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AH chem > unit 1 > Flashcards

unit 1 Flashcards

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1
Q

Electromagnetic radiation

A

A wave said to have dual nature (can behave as a wave and a particle)

waves with wavelength between 10 ^-14 and 10 ^4 m.

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2
Q

Speed of light

A

2x10 ^8 m/s

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3
Q

forms of electromagnetic radiation

A

Radio Waves, Microwaves, infrared, ultraviolet, x rays, gamma rays

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4
Q

What is the equation that links the speed of light, frequency and wavelength

A

c=fλ

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5
Q

EMR can exist as a stream of particles called

A

photons

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6
Q

equation which links energy and frequency of a photon

A

E=hf

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7
Q

equation used to calculate the energy of 1 mole of photons

A

E=Lhf

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8
Q

Combination of c=λf and E=Lhf

A

E= Lhc/λ

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9
Q

Lhc =

A

0.12

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10
Q

flame colours

A

when some metals burn they emit bright colours. This is because excited electrons are falling to lower energy levels and emitting light

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11
Q

Atomic Emission Spectra stage 1

A

The gaseous element is excited using high temperatures. This causes electrons to be promoted to higher energy levels. Electrons are said to be “excited”

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12
Q

Atomic Emission Spectra stage 2

A

Electrons fall down from these higher levels and in doing so emit energy in the form of light. (photon)

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13
Q

Atomic Emission Spectra stage 3

A

the electrons which make a big transition have high energy, high frequency and small wavelength.

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14
Q

Atomic Emission Spectra stage 4

A

the electrons which make a big transition have low energy, low frequency and large wavelength.

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15
Q

Atomic Emission Spectra stage 5

A

the different wavelengths of light given off pass into a prism where they are refracted. the light with the shortest wavelength are refracted most. A line spectrum is produced

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16
Q

Atomic Emission Spectra stage 6

A

each element has its own characteristic line spectrum. This line spectrum is concrete evidence for the existence of fixed energy levels in atoms.

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17
Q

Atomic Emission Spectra stage7

A

each line in the spectrum corresponds to the energy given out when an excited electron falls to a lower energy level.

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18
Q

what are the different transitions within an atom

A

Lyman : energy level which the electron falls=1
(high f, low λ) Part of the spectrum where lines are seen= UV

Balmer: energy level which the electron falls=2
Part of the spectrum where lines are seen= visible

Paschen :energy level which the electron falls=3
(low f, high λ) Part of the spectrum where lines are seen= IR

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19
Q

Atomic emission spectroscopy (AES)

A

Each element has its own characteristic emission spectrum, which can be matched to an unknown sample.
The position of spectral lines identify which element is present.
The Intensity tells you the concentration of the element

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20
Q

Atomic absorption spectroscopy (AAS)

A

Similar to AES however the wavelengths and intensity of the radiation are absorbed, which are measured during the promotion of electrons.

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21
Q

atomic orbital

A

a region of space where an electron is likely to be found

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22
Q

Principle quantum number, n

A

relates to the overall size/ energy or the orbitals.

for example sodium has electron arrangement 2,81, the value of n would be 1,2,3 respectively.

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23
Q

angular momentum quantum number, l

A

each electron shell is further divided into subshells.
values range form 0,1,2,3

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24
Q

magnetic quantum number, mℓ

A

tells us the multiplicity(how many) and orientation or the orbitals

values range from -3,…..,3

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25
Q

Values of n and corresponding letter

A

0 - s
1- p
2- d
3- f

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26
Q

shapes of orbitals: s orbitals

A

spherical,
-ℓ= 0
-mℓ = 0
-1 value = 1 orbital

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27
Q

shapes of orbitals: p orbitals

A

figure of 8,
- ℓ = 1
-mℓ = -1,0,1
3 values = 3 orbitals

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28
Q

shapes of orbitals: D orbital

A

ℓ =2
mℓ = -2,-1,0,1,2
5 values = 5 orbitals

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29
Q

spin quantum number

A

values of +1/2 or -1/2

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30
Q

Pauli exclusion principle

A

No two electrons can have the same set of 4 quantum numbers. meaning that the number of electrons in an given orbitals cannot exceed 2 and if there are two electrons in an orbital they must have opposite spin

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31
Q

Hund’s rule

A

electrons fill orbitals singly with parallel spins. before spin pairing.

