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Bonding And Structure Flashcards

1
Q

Factors affecting ionic bond strength

A

Nuclear charge/charge of ions
Ionic radii

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2
Q

Effect of nuclear charge / charge of ions on ionic bond strength

A

Higher charge= stronger

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3
Q

Effect of ionic radii on ionic bond strength

A

Smaller ionic radii= stronger

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4
Q

Describe strength of ionic bind across a period

A

Cationic radii decreases
Charge increases
Therefore for cations the ionic bond strength increases across a period

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5
Q

2 different types of substance

A

Molecular= simple or macro (covalent)
Giant= giant ionic or giant covalent or giant metallic

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6
Q

Describe giant ionic crystal lattice

A

Oppositely charged ions held in a regular 3D lattice by electrostatic attraction. E.g NaCl
Ionic bonds= multidirectional

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7
Q

Properties of ionic

A

High mp and bp= many multidirectional IB to break
Only conduct when molten or aq
Brittle= if arrangement is disrupted then ions repel
Often soluble in polar solvents e.g water but insoluble in non polar

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8
Q

What is stronger, single or triple bonds

A

Single= longer and weaker
Triple= shorter and stronger
Nuclei can remain closer if the shared electron density contains more bonding electrons to overcome repulsion.

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9
Q

Covalent defintion

A

Consists of shared pairs of electrons normally with one electron being supplied by each atom either side of bond.

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10
Q

Where can a covalent bond form

A

Atoms of same element e.g N2 O2 Cl2
Atoms of different non-metal elements, E.g CO2 SO2
When one of the elements is in middle of table e.g CCl4 SiCl4
Head of the group elements with high ionisation energies e.g BeCl2 AlCl3

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11
Q

With what electronegativity difference is ionic vs polar cov vs non polar cov

A

<0.5= non polar covalent
0.5<x>1.7= polar covalent
>1.7=ionic</x>

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12
Q

Factors affecting covalent bond strength

A

Sum of atomic radii/size. Smaller=stronger
Number of bonding electron pairs. Higher neg charge attracted to nucleus.
Bond length varies bond strength. Shorter=stronger longer=weaker

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13
Q

What does VSEPR stand for

A

Valence
Shell
Electron
Pair
Repulsion

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14
Q

Rules for shape of molecules

A

Draw lewis model
Count electron groups
Count bonded pairs
Count lone pairs

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15
Q

State the different shapes

A

Linear 180
Bent 104.5
Trigonal planar 120
Tetrahedra 109.5
Octahedral 90 & 180
Trigonal pyramid 107
Trigonal bypyramid

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16
Q

State bond angle of methane

A

Tetrahedra 109.5

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17
Q

State bond angle of boron trifluoride

A

BF3
Trigonal planar 120

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18
Q

Whats more electron dense, lone pair or bonded pair
What is the greatest repulsion between

A

Lone pair
LP is pulled to one nucleus therefore more compact and provide more repulsion.
BP evenly pulled towards each nuclei

2 lone pairs

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19
Q

What is a lewis diagram

A

Involves lone pairs and represents bonding electron pairs with a stick

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20
Q

What is a sigma

A

sigma bonds (σ bonds) are the strongest type of covalent chemical bond. They are formed by head-on overlapping between atomic orbitals.

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21
Q

Dative covalent bond

A

Shared pair of electrons come from only one atom
Both electrons supplied by one atom in a covalent bond

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22
Q

Examples of dative covalent

A

CO carbon monoxide
NH4+ ammonia
H30+
NH3BF3

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23
Q

Al2Cl6

A

Two AlCl3 molecules join together through two dative covalent bonds to form the dimer Al2Cl6

