Chemistry 3A - Chemical Bonding 1 (Ionic & Covalnent Bonds)

Front 1

Octet rule

Back 1

Atoms TEND to lose, gain, or share electrons till they are surrounded by eight valence elctrons

Front 2

Characteristics of Noble gases ( 3)

Back 2

1) High I.E

2) Low affinity for additional electrons

3) General lack of reativity

Front 3

Define ionic bond

Back 3

Electrostatic force of attraction between 2 oppositely charged ions ( Between metals and non-metals)

Front 4

Covalent Bond

Back 4

The attraction between the localised shared electrons and the nuclei (between non-metals)

Front 5

Define dative bond

Back 5

Shared pair of electrons is provided by only one of the bonding atoms

Front 6

Conditions for co-ordinate bond (2)

Back 6

1) Lone pair of electrons for the donor

2) Vacant low-lying orbital for the acceptor

Front 7

Relationship between cat/anions and electronegativity

Back 7

For cations (e.g. NH4+) electrons are lost from the less electronegative atom (usually central atom)

For anions (e.g. S04 2- ) electrons are gained by the more electronegative ion

Front 8

Less than 8 electrons in valence shell = ?

Back 8

Electron deficient

Front 9

Molecules with unpaired electrons (e.g. n02) = ?

Back 9

Radicals

Front 10

Why can Pcl5 exist but not Ncl5?

Back 10

Expansion of octet is only available for elements in period 3 or more , as they will have the low-lying vacant orbitals to accommodate the expansion.

N is from period 2 .

Front 11

What is the VSEPR Theory and what is it used for

Back 11

Valence Shell Electron Pair repulsion Theory

To predict the molecular geometry or shape of a species

Front 12

2 Principles of the VSEPR Theory

Back 12

1) Electron pairs around the central atom of a molecule are arranged as far apart as possible to minimise inter-electronic repulsion - to have the greatest stability

2) Lp - lp repulsion > bp-lp repulsion > bp-bp repulsion

Front 13

Is a double bond 2 regions of electron density?

Back 13

NO

Front 14

2 regions of electron density

Back 14

0 lp - Linear - 180

Front 15

3 regions of electron density

Back 15

0lp - trigonal planar - 120

1lp - bent - <120

Front 16

4 regions of electron density

Back 16

0lp - tetrahedral - 109.5 / 109

1lp - trigonal bypyramidal - 107

2lp - bent - 105

Front 17

5 regions of electron density

Back 17

0lp - Trigonal bipyramidal - 120 + 90

1lp - Seesaw - NA

2lp - T-Shape - 90

3lp - Linear - 180

Front 18

6 regions of electron density

Back 18

0lp - Octahedral - 90

1lp - square pyramidal - NA

2lp - Square planar - 90

Front 19

Effect of electronegativity of central atom on bond angles

Back 19

More electronegative - draws electron density to itself - bond pair is nearer to the nucleus - exert more repulsion

Front 20

Covalent bonding is the overlapping of valence orbitals - why does the overlapping need to be effective ?

Back 20

To ensure that the electron density in the region between the nuclei is a maximum

Front 21

What is the condition for the optimum bonding distance?

Back 21

Molecule must be most stable - when the energy level is at a minimum

Front 22

Two types of Covalent bond depending on the way the orbital overlap

Back 22

1) Sigma bond

2) pi bond

Front 23

Characteristics of sigma bonds

1) Type of overlap

2) Type of orbitals

3) Electron density concentration

Back 23

1) Head to head overlap of 2 atomic orbitals

2) Can be s,p, or d orbitals

3) Electron density is concentrated between the nuclei of the two bonding atoms

Front 24

Characteristics of pi bonds

1) Type of overlap

2) Type of orbitals

3) Electron density concentration

Back 24

1) Collateral , side-on overlap

2) p or d orbitals

3) Electron cloud above an below the nuclear axis but 0 electron density along the axis

Front 25

Comparison between the two types

Back 25

1) pi bond is weaker as the side - on overlap is less effective

Front 26

Can pi bonds exist in single bonds ?

