Electro Flashcards
(118 cards)
1
Q
electromagnetic Spectrum
A
a range of frequencies that covers all electromagnetic radiation and their respective wavelengths and energy
2
Q
what is the electromagnetic spectrum divided into?
A
it is divided into bands and regions
3
Q
which type of chemistry is the electromagnetic spectrum important for?
A
analytical chemistry
4
Q
what relationship does the spectrum show
A
the relationship between frequency, wavelength and energy
5
Q
frequency
A
how many waves pass per second
6
Q
wavelength
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the distance between two consecutive peaks on the wave
7
Q
which types of waves are dangerous and why?
A
gamma rays, x rays, and UV radiation are all dangerous. they have a high frequency and high energy which is damaging to your health
8
Q
order of waves from lowest energy to highest
A
radio waves, microwaves, infrared, visible light, ultraviolet, x rays and gamma rays
9
Q
a lower energy results in
A
a long wavelength, a low frequency
10
Q
a higher energy results in
A
a shorter wavelength, a high frequency
11
Q
all light waves travel at the same
A
speed
12
Q
what distinguishes light waves?
A
their diffrent frequencies
13
Q
speed of light symbol and its value
A
the speed of light symbol is a constant “c” and
has a value of 3.00 x 10^8 ms^-1
14
Q
what is frequency inversely proportional to?
A
frequency “f” is inversely proportional to wavelength ( λ)
the higher the frequency, the shorter the wavelength
15
Q
equation that links constant, frequency and wavelength
A
c=fλ
16
Q
continous spectrum
A
a visible region that contains all the colours of the spectrum
what you see in a rainbow, which is formed by the refraction of white light through a prism or water droplets in air
17
Q
a line specrum only shows
A
certain frequencies
18
Q
what does the helium spectrum diagram show?
A
the line spectrum of helium shows only certain frequencies of light.
it tells us that the emitted light from the atoms can only be certain fixed frequencies- it is quantised
quanta = “little packet”
19
Q
electrons can only possess…. amount of energy
A
certain amounts of energy, they cannot have any energy value
20
Q
where and how do electrons move
A
electrons move rapidly around the nucleus in energy shells
21
Q
what happens to electrons when energy is increased?
is the process reversible?
A
If their energy is increased, then they can jump to a higher energy level
the process is reversible, so electrons can return to their original energy levels, when this happens they emit energy
the frequency of enrgy is exactly the same, it just gets emitted instead of absorbed
22
Q
what does the diffrence between absorption and emission depend on
A
it depends on whether the electrons are jumping from lower to higher energy levels or the other way around
23
Q
the energy electrons emit is
what does it correspond to
what to do when energy is emitted in the visible region, what does it result in
A
a mixture of diffrent frequencies
which is thought to correspond to the many posssibilities of electron jumps between energy shells
if the the energy is emitted in the visible region, it can be analysed by passing it through a driffraction grating
the result is a lien emission spetrum
24
Q
spectrum of hydrogen diagram
what does each line have and what does this suggest
A
each line is a specific energy value
this suggests that electrons can only possess a limited choice of allowed energies
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what are the packets of energy in electrons of lines called
these packets of energy are called quanta or (plural "quantum")
26
where do the lines get closer in a spectrum of hydrogen
what is this called?
the lines get closer towards the blue end of the spectrum
this is called convergence
27
what is convergence
when lines get closer towards the blue end of the spectrum in a spectrum of hydrogen. The set of lines is converging towards the higher energy end, so the electron is reaching maximum amount of energy
28
what is the maximum amount of energy of an electron?
their maximum corresponds to the ionisation energy of the electron
29
who were the first lines in a spectrum of hydrogen observed by
these lines were first observed by swiss school teacher Johannes Balmer, they were named after him
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what do these lines correspond to
we now know that these electrons correspond to the electron jumping from higher levels down to the second or n=2 energy level
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large version of hydrogen spectrum from infrared to ultraviolet region
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what lines can we see in the full version of hydrogen spectrum
we can see sets or families of lines
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Balmer couldn't explain why the lines were formed
who's theory explained it?
