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Chemistry > 3.1 Physical Chemistry > Flashcards

3.1 Physical Chemistry Flashcards

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1
Q

What information does a mass spectrometer give?

A

Relative isotopic mass + relative abundances

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2
Q

What happens during TOF mass spectrometry?

A

1) ionisation into +1 ions
2)accelerated through electric field to have same KE
3)ion drift: ions travel at different speeds based on mass (TOF)
4) detection=current produced (abundance)

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3
Q

What is electron impact ionisation?

A

High energy e- fired at vaporised element or low Mr compound.
Loses e- to become +1

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4
Q

What is electronspray ionisation?

A

•dissolved in volatile liquid (methanol/water)
•injected through hypodermic needle
•needle is charged so it gains a proton (H+)

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5
Q

When is electron spray ionisation used?

A

For high Mr compounds

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6
Q

What is the definition of relative atomic mass? (Ar)

A

Average mass of atoms compared to 1/12th mass of C-12 atoms

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7
Q

How do you work out the Mr from a mass spectrum for electro spray ionisation?

A

Main peak= Mr + 1 so take away 1

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8
Q

Why is there a small peak after the main peak?

A

In organic molecules- due to small no. of molecules containing 13C or 2H

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9
Q

Why is ionisation necessary to be analysed by TOF mass spectrometry?

A

•ions are accelerated by an electric field (mass)
•ions create current when hitting detector (abundance)

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10
Q

What are the different orbitals and how many e- can they hold?

A

S=2
P=6
D=10

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11
Q

How is the electron configuration for Cr and Cu different?

A

They have 4s1 instead of 4s2 which makes it more stable as e- repel each other in the same suborbital

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12
Q

What is ionisation energy?

A

Energy needed to remove one e- from each atom in 1 mole of gaseous atoms to form 1 mole of gaseous ions

X(g) = X+ (g) + e-

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13
Q

What are the factors affecting ionisation energy?

A

•atomic radius
•charge of nucleus
•shielding

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14
Q

What happens to the ionisation energy down a group and why?

A

It decreases the ionisation energy as there is more shielding + increased atomic radius so weaker attraction

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15
Q

What happens to ionisation energy across a period and why?

A

I.E decreases as there is the same shielding but increased nucleus charge so atomic radius decreases. Increasing attraction

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16
Q

What are exceptions which cause decrease of ionisation energy across a period?

A

• e- in e.g 2p orbital not 2s as 2p is in higher energy levels
• 2e- in a suborbital of an orbital as they repel, increasing energy

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17
Q

Why is the second successive ionisation energy of an element higher than 1st?

A

2nd e- is being lost from a positive ion which increases its attraction to the nucleus. + less repulsion

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18
Q

How is % error calculated?

A

(+/- uncertainty divided by measurement) *100

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19
Q

What is an orbital?

A

A region around the nucleus that can hold upto 2 e-

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20
Q

Why should % yield be as high as possible?

A

To maximise the amount of products made

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21
Q

Why should atom economy be as high as possible?

A

•To maximise the mass of reactants that end up as useful products
•To minimise the amount of waste by product

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22
Q

What is the ideal gas equation? (With units)

A

PV=nRT
R: 8.31
T: K (•C + 273 = K)
P: Pa
V: m3

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23
Q

What do empirical and molecular formula mean?

A

Empirical: simplest whole no. ratio of atoms of each element in a compound
Molecular: actual no. of atoms of each element in a compound

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24
Q

How is the no. of molecules calculated?

A

Avogadro’s constant X moles

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25
Q

What are the steps to making a standard solution?

A

1.weight the mass on digital balance
2.add distilled water to beaker and stir to dissolve
3.tip solution with funnel into volumetric flask
4.rinse beaker and stirring rod and add to flask
5. Top flask to correct volume until bottom of meniscus on line
6.stopper flask and turn upside down to mix

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26
Q

What are the indicators used in titration? (+ colour change)

A

•methyl orange: red in acid, yellow in alkali
•phenolphthalein: colourless in acid, pink in alkali

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27
Q

What is a coordinate (dative) covalent bond? (Arrow bond)

A

It contains a shared pair of e- (lone pairs) with both e- supplied by one atom to form a covalent bond

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28
Q

What are bonding and e- pairs?

