Return to AS and A2 GCE chemistry
1.1a A model of the atom
An atom contains protons, neutrons and electrons. The protons and neutrons are found in the nucleus of the atom while the electrons exist in the shells, energy levels or quantum shells around the nucleus. The existence of the nucleus was proved by Geiger and Marsden in their experiment when alpha particles were fired at gold foil. Deflection of the alpha particles could only be explained by the existence of a nucleus containing all of the atom's mass.
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1.1b Relative masses
The unit of mass is 1/12 of the mass of an atom of the isotope Carbon-12 (12C=12 exactly)
The Relative Atomic Mass (Ar)of an element is the ratio of the mass of an average atom of that element to 1/12 of the mass of an atom of the nuclide Carbon-12.
The Relative Isotopic Mass of a nuclide is the ratio of the mass of one atom of that nuclide to 1/12 of the mass of a Carbon-12 atom.
The Relative Molecular Mass (Mr)of a substance is the ratio of the mass of an average molecule of that substance to 1/12 of the mass of an atom of the nuclide Carbon-12.
1.1c Atomic number Z, mass number and relative isotopic mass
The Atomic Number of an element is the number of protons in the nucleus of one of the atoms of that element.
The Mass Number of a nuclide is the sum of the Protons and Neutrons in the nucleus of an atom of that nuclide.
The Relative Isotopic Mass and Mass Number of a nuclide can be taken to have the same value.
1.1d Isotopes and relative atomic mass
Isotopes are atoms of the same element with different numbers of neutrons in the nucleus
The Relative Atomic Mass of an element is the weighted (to take account of relative abundance) average of the the Relative Isotopic Masses of all of the isotopes of that element.
E.g. Chlorine has two isotopes with mass numbers (and relative isotopic masses) 35 and 37 35 37
75% is Cl and 25% is Cl
Let there be 100 atoms
Total mass of 100 atoms = (75 * 35) + (25 * 37) = 3550
Average mass of an atom (relative atomic mass of chlorine) = Total mass /Number of atoms
so relative atomic mass of chlorine = 35.5
1.1e Use of the low-resolution mass spectrometer
In a mass spectrometer the stages are 1. vaporisation by heating to give gas atoms or molecules, 2. ionisation by electron bombardment to form positive ions, 3.acceleration of positive ions, 4.deflection by a variable magnetic field, 5. detection. The mass spectrum of an element shows the relative isotopic mass and relative abundance of each isotope of the element being tested. In a mass spectrum the height of each peak = the relative abundance. The Relative Atomic Mass of an element can be found by finding the sum of the products of the relative abundance of each isotope and its relative isotopic mass and the dividing by the total relative abundance.
For 100 neon atoms the total mass is (90.5*20)+(0.3*21)+(9.2*22) = 2018.7
relative atomic mass of neon 2018.7/100 = 20.2
Very accurate masses can be read from the spectrum if needed e.g. 20.994 for neon-21.
The peak in the spectrum on the far
right has the highest value of m/e and is called the molecular ion. This
peak gives the Relative Molecular Mass of a compound. Below ethanol can be seen
to have a relative molecular mass of 46.
Some organic molecules (e.g. Ethanol ) are fragmented in a mass spectrometer. These fragments (which are parts or fragments of the molecular ion) can be seen in the peaks. The pattern of fragments can be used to work out the structure of original molecule. Above ethanol can be seen to contain a CH3 group, a C2H5 group and an OH group.
1.1f Defining ionisation energies
The first ionisation energy is the enthalpy change when 1 mole of gaseous atoms of an element each lose an electron to form gaseous ions each with a single positive charge. It is the enthalpy change for the reaction:
M(g) ---> M+(g) + e-
The second ionisation energy is the enthalpy change when 1 mole of gaseous ions of an element each with a single positive charge each lose an electron to form gaseous ions each with a double positive charge. It is the enthalpy change for the reaction:
M+(g)---> M2+(g) + e-
Task1.1f.1 Define the third ionisation enthalpy for the element M and give an equation.
Task1.1f.2 Define the first ionisation enthalpy for sodium, magnesium, chlorine and neon.
