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chemistry > Bonding And Structure > Flashcards

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
Q

Describe trigonal planar shape of molecules

A

N. Bonding pairs 3
N. Lone pairs 0
Bond angle 120
Examples BF3, AlCl3, SO3, NO3-, CO3 2-

26
Q

Describe tetrahedra shape of molecules

A

N. Bonding pairs 3
N. Lone pairs 0
Bond angle 109.5
Examples SiCl4, SO4 2-, NH4+, ClO4-

27
Q

Describe trigonal pyramidal shape of molecules

A

N. Bonding pairs 3
N. Lone pairs 1
Bond angle 107
Examples NCl3, PF3, ClO3,

28
Q

Describe bent shape of molecules

A

N. Bonding pairs 2
N. Lone pairs 2
Bond angle 104.5
Examples OCl2, H2S, OF2, SCl2

29
Q

Describe trigonal bipyramidal shape of molecules

A

N. Bonding pairs 5
N. Lone pairs 0
Bond angle 120 & 90
Examples PCl5

30
Q

Describe octahedral shape of molecules

A

N. Bonding pairs 6
N. Lone pairs 0
Bond angle 90&180
Examples SF6

31
Q

How to explain shape of molecules

A
  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
Q

Lone pairs effect on bond angle

A

Lone pairs repel more than bonding pairs and so reduce bond angle, approx 2.5 degrees per lone pair

33
Q

Electronegativity definition

A

Relative tendency of an atom in a covalent bond in a molecule to attract electrons kn a covalent bond to itself

34
Q

How is electronegativity measured

A

Pauling scale 0-4

35
Q

Where are the most electronegative elements found

A

Top right, N O F Cl

36
Q

Where are least electronegative elements found

A

Bottom left, Fr Cs Rb

37
Q

Factors affecting electronegativity

A

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
Q

When is a polar covalent bond formed

A

Electronegativity difference of between 0.3-1.7

39
Q

What is polar covalent

A

Unequal distribution of electrons in the bind so produces charge separation (dipole)

40
Q

When is ionic bond formed difference if EN

A

EN difference of above 1.7

41
Q

What is a symmetric molecule

A

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
Q

What happens if a charged rod is brought close to a jet of polar liquid flowing from a burette

A

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
Q

What happens if a charged rod is brought close to a jet of non polar liquid flowing from a burette

A

Not be deflected

44
Q

What are london forces

A

Occur between all molecular substances and noble gases, do not occur in ionic substances

45
Q

Main factor affecting size of london forces

A

More electrons= higher chance that temporary dipoles will form= make london forces stronger= more energy required to break= higher bp

46
Q

miscible

A

2 liquids dissolvig/mixing

47
Q

immiscible

A

2 liquids don’t mix/ separate

48
Q

evidence for polar molecules method

A

-liquid in a burette
-allow to run out
-charged rod alongside stream of liquid
-polar molecules=detected by electrostatic attraction
-non-polar=unaffected

49
Q

what is stronger intermolecular or intramolecular

A

intramolecular

50
Q

3 types of intermolecular bond

A

london dispersion forces
permanent dipole-dipole
hydrogen bonds

51
Q

3 types of intermolecular bond

A

london dispersion forces
permanent dipole-dipole
hydrogen bonds

52
Q

list bonds from strongest to weakest

A

metallic
ionic
covalent
hydrogen
permanent dipole
induced dipole (london)

53
Q

how is an induced dipole formed

A

as the temporary dipole approaches the non-polar molecule it induces a dipole in the neighboring non-polar molecule

54
Q

how is a temporary dipole formed

A

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
Q

how is a temporary dipole formed

A

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
Q

what IMFs do non polar molecules have

A

london

56
Q

what IMFs do polar molecules have

A

london, permanent dipole-dipole, can have hydrogen

57
Q

trend in BP down group 4. why

A

increases due to n. electrons increasing so LF stronger, n. protons/nuclear charge increases

58
Q

when do hydrogen bonds occur

A

when H is covalently bonded to N O or F (highly electronegative)

59
Q

how do hydrogen bonds occur

A

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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