Topic 2: Atomic Structure and the Periodic Table

Question 1

Fundamental particles:

Answer 1

-the atomic structure model has evolved over time, as knowledge and scientific understanding changes
-the current, accepted model of the atom consists of a small, dense central nucleus surrounded by orbiting electrons in electron shells
-this was discovered in the Rutherford scattering experiment in 1911
-the nucleus consists of protons and neutrons giving it an overall positive charge
-it contains almost the entire mass of the atom
-in a neutral atom, the number of electrons is equal to the number of protons due to the relative charges

Question 2

Protons, Neutron and Electrons:

Answer 2

-relative charge: proton -> +1, neutron -> 0, electron -> -1
-relative mass: proton -> 1, neutron -> 1, electron -> 1/1840
-the maximum number of orbiting electrons that can be held by any single shell, depends on the number of the shell, this can be calculated using 2n^2 where n is the number of the shell
-example: Electrons in shell 2 = 2(2^2) = 8 electrons
-each electron shell must fill before the next one can hold any electrons

Question 3

Atomic Number and Mass Number:

Answer 3

-mass number is represented using A and can be calculated as the sum of protons and neutrons in an atom
-atomic number is represented using Z and is equal to the number of protons in an atom, hence it can be referred to as proton number

Question 4

Isotopes:

Answer 4

-> isotopes are atoms of the same element with the same atomic number, but with a different number of neutrons, resulting in a different mass number
-neutral atoms of isotopes will react chemically in the same way because their proton number and electron configuration is the same
-the sharing and transfer of electrons is unaffected
-however, the different mass numbers means they have different physical properties
-example: the following are all isotope of hydrogen:
Hydrogen = 1 proton and 1 neutron
Deuterium = 1 proton and 2 neutrons
Tritium = 1 proton and 3 neutrons

Question 5

Relative Masses:

Answer 5

-Relative atomic mass (Ar) is defined as:
-> the mean mass of an atom of an element, relative to 1/12 of the mean mass of an atom of the carbon-12 isotope
-this takes the relative abundances of the different isotopes of an element into account
Ar = means mass of an atom of an element/ 1/12 x mean mass of C-12 isotope

Question 6

Relative isotopic mass:

Answer 6

-> the isotopic mass of an isotope relative to 1/12 of the mean mass of an atom of the carbon-12 isotope

Question 7

Relative molecular mass (Mr):

Answer 7

-> the mean mass of a molecule of a compound, relative to 1/12 of the mean mass of an atom of the carbon-12 isotope
-it can be calculated for a molecule by adding together the separate Ar values of the component elements
-example: Mr of H2O = (2x1.0) + (1x16.0) = 18.0
-calculating the relative molecular mass can be used to identify the molecules present in a sample
-relative formula mass is similar to Mr but is used for compounds with giant structures

Question 8

Ions and Mass Spectrometry:

Answer 8

-ions are formed when an atom loses or gains electrons meaning it is no longer neutral and will have an overall charge
-they are very useful in the analytical technique of mass spectrometry
-it is used to identify different isotopes and find the overall relative atomic mass of an element

Question 9

Time of Flight (TOF): Mass Spectrometry

Answer 9

-this form of mass spectrometry records the time it takes for ions of each isotope to reach a detector, using this spectra can be produces showing each isotope present
1. Ionisation - a sample of an element is vaporised and injected into the mass spectrometer where a high voltage is passed over the chamber
-this causes electrons to be removed from the atoms (they are ionised) leaving +1 charged ions in the chamber
2. Acceleration- these positively charged ions are then accelerated towards a negatively charged detection plate
3. Ion drift - the ions are then deflected by a magnetic field into a curved path
-the radius of the path is dependent on the charge and mass of the ion
4. Detection - when the positive ions hit the negatively charged detection plate, they gain an electron, producing a flow of charge
-the greater the current produced, the greater the abundance of that particular ion
5. Analysis - these current values are then used in combination with the flight times to produce a spectra print-out with the relative abundance of each isotope displayed
-during the ionisation process, a 2+ charged ion may be produces which means it will be more affected more by the magnetic field producing a curved path of smaller radius
-as a result, this mass to charge ratio (m/z) is halved which can be seen on spectra as a trace at half the expected m/z value

