeF 4 AsI 5 SF 3 + I 3 -

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