A

Charge cloud that repel each other

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29
Q

What effects the shape of a molecule?

A

Pairs of e- in outer shell of atoms arrange as far apart to minimise repulsion= stable
•lone pairs will repel most= reduce bond angle by 2.5

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30
Q

What are all the shapes of molecules names and bond angles?

A

•2 e- pair: linear, 180
•3: trigonal planar, 120
•4: tetrahedral, 109.5
•5: trigonal bipyramidal, 120 + 90
•6: octahedral, 90

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31
Q

What is electronegativity?

A

The ability of atom’s nucleus to attract the pair of e- in a covalent bond to form dipoles with the other atom

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32
Q

Why do some molecules with polar bonds not have a permanent dipole?

A

•The polar bonds are due to different electronegativities
•these bonds are arranged symmetrically around central atoms so will have their charges cancel out= no dipole moments

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33
Q

What is a dipole?

A

Difference in charge between 2 atoms with different electronegativities caused by shift in e- density

34
Q

What is an induced dipole dipole force? (Van der waals)

A

•e- in charge clouds always move which forms temporary dipoles
•this can induce another temporary dipole in a neighbouring molecule
•the 2 dipoles are attracted to each other

35
Q

What is an permanent dipole dipole force?

A

•permanent dipoles between molecules with atoms of different electronegativities
•they will attract neighbouring dipoles

36
Q

What is a hydrogen bond?

A

•only happens between N O F, as they have lone pairs and very high electronegativities
•bond is so polarised and H has very high charge density = H-bonds form between lone pairs and H atoms of other molecules

37
Q

Why does S have a higher boiling point than P?

A

S8 molecules compared to P4 molecules have more e- (bigger molecules) so stronger van der waals

38
Q

Which of Na+ and Mg2+ is the smaller ion?

A

Mg2+ as it will have same shielding but a higher nuclear charge

39
Q

What is an advantage of using a conical flask instead of a beaker in titration?

A

•Less chance of splashing
•less chance of losing solution

40
Q

What is enthalpy change and standard conditions?

A
  1. Heat energy change measured under constant pressure
  2. 100kPa and 298K
41
Q

What is standard enthalpy of combustion?

A

Enthalpy change when 1 mole of a substance is completely burned in oxygen under standard conditions

42
Q

What is standard enthalpy of formation?

A

Enthalpy change when 1 mole of a compound is formed from its elements in their standard states under standard conditions

43
Q

What are the 2 equations needed for calorimetry experiments?

A

q=mc^T
^T= change in temperature
c= 4.18 J/g/K
m= mass of the substance that has the temp. change + c

Enthalpy change= q/mol
-ve for exothermic
+ve for endothermic

44
Q

What is Hess’s law?

A

Total enthalpy change is independent of the route taken

45
Q

Why do values of mean bond enthalpy differ from those in Hess’s law?

A

Mean bond enthalpies are mean values calculated across a range of compounds containing that bond whereas Hess’s law find the actual values of the bond enthalpies in that reaction.

46
Q

What is bond enthalpy?

A

Average energy needed to break 1 mole of covalent bonds, averaged over a range of molecules

47
Q

When will a reaction occur?

A

1.collisions must occur in the right orientation
2.collide at their specific activation energy

48
Q

What is activation energy?

A

Minimum amount of energy needed to break bonds and start a reaction

49
Q

What happens to rate of reaction when temperature increases? (+to maxwell-Boltzmann distribution?)

A

•More KE so more particles can collide frequently and successfully with their activation energy
•curve= same AE but more molecules reacting

50
Q

What happens to rate of reaction as concentration and pressure increases?

A

More crowded so collide more frequently, increasing rate of reaction

51
Q

How does a catalyst increase rate of reaction?

A

Provides an alternative reaction pathway with lower activation energy, without being chemically used up

52
Q

How can reaction rates be calculated?

A

•timing how long a precipitate takes to form (mark obscured but can be subjective)
•measure decrease in mass due to gas released so use fume cupboard
•measure volume of gas given off with gas syringe

53
Q

What is dynamic equilibrium and when does it occur?

A

•forward and reverse reactions at an equal rate
•products and reactant concentration remain constant
•happen only in closed system

54
Q

What is Le chatlier’s principle and in which type of reactions does it occur in?