Task1.1f.3 Define the second ionisation enthalpy for lithium, aluminium and oxygen.
1.1g ionisation energies and quantum shells
The graph of successive ionisation energies against ionisation number shows electrons grouped into three energy levels or quantum shells with similar energies. Electron 1 (ionisation number = 1, removed first) from third shell, electrons 2-9 in second shell, electrons 10 and 11(removed last) from first shell closest to the nucleus.
Sketch similar graphs for carbon, silicon and potassium (atomic numbers 6, 14 and 19)
1.1h Evidence for s, p and d orbitals
The graph of first ionisation energy against atomic number shows the grouping of electrons into s, p (and d) sub-shells (orbitals) within the energy levels. The general trend is for ionisation energy to increase with increasing atomic number across a period but B < Be and O < N.
< Be because Be completes the 2s shell. B has 1 electron in the 2p shell. 2s
shields better than 2p so therefore the 2p electron is easier to remove.
< N because of the stability of the half-filled p shell for N. For O an
electron is being added to an occupied orbital
Task1.1h.1 Label the diagram with symbols for the elements.
Task1.1h.2 Label electrons removed as s or p.
Task1.1h.3 Write two sentences using the frame: The ionisation energy (increases/decreases) (from left to right across a period/down a group) because (distance from outer electrons to nucleus is increasing/the nuclear charge in increasing)
1.1i s,p and d-block elements
All s-block elements have their outer electrons in s-orbitals.
All p-block elements have their outer electrons in p-orbitals.
All d-block elements have electrons in their d-orbitals which are in the process of being filled.
1.1j Electronic configurations from hydrogen to krypton
Electronic Configuration can be predicted. If the following rules are followed:
Electrons go into lower energy levels before higher ones.
Electrons go into lower subshells before higher ones.
Electrons occupy orbitals with 1 electron rather than 2 if possible.
Electrons can only occupy the same orbital if they have opposite spins.
When filling d-orbitals electrons create 3d5 and 3d10 by losing a 4s electron as half filled or filled d subshells are more stable than other arrangements.
An orbital is represented by a box
The orbital order 1s, 2s, 2p, 3s, 3p, 4s, 3d, 4p
Max electrons in shell s=2, p=6( 3 orbitals/squares), d=10( 5 orbitals/squares)
e.g Kr (the biggest you need)
1s2 2s2 2p6 3s2 3p6 3d10 4s2 4p6
1s2 2s2 2p6 3s2 3p6 3d1 4s2
1s2 2s2 2p6 3s2 3p6 3d5 4s1
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Task 1.1j.1 Draw s,p,d electronic
configurations for Be, F, Mg, S and K.
Task 1.1j.2 Draw electrons in boxes diagrams for N, Ne, Na, Ar and Ca.
Task 1.1j.3 Draw s,p,d electronic configurations for Sc, V, Mn, Ni, Zn and Cu.
Task 1.1j.4 Draw electrons in boxes diagrams for Ti, Fe, Ni, Se, Kr and Cr
1.1k Electronic structure and chemical properties
s block elements which easily lose 1 or two electrons to form stable ions. s block elements are strongly metallic.
p block elements that are strongly non metallic. They may gain 1,2 or 3 electrons to form -ve ions with a stable inert gas configuration. Electron sharing is also common. Group 0 have complete outer shell, therefore they are chemically inert.
d-block elements lose electrons and therefore have metallic properties.
1.1l First and second electron affinities
The 1st Electron Affinity of an element is the enthalpy change when 1 mole of gaseous atoms of that element each gain an electron to form gaseous ions each with a single negative charge. Often exothermic as the negative electron is attracted to positive nucleus. It is the enthalpy change for the following reaction:
X(g) + e- --> X-(g)
The 2nd Electron Affinity of an element is the enthalpy change when 1 mole of gaseous ions each with a single negative charge each gain an electron to form gaseous ions each with a charge of -2. Often endothermic as negative electron repelled by the negative ion. It is the enthalpy change for the following reaction:
X-(g) + e- --> X2-(g)