Electronic Structure Of Atoms Flashcards
What is an energy level?
It’s a region of definite energy within an atom that electrons occupy
What elements have the same number of electrons in the outer most energy level
Are in the same group
How do elements emit light and give an example
When the atoms of the elements are supplied with energy under certain
conditions
Energy can be supplied by a flame
What flame colour does barium have?
Yellow-green
What flame colour does copper have?
Blue-green
What flame colour does lithium have?
Deep red
What flame colour does potassium have?
Lilac
What flame colour does sodium have?
Yellow
What flame colour does strontium have?
Red
Flame test
- crush salt to be tested with mortar and pestle
- dip soaked splint in salt and put splint in flame
- note colour
- repeat with other salts and note colour
Give example of flame tests in everyday life
Firework displays - strontium and barium
What is the discharge tube used for?
If an element is easily vaporised, it emits light of a characteristic colour when placed in discharge tube at low pressure and high voltage
What is a spectroscope?
It analyses light emitted by elements
What can flame tests be used for
To distinguish between different metals in their compounds
What is white light?
Mixture of visible light of all wavelengths
What occurs in a spectroscope when analysing white light?
The light waves are bent to different extents forming a continuous brand of colours or continuous spectrum [red to violet]
What is viewed when light emitted from a discharge tube containing hydrogen is analysed using a spectroscope?
A series of coloured lines of definite wavelength against a dark background is observed = line spectrum
How could you differ different elements?
By their emission spectrum. The emission spectrum of an element is characteristic of that element and is different from that of any other element
Why are certain colours not seen in the hydrogen/line spectrum?
Certain colours not able to be absorbed
What kind of light do elements emit?
Visible light and light in the ultraviolet and infrared regions
What does the hydrogen spectrum include?
Series of lines in both the visible and ultraviolet regions
What contains all the lines in the visible region of the hydrogen spectrum?
The Balmer series
What contains the lines in the ultraviolet region of the hydrogen spectrum?
Lyman series
What contains the lines in the infrared region of the hydrogen spectrum?
Paschen series
What does each element have?
A characteristic emission spectra and an absorption spectra
What is the absorption spectrum?
It’s the spectrum that is observe after the white light has been passed through an element
What does an absorption spectrum consist of?
A series of dark lines against a coloured background. The dark lines are at exactly the same wavelengths as the coloured lines in the emission spectrum of that element.
How can you measure the amount of an element in a sample
From the amount of light absorbed - atomic absorption spectrometer
The amount of light absorbed depends on what?
The amount of element in the sample
How is the atomic absorption spectrometry useful?
Analysis of heavy metals - estimate the amount of lead in a blood sample
What theory did Neil’s Bohr put forward to do with the line spectrum ?
- Hydrogen electron is restricted to those regions of the atom that have certain energy values = energy levels and it cannot have an energy value lying between energy levels
- when an electron moves from a higher energy level [E2] to a lower energy level [E1] a definite amount of energy is emitted, amount of energy emitted is equal to the energy difference = E2 - E1 = hf
- atom absorbs amount of energy equal to the energy difference between e1 and e2 the electron will move to the higher level
- each energy level has associated with it an integer called the principal quantum number
What is the ground state?
It is the lowest energy state for the hydrogen atom
What is the excited state?
It is a higher energy state for the hydrogen atom
Describe the Bohr theory which accounts for the line spectrum of the hydrogen atom
- Normal circumstances, the hydrogen electron is in n=1 energy level, the ground state -the lowest energy state- for the hydrogen atom.
- if receives enough energy [eg, discharge tube], it moves to n=2, excited state -higher energy state- for the hydrogen atom.
- it is unstable in this state and eventually will drop back to ground state, emitting energy equal to the difference between energies of n=2 and n=1 energy levels.
How are the ultraviolet [lyman series] lines obtained?
When the electron falls from the n=3 to the n=1 level
How do the visible lines in the hydrogen spectrum -the balmer lines- arise?
Electrons in n=3 energy level falls to n=2
When electrons fall they emit energy or light at definite wavelength, what does the light appear as?
One of the line on the line spectrum of the element
How is the line spectrum unique to every element?
In atoms of the element, the spacings between the energy levels are unique which give rise to unique electronic transitions
What is the Heisenberg uncertainty principle?
It states that it is not possible to determine at the same time the exact position and velocity of an electron
What is an atomic orbital?
It is a region in space where the probability of finding an electron is relatively high
What are the three types of orbitals and what are their shapes
S orbitals -spherical shape
P orbitals - dumbbell shape
D orbitals - too complex
How many electrons can all orbitals hold?
