notes unit 4- bonding ii · mono-1 hexa-6 di-2 hepta-7 tri-3 octa-8 tetra-4 nona-9 penta-5 deca-10...

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Unit 4‐ Bonding IIReview

Unit 4‐ Bonding II

Compound Bond Type Compound Bond Type

NaCl Ionic NCl3 CovalentCO Covalent PF3 CovalentFeNi Metallic CaCl2 IonicSiS2 Covalent Fe2O3 Ionic

Determine the type of bond (Ionic, Covalent or Metallic) in the following compounds:

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Draw the Lewis Structures•

1) PBr3 4) NO2‐1

2) N2H2 5) C2H4

3) CH3OH 6) HBr

Draw the Lewis Structures•

1) PBr3 4) NO2‐1 [ ] ‐1

2) N2H2 5) C2H4

3) CH3OH 6) HBr

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1. Positive ion is written first … this is usually a metal2. Negative ion is written second … this is usually a nonmetal3. Subscripts are used to show how many ions of each part are in the compound. They are used to balance the charge of the ions. The overall charge should be “0”

Criss-cross method: Examples: 1. Sodium oxide

* sodium is the positive ion = +1* oxide is the negative ion = -2 * therefore … it takes 2 sodium ions to balance

the charge of the oxideFormula = Na2O

Rules of writing Ionic formulas:

2. Calcium nitrate calcium is the positive ion = +2nitrate is the negative ion = -1therefore … it takes 2 nitrates to balance the charge of

calciumFormula = Ca(NO3)2

3. Aluminum sulfide aluminum is the positive ion = +3sulfide is the negative ion = -2therefore … it takes 2 aluminum ions and 3 sulfide

to balance the chargeFormula = Al2S3

Rules of writing formulas:

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Prefix Number Prefix Number

mono- 1 hexa- 6

di- 2 hepta- 7

tri- 3 octa- 8

tetra- 4 nona- 9

penta- 5 deca- 10

Covalent NamingBinary covalent compounds are characterized by having two nonmetals. Naming these compounds involves the use of numerical prefixes:

Write the correct formulas for each covalent compound:Compound Name Oxidation States Covalent Formula

WaterO (2)

H (1)

Carbon DioxideC (4)

O (2)

Chlorine (Diatomic Element) Cl (1)

Methane (5 total atoms)C (4)

H (1)

Ammonia (4 total atoms)N (3)

H (1)

Carbon tetrabromide (5 total

atoms)

C (4)

Br (1)

Phosphorous trichloride (4

total atoms)

P (3)

Cl (1)

Diphosphorous trioxide (5

total atoms)

P (3)

O (2)

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WaterO (2)

H (1) H20

Carbon DioxideC (4)

O (2) CO2

Chlorine (Diatomic Element) Cl (1)


H (1)


H (1)


atoms)

C (4)

Br (1)


total atoms)

P (3)

Cl (1)


total atoms)

P (3)

O (2)


WaterO (2)

H (1) H20

Carbon DioxideC (4)

O (2) CO2

Chlorine (Diatomic Element) Cl (1) Cl2


H (1) CH4


H (1)


atoms)

C (4)

Br (1)


total atoms)

P (3)

Cl (1)


total atoms)

P (3)

O (2)

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WaterO (2)

H (1) H20

Carbon DioxideC (4)

O (2) CO2

Chlorine (Diatomic Element) Cl (1) Cl2


H (1) CH4


H (1) NH3


atoms)

C (4)

Br (1) CBr4


total atoms)

P (3)

Cl (1) PCl3


total atoms)

P (3)

O (2) P2O3


Compound Bond Type Compound Bond Type

NaCl Ionic NCl3 CovalentCO Covalent PF3 CovalentFeNi Metallic CaCl2 IonicSiS2 Covalent Fe2O3 Ionic

Determine the type of bond (Ionic, Covalent or Metallic) in the following compounds:

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Criss‐Cross rule* Write out symbols and charge of elements* Criss‐Cross charges as subscripts (Swap and Drop)* Combine as a formula unit