Planck's Quantum Theory in 1900
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who applied the quantum theory to electrons, and what did they propose?
Niels Bohr applied the quantum theory to electrons in 1913 and proposed that electrons could only exist in fixed energy levels
35
what evidence did the line emission spectrum of hydrogen provide, what was then deduced from this?
The line emission spectrum of hydrogen provided evidence of energy levels
and it was deduced that the families of lines corresponded to electrons jumping from higher levels to lower levels
36
diagram to show the energy transitions for the hydrogen atom.
what did the findigs about this help scientists with
the finding helped scientists to understand how electrons work and provided the backbone to our knowledge of energy levels, sublevels and orbitals
37
what is the electron jump and energy of infrared region?
the jump is n∞→ n, the energy is low
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what is the electron jump and energy for the visible region
n∞ → n the energy is ↓
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electron jump and energy for ultraviolet region
n∞ → n energy is high
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ex.
which electron transition in the hydrogen atom emits visible light?
emission in the visible region occurs for a electron jumping from any energy level to n=2
ex. n = 3 to n = 2
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what is the electronic configuration
the arrangement of electrons in an atom
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how are electrons arranged around the nucleus
in principal energy levels or principal quantum shells
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how are principal quantum numbers used
what happens when the principal quantum number is lower/higher
principal quantum numbers (n) are used to number the energy levels or quantum shells
the lower the principal quantum number, the closer the shell is to the nucleus
the higher the principal quantum number, the greater the energy of the electron within that shell
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how many electrons can the principal quantum number n=1 hold
up to 2 electrons
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how many electrons can the principal quantum number n=2 hold
up to 8 electrons
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how many electrons can the principal quantum number n=3 hold
up to 18 electrons
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how many electrons can the principal quantum number n=4 hold
up to 32 electrons
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what is the mathematical relationship between the number of electrons and the principal energy level
the mathematical relationship between the number of electrons
and the principal energy level is 2n^2
ex. in the third shell n = 3 and the number of electrons is 2 x (3^2 ) = 18
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electrons are arranged in..... which are numbered by....
electrons are arranged in principal quantum shells which are numbered by principal quantum numbers
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The principal quantum shells are split into sub shells which are given by
The letters s,p and d
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subshells when elements have more than 57 electrons
they also have an f subshell
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the energy of the electrons in the subshells increases in order
The energy of the electrons in the subshells increases in the order s < p < d
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which shells do the order of subshells overlap for?
the order of subshells overlaps for the higher principal quantum shells
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What do subshells contain
they contain one or more atomic orbitals
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where do orbitals exist as and where can electrons be found
orbitals exist at specific energy levels and electrons can only be found at these specific levels, not in between
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each atomic orbital can be occupied by a maximum of how many electrons
Each atomic orbital can be occupied by a maximum of two electrons
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shape of s and p orbitals
orbitals have specific 3D shapes
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difference between s and p orbitals
S orbitals have a spherical shape while p orbitals are dumb bell shaped structures with two lobes along the x, y and z axis
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summary diagrams of s and p orbitals
overview of shells, subshells and atoms
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what is the ground state
the most stable electronic configuration of an atom which has the lowest amount of energy
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with what principle is ground state achieved
the AUFBAU PRINCIPLE, by filling the subshells of energy with the lowest energy first (1s).