Two
Of the orbitals, which have the highest and lowest energy?
3px, 3py, 3pz = equal energy
3s = lower energy
3d = highest energy
In relation to orbitals, how does the energy impact the position?
The lower the energy, the closer it is to the nucleus
What is an energy sublevel?
It is a subdivision of an energy level containing one or more orbitals, all of which have the same energy
How many electrons can s sublevel hold?
2
How many electrons can the p sublevel hold?
6
How many electrons can the d sublevel hold?
10
How many electrons can the f sublevel hold?
14
What rules are used to weigh electrons to various sublevel and orbitals?
Aufbau principle - electrons will occupy the lowest energy sublevel available
Not more than two electrons can occupy an orbital at one time
Electrons occupy orbitals of equal energy singly where possible
What are the limitations of the Bohr theory?
- It only worked well for hydrogen
- didn’t take into account that electrons have the properties of waves as well as particles
- did not allow for Heisenberg uncertainty principle
- did not explain discovery of sublevel
- did not account for existence of orbitals
What is the atomic radius of an element?
It is half the distance between the nuclei of two atoms of the element that are joined together by a single covalent bond
What are the trends in the size of the atomic radius going down the group
- electrons in inner levels partially neutralise attractive force of nucleus by repelling outer electrons
- in group 1, an extra energy level is added and effected nuclear charge experienced by outermost electron is much less than full nuclear charge
- atomic radius increases, due to the addition of extra energy levels, resulting in a extra screening by energy levels
Trends in size of atomic radius going across a period
- Nuclear charge - exerts a greater attractive force on the outer electrons
- Screening effect - extra electrons are added to same energy level so there is no screening effect
What is the first ionisation energy of an element?
It’s the minimum energy in kilojoules required to remove the most loosely bound electron from each isolated atom in a mole of the element in its ground state
How is the first ionisation energy represented?
X = X+ + e-
What is the first ionisation energy measured in
Kilojoules per mole
Does the first ionisation energy increase or decrease across a period? And which is easier/harder?
The values increase across a period making it harder to remove the loosely bound electron
- nuclear charge increases - strong force of attraction for the electrons from the nucleus and therefore electrons are held more tightly and so it’s harder to remove an electron
- decrease in atomic radius due to the nuclear charge exerting an attractive force on outer electrons + no screening effect makes it easier to remove electron from an atom
Does the ionisation energy increase or decrease going down a group? Is it harder or easier to remove loosely bound electron?
The energies decrease, easier
- increase in atomic radius, easier to remove an electron despite increased nuclear charge
- screening effect of inner energy levels, causing the nuclear charge experienced by the outermost electron being much less than the full nuclear charge
Exception - beryllium and boron
Beryllium has higher ionisation energy value than boron. [full outer sublevel]
Extra electron is added to boron 2px orbital which has higher energy than 2s orbital. Electron is more readily removed than the 2s electron in beryllium, beryllium has a full outer sublevel and is stable
Nitrogen and oxygen
What is the second ionisation energy?
It’s the minimum energy required to remove the most loosely bound electron from each singly charged positive ion in a mole of these ions
How is the second ionisation energy represented
X+ = x2+ + e-
Elements with similar outer electronic configurations but inner orbitals incomplete
Will have similar chemical properties
Elements with different outer electronic configurations
Have different chemical properties
What is the aufbau principle?
It states that electrons will occupy the lowest energy sublevel available
What are the exceptions to electronic configurations?
Copper and chromium
[just take one from second last one and give to last]
What are large and small atoms?
Large atomic radius - large atom
Small atomic radius - small atom
Identify the main energy levels involved by in the electron transition that gives ride to Balmer series
3 and 2
Why is it difficult to specify the absolute boundary of an atom?
Heisenberg uncertainty principle
Impossible measure position and velocity of electron in an atom
Account for general trend in atomic radii going across the second period
Decreases
- electrons in inner levels partially neutralise attractive force of nucleus by repelling outer electrons
- the increasing nuclear charge exerts a greater attractive force on the outer electrons and pulls them in
- electrons are added within the same energy levels, screening effect by inner levels is the same
Set of successive ionisation energy values for electrons in carbon
[1086,2353,4620,6223,37831,47277]
How provide evidence for
- no. electrons in a carbon atom
- no. electrons in each main energy level in a carbon atom
- each value = one electron
- gradual increase for first four values = four electrons in outer shell [easier to remove]
Why is the first ionisation energy of oxygen lower than that of nitrogen despite the general increase
Nitrogen is relatively stable
-3 half filled 2p orbitals