A

•If reversible reaction has a change in concentration, pressure or temperature, equilibrium position will move to counteract the change.
•homogeneous equations: states are the same for both reactants and products

55
Q

What doesn’t affect the position of equilibrium?

A

Catalysts, but equilibrium is reached faster

56
Q

Why are high pressures and temperatures compromised?

A

•too expensive e.g equipment
•rate of reaction must be considered as-well as yield

57
Q

What is an oxidising and reducing agent?

A

•oxidising agent will accept e-
•reducing agents will donate e-

58
Q

What is an oxidation number?

A

E- needed to gain or lose for an atom to form a chemical bond

59
Q

Lattice dissociation enthalpy

A

The standard enthalpy change when 1 mole of an ionic solid is separated into its gaseous ions

60
Q

Lattice formation enthalpy

A

The standard enthalpy change when 1 mol of solid lattice is formed from its gaseous ions

61
Q

Enthalpy of solution

A

The standard enthalpy change when 1 mole of an ionic solid dissolves in a large enough amount of water that the dissolved ions are well separated and do not interact with each other

62
Q

Enthalpy of atomisation

A

The standard enthalpy change when 1 mole of gaseous atoms is formed

63
Q

Enthalpy of hydration

A

The enthalpy change when one mole of separated gaseous ions is dissolved completely in water to form a mole of aqueous ions

64
Q

1st electron affinity

A

The enthalpy change when one mole of gaseous atoms each gains an electron to form a mole of gaseous 1- ions

65
Q

Bond dissociation enthalpy

A

The enthalpy change when one mole of bonds of the same type in gaseous molecules is broken

66
Q

What does the perfect ionic model assume and why is it not true?

A

•Ions are perfectly spherical
•e- clouds around ions are attracted towards each ion= distortion which may cause covalent characters when polarisation is large

67
Q

What makes each ion more polarising? (+ and -)

A

+ve=
•higher charge density when smaller ionic radius and bigger charge

-ve=
•larger ionic radius and smaller charge

68
Q

What happens to the lattice dissociation energy as we move down group 7?

A

It will decrease as the size of the negative ion increases

69
Q

What does the difference between experimental and theoretical LDE tell you?

A

Bigger the difference between the LDR’s = more covalent character which causes strength of bond to increase

70
Q

How is solubility of an ionic compound predicted by its enthalpy change of solution?

A

If enthalpy change of solution is negative, the ionic compound is likely soluble

71
Q

What is entropy and its units?

A

The measure of disorder (how many different arrangements a molecule can make)
J K-1 mol-1

72
Q

How can you find out if a reaction is feasible?

A

-ve gibbs free energy change
•but reaction may not occur as the activation energy may be too high

Gibbs = enthalpy change - T(entropy change)

73
Q

What is a Bronsted Lowry acid?

A

Proton donors

74
Q

What is a Bronsted Lowry base?

A

Proton acceptors

75
Q

What is the expression for Kw and why?

A

Kw= [H+][OH-]
Kw = 10-14 mol2dm-6 at given temp.

H20 = H+ and OH- , water dissociates into its ions very weakly, so little OH- and H+ compared to H20 that we assume conc. of water is constant

76
Q

What is kw of pure water and why?

A

Kw= [H+]2
As [OH-] = [H+]

77
Q

What is Ph and equations?

A

Ph = conc. of H+ ions (2 d.p)
Ph= -log(H+)
[H+]= 10 -ph

78
Q

What is Ka and equations?

A

•weaks acids only dissociate slightly in aq so we use ka
Ka= [H+][A-]/[HA]

Ka= [H+]2/[HA]

79
Q

What is Pka and what does it tell you?

A

Pka=-log(Ka)
Ka= 10-pka
The lower the value of Pka, the stronger the acid

80
Q

What can we assume at the half neutralisation point?

A

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

81
Q

What is a suitable indicator?

A

Must change colour entirely within the vertical part of the titration curve= end point

82
Q

What is a buffer? (Acidic/basic)

A

A buffer is a chemical that resists the change in Ph when small amounts of acid or base are added

Acidic: resist the change in ph to keep below ph 7 (made from weak acid and its salt)
Basic: resist the change in ph to keep above ph 7 (made from weak base and its salt)

Chemistry (3 decks)

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