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the order of the subshells in terms of increasing energy does not follow a regular pattern at n=3 and higher
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when do the principal quantum shells increase in energy
the principal quantum increase in energy with increasing principal quantum number
eg. n=4 is higher than n=2
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the subshells increase with order s < p < d < f, with expection to
the only exception is on the 3d orbital which has slightly higher energy than the 4s orbital, so the 4s orbital is filled before the 3d orbital
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energy levels
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each shell can be divided into
each shell can be divided into further subshells, s p d f
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how many orbitals can subshell s hold
1 orbital
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how many orbitals can subshell p hold
3 orbitals labeled px , py and pz
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how many orbitals can subshell d hold
5 orbitals
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how many orbitals can subshell f hold
7 orbitals
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maximum number of electrons each orbital can hold
maximum number of 2 electrons
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maximum number of electrons subshell s
1x2 = 2 electrons
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maximum number of electrons subshell p
3x2= 6 electrons
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maximum number of electrons subshell d
5x2=10 electrons
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maximum number of electrons subshell f
7x2= 14 electrons
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what happens to the energy in orbitals of the same subshell at ground state
In the ground state, orbitals in the same subshell have the same energy
and are said to be degenerate, so the energy of a px orbital is the same as a py orbital
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division of shells diagram
shells are divided into subshells which are then divided into orbitals
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what does the principle quantum number tell us
The principal quantum number (n) tells you the energy level of an atom and how many types of orbitals are available.
The types of orbitals available depend on n, and they are labeled s, p, d, and f.
Principal Quantum Number (n) | Types of Orbitals Available
1 | s
2 | s, p
3 | s, p, d
4 | s, p, d, f
n = 1 → only 1s orbital
n = 2 → 2s and 2p orbitals
n = 3 → 3s, 3p, and 3d orbitals
n = 4 → 4s, 4p, 4d, and 4f orbitals
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what is the shape of an s orbital
how does the size of the s orbital increase
the size of the s orbital increases with increasing shell number
eg. the orbital s of the third quantum shell (n=3) is bigger than the s orbital of the first quantum shell (n=1)
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what is the shape of an p orbital
how does the size of the p orbital increase
The p orbitals are dumbbell shaped
every shell has three p orbitals excepts for the first one (n=1)
The p orbitals occupy the x, y and z axes and point atright angles to each other, so are
oriented perpendicular to one another
The lobes ofthe p orbitals become larger and longer with increasing shell number
83
the electron configuration gives us information about
about the number of electrons in each shell, subshell and orbital of an atom
the subshells are filled in order of increasing energy
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what does the electron configuration show
the number of electrons occcupying a subshell in a specific shell
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electrons can be imagined as small spinning charges which
rotate around their own axid in either clockwise or anticlockwise direction
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electrons with the same spin
repel eachother which is also called spin pair repulsion
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due to the spin repel repulsion of electrons in the same spin, elecrons will
electrons will occupy separate orbitals in the same subshell first to minimise this repulsion and have their spin in the same direction
They will then pair up, with a second electron being added to the
first p orbital, with
its spin in the opposite direction
This is known as Hund's Rule
E.g. if there are three electrons in a p subshell, one electron will go into each px , py
and pz orbital
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hunds rule
electrons occupy the same region of space in orbitals even though there is ....
the principle quantum number indicates the energy level of a particular shell but also indicates the energy of the electrons in that shell
A 2p electron is in the second shell and therefore has an energy corresponding to n=2
electrons occupy the same region of space in orbitals even though there is repulsion between them
89
according to pauli exclusion principle, an orbital can only hold how many electrons, why?
