OCR CHEMISTRY ANotes for Topic 11.1.1 Atoms1.2.1 Electron Structure1.3.1 Periodicity1.1.1 ATOMSa) Protons, neutrons and electronsAtoms are made up of three fundamental particles: protons, neutrons and electrons.Protons and neutrons are found in the nucleus and are collectively called nucleons. Electrons orbit the nucleus in a similar way to that in which planets orbit a sun. In between the electrons and nucleus there is nothing (empty space).The nucleus is very small; if an atom were the size of a football pitch, the nucleus would be the size of a drawing pin.The basic properties of these three particles can be summarized in the following table:ParticleChargeMassProton+1 unitApprox 1 unitNeutronNo chargeApprox 1 unitElectron-1 unitApprox 1/1840 units (very small)1 unit of charge is 1.602 x 10-19 coulombs. A proton is given a charge of +1 and an electron a charge of -1. All charges are measured in these units.1 unit of mass is 1.661 x 10-27 kg. This is also not a convenient number, so we use “atomic mass units〞. Since the mass of protons and neutrons varies slightly depending on the nucleus, then in order to define an “atomic mass unit〞 we need to choose one nucleus as a standard. For this purpose 126C , or “carbon-12〞, was chosen because its mass per nucleon (1.661 x 10 –27 kg) is around average, which means all the other nuclei have masses close to whole numbers. An atomic mass unit is thus defined as 1/12th of the mass of one atom of carbon-12. Everything else is measured relative to this quantity.b) Atomic numbers, mass numbers and isotopesAn atom is named after the number of protons in its nucleus. If the nucleus of an atom has 1 proton, it is hydrogen; if it has two protons, it is helium; if it has 3, it is lithium etc. The number of protons in the nucleus of an atom is called the atomic number. It has the symbol Z.The atomic number is the number of protons in the nucleus of an atomNot all atoms of the same element have equal numbers of neutrons; this may vary slightly. The sum of the number of protons and neutrons in the nucleus of an atom is called its mass number. It is represented by the symbol A.The mass number is the sum of the number of protons and neutrons in the nucleus of an atomThe nucleus of an atom can thus be completely described by its mass number and its atomic number. It is generally represented as follows:AZEEg. 94Be, 126C, 2412MgAtoms with the same atomic number but with different mass numbers (ie different numbers of neutrons) are called isotopes.Isotopes are atoms with the same atomic number but with different mass numbersEg magnesium (atomic number 12) has 3 naturally occurring isotopes: 2412Mg: 12 protons, 12 neutrons2512Mg: 12 protons, 13 neutrons2612Mg: 12 protons, 14 neutronsIn a neutral atom, the number of protons and electrons are the same. However, many elements do not exist as neutral atoms, but exist as ions. Ions are species in which the proton and electron numbers are not the same, and hence have an overall positive or negative charge. The number of electrons in a species can be deduced from its charge:Eg2412Mg2+: 12p, 12n, 10e2412Mg+: 12p, 12n, 11e2412Mg 12p, 12n, 12e2412Mg-: 12p, 12n, 13eb) Relative atomic massThe mass of an atom is measured in atomic mass units, where one unit is 12th of the mass of one atom of carbon-12.The relative isotopic mass of an isotope is the ratio of the mass of one atom of that isotope to 1/12th of the mass of one atom of carbon-12. It is usually very close to a whole number ratio:IsotopeMass numberRelative isotopic mass11H11.00642He44.00394Be99.0122713Al2726.9195927Co5958.933The masses of protons and neutrons vary slightly from isotope to isotope, so the relative isotopic mass is not exactly a whole number.The relative atomic mass of an atom is the ratio of the average mass of one atom of that element to 1/12th of the mass of one atom of carbon-12.The RAM is the average mass of all the isotopes, and is often not close to a whole number:ElementCommon mass numbersRelative atomic massMg24, 25, 2624.32Cl 35, 3735.45Br79, 8179.91Ba134, 135, 136, 137, 138137.33Species consisting of more than one atom also have a characteristic mass:The relative molecular mass of a molecule is the ratio of the average mass of that molecule to 1/12th of the mass of an atom of carbon-12. The relative formula mass of a compound is the ratio of the average mass of one formula unit of that compond to 1/12th of the mass of an atom of carbon-12. The relative molecular/formula mass of a species is the sum