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Biochem Y1 > FoC : Structure, Bonding And Equalibria > Flashcards

FoC : Structure, Bonding And Equalibria Flashcards

1
Q

Covalent bonding

A

The formal sharing of electrons between two atoms. In heteroatomic molecules this is due to differences in electronegativity

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2
Q

Ionic bonding

A

The formal transfer of an electron from one atom to another

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3
Q

Metallic bonding

A

Low electronegativity elements, electrons are delocalised in a sea of electrons

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4
Q

Hydrogen bonding

A

Electrostatic interaction between a (H bonding to an electronegative atom )and an electronegative atom

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5
Q

Hydrogen bonding is an example of what type of interaction

A

Non covalent or electrostatic

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6
Q

Electrostatic

A

Positive charge is attracted to negative charge

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7
Q

Pauling definition of electronegativity

A

Ability of an atom to attract and electron to itself

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8
Q

Whose definition of electronegativity is based upon physical data

A

Allen

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9
Q

What are electrons with low electronegativity sometimes called?

A

Electropositive (meaning only very very weak electronegativity)

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10
Q

What is the van Arkel-Ketelaar triangle

A

Triangle that allows us to predict the dominant bonding in a compound using electronegativity and the different in electronegativity between the bonded atoms

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11
Q

If the reaction is exothermic what does that mean about the strength of the bond in the product

A

Stronger then the starting material
Thermodynamically favoured

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12
Q

If the charge in enthalpy is positive what does this mean about the thermodynamic favourability

A

It’s not

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13
Q

Trends in ionisation energy

A

Generally increases as protons increases
Shielding same
Therefore greater nuclear charge (Zeff)

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14
Q

What is the effect of d orbitals on ionisation energy

A

D orbitals being filled has a small increase in Zeff.
The 10 electrons don’t effectively shield the 10 added protons
Zeff is higher

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15
Q

Why is further IE’s higher then first IE

A

It’s harder to remove and electron from a cation
The attraction to the electron to positively charged ion is greater

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16
Q

Trends in electron gain enthalpy

A

Normally exothermic
Typically decreases with more protons
Sufficient attraction between the nucleus and new electron
Overcome electron pair repulsion
Exception at N
Electron pair repulsion means more energy required to gain an electron

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17
Q

Trends in lattice enthalpy

A

Lattice enthalpy decreases as ion size increases
Greater charge = greater lattice enthalpy
When formed from charges greater then +1 greater lattice enthalpy

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18
Q

Why don’t noble gases form ions

A

The IE is so high
as the nuclear charge is at maximum
Energy can’t be recouped forming bonds
Isn’t thermodynamically favourable

19
Q

General strength of covalent bonds as opposed to ionic and h bonding

A

Weaker then ionic
Stronger then H

20
Q

What is the shape of the p and s orbitals

A

P = dumbbell
S = spherical

21
Q

Difference between multipul bonds and single bonds

A

Multiple bonds are shorter and stronger

22
Q

Covalent bond dissociation energy trends

A

Generally bonds to large atoms are weak and longer
Orbitals increase in size and more diffuse
Can’t overlap as efficiently
Bonds to electronegative elements are shorter
Elec neg have smaller orbitals
E- more strongly attrac to nuc
So more overlap of orbitals

23
Q

What are the limitations of VSEPR

A

•In some compounds (such as [TeBr6]2-) predicts inaccurately
As lp is stereochemically inactive
As lp occupies s orbital
•can’t use for radicals
•can’t use for transition metals
•can’t use for F block elements

24
Q

what does the valence bond theory assume

A

bonds are located between pairs of atoms

25
Q

what does hybridisation suggest will happen to 2s and 2p orbitals

A

combine to form 4 new sp3 hybrid orbitals so all bonds are identical and each contain an electron

26
Q

what does hybridisation suggest about the formation of bonds

A

bond is formed through the overlap sp3 orbitals and the 1s orbital. overlap causes a 2 electron covalent bond

27
Q

why do double bonded compunds not have sp3 bonding on the double bond

A

they wouldnt align with the geometry of the molecule

28
Q

what hybridisation do double bonds have

A

sp2 created through the overlap of 2 p orbitals and 1 s orbital , 1 p orbital is not hybridised
c-c sigma bond formed from 2 sp2 overlap and pi bond formed from unhybridised p orbital overlap

29
Q

when are sp orbitals found

A

in triple bonds the overlap of1 s and 1 pz orbital

30
Q

how are all the bonds formed in a triple bond

A

2 pi bonds formed at 90 degrees from one another from the 2px overlap and 2py overlap (unhybridised) ans sp hybrisation 180 degrees from one another

31
Q

bronstedlowry definition of acid & base

A

acid : proton donar
base: proton acceptor

32
Q

lewis definition of acid & base

A

acid: electron pair acceptor
base: electron pair donar

33
Q

how to calculate Ka

A

[A-][H3O+]/ [HA]

34
Q

if the acid is stronger what does this mean in terms of conjugate acid, equalibrium and Ka

A

conjugate acid increases
equalibrium shifts further right
Ka increases

35
Q

pKa calculation

A

pKa = -log10 Ka

36
Q

what factors control acidity

A

strength of HA (acid)
electronegativity of A
stabilisation of A- (conjugate base)
the solvent

37
Q

whats a polybasic acid

A

may lose more then one proton

38
Q

how to calculate Kb

A

[BH+][OH-]/[B]

39
Q

what is solvent leveling

A

strongest acids appear to have the same strength when fully dissociated because water is a good proton acceptor

40
Q

what can water do that other molecules cannot

A

self ionise

41
Q

how to calculate the degree of dissociation

A

( [H3O+]/[HA] ) x 100

42
Q

how to calculate kw

A

[H3O+][OH-]

43
Q

how to calculate the pH of buffer solutions

A

pH = pKa + log10[A-]/[HA]

Biochem Y1 (11 decks)

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