Question 10

Ar and Mass Spectrometry:

Answer 10

-using this print-out spectra, the Ar (relative atomic mass) can be calculated my multiplying each m/z value by its abundance and adding each of these together, before dividing by the total abundance of all species present
Ar = (m/z x abundance)/total abundance
-using this calculated value of Ar, the element can be identified by referring to the Periodic Table

Question 11

Predicting Mass Spectra:

Answer 11

-if you know the abundance of an isotope you can generate the mass spectra for its molecules, including relative speak heights
-example: The relative abundance of ^35Cl atoms is 75% and ^37Cl atoms is 25%. In other words, for every 100 atoms of chlorine, 25 would be ^35Cl and 75 would be ^37Cl. Spectra produced by the mass spectrometry of chlorine display a characteristic pattern in a 3:1 ratio for Cl^+ ions and a 3:6:9 ratio for Cl2^+ ions. This is because one isotope is more common than the other and the chlorine molecule can form in different combinations
^70CL2^+ = 35 + 35
^72Cl2^+ = 35 + 37 OR 37 + 35
^74Cl2^+ = 37 + 37

Question 12

Ionisation Energy: first ionisation

Answer 12

-> the minimum energy required to remove one mole of electrons from one mole of atoms in a gaseous state, it is measured in kJmol^-1
Na (g) -> Na^+ (g) + e^-
-successive ionisation energies occur when further electrons are removed
-this usually requires more energy because, as electrons are removed, the electrostatic force of attraction between the positive nucleus and the negative outer electron increases
-more energy is therefore needed to overcome this attraction, causing ionisation energy to increase
-all ionisation energies are endothermic, as the removal of electrons requires an energy input

Question 13

Ionisation energy: second ionisation

Answer 13

-> the minimum energy required to remove one mole of electrons from one mole of +1 ions in a gaseous state, it is measured in kJmol^-1
Na^+ (g) -> Na^2+(g) + e^-
-the definition can be used to define successive ionisation energies
-ionisation energy is influenced by 3 factors:
1. The number of protons
2. The electron shielding
3. The sub shell from which the electron is removed

Question 14

First ionisation energy: trends

Answer 14

-first ionisation energy follows trends within the Periodic Table as they are influenced by proton-electron forces of attraction, and electron shielding
-along a period: first ionisation energy increases due to a decreasing atomic radius and greater electrostatic forces of attraction, electrostatic forced of attraction increase since there is an increasing number of protons
-down a group: first ionisation energy decreases due to an increasing atomic radius and electron shielding which reduces the effect of the electrostatic forces of attraction
-when successive ionisation energies are plotted on a graph, a sudden large increase indicates a change in energy level which is because the electron is being removed from an orbital closer to the nucleus so more energy is required to do so
-this large energy increase provides supporting evidence for the atomic orbital theory
-the first ionisation energy of Aluminium is lower than expected due to a single pair of electrons with opposite spin
-as a result there is a natural repulsion which reduces the amount of energy needed to be put in to remove the outer electron

Question 15

Electron orbitals:

Answer 15

-electrons are held in clouds of negative charge called orbitals
-there are different types of orbital: s,p, d and f, each orbital can hold 0,1 or 2 electrons
-when an orbital holds 2 electrons the electrons must have opposite spins which have different shapes

Question 16

Electron Orbitals: spin

Answer 16

-within an orbital, electrons pair up with opposite spin so that the atom is as stable as possible
-electrons in the same orbital must have opposite spins, spin is represented by opposite arrows
-overall there are three rules for writing out electron configuration:
1. The lowest energy orbital is filled first
2. Electrons with the same spin fill up an orbital first before pairing begins
3. No single orbital holds more than 2 electrons

