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Fundamnetals of Chemistry for Biochemsists CHE00031C > Structure, Bonding and Equilibria > Flashcards

Structure, Bonding and Equilibria Flashcards

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1
Q

Covalent bonding

A

The formal sharing of electrons between two atoms

  • Found in elements with high electronegativity and in heteroatomic molecules with a small difference in electronegativity
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2
Q

Ionic bonding

A

In compounds where constituent elements have a high difference in elelctronegativity

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3
Q

Metallic bonding

A

Found in elements with low electronegtaivity (bottom left of periodic table) electrons delocalised into a “sea of electrons”

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4
Q

Hydrogen bonding

A

Primarily electrostatic, traditionally between a H atom attached to an electronegative atom (non-covalent interaction)

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5
Q

Define elctronegtivity

A

ability of an atom to attract an electron to itself (Pauling)

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6
Q

Trend of electrongativity down the periodic table

A

decreases down the table (highest in top right, lowest in bottom left)

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7
Q

Define electropositive

A

not very electronegative

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8
Q

ΔH of atomisation

A

The amount of enthalpy change when a compound’s bonds are broken and the component atoms are separated into single atoms

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9
Q

ΔH of reaction

A

The enthalpy change that accompanies a chemical reaction

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10
Q

1st ionisation enthalpy

A

The enthalpy change associated with the removal of the first electron from an isolated gaseous atom

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11
Q

2nd ionisation enthalpy

A

The enthalpy change associated with the removal of the second electron from an isolated gaseous ion

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12
Q

Electron gain enthalpy

A

The amount of energy released when an electron is added to an isolated gaseous atom

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13
Q

Lattice enthalpy

A

A measure of the strength of the forces between the ions in an ionic solid

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14
Q

In terms of the van Arkel-Katelaar triangle, what is an ionic bond?

A

A high diff. in electrongativity

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15
Q

In terms of the van Arkel-Katelaar triangle, what is an covalent bond?

A

A high electronegativity and a low difference between the two constituents

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16
Q

In terms of the van Arkel-Katelaar triangle, what is an metallic bond?

A

A low electrongativity with a low difference between constituents

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17
Q

Describe the trends in 1st ionization energy across period 2

A
  • Moving left to right protons increase (nuclear charge increases)
  • The outer electrons are all in the same area of space so do not shield each other well (Zeff increases)
  • Attraction between nucleus and outer e- increase
  • IE increases
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18
Q

Explain the dip in 1st IE at B

A

New e- in p-orbital does no penetrate 2s very well so increase in Zeff is not as much as expected

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19
Q

Explain the dip in 1st IE at O

A

Two electrons in same p-orbital, e- repulsion (less energy to remove)

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20
Q

Describe and explain the trend in 1st IE as you move down the periodic table

A
  • IE decreases
  • Outer e- further from nucleus
  • Less energy required to remove outer e-
  • NOT AN EVEN TREND - d-orbitals being filled have a small effect on Zeff due to the 10 e- not effectively shielding the 10 new protons (Zeff increases)
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21
Q

Describe and explain the trend in IE beyond 1st as you move down the periodic table

A
  • 2nd and higher IE is always higher (e- attraction greater to + cation)
  • Significant increase in IE when removing from a lower shell (closer to nucleus)
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22
Q

Describe trends in lattice enthalpy

A
  • Lattice enthalpy decreases as ion size increases
  • Greater charge on either ion leads to greater lattice enthalpy
  • Gain in lattice enthalpy is normally more than economical for amount of energy entering system in IE
  • Lattice enthalpy decreases down a group
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23
Q

Describe and explain the trend in covalent bond dissociation energies

A

As atoms get larger their orbitals increase in size and become more diffuse (BDE decreases) therefore, weaker and longer bond
- Bonds to electronegative atoms are stronger and shorter

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24
Q

Why do covalent bonds form?

A
  • both e- attracted to both positive nuclei
  • 2 e- repel eachother
  • 2 nuclei repel eachother
  • equilibrium in terms of internucleur distance and energy found
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25
Q

What is hybridization?

A

the combining of s and p orbitals to form sp orbitals (which point to the four corners of the tetrahedron)

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26
Q

What hybridization occurs in an alkane?

A

sp3 (1s and 3p orbitals) (4 sigma bonds)

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27
Q

What hybridization occurs in an alkene?

A

sp2 (1s and 2p orbitals) (final p orbital forms pi bond) (3 sigma bonds formed by sp orbitals)

27
Q

What hybridization occurs in an alkyne?

A

sp (1s and 1p orbital) (2 sigma bonds forrmed by sp orbitals) (2p orbitals form the two pi bonds at 90 degrees to eachother)

28
Q

What does an increase in s character in a given bond do to bond length?