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32
Q

Aufbau principle

A

electrons fill lowest energy orbitals first

1s, 2s, 2p, 3s, 3p, 4s, 3d

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33
Q

degenerate orbitals

A

orbitals with the same energy

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34
Q

Ionising energy

A

ionising energy increases across a period as more protons mean more pull

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35
Q

why ionising energy differs

A

if you are taking an electron form a full stable sunshell
if you are taking it from a half full subshell

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36
Q

the periodic table is split into 4 blocks

A

s d p

  f

this is because the last letter/ orbital filled defines what black the element is in.

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37
Q

shapes of molecules

A

the shape of a molecule is determined by the number of bonding and nonbonding electron pairs.

given that a bond is a share of a pair of electrons, both being negatively charged they repel each other. the molecule adopts a shape which will minimise repulsion and maximise separation.

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38
Q

Acronym

how is the shape of a molecules governed

A

VSEPR, valence shell electron pair repulsion

39
Q

Formula used to calculate number of non/bonding pairs

A

number of electron pairs =

no of outer electrons on central atom + no of bonds it makes/2

40
Q

linear shape

A

2 electron pairs all bonding

bond angle of 180

41
Q

trigonal

A

3 pairs all bonding

bond angle of 120

42
Q

tetrahedral

A

4 electron pairs all bonding

bond angle of 109.5

43
Q

trigonal pyramidal

A

4 electron pairs 3 bonding 1 non

44
Q

tetrahedral angular

A

4 pairs 2 bonding 2 non

45
Q

trigonal bipyramidal

A

5 electron pairs all bonding
or
3 bonding, 2 non

bond angle of 120 and 90

46
Q

octahedral

A

6 electron pairs all bonding

bond angle 90

47
Q

Calculating Bonding in polyatomic ions

A

the same formula however +1 for a negative charge and -1 for a positive one.

48
Q

Non bonding electrons

A

lone pairs

49
Q

dative covalent bond

A

a covalent bond in which both electrons come from the same atom

50
Q

transition metal

A

a mental with an incomplete d subshell in at least one of its ions.

51
Q

what principle do copper and chromium violate

A

Aufbau in favour of a stable d shell

52
Q

orbitals

what happens when a transition metal forms an ion

A

the 4s electrons are lost first

53
Q

Scandium and zinc

A

are not considered as true transition metals and sc3+ has no d subshell and Zn2+ has a full d subshell

54
Q

Transition metal complexes

A

A complex consist of a central ion surrounded by ligands

55
Q

ligands

A

negatively charged ions or neutral molecules which can bind to the central metal ion by donating a electron pair forming a dative covalent bond

common charges lingands - f-, cl- ,CN-

56
Q

monodentate ligand

A

binds to the central atom via the donation of one electron pair.

57
Q

Bidentate ligands

A

each ligand donates 2 electron pairs.

58
Q

polydentate ligands

A

donates more than 2 electron pairs.

59
Q

coordination number

A

how many bonds the central atom is making.

60
Q

Colour in transition metal complexes

A

In a free transition metal atom/ion, the 5 d orbitals are degenerate, but when ligands approach degeneracy is lost due to repulsion and d orbital splitting. giving rise to colour.

61
Q

why are transition metals coloured 1

A

ligands cause loss of degeneracy / d orbital spilling through repulsion of electrons.

62
Q

why are transition metals coloured 2

A

electrons in the lower energy d orbitals can absorb light and be promoted to the higher energy level. (the energy/ frequency/ wavelength/ colour of light absorbed depends on the energy gap)

63
Q

why are transition metals coloured 3

A

white light is a combination of 3 major colours red green and blue.
if red is absorbed to promote the electron, green and blue are transmitted and shows cyan in colour.
If green, magenta and if blue, yellow.

64
Q

spectrochemical series

A

different ligands have different field strengths, and split d orbitals to different extents.

