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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
Values of n and corresponding letter
0 - s 1- p 2- d 3- f
26
shapes of orbitals: s orbitals
spherical, -ℓ= 0 -mℓ = 0 -1 value = 1 orbital
27
shapes of orbitals: p orbitals
figure of 8, - ℓ = 1 -mℓ = -1,0,1 3 values = 3 orbitals
28
shapes of orbitals: D orbital
ℓ =2 mℓ = -2,-1,0,1,2 5 values = 5 orbitals
29
spin quantum number
values of +1/2 or -1/2
30
Pauli exclusion principle
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
31
Hund's rule
electrons fill orbitals singly with parallel spins. before spin pairing.
32
Aufbau principle
electrons fill lowest energy orbitals first 1s, 2s, 2p, 3s, 3p, 4s, 3d
33
degenerate orbitals
orbitals with the same energy
34
Ionising energy
ionising energy increases across a period as more protons mean more pull
35
why ionising energy differs
if you are taking an electron form a full stable sunshell if you are taking it from a half full subshell
36
the periodic table is split into 4 blocks
s d p f this is because the last letter/ orbital filled defines what black the element is in.
37
shapes of molecules
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.
38
# Acronym how is the shape of a molecules governed
VSEPR, valence shell electron pair repulsion
39
Formula used to calculate number of non/bonding pairs
number of electron pairs = no of outer electrons on central atom + no of bonds it makes/2
40
linear shape
2 electron pairs all bonding bond angle of 180
41
trigonal
3 pairs all bonding bond angle of 120
42
tetrahedral
4 electron pairs all bonding bond angle of 109.5
43
trigonal pyramidal
4 electron pairs 3 bonding 1 non
44
tetrahedral angular
4 pairs 2 bonding 2 non
45
trigonal bipyramidal
5 electron pairs all bonding or 3 bonding, 2 non bond angle of 120 and 90
46
octahedral
6 electron pairs all bonding bond angle 90
47
Calculating Bonding in polyatomic ions
the same formula however +1 for a negative charge and -1 for a positive one.
48
Non bonding electrons
lone pairs
49
dative covalent bond
a covalent bond in which both electrons come from the same atom
50
transition metal
a mental with an incomplete d subshell in at least one of its ions.
51
what principle do copper and chromium violate
Aufbau in favour of a stable d shell
52
# orbitals what happens when a transition metal forms an ion
the 4s electrons are lost first
53
Scandium and zinc
are not considered as true transition metals and sc3+ has no d subshell and Zn2+ has a full d subshell
54
Transition metal complexes
A complex consist of a central ion surrounded by ligands
55
ligands
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
monodentate ligand
binds to the central atom via the donation of one electron pair.
57
Bidentate ligands
each ligand donates 2 electron pairs.
58
polydentate ligands
donates more than 2 electron pairs.
59
coordination number
how many bonds the central atom is making.
60
Colour in transition metal complexes
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
why are transition metals coloured 1
ligands cause loss of degeneracy / d orbital spilling through repulsion of electrons.
62
why are transition metals coloured 2
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
why are transition metals coloured 3
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
spectrochemical series
different ligands have different field strengths, and split d orbitals to different extents.
65
2 equations showing how water molecules dissociate or ionise. Full and simplified
H2O + H2O ⇌ H3O+ + OH- : full H2O ⇌ H+ + OH- : shortened
66
Formula for the hydronium ion
H30+ (hydrated proton) can also be written as H+
67
the ionic product of water
kw = [H+] [OH-] = 1x10-14 or kw = [H30+] [OH-] = 1x10-14
68
amphoteric
can behave as an acid or a base
69
what is the difference between a strong and a weak acid
Strong acids fully dissociate into ions. weak acids partially dissociate
70
all Ka values for weak acids are significantly less than 1, what does this mean for equilibrium
lies to the left
71
diprotic acid
2 protons (2H+ ions) H2so3
72
monoprotic
1 proton (HCl)
73
what is the difference between a strong and a weak base
Strong bases fully dissociate into ions. weak bases partially dissociate
74
all Kb values for weak bases are significantly less than 1, what does this mean for equilibrium
lies to the left
75
Strong acids are better conductors of electricity as
it is fully dissociated it has more moving ions.
76
why do two equimolar acids need the same volume of alkali for neutralization
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
Strong bases are better conductors of electricity as
it is fully dissociated it has more moving ions.
78
why do two equimolar bases need the same volume of acid for neutralization
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
producing salt from acid and base
acid + alkali ---> salt + water
80
types of strong acids
Hydrochloric acid (HCl) Nitric acid (HNO3) Sulfuric acid (HSO4)
81
types of weak acid
Sulphurous acid (H2SO3) ethanoic acid (CH3COOH) carbonic acid (H2CO3)
82
types of strong alkali
Sodium hydroxide (NaOH) Potassium hydroxide (KOH) lithium hydroxide (LiOH)
83
weak alkali's
Ammonia (NH3)
84
how Ph of a salt solution is determined
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
why is a salt solution acidic for example (ammonium chloride)
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
lowry -bronsted theory
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
Buffer solutions
a solution which contains a constant pH when small amounts of acid/alkali are added
88
to make an acidic buffer
a mixture of a weak acid and a salt of that weak acid
89
to make an basic buffer
ma mixture of a weak base and and salt of that weak base
90
how acid buffers work - ethanoic acid and sodium ethanoate
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
indicators
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
titration curves
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
equivalence point
the point at which the number of moles of acid = the number of moles of alkali to produce a salt.
94
choosing an indicator
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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