A

Shortens it as s-electrons held closer to the nucleus

29
Q

Rules for using VSEPR to determine shape

A

1) Determine which atom you’re determining geometry around
2) Take the group number for this atom and add electrons for substituents (1 for single bond,2 for double etc.)
3) Add/remove e- for charge
4) Divide by 2 to determine no. PAIRS
5) Deduce no. lone pairs and bonding pairs
6) Remember to consider any double or triple bonds before determining the geometry

30
Q

How does EPG affect hybridization?

A

Linear - sp
Triagnol planar - sp2
Tetrahedral - sp3
Trigonal bipyramid - sp3d
Octahedral - sp3d2

31
Q

Steps for drawing a resonance form

A

1) draw a Lewis dot structure with charges
2) Identify which groups can donate pi e-/lone pairs and those that can accept them
3) Draw curly arrows between the two groups (must be conjugated)
4) ensure nothing has too many bonds
5) draw the new structure
6) Check stability

32
Q

Define Bronsted-Lowry acid

A

An acid is a substance which is a proton donor

33
Q

Define Bronsted-Lowry base

A

A base is a substance that accepts protons

34
Q

Define Lewis acid

A

e- pair acceptor

35
Q

Define Lewis base

A

e- pair donor

36
Q

The formula for the equilibrium constant

A

Kc = [A-][H3O+]/[HA][H2O]

37
Q

The formula for acid dissociation constant

A

Ka = [H+][A-]/[HA]

38
Q

What does a high Ka indicate in terms of acid strength

A

High Ka = high acid strength

39
Q

pKa formula

A

pKa = -log(Ka)

40
Q

What does a high pKa indicate in terms of acid strength

A

low acid strength

41
Q

Factors affecting acidity

A

1) strength of H-A bond
2) Electronegativity of A
3) stabilization of A-
4) Solvent

42
Q

What is an indication of a high stability of A-

A
  • A large no. resonance forms (- charge more dispersed across mol) less likely to reform acid
  • If the resonance forms place - on an electronegative atom, increases the stability of the conjugate base
43
Q

Kb formula

A

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

44
Q

pKb formula

A

pKb = -log(Kb)

45
Q

Ka of the conjugate base formula

A

[B][H3O+]/[BH+]

46
Q

What does a high Ka of the conjugate base indicate about the base strength?

A

Weak base

47
Q

What does a high pKa of the conjugate base indicate about the base strength?

A

Strong base

48
Q

What does solvent leveling mean?

A

The effect of strong acids appearing to be equally strong because they all fully dissociate in a good base

49
Q

Example of a worse base than water to test differences in acidity between strong acids

A

Ethanoic acid CH3COOH

50
Q

What will a base stronger than water do in water?

A

It will deprotonate water to OH- (water acts as an acid)

51
Q

Kc formula describing self-ionisation of water

A

Kc = [H3O+][OH-]/[H2O]

52
Q

define amphiprotic

A

A substance that can act as both an acid and a base

53
Q

Self ionisation constant

A

Kw

54
Q

Kw expression for water

A

Kw = [H3O+][OH-] (units: mol2dm-6)

55
Q

prove water’s pH is 7

A

at 298K Kw = 10^-14mol2dm-6
In pure water [H3O+]=[OH-]
Therefore [H3O+] = 10^-7moldm-3

pH=-log[H3O+]
pH=-log(10^-7)
pH=7

56
Q

How does pH of water vary with temperature?

A
  • H2O dissociation is endothermic
  • As temp increases pH decreases
  • This is always neutral as pure water is always considered “neutral”
57
Q

pH formula

A

pH=-log[H3O+]

58
Q

What defines a strong acid in terms of pKa

A

pKa<0

59
Q

How to determine pH of a strong acid (pKa<0)

A

Because acid is completely dissociated, you can assume that [H3O+(aq)]=[HA]

pH = -log[H3O+]
pH = -log[HA]

60
Q

Define buffer

A

Solutions able to resist changes in pH on the addition of an acid or base or on dilution

61
Q

What is a buffer made up of

A

A buffer is made up of weak acids and their conjugate bases or weak bases and their conjugate acids

e.g. CH3COOH + CH3COONa, NH3 + NH4Cl

62
Q

What is the Henderson-Hasslebach equation used for

A

to find the pH of a buffer

63
Q

What is the Henderson-Hasslebach equation?

A

-log(Ka) = -log(pKa) + log([A-]/[HA])

64
Q

Henderson hasslebach equation

A

pH = pKa + log([A-]/[HA])

Fundamnetals of Chemistry for Biochemsists CHE00031C (5 decks)

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