An orbital can only hold two electrons and they must have opposite spin - the is known
as the Pauli Exclusion Principle
This is because the energy required to jump to a higher empty orbital is greater than the inter-electron repulsion-- for this reason, they pair up and occupy the lower energy levels first
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electron configurations can also be represented using
where each box represents
the orbital spin diagrams
Each box represents an atomic orbital
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how are the boxes arranged in orbital spin diagrams
The boxes are arranged in order ofincreasing energy from lowerto higher (i.e. starting
from closestto the nucleus)
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how are electrons represented in orbital spin diagrams
The electrons are represented by opposite arrows to show the spin of the electrons
E.g.the box notation fortitanium is shown below
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writing out electronic configurations tells us
how the electrons in an atom or ion are arranged in their shells, subshells and orbitals
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how can electron configuration be done
This can be done using the full electron configuration or the shorthand version
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full electron configuration
The full electron configuration describes the arrangement of all electrons from the 1s subshell up
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shorthand electron configuration
The shorthand electron configuration includes using the symbol of the nearest
preceding noble gas to account for however many electrons are in that noble gas, followed by the rest of the electron con
guration
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Negative ions are formed by
adding electrons to the outer subshell
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ex. full and shorthand electron configuration of gallium
Gallium has 31 electrons so the full electronic configuration is:
Full:1s^2 2s^2 2p^6 3s^2 3p^6 3d^10 4s^2 4p^1
Shorthand: [Ar] 3d^10 4s^2 4p^1
98
transition metals fill the.... subshell, they also lose electrons from the... rather than from the .... subshell
The transition metals fill the 4s subshell before the 3d subshell, but they also lose electrons from the 4s first rather than from the 3d subshell
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positive ions are formed by
removing electrons from the outer subshell
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the periodic table is split up into four main blocks depending on their electron configuration
s block elements (valence electron(s) in s orbital)
p block elements (valence electron(s) in p orbital)
d block elements (valence electron(s) in d orbital)
f block elements (valence electron(s) in f orbital)
elements can be divided into four blocks according to their outer shell electron configuration
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exceptions to the Aufbau Principle
Chromium and copper have the following electron con
gurations:
Cr is [Ar] 3d^5 4s^1 not [Ar] 3d^4 4s^2
Cu is [Ar] 3d^10 4s^1 not [Ar] 3d^9 4s^2
This is because the [Ar] 3d^5 4s^1 and [Ar] 3d^10 4s^1 configurations are energetically avourable
By promoting an electron from 4s to 3d,these atoms achieve a half full or full d-subshell,respectively
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ex. full and shorthand electron configuration of potassium
Potassium has 19 electrons so the full electronic configuration is:
1s^2 2s^2 2p^6 3s ^2 3p^6 4s^1
The 4s orbital is lower in energy than the 3d subshell and is therefore
filled first
The nearest preceding noble gas to potassium is argon which accounts for18
electrons
so the shorthand electron configuration is:
[Ar] 4s^1
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ex. full and shorthand electron configuration of calcium
Calcium has 20 electrons so the full electronic configuration is:
1s^2 2s^2 2p^6 3s^2 3p^6 4s^2
The 4s orbital is lower in energy than the 3d subshell and is therefore
filled first
The shorthand version is [Ar] 4s^2 since argon is the nearest preceding noble gas
to calcium which accounts for18 electrons
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electrons move .... around.... in......
Electrons move rapidly around the nucleus in energy shells
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ex. full and shorthand electron configuration of Ca2+
If you ionise calcium and remove two of its outer electrons,the electronic
configuration of the Ca2+ ion is identical to that of argon:
Ca2+ is 1s^2 2s^2 2p^6 3s^2 3p^6
Ar is also 1s^2 2s^2 2p^6 3s^2 3p^6 so the shorthand version is [Ar]
99
Heat or electricity can be used to do what?
Heat or electricity can be used to excite an electron to a higher main energy level
These range from
n =1 (ground state) to n = ∞
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when electrons fall back down, they must
When the electrons 'fall' back down they mustlose the energy difference between the two energy levels.
This loss of energy is performed by releasing electromagnetic energy
in the form of infrared, visible light or ultraviolet radiation.
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when the electron dalls back to n=1 (ground state). The energy released is in which spectrum?
When the electron falls back to n = 1(ground state) the energy released is in the
ultraviolet region of the spectrum
This corresponds to the Lyman series
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Bohr's model
electrons exist in discrete energy levels so an exact amount of energy is required for an electron to "jump" an energy level, a little like a ladder
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diagram to show the release of phtoton when an electron is promoted
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limitations to Bohr's model
Assumes positions of electrons are fixed
Assumes energy levels are spherical in nature
Bohr limited calculations to hydrogen only, so does not explain the line spectra of
other elements containing more than one electron
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As the line spectra is produced, the lines will become closer together. the point where lines appear to meet is
the limit of convergence
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the convergence limit is the
frequency at which the spectral lines converge
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the energy required for an electron to escape an atom or reach the upper limit of convergence
is the ionisation energy
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the frequency of radiation in the emission spectrum at the limit of convergence can be used to determine
the first ionisation energy or IE1
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