Back 26

No. pi bonds occur when atoms undergo multiple bonding - one of the bonds must be a sigma bond, and the rest are pi bonds

E.g. triple bond consists of one sigma and 2 pi bonds

Front 27

Relationship between covalent bond length and strength of covalent bond

Back 27

Stronger the covalent bond - shorter the bond length

Front 28

Bond energy = bond enthalpy ?

Back 28

YEP

Front 29

Define bond energy

Back 29

The average energy absorbed when 1 mol of a particular bond is broken in the gaseous state

Front 30

Factors affecting covalent bond strength (3)

Back 30

1) Number of bonds between atoms

2) Effectiveness of overlap between orbitals

3) Difference in electronegativities of the bonding atoms (bond polarity)

Front 31

1) Number of bonds between atoms

Back 31

Increased number of bonds - increased number of shared electrons btw the 2 atoms - increased electrostatic attraction btw bond pairs and nuclei - bond strength increases

Front 32

2) Effectiveness of overlap of the orbitals

Back 32

Larger orbitals - more diffuse - overlap is less effective

Front 33

3) Difference in electronegativities in the bonding atoms

Back 33

Different electronegativities - electron density distribution is asymmetrical - partial charges arise on the bonded atoms - polar covalent bond - increase in electrostatic attraction ( between delta + and data - ) - increased bond strength

Simply put : Greater the difference in electronegativities - higher the bond parity - stronger the bond

Front 34

What is hybridisation?

Back 34

Term applied to the mixing of atomic orbitals in an atom to generate a set of hybrid orbitals

Front 35

What does sp3 hybridisation lead to ?

Back 35

tetrahedral electron pair geometry

Front 36

What does sp2 hybridisation lead to ?

Back 36

Trigonal planar electron pair geometry

Front 37

What does sp hybridisation lead to

Back 37

linear electron pair geometry

Front 38

What about the unhybridized orbitals?

Back 38

They remain ; perpendicular to each other and the line containing the hybridised orbital

Front 39

Effect of hybridisation one bond strength and length

Back 39

Since s orbital is spherical and closer to the nucleus , the higher the s character of the hybrid orbital - more effective overlap with orbital of other atoms - increase in bond strength and decrease in bond length

E.g. sp has more s character than sp3

Front 40

What is in a resonance hybrid(e.g. O3)

Back 40

pi bonding electrons are delocalised over the atoms via the overlapping of p orbitals of atoms , which are perpendicular to the plane of the molecule ( containing sigma bonds)

Front 41

Why resonance hybrid?

Back 41

More stable than any of the other resonance structures

Front 42

Range of bond types

Back 42

Ionic - Ionic with partial covalent character - polar covalent bond - non polar / pure covalent bond

Front 43

Formation of ionic bonds with covalent character

Back 43

Attraction of the cation for the valence shell electrons of the anion - distortion / polarisation of the electron cloud - pulling it into the region between the two nuclei - electron sharing , ergo, covalency

Front 44

Degree of covalent character depends on (2)

Back 44

1) Polarising power of the cation - in turn dependent on the charge density

2) Polarisability of the anion depends on the size of the anion and charge

Front 45

Formation of Polar covalent bond

Back 45

Two atoms with different electronegativities - electron pair is not equally shares - bond becomes polar

Front 46

Degree of polarity depends on

Back 46

Dipole moment - Dependent on difference in electronegativities.

The product of the charge of the distance between the charges

Indicated by an arrow with a line striking close to the end

Front 47

What is the overall dipole moment of a molecule ?

Back 47

The vector sum of all the bond dipol moments

Front 48

Is Co2 a polar molecule ?

Back 48

No. Although the CO bonds are polar covalent bonds, the bond dipole moments cancel each other out vectorially - overall dipole moment is 0

Front 49

Which molecular shapes cancel out the bond dipole moments of polar covalent bonds?

Back 49

Linear , trigonal planar, tetrahedral, trigonal bipyramidal, octahedral

* all shapes without any lone pairs

Front 50

C-H bond ; is it polar?

Back 50

No . the electronegativity difference is almost zero - thus it is negligible