Atomic Structure - Unit 1, Section 1 Flashcards
The structure of the atom
made up of protons, neutrons and electrons.
protons - 1, +1
neutrons - 1, 0
electrons - 1/2000, -1
most of the mass is concentrated in the nucleus - protons and neutrons found.
what is the mass number? (A)
total number of protons and neutrons in the nucleus of an atom.
what is the atomic number? (Z)
the number of protons in the nucleus of an atom.
it identifies the element.
what is an ion?
an atom that has become positively or negatively charged by either gaining or losing electrons.
what is an isotope?
atoms with the same number of protons but a different number of neutrons.
similarities and differences between isotopes
they have the same electron configuration so therefore have the same chemical properties.
have different physical properties as this tends to depend more on the mass of the atom.
John Dalton’s atomic model - 19th Century
described atoms as solid spheres
different spheres made up the different elements
J.J Thomson’s atomic model - 1897
concluded that an atom must contain even smaller, negatively charged particles - electrons.
made the ‘plum pudding model’ - solid sphere with negative electrons spread throughout.
Ernest Rutherford’s atomic model - 1909
gold foil experiment with Geiger and Marsden.
they fired positively charged alpha particles at a thin sheet of gold.
the plum pudding model would mean that most of the alpha particles would be deflected slightly by the positive ‘pudding’.
however, most of the particles passed straight through and only a small number were deflected, meaning the plum pudding model was incorrect.
made the nuclear model of the atom - a tiny, positively charged nucleus at the centre, surrounded by a cloud of negative electrons. Most of the atom is empty space.
Neils Bohr’s atomic model
scientists realised that electrons in a cloud would spiral into the nucleus and cause the atom to collapse, making Rutherford’s model incorrect.
Bohr proposed new model with 4 principles:
- electrons only exist in fixed orbits (shells)
- each shell has a fixed energy
- when an electron moves between shells, electromagnetic radiation is emitted or absorbed.
- because the energy of shells is fixed, the radiation will have a fixed frequency.
the frequencies were already known from experiments and the Bohr model fitted these observations.
Other atomic models - Quantum model
Bohr model wasn’t 100% right as scientists discovered that not all electrons in a shell had the same energy.
they refined the model and included sub-shells.
Today’s model is based on quantum mechanics. The quantum model explained observations that cannot be accounted for by the Bohr model.
what is Relative Atomic Mass? (Ar)
the average mass of an atom of an element relative to 1/12th the mass of an atom of carbon-12.
what is relative isotopic mass?
the mass of an atom of an isotope of an element relative to 1/12th the mass of an atom of carbon-12.
calculating relative atomic mass
isotopic masses x percentages / 100
calculating relative molecular mass (Mr)
the average mass of a molecule relative to 1/12th the mass of an atom of carbon-12.
add up all relative atomic mass values of all the atoms in the molecule.
calculating relative formula mass
the average mass of a formula unit relative to 1/12th the mass of an atom of carbon-12.
add up all the relative atomic masses of all the ions in the formula unit.
what is a mass spectrometer?
a machine used to analyse elements or compounds. it can provide information about the relative atomic mass of an element and the relative abundance of its isotopes, or the relative molecular mass of a molecule if you use it to analyse a compound.
how a TOF mass spectrometer works
IONISATION x2
- the sample needs to be ionised before it enters the mass spectrometer.
-Electrospray ionisation = the sample dissolved in a solvent and pushed through a small nozzle at high pressure. A high voltage is applied to it, causing each particle to gain a H+ ion. The solvent is then removed, leaving a gas made up of positive ions.
- Electron impact ionisation = the sample is vaporised and an electron gun is used to fire high energy electrons at it. This knocks one electron off each particle so they become 1+ ions.
ACCELERATION
- the positive ions are accelerated by an electric field. The electric field gives the same kinetic energy to all the ions. The lighter ions experience a greater acceleration than heavier ions as they accelerate more easily.
ION DRIFT
- the ions enter a region with no electric field. They drift through at the same speed as they left the electric field, so the lighter ions will be drifting at higher speeds.
DETECTION
- because the lighter ions travel through the drift region st higher speeds, they reach the detector first. The detector detects the current created when the ions hit it and records how long they took to pass through the spectrometer. This data is then used to calculate the mass/charge values needed to produce a mass spectrum.
what is a mass spectrum?
a type of chart produced by a mass spectrometer. it shows information about the sample that was passed through the spectrometer.