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24
Q

Describe the linear shape of molecules

A

N. Bonding pairs 2
N. Lone pairs 0
Bond angle 180
Examples CO2, BeF2, HCN, CS2

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25
Describe trigonal planar shape of molecules
N. Bonding pairs 3 N. Lone pairs 0 Bond angle 120 Examples BF3, AlCl3, SO3, NO3-, CO3 2-
26
Describe tetrahedra shape of molecules
N. Bonding pairs 3 N. Lone pairs 0 Bond angle 109.5 Examples SiCl4, SO4 2-, NH4+, ClO4-
27
Describe trigonal pyramidal shape of molecules
N. Bonding pairs 3 N. Lone pairs 1 Bond angle 107 Examples NCl3, PF3, ClO3,
28
Describe bent shape of molecules
N. Bonding pairs 2 N. Lone pairs 2 Bond angle 104.5 Examples OCl2, H2S, OF2, SCl2
29
Describe trigonal bipyramidal shape of molecules
N. Bonding pairs 5 N. Lone pairs 0 Bond angle 120 & 90 Examples PCl5
30
Describe octahedral shape of molecules
N. Bonding pairs 6 N. Lone pairs 0 Bond angle 90&180 Examples SF6
31
How to explain shape of molecules
1. State number of bonding pairs and lone pairs of electrons. 2. state that electron pains repel and try to get as far apart as possible (or to a position of minimum repulsion) 3. If there are no lone pairs, state that the electron pairs repel equally. 4. If there are lone pairs of electrons, then state that lone pains repel more than bonding pairs. 5. state actual shape and bond angle.
32
Lone pairs effect on bond angle
Lone pairs repel more than bonding pairs and so reduce bond angle, approx 2.5 degrees per lone pair
33
Electronegativity definition
Relative tendency of an atom in a covalent bond in a molecule to attract electrons kn a covalent bond to itself
34
How is electronegativity measured
Pauling scale 0-4
35
Where are the most electronegative elements found
Top right, N O F Cl
36
Where are least electronegative elements found
Bottom left, Fr Cs Rb
37
Factors affecting electronegativity
Increases across a period as protons increase and atomic radii decreases Decreases down a group because distance between nucleus and outer electrons increases and shielding by inner electrons increases
38
When is a polar covalent bond formed
Electronegativity difference of between 0.3-1.7
39
What is polar covalent
Unequal distribution of electrons in the bind so produces charge separation (dipole)
40
When is ionic bond formed difference if EN
EN difference of above 1.7
41
What is a symmetric molecule
All bonds are identical with no lone pairs Will not be polar even if individual bonds are e.g CO2 Individual dipoles of bonds cancel out due to symmetrical shape If there is no NET dipole moment then the molecule is non polar
42
What happens if a charged rod is brought close to a jet of polar liquid flowing from a burette
The jet of liquid will be attracted to the electrostatic force of the rod. The dipoles will align, the stronger the dipole the more deflection of the jet Deflected
43
What happens if a charged rod is brought close to a jet of non polar liquid flowing from a burette
Not be deflected
44
What are london forces
Occur between all molecular substances and noble gases, do not occur in ionic substances
45
Main factor affecting size of london forces
More electrons= higher chance that temporary dipoles will form= make london forces stronger= more energy required to break= higher bp
46
miscible
2 liquids dissolvig/mixing
47
immiscible
2 liquids don't mix/ separate
48
evidence for polar molecules method
-liquid in a burette -allow to run out -charged rod alongside stream of liquid -polar molecules=detected by electrostatic attraction -non-polar=unaffected
49
what is stronger intermolecular or intramolecular
intramolecular
50
3 types of intermolecular bond
london dispersion forces permanent dipole-dipole hydrogen bonds
51
3 types of intermolecular bond
london dispersion forces permanent dipole-dipole hydrogen bonds
52
list bonds from strongest to weakest
metallic ionic covalent hydrogen permanent dipole induced dipole (london)
53
how is an induced dipole formed
as the temporary dipole approaches the non-polar molecule it induces a dipole in the neighboring non-polar molecule
54
how is a temporary dipole formed
electron charge cloud in non-polar molecules is constantly moving during this movement the cloud can be more on one side of the atom or molecule causes temp dipole to arise
54
how is a temporary dipole formed
electron charge cloud in non-polar molecules is constantly moving during this movement the cloud can be more on one side of the atom or molecule causes temp dipole to arise
55
what IMFs do non polar molecules have
london
56
what IMFs do polar molecules have
london, permanent dipole-dipole, can have hydrogen
57
trend in BP down group 4. why
increases due to n. electrons increasing so LF stronger, n. protons/nuclear charge increases
58
when do hydrogen bonds occur
when H is covalently bonded to N O or F (highly electronegative)
59
how do hydrogen bonds occur
H only has 2 electrons shared in covalent bond, if pulled away from H atom, the H nucleus is left exposed and very eectron deficient (delta positive). this delta positive H in this extremely polar covalent bond can then interact with delta negative N O F in another molecule.

chemistry (7 decks)

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