of the relative atomic masses of its constituent atoms.Eg The relative molecular mass of CO2 is 12.0 + 16.0 + 16.0 = 44.0The relative atomic mass can be calculated by the formula:Σ (perentage abundance of each isotope x mass of each isotope) 100eg Neon consists of 90% neon-20 and 10% neon-22. What is its relative atomic mass?ram = (90 x 20 + 10 x 22)/100 = 20.21.2.1 ELECTRON STRUCTUREa) Energy levelsElectrons do not orbit the nucleus randomly; they occupy certain fixed energy levels. Each atom has its own unique set of energy levels, which are difficult to calculate but which depend on the number of protons and electrons in the atom.Energy levels in an atom can be numbered 1,2,3,…. To infinity. 1 is the lowest energy level (closest to the nucleus) and energy level infinity corresponds to the energy of an electron which is not attracted to the nucleus at all. The energy levels thus converge as they approach infinity:b) Orbitals and sub-levelsElectrons do not in fact orbit the nucleus in an orderly way. In fact they occupy areas of space known as orbitals. The exact position of an electron within an orbital is impossible to imagine; an orbital is simply an area of space in which there is a high probability of finding an electron.Orbitals can have a number of different shapes, the most common of which are as follows:s-orbitals: these are spherical.Every energy level contains one s-orbital.An s-orbital in the first energy level is a 1s orbital.An s-orbital in the second energy level is a 2s orbital, etcp-orbitals: these are shaped like a 3D figure of eight. They exist in groups of three: Every energy level except the first level contains three p-orbitals. Each p-orbital in the same energy level has the same energy but different orientations: x, y and z.A p-orbital in the second energy level is a 2p orbital (2px, 2py, 2pz)A p-orbital in the third energy level is a 3p orbital (3px, 3py, 3pz), etcIn addition, the third and subsequent energy levels each contain five d-orbitals, the fourth and subsequent energy levels contain seven f-orbitals and so on. Each type of orbital has its own characteristic shape.S, p and d orbitals do not all have the same energy. In any given energy level, s-orbitals have the lowest energy and the energy of the other orbitals increases in the order p < d < f etc. Thus each energy level must be divided into a number of different sub-levels, each of which has a slightly different energy.The number and type of orbitals in each energy level can thus be summarised as follows: Energy levelNumber and type of orbital1st sub-level2nd sub-level3rd sub-level4th sub-level5th sub-level11 x 1s21 x 2s3 x 2p31 x 3s3 x 3p5 x 3d41 x 4s3 x 4p5 x 4d7 x 4f51 x 5s3 x 5p5 x 5d7 x 5f9 x 5gSince the different sub-levels have different energies, and the energies of the different levels get closer together with increasing energy level number, the high energy sub-levels of some energy levels soon overlap with the low energy sub-levels of higher energy levels, resulting in a more complex energy level diagram:Starting with the lowest energy, the orbitals can thus be arranged as follows:1s 2s 2p 3s 3p 4s 3d 4p 5s 4d 5p 6s 4f 5d 6p 7s 5f 6d c) ElectronsElectrons repel each other. In a small space such as an orbital, it is impossible to put more than two electrons.Since electrons are charged particles, and moving charges create a magnetic field, it is possible to create a small magnetic attraction between two electrons if they are spinning in opposite directions in the same orbital. This is the reason two electrons, and not one, are permitted in the same orbital.An orbital is therefore a region of space which can hold up to two electrons, with opposite spins.It is thus possible to calculate the maximum possible number of electrons in each sub-level, and thus in each energy level:Energy LevelNumber of electrons in each sub-levelMax. no of electrons12 x 1s222 x 2s, 6 x 2p 832 x 3s, 6 x 3p, 10 x 3d1842 x 4s, 10 x 3d, 6 x 4p, 7 x 4f32Note, however, that the 4s electrons in energy level 4 are filled before the 3d electrons in energy level 3.d) Electron arrangement in orbitalsThere are three rules which determine the way in which electrons fill the orbitals1. Aufbau/building principle: electrons always fill the lowest energy orbitals first.2. Hund's rule: electrons never pair up in the same orbital until all orbitals of the same energy are singly occupied, and all unpaired electrons have parallel spin.3. Pauli exclusion principle: only two electrons may occupy