atomic structure Flashcards
(21 cards)
describe the development of atoms?
lead scientist - John Dalton
when - early 1800s
description of the model - solid sphere, atom smallest particle, elements are different as they have different weights.
describe the development of the plum pudding model?
lead scientist - JJ Thompson
when - 1897
description of the model - positive sphere with electrons randomly scattered in.
what did they discover - electrons.
describe the development of the first nuclear model?
lead scientist - Ernest Rutherford
when - 1911
description of model - small, dense positively charged nucleus with electrons surrounding it.
what did they discover - protons, most of an atom’s mass is in its nucleus, rest of atom is empty space.
what did they do - alpha particle experiment.
describe the development of the second nuclear model?
lead scientist - Neils Bohr
when - 1913
description of model - electrons move around the nucleus in energy levels.
what did they discover - the way electrons move around the nucleus.
describe the development of the third nuclear model?
lead scientist - James Chadwick
when - 1932
description of model - discovered isotopes which helped him discover neutrons.
what did they discover - neutrons.
what can mass spectrometry be used for?
To find the abundance and mass of each isotope in an element which allows us to determine its relative atomic mass (and thus identify the element).
To find the relative molecular mass of substances made of molecules.
Outline the first stage time of flight mass spectrometry (ionisation)?
Electron impact
- High energy electrons are fired at the sample from an electron gun.
- This knocks off one electron from each atom/molecule to form 1+ ion.
Electrospray
- The sample is dissolved in a volatile solvent and injected through a fine hypodermic needle as a fine spray into a vacuum in the ionisation chamber.
- A very high voltage is applied to the end of the needle where the spray emerges (the needle is positively charged).
- The particles gain a proton and become ions as a fine mist.
- The solvent evaporates leaving 1+ ions.
Outline the second stage of time of flight mass spectrometry (acceleration of ions)?
The ions are accelerated using an electric field so that all the ions have the same kinetic energy.
Outline the third stage of time of flight mass spectrometry (separation of charged ions)?
- Ion drift (the ions then enter the flight tube).
- Ions with different masses have a different time of flight.
- The lighter ions travel faster and take less time to reach the detector.
Outline the fourth stage of time of flight mass spectrometry (detection)?
- The detector is a negatively charged plate.
- A current is produced when the ions hit the plate as electrons flow from the plate to the positive ions.
- The size of the current is proportional to the number of ions.
What equations do you need to know?
m₁/t₁² = m₂/t₂² —-> t₂=√ m₂/m₁÷t₁²
v = √2KE/m
t = d/v
Give information about electron shells, sub-shells and orbitals?
In the modern model of the atom, electrons are found in specific energy levels known as shells surrounding the nucleus.
Shells that are further from the nucleus hold electrons with higher energy compared to those closer.
Each shell is defined by a principal quantum number (n = 1, 2, 3…), indicating its relative distance from the nucleus.
These shells are further divided into sub-shells, which have distinct energy levels and are labelled as s, p, d, and f.
Within sub-shells, electrons are located in orbitals, which are regions with a high probability of finding an electron.
An orbital is defined as a region around the nucleus that can accommodate up to two electrons with opposite spins.
What are the guidelines for deducing an atoms electron configuration?
- Electrons fill the lowest energy orbitals first (for example, calcium’s electron configuration is 1s2, 2s2 ,2p6, 3s2, 3p6, 4s2). The 4s sub-shell is filled before the 3d sub-shell because the 4s orbital has a lower energy than the 3d orbitals in neutral atoms.
- Electrons first occupy orbitals of equal energy singly before pairing up.
- When two electrons occupy the same orbital, they must have opposite spins (up and down) to minimise electron-electron repulsion.
- For ions in the s and p blocks, electrons are added to or removed from the highest occupied sub-shell, (e.g. the electronic configuration of Mg²⁺ is 1s2, 2s2, 2p6 and the electronic configuration of Cl⁻ is 1s2 2s2 2p6 3s2 3p6).
Describe the electron configuration of unusual transition metals.
Chromium (Cr) and copper (Cu) exhibit unusual electron configurations: Cr: 1s2, 2s2, 2p6, 3s2, 3p6, 3d5, 4s1 (instead of 3d4, 4s2) and Cu: 1s2, 2s2, 2p6, 3s2, 3p6, 3d10, 4s1 (rather than 3d9, 4s2).
These exceptions occur because configurations with a half-filled (d5) or fully filled (d10) d sub-shell are energetically more favourable.
Describe the electron configuration of transition metals that form ions.
For transition metals forming ions, the loss of electrons happens from the 4s orbital before the 3d orbital.
For example, consider an iron atom (Fe) with 26 electrons: the electron configuration of the iron atom is: 1s2, 2s2, 2p6, 3s2, 3p6, 3d6, 4s2.
When an iron atom loses three electrons to form an Fe3+ ion: the 4s electrons are lost first. then, one 3d electron is lost.
The resulting electron configuration of the Fe3+ ion with 23 electrons is: 1s2, 2s2, 2p6, 3s2, 3p6, 3d5.
What is the definition of the 1st ionisation energy?
The energy required to remove one electron from each atom in a mole of gaseous atoms, producing one mole of 1+ gaseous ions.
What is the definition of the 2nd ionisation energy?
The energy required to remove the second electron (not both electrons).
Why does the 1st ionisation energy decrease as you go down a group?
- Atomic radius increases.
-Shielding by inner electrons increases. - So less energy required to remove an electron.
Why does the 1st ionisation energy increase as you go across a period?
- Atomic radius decreases.
- Nuclear charge increases.
- Shielding by the inner electrons stays the same.
- So more energy is required to remove one electron.
Why does the 1st ionisation energy dip from group 2 to group 3?
- Group 3 atoms outer electron is in a p orbital whereas group 2 is from an s orbital.
- The outer p1 electron is further away from the nucleus.
- The inner s2 electrons increase the electron shielding.
- So less energy is required to remove the outer p1 electron.
Why does the 1st ionisation energy dip from group 5 to group 6?
- Group 6 atoms outer electron is in an orbital with another electron whereas group 5’s is the only electron in its orbital.
- The repulsion of the two electrons in the same p orbital leads to less energy being required to remove the outer electron.