65
Q

2 equations showing how water molecules dissociate or ionise. Full and simplified

A

H2O + H2O ⇌ H3O+ + OH- : full

H2O ⇌ H+ + OH- : shortened

66
Q

Formula for the hydronium ion

A

H30+ (hydrated proton)

can also be written as H+

67
Q

the ionic product of water

A

kw = [H+] [OH-] = 1x10-14
or
kw = [H30+] [OH-] = 1x10-14

68
Q

amphoteric

A

can behave as an acid or a base

69
Q

what is the difference between a strong and a weak acid

A

Strong acids fully dissociate into ions. weak acids partially dissociate

70
Q

all Ka values for weak acids are significantly less than 1, what does this mean for equilibrium

A

lies to the left

71
Q

diprotic acid

A

2 protons (2H+ ions) H2so3

72
Q

monoprotic

A

1 proton (HCl)

73
Q

what is the difference between a strong and a weak base

A

Strong bases fully dissociate into ions. weak bases partially dissociate

74
Q

all Kb values for weak bases are significantly less than 1, what does this mean for equilibrium

A

lies to the left

75
Q

Strong acids are better conductors of electricity as

A

it is fully dissociated it has more moving ions.

76
Q

why do two equimolar acids need the same volume of alkali for neutralization

A

although the weak acid has fewer H+ ions, initially once alkali is added the concentration of H+ drops and equilibrium shifts to the right and continues to do so until it fully dissociates

77
Q

Strong bases are better conductors of electricity as

A

it is fully dissociated it has more moving ions.

78
Q

why do two equimolar bases need the same volume of acid for neutralization

A

although the weak acid has fewer OH+ ions, initially once H+ is added the concentration of OH+ drops and equilibrium shifts to the right and continues to do so until it fully dissociates

79
Q

producing salt from acid and base

A

acid + alkali —> salt + water

80
Q

types of strong acids

A

Hydrochloric acid (HCl)
Nitric acid (HNO3)
Sulfuric acid (HSO4)

81
Q

types of weak acid

A

Sulphurous acid (H2SO3)
ethanoic acid (CH3COOH)
carbonic acid (H2CO3)

82
Q

types of strong alkali

A

Sodium hydroxide (NaOH)
Potassium hydroxide (KOH)
lithium hydroxide (LiOH)

83
Q

weak alkali’s

A

Ammonia (NH3)

84
Q

how Ph of a salt solution is determined

A

if the parent acid is strong and the parent base is weak = acid
if parent acid is weak and parent base is strong = base
if both equal= neutral

85
Q

why is a salt solution acidic for example (ammonium chloride)

A

ammonium ions like reacting with hydroxide ions (equilibrium is to the left)
There are lots of ammonium ions from the fully dissociated ammonium chloride which react with the OH- ions causing the concentration of OH- to fall.
Equilibrium shifts right creating a surplus of H+, hence acidic

86
Q

lowry -bronsted theory

A

an acid is - a proton (H+) donor
a base is - a proton acceptor

conjugate base- what is formed when the acid has lost its proton
conjugate acid -what is formed when the base has gained a proton

87
Q

Buffer solutions

A

a solution which contains a constant pH when small amounts of acid/alkali are added

88
Q

to make an acidic buffer

A

a mixture of a weak acid and a salt of that weak acid

89
Q

to make an basic buffer

A

ma mixture of a weak base and and salt of that weak base

90
Q

how acid buffers work - ethanoic acid and sodium ethanoate

A

in the addition of H+:
ethanoate ion reacts readily with H+ ions so that when H+ is added to the buffer the ethanoate ions from the salt will quickly react with them.

In the addition of OH- :
OH- ions will react with H+ ions in eqn 2. H+ will fall but equilibrium shifts right to replace.

91
Q

indicators

A

indicators are weak acids in which the colour of the acid is different to that of the conjugate base. in the case of litmus, red and blue.

for the colour to change of an indicator to be distinguishable the [HIn] and [Lm-] must differ approx by a factor of 10.

92
Q

titration curves

A

strong base+strong acid: equivalence point 7
weak acid+strong base: equivalence point above 7
strong acid+weak base: equivalence point below 7
weak acid+weak base: equivalence point 7

93
Q

equivalence point

A

the point at which the number of moles of acid = the number of moles of alkali to produce a salt.

94
Q

choosing an indicator

A

The pH range of the indicator should be coincide with the range of rapid pH change on the graphs.

AH chem (4 decks)

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