the same orbital, and they must do so with opposite spin.The arrangement of electrons in an atom is known as its electronic configuration. It can be represented in two ways:The arrow and box method represents each orbital as a box and each electron as an arrow. The direction of spin is shown by the orientation of the arrow.The electronic configuration of the first 18 elements using the arrow in box method is as follows: 1s 2s 2p 3s 3pH↑He↑↓Li↑↓↑Be↑↓↑↓B↑↓↑↓↑C↑↓↑↓↑↑N↑↓↑↓↑↑↑O↑↓↑↓↑↓↑↑F↑↓↑↓↑↓↑↓↑Ne ↑↓↑↓↑↓↑↓↑↓Na↑↓↑↓↑↓↑↓↑↓↑Mg↑↓↑↓↑↓↑↓↑↓↑↓Al↑↓↑↓↑↓↑↓↑↓↑↓↑Si↑↓↑↓↑↓↑↓↑↓↑↓↑↑P↑↓↑↓↑↓↑↓↑↓↑↓↑↑↑S↑↓↑↓↑↓↑↓↑↓↑↓↑↓↑↑Cl↑↓↑↓↑↓↑↓↑↓↑↓↑↓↑↓↑Ar↑↓↑↓↑↓↑↓↑↓↑↓↑↓↑↓↑↓The orbital method indicates the number of electrons in each orbital with a superscript written immediately after the orbital.The electronic configurations of the first eighteen elements can be shown with the orbital method as follows:H: 1s1He: 1s2Li: 1s22s1Be: 1s22s2B: 1s22s22p1C: 1s22s22p2 or 1s22s22p63s23px13py1N: 1s22s22p3 or 1s22s22p63s23px13py13pz1O: 1s22s22p4 or 1s22s22p63s23p23px23py13pz1F: 1s22s22p5Ne: 1s22s22p6Na: 1s22s22p63s1Mg: 1s22s22p63s2Al: 1s22s22p63s23p1Si: 1s22s22p63s23p2 or 1s22s22p63s23px13py1P: 1s22s22p63s23p3 or 1s22s22p63s23px13py13pz1S: 1s22s22p63s23p4 or 1s22s22p63s23px23py13pz1Cl: 1s22s22p63s23p5Ar: 1s22s22p63s23p6 A shorthand form is often used for both the above methods. Full shells are not written in full but represented by the symbol of the element to which they correspond, written in square brackets.Eg. 1s22s22p6 is represented as [Ne] and 1s22s22p63s23p6 is represented as [Ar].The shorthand electronic configuration of the elements with atomic numbers 18 to 36 can be written as follows: 4s 3d 4pK[Ar]↑Ca[Ar]↑↓Sc[Ar]↑↓↑Ti[Ar]↑↓↑↑V[Ar]↑↓↑↑↑Cr[Ar]↑↑↑↑↑↑Mn[Ar]↑↓↑↑↑↑↑Fe[Ar]↑↓↑↓↑↑↑↑Co[Ar]↑↓↑↓↑↓↑↑↑Ni[Ar]↑↓↑↓↑↓↑↓↑↑Cu[Ar]↑↑↓↑↓↑↓↑↓↑↓Zn[Ar]↑↓↑↓↑↓↑↓↑↓↑↓Ga[Ar]↑↓↑↓↑↓↑↓↑↓↑↓↑Ge[Ar]↑↓↑↓↑↓↑↓↑↓↑↓↑↑As[Ar]↑↓↑↓↑↓↑↓↑↓↑↓↑↑↑Se[Ar]↑↓↑↓↑↓↑↓↑↓↑↓↑↓↑↑Br[Ar]↑↓↑↓↑↓↑↓↑↓↑↓↑↓↑↓↑Kr[Ar]↑↓↑↓↑↓↑↓↑↓↑↓↑↓↑↓↑↓Note the unusual structures of chromium and copper. e) Electron arrangement in ionsThe electronic configuration of ions can be deduced by simply adding or removing the appropriate number of electrons. Electrons are removed in reverse order to which they are added.f) Effect of electronic configuration on chemical propertiesThe chemical properties of an atom depend on the strength of the attraction between the outer electrons and the nucleus. These in turn depend on the number of protons and on the electronic configuration, and so it follows that these two factors are instrumental in determining the chemical properties of an atom.This is in contrast with the neutron number however, which has no effect on the chemical properties of an atom. Neutrons have no charge and hence exert no attractive force on the nucleus.Isotopes, therefore, tend to have very similar chemical properties since they have the same atomic number and the same electronic configuration. They differ only in number of neutrons, which do not directly influence the chemical properties of an element.g) Ionisation EnergiesThe first ionisation energy of an element is the energy required to remove one electron from each of a mole of free gaseous atoms of that element.It can also be described as the energy change per mole for the process:M(g) à M+(g) + eThe amount of energy required to remove an electron from an atom depends on the number of protons in the nucleus of the atom and on the electronic configuration of that atom.The second ionisation energy of an atom is the energy required to remove one electron from each of a mole of free gaseous unipositive ions.M+(g) à M2+(g) + eOther ionisation energies can be defined in the same way:The third ionisation energy of an atom is the energy required to remove one electron from each of a mole of bipositive ions.M2+(g) à M3+(g) + eThe nth ionisation energy can be defined as the energy required for the processM(n-1)+(g) à Mn+(g) + eA number of factors must be considered when explaining ionisation energies:- Energy is required to remove electrons from atoms in order to overcome their attraction to the nucleus. The greater the number of protons, the greater the attraction of the electrons to the nucleus and the harder it is to remove the electrons. The number of protons in the nucleus is known as the nuclear charge.- The effect of this nuclear charge, however, is cancelled out to some extent by the other electrons in the atom. Each inner shell and inner sub-shell electron effectively cancels out one unit of charge from the nucleus. This is known as shielding.