Equation Form of Balancing Charges(Number of Cations)x(Cation Charge) + (Number of Anions)x(Anion Charge) = 0

EX: Aluminum and Oxygen EX: Barium and Oxygen

Al3+ O2‐ Ba2+ O2‐

AL2O3 BaO

Lithium Iodide (LiI) Strontium Chloride (SrCl2) Sodium Sulfide (Na2S)

Balancing Charges:

Balancing Charges Practice

• Lithium Iodide

Li +1 I‐1 so LiI

• Strontium Chloride

Sr+2 Cl‐1 so SrCl2• Sodium Sulfide

Na +1 S‐2 so Na2S

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Cl‐ S‐2 F N‐3 O P‐3

Mg+2

Cs+

Cr+3

Na

Zn+2

Al+3

K


Cl‐ S‐2 F N‐3 O P‐3

Mg+2MgCl2 MgS MgF2 Mg3N2 MgO Mg3P2

Cs+CsCl Cs2S CsF Cs3N Cs2O Cs3P

Cr+3CrCl3 Cr2S3 CrF3 CrN Cr2O3 CrP

NaNaCl Na2S NaF Na3N Na2O Na3P

Zn+2ZnCl2 ZnS ZnF2 Zn3N2 ZnO Zn3P2

Al+3 AlCl3 Al2S3 AlF3 AlN Al2O3 AlP

K KCl K2S KF K3N K2O K3P

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Valence Shell Electron Pair Repulsion Theory

• This is the way that we predict the geometry shape of molecules, A model was developed a qualitative model known as Valence Shell Electron Pair Repulsion Theory (_VSEPR__ Theory).

• The basic assumptions of this theory are summarized below.

– 1) The electron pairs in the valence shell around the central atom of a molecule repel each other and tend to orient in space so as to minimize the repulsions and maximize the distance between them.

– 2) There are two types of valence shell electron pairs: __Bonded____ pairs and __Unbonded____ pairs

• Bond pairs are __Electrons shared__ by two atoms and are attracted by two nuclei. Hence they occupy less space and cause less repulsion.

• Lone pairs are pairs of electrons not involved in bond formation and are in attraction with only one nucleus. Hence they occupy more space. As a result, the lone pairs cause more repulsion.

• Note: The bond pairs are usually represented by a _line drawn between the two atoms_, whereas the lone pairs are represented by a lobe with two electrons.

Valence Shell Electron Pair Repulsion Theory

• 3) In VSEPR theory, the _double and triple_ bonds are treated as if they were single bonds. The electron pairs in multiple bonds are treated collectively as a single super pair.

• 4) The shape of a molecule can be predicted from the number and type of valence shell electron pairs around the central atom.

• When the valence shell of central atom contains only bond pairs, the molecule assumes symmetrical geometry due to even repulsions between them.

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Unit 4‐ Bonding

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Unit 4‐ Bonding

• Steric number is the total number of atoms bonded to the central atom and plus the number of lone pairs on the central atom.

Molecule Structural Diagram Oxidation State of each

element

Molecular Geometry

CClF3C (4) CL (1) F (1)

SF2

BF3

SiBr4

NH3

VSEPR PracticeComplete the table with the requested information.

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Molecule Structural Diagram Oxidation State of each

element

Molecular Geometry

CClF3C (4) CL (1) F (1)

SF2 S (2) F (1)

BF3B (3) F (1)

SiBr4

Si (4) Br(1)

NH3N (3) H (1)

VSEPR PracticeComplete the table with the requested information.

Polyatomic Ions

• Polyatomic ions are groups of atoms that are covalently bonded, but carry an overall net charge.

• The names of polyatomic ions must be memorized to appropriately name these compounds.

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Polyatomic Ions

ClO3‐ Chlorate

Polarity• Bond Polarity• Electronegativity• Ionic bonds have an electronegativity difference that is greater than 1.7.

Covalent bonds have an electronegativity difference less than (or equal to) 1.7. Electronegativity differences between 0 and 0.4 indicate non‐polar covalent bonds. Electronegativity differences between 0.4 and 1.7 indicate polar covalent bonds.