Question 17

Electron orbitals: Exceptions to the rules

Answer 17

-if electron spins are unpaired and therefore unbalanced, it produces a natural repulsion between the electrons making the atom very unstable
-if this is the case, the electrons may take on a different arrangement to improve stability
-example: the 3p^4 orbital contains a single pair of electrons with opposite spins, making it unstable: therefore the electron configuration changes to become 3p^34s^1 which is a much more stable arrangement

Question 18

Electron Configurations:

Answer 18

-scientific ideas on electronic configuration have developed over time as new discoveries are made, the current, accepted model is based on the following evidence:
1. Emission spectra provide evidence for the existence of quantum shells
2. Successive ionisation energies provide evidence for quantum shells within atoms and suggest the group to which the element belongs
3. First ionisation energy of successive elements provides evidence for electron sub shells
-these orbitals correspond with blocks on the Periodic Table, each element in the block has outer electrons in that orbital

Question 19

Electron configuration: different sub-shells

Answer 19

-each sub shell has a different number of orbitals and therefore can hold a different number of electrons before the next one is filled: s-sub shell = 2 electrons, p-sub shell = 6 electrons, d-sub shell = 10 electrons
-the energy of the orbitals increases from s to d meaning the orbitals are filled in this order, each orbital is filled before the next one is used to hold electrons
-example: Sodium has 11 electrons, these would be written in the following configuration -> Na = 1s^2, 2s^2, 2p^6, 3s^1
-it has 3 energy levels and 4 orbitals holding the 11 electrons
-there are two main exceptions to electron configuration, a completely full or half-dull d sub level is more stable than a partially filled d sub level, so an electron from the 4s is excited to the 3d orbital
Chromium: 1s^2, 2s^2, 2p^6, 3s^2, 3p^6, 3d^5, 4s^1
Copper: 1s^2, 2s^2, 2p^6, 3s^2, 3p^6, 3d^10, 4s^1

Question 20

Periodicity:

Answer 20

-> refers to the study of patterns of physical, atomic and chemical properties within the Periodic Table that repeat regularly
-the Periodic Table arranges the known elements according to proton number,
-all the elements along a period have the same number of electron shells
-all the elements down a group have the same number of outer electrons, this number is indicated by the group number
-elements are classified into blocks within the Periodic Table that show electron configuration:
S-block = groups 1 and 2
P-block = groups 3 and 0
D- block = transitions metals
F- block = radioactive elements
-these different electron configurations are often linked to other trends within the Periodic Table, periodicity is the study of these trend

Question 21

Periodicity Data:

Answer 21

-the trends of properties like melting point, atomic radius and ionisation energy can be represented on a graph, these trends are known as periodic properties
-ionisation energies may be shown on a logarithmic graph, which has a scale where a one unit increase represents an increase by a factor of 10
-for instance, the numbers 10 and 100 would be the same distance apart as the numbers 100 and 1000, this enable you to plot data that has a large range

Question 22

Atomic radius:

Answer 22

-along a period, atomic radius decreases
-this is due to an increased nuclear charge for the same number of electron shells
-the outer electrons are pulled in closer to the nucleus as the increased charge produces a greater attraction, as a result, the atomic radius for that element is reduced
-down a group, atomic radius increases, with each increment down a group, an electron shell is added each time
-this increases the distance between the outer electrons and the nucleus, reducing the power of attraction
-more shells also increases electron shielding where the inner shells create a ‘barrier’ that blocks the attractive forces
-therefore, the nuclear attraction is reduced further and atomic radius increases

Question 23

Ionisation energy: as a trend

Answer 23

-along a period, ionisation energy increases, the decreasing atomic radius and increasing nuclear charge means that the outer electrons are held more strongly and therefore more energy is required to remove the outer electron and ionise the atom
-down a group, ionisation energy decreases, the nuclear attraction between the nucleus and outer electrons reduces and increasing amount of shielding means less energy is required to remove the outer electron