- The outermost electrons in the atom thus only feel the residual positive charge after all inner shell and inner sub-shell electrons have cancelled out much of the nuclear charge. This residual positive charge is known as the nuclear attraction.- Electrons repel each other, particularly when they are in the same orbital. The degree of repulsion between the outermost electrons affects the ease with which electrons can be moved.When considering trends in ionisation energies, it is thus necessary to consider 4 factors:- nuclear charge- shielding- nuclear attraction- electron repulsionIt always becomes progressively more difficult to remove successive electrons from an atom; the second ionisation energy is always greater than the first, the third always greater than the second and so on. There are two satisfactory explanations for this:As more electrons are removed from an atom, the number of electrons remaining in the atom decreases. The repulsion between these electrons therefore decreases, while the number of protons remains the same. The remaining electrons are thus more stable and increasingly difficult to remove.The difference in successive ionisation energies, however, varies widely and depends on the electronic configuration of the atom in question. The difference in successive ionisation energies of an atom can be predicted qualitatively by consideration of the effective nuclear charge on the electron to be removed and the shielding of that electron by the inner shell and inner sub-shell electrons.Consider the successive ionisation energies of aluminium, 1s22s22p63s23p1: The 1st, 2nd and 3rd ionisation energies are fairly low because the 3p and 3s electrons are shielded by all the other electrons in inner shells, so the nuclear attraction is low.1st: 578 kJmol-1, 2nd: 1817 kJmol-1, 3rd: 2745 kJmol-1There is a huge jump to the 4th ionisation energy, since a 2p electron is now being removed (from an inner energy level), resulting in a large drop in shielding. 4th: 11578 kJmol-1, 5th: 14831 kJmol-1, 6th: 18378 kJmol-1, 7th: 23296 kJmol-1, 8th: 27460 kJmol-1, 9th: 31862 kJmol-1; 10th: 38458kJmol-1, 11th: 42655 kJmol-1There is a huge jump to the12th ionisation energy, since a 1s electron is now being removed (from an inner energy level), resulting in a large drop in shielding.12th: 202176kJmol-1, 13th: 222313kJmol-1.These ionisation energies could be plotted on a graph as follows:Note that the largest jumps by far occur between the 3rd and 4th ionisation energies, and between the 11th and 12th ionisation energies. In practice only large jumps such as this are visible on such a graph.The relative values of successive ionisation energies are therefore a direct indicator of the electronic configuration of the atom in question.The trends can be summarised as follows:1. The successive ionisation energies of an atom always increase. The more electrons that are removed, the fewer the number electrons that remain. There is therefore less repulsion between the electrons in the resulting ion. The electrons are therefore more stable and harder to remove.2. By far the largest jumps between successive ionisation energies come when the electron is removed from an inner shell. This causes a large drop in shielding, a large increase in effective nuclear charge and a large increase in ionisation energyBy applying the above principles in reverse, it is also possible to predict the electronic structure of a species by analysis of the successive ionisation energy data:Eg Si:Large jumps occur between 4th and 5th and between 12th and 13th.Therefore there are three shells: The first contains 2 electrons, the second 8 and the third 4. 1.3.1 PERIODICITYa) Structure of the Periodic TableThe periodic table is a list of all known elements arranged in order of increasing atomic number, from 1 to 106. In addition to this, the elements are arranged in such a way that atoms with the same number of shells are placed together, and atoms with similar electronic configurations in the outer shell are also placed together. This is achieved as follows:The elements are arranged in rows and columns. Elements with one shell are placed in the first row (ie H and He), Elements with two shells are placed in the second row (Li to Ne) and so on.A row of elements thus arranged is called a period. In addition, the elements are aligned vertically (in columns) with other elements in different rows, if they share a similar outer-shell electronic configuratio。