• Polar Covalent Bond‐ a covalent bond in which the electrons are not shared equally because one atom attracts them more strongly than the other.

• Non‐polar Covalent Bond‐ a covalent bond in which the electrons are shared equally.

• Use the periodic table of electronegativities to answer the questions on electronegativity differences.

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Electronegativity

H

2.1

Li

1.0

Be

1.5 ELECTRONEGATIVITY

(electron attraction!)

B

2.0

C

2.5

N

3.0

O

3.5

F

4.0

Na

0.9

Mg

1.2

Al

1.5

Si

1.8

P

2.1

S

2.5

Cl

3.0

K

0.8

Ca

1.0

Sc

1.3

Ti

1.5

V

1.6

Cr

1.6

Mn

1.5

Fe

1.8

Co

1.9

Ni

1.9

Cu

1.9

Zn

1.6

Ga

1.6

Ge

1.8

As

2.0

Se

2.4

Br

2.8

Rb

0.8

Sr

1.0

Y

1.2

Zr

1.4

Nb

1.6

Mo

1.8

Tc

1.9

Ru

2.2

Rh

2.2

Pd

2.2

Ag

1.9

Cd

1.7

In

1.7

Sn

1.8

Sb

1.9

Te

2.1

I

2.5

Cs

0.7

Ba

0.9

La‐Lu

1.0‐1.2

Hf

1.3

Ta

1.5

W

1.7

Re

1.9

Os

2.2

Ir

2.2

Pt

2.2

Au

2.4

Hg

1.9

Tl

1.8

Pb

1.9

Bi

1.9

Po

2.0

At

2.2

Fr

0.7

Ra

0.9

Ac

1.1

Th

1.3

Pa

1.4

U

1.4

Np‐No

1.4‐1.3

Electronegativity

• Determine the type of bond that would form between the following two elements using differences in electronegativity.

Example: Mg – O– O is 3.5 and Mg is 1.2, therefore, the difference is 3.5 – 1.2 = 2.3 IONIC

Example: Cl – Cl– Cl is 3.0. The difference is 3.0 – 3.0 = 0

NON‐POLAR COVALENT

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ElectronegativityBond Electronegativity Difference Bond Type

1. C – N

2. Li – F

Bond Electronegativity Difference Bond Type

3. N – Cl

4. Na ‐ Cl

5. O – F

6. B – H

7. Ba – F

8. C – H

ElectronegativityBond Electronegativity Difference Bond Type

1. C – N 0.5 Polar Covalent

2. Li – F 3.0 Ionic

Bond Electronegativity Difference Bond Type

3. N – Cl 0.0 Covalent

4. Na ‐ Cl 2.1 Ionic

5. O – F 0.5 Polar Covalent

6. B – H 0.1 Covalent

7. Ba – F 3.1 Ionic

8. C – H 0.4 Covalent

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Molecular Polarity• Dipole moment‐ a property of a molecule whereby the charge distribution can be represented by a center of positive charge and a center of negative charge.

•• Polar Molecule‐ a molecule that has a permanent dipole moment.

• Determining if a molecule is polar.• If ALL of the bonds are non‐polar, then the molecule is non‐polar.

• If some or all of the bonds are polar, you can consider the vectors. Vectors are arrows that point in the direction of the negative charge (the direction the electrons are pulled).

Molecular Polarity

• Examples:

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Molecular Polarity

Structural formula Polar or non‐polar

Formula *Diagram Formula Polar/Non‐polar

NH3

SCl2

CF4

PCl3

H2S

C2H2

* Bent must be draw as bent

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Structural formula Polar or non‐polar

Formula *Diagram Formula Polar/Non‐polar

NH3

Polar

SCl2Polar

CF4 Non Polar

PCl3Polar

H2SPolar

C2H2NonPolar

* Bent must be draw as bent

notes unit 4- bonding ii · mono-1 hexa-6 di-2 hepta-7 tri-3 octa-8 tetra-4 nona-9 penta-5 deca-10...

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