Interpreting a mass spectrum
- if the sample is an element, each line will represent a different isotope of the element.
- the y-axis gives the abundance of ions as a percentage.
- for an element, the height of each peak shows the relative isotopic abundance.
- the x-axis units are given as a ‘mass/charge’ ratio. you can assume that the x-axis is the relative isotopic mass.
identifying elements
mass spectrometry can be used to identify elements.
elements with different isotopes produce several lines in the mass spectrum as they all have different masses.
this produces characteristic patterns which can be used as ‘fingerprints’ to identify certain elements.
calculating relative atomic mass (Ar) using mass spectra
- for each peak, read the % relative isotopic abundance from the y-axis and the relative isotopic mass from the x-axis. Multiply them together to get the total relative mass for each isotope.
- add up these totals
- divide by 100
electron configuration order
1s, 2s, 2p, 3s, 3p, 4s, 3d, 4p, 5s, 4d, 5p, 6s, 4f, 5d, 6p, 7s, 5f.
electron configuration rules
- fill up lowest energy sub-shells first.
- fill orbitals in a sub-shell singly before they start sharing.
Electron configuration of Chromium (Cr)
1s2, 2s2, 2p6, 3s2, 3p6, 3d5, 4s1.
donates one of its 4s electrons to the 3d sub shell as it is more stable.
Electron configuration of Copper (Cu)
1s2, 2s2, 2p6, 3s2, 3p6, 3d10, 4s1.
donates one of its 4s electrons to the 3d sub shell as it is more stable.
electron configuration of transition metals - odd thing
they lose their 4s electrons before their 3d electrons.
electronic structure and chemical properties
the number of outer shell electrons decides the chemical properties of an element.
the S block elements (Group 1 and 2) have 1 or 2 outer shell electrons which are easily lost to form positive ions with a noble gas configuration.
the P block elements (Groups 5,6 and 7) can gain 1,2 or 3 electrons to form negative ions with a noble gas configuration.
Groups 4-7 can also share electrons when they form covalent bonds.
Group 0 elements have completely filled S and P sub shells so don’t gain, lose or share electrons.
Ionisation
when electrons have been removed from an atom or molecule.
first ionisation energy
the energy needed to remove 1 electron from each atom in 1 mole of gaseous atoms to form 1 mole of gaseous 1+ ions.
Important points about ionisation energies
- must use the gas state symbol
-always refer to 1 mole of atoms - the lower the ionisation energy, the easier it is to form a positive ion
factors affecting ionisation energy
- nuclear charge - the more protons there are in the nucleus, the more positively charged the nucleus is and the stronger the attraction for the electrons.
- distance from nucleus - an electron close to the nucleus will be much more strongly attracted than one further away.
- shielding - as the number of electrons between the outer electrons and nucleus increases, the outer electrons feel less attraction to the nucleus. this lessening of the pull of the nucleus is as a result of shielding.
second ionisation energy
the energy needed to remove one electron from each ion in 1 mole of gaseous 1+ ions to form 1 mole of gaseous 2+ ions.
general equation for successive ionisation energies
X(n-1)+ (g) -> Xn+ (g) + e-
Ionisation trends down Group 2
- 1st ionisation energy decreases down Group 2.
- each element going down the group has one more electron shell than the one above. The extra shell will shield the outer electrons from the attraction of the nucleus. Also, the extra shell means the outer electrons will be further from the nucleus so the nucleus’s attraction will be reduced. Both these factors make it easier to remove outer electrons, resulting in lower ionisation energies.
Ionisation trends across periods
as you move across a period, the general trend is for ionisation energies to increase. this can be explained by the fact that the number of protons is increasing, which means a stronger nuclear attraction.
Drop between Group 2 and 3 - ionisation trends
Aluminium’s (G3) outer electron is in a 3p orbital rather than a 3s. The 3p orbital has a slightly higher energy than the 3s orbital, so the electron is to be found further from the nucleus, as the 3p orbital has additional shielding provided by the 3s electrons - Lower ionisation energy than Mg when it is expected to be higher.
Drop between Group 5 and 6 - ionisation trends
The shielding is identical in the phosphorus (G5) and sulfur (G6) atoms and the electron is being removed from identical orbitals.
In phosphorus’s case, the electron is being removed from a singly occupied orbital, whereas in sulfurt it is being removed from a fully occupied orbital. The repulsion between the 2 electrons in this orbital means that the electrons are easier to remove, therefore sulfur has a lower ionisation energy.
