GCE Chemistry 1.2 Atomic Structure PDF

FACTFILE: GCE CHEMISTRY 1.2 ATOMIC STRUCTURE Atomic Structure Learning Outcomes 1.2.9 demonstrate understanding that an orbital is a region within an atom that can hold up to two Students s...

FACTFILE: GCE CHEMISTRY 1.2 ATOMIC STRUCTURE Atomic Structure Learning Outcomes 1.2.9 demonstrate understanding that an orbital is a region within an atom that can hold up to two Students should be able to: electrons with opposite spins and describe the 1.2.1 describe the properties of electrons, protons shape of s and p-orbitals; and neutrons in terms of their location in the atom and their relative masses and charges; 1.2.10 classify an element as belonging to the s, p, d or f block according to its position in the 1.2.2 explain the terms atomic number and mass Periodic Table; number and use them to deduce the numbers of protons, neutrons and electrons in an atom 1.2.11 define and write equations for the first and or ion; successive ionisation energies of an element in terms of one mole of gaseous atoms and 1.2.3 define the terms relative atomic mass ions; and relative isotopic mass in terms of the carbon-12 standard; 1.2.12 demonstrate understanding that successive ionisation energies can be used to predict 1.2.4 define and demonstrate understanding of the the group of an element, and that graphs term isotopes; of successive ionisation energies against number of electrons removed, for an element, 1.2.5 define the terms relative molecular mass (for give evidence for the existence of shells; molecules) and relative formula mass (for ionic compounds) in terms of the carbon-12 1.2.13 explain the trend in the first ionisation standard and calculate their values from energies of atoms down Groups, and across relative atomic masses; Periods in terms of nuclear charge, distance of outermost electron from the nucleus, 1.2.6 interpret mass spectra of elements by shielding and stability of filled and half- calculating relative atomic mass from isotopic filled sub-shells; abundances and vice versa; 1.2.14 demonstrate understanding that graphs of 1.2.7 predict the mass spectra of diatomic elements, first ionisation energies of elements up to for example chlorine; krypton provide evidence for the existence of shells and sub-shells 1.2.8 deduce the electronic configuration of atoms and ions up to krypton in terms of shells and sub-shells using the building up principle (s, p and d notation and electrons in boxes notation); 1 FACTFILE: GCE CHEMISTRY / 1.2 ATOMIC STRUCTURE Atomic Structure Isotopes All atoms consist of a number of fundamental, Isotopes are atoms which have the same atomic sub-atomic particles. There are three, the electron, number but a different mass number (contain the proton and the neutron. Different atoms (and the same number of protons but a different therefore elements) have different numbers of these number of neutrons). three fundamental particles. For example, 12C contains 6 neutrons, whereas Particle Relative Relative Position 13 C contains 7 neutrons. Both isotopes contain Charge Mass in Atom 6 protons and 6 electrons and have a different protons +1 1 Nucleus number of neutrons, and a different mass number. neutrons 0 1 Nucleus A mass spectrometer, is used to obtain accurate electrons -1 1/1840 Shells atomic masses by measuring the masses and relative abundances of the isotopes of an atom. In a mass The atomic number and mass number give us spectrometer, particles are turned into positive ions, important information about an atom and are accelerated and then deflected by an electromagnet. particularly useful in distinguishing one isotope of The resulting path of ions depends on their ‘mass to an element from another. charge’ ratio (m/z). Whilst initially used to determine accurate relative atomic masses, mass spectrometry Atomic number is now widely used to determine the relative formula masses of unknown compounds (e.g. in forensic The atomic number is the number of protons science) in order to help identify them. (in the nucleus) of an atom. For an atom (which is neutral), this number also corresponds to the For example, consider the mass spectrum of zirconium. number of electrons. 100 Mass number relative abundance (%) The mass number is the total number of protons and neutrons (in the nucleus) of an atom. For example: 50 Mass number PROTONS + NEUTRONS 23 Atomic number 11 Na 88 90 92 94 96 98 PROTONS m/z The height of each peak is proportional to the There are 11 protons, 11 electrons and 12 neutrons amount of each isotope present (i.e. it’s relative in a sodium atom. abundance). Most ions have a 1+ charge so that the m/z ratio is numerically equal to mass m of the ion. The five peaks in the mass spectrum shows that there are 5 isotopes of zirconium - with relative isotopic masses of 90, 91, 92, 94 and 96 on the C-12 scale. The relative sizes of the peaks gives you a direct measure of the relative abundances of the isotopes. In this case, the 5 isotopes (with their relative percentage abundances) are: zirconium-90 51.5 zirconium-91 11.2 zirconium-92 17.1 zirconium-94 17.4 zirconium-96 2.8 2  FACTFILE: GCE CHEMISTRY / 1.2 ATOMIC STRUCTURE The relative atomic mass of zirconium is Electronic configuration calculated by: Electrons are arranged in energy levels. The energy level n=1 is closet to the nucleus. Energy level are (90 x 51.5)+(91 x 11.2)+(92 x 17.1)+(94 x 17.4)+(96 x 2.8) made of subshells which are made of orbitals.An RAM = orbital is a region within an atom that can hold 100 up to two electrons with opposite spins. Examples = 4635 + 1019.2 + 1573.2 + 1635.6 + 268.8 include s orbitals which are spherical shaped and p orbitals which are dumbbell shaped: 100 An s subshell holds up to 2 electrons, a p up to 6 , a = 9131.8 d up to 10 and an f up to 14. 100 = 91.3 The relative atomic mass (RAM) is the average s orbital p orbital (weighted mean) mass of an atom of an element relative to one-twelfth of the mass of an atom of carbon-12. Quantum shell Sub-shells Number of Number of (Energy Level) electrons orbitals The relative isotopic mass (RIM) is the mass of n=1 One sub-shell 2 1 an atom of an isotope of an element relative to 1s one-twelfth of the mass of an atom of carbon-12. n=2 Two sub-shells RFM is for ionic compounds and for giant covalent 2s 2 1 compounds. 8 2p 6 3 n =3 Three sub-shells You should be able to predict the mass spectra of 3s 2 1 diatomic elements, for example chlorine. Chlorine 18 3p 6 3 has 5 peaks in the mass spectrum of Cl2 one at 35, 3d 10 5 37, 70, 72 and 74. n=4 Four sub-shells The relative formula mas (RFM) is the average 4s 2 1 (weighted mean) mass of a formula unit relative to 4p 6 32 3 one-twelfth of the mass of an atom of carbon-12. 4d 10 5 4f 14 7 The relative molecular mass (RMM) is the average (weighted mean) mass of a molecule relative to one-twelfth of the mass of an atom of carbon-12. RMM is used for molecular covalent substances. 3 FACTFILE: GCE CHEMISTRY / 1.2 ATOMIC STRUCTURE When writing electron configurations, electrons are placed in the first energy level (the energy level closest to the nucleus) and subsequent increasing energy levels. Electrons are not paired until a subshell is half filled. When two electrons are placed in the same orbital, they have opposite spin and are shown as vertical arrows in opposite directions. Sodium (11 electrons) Na 1s2 2s2 2p6 3s1 Phosphorus (15 electrons) P 1s22s22p63s23p3 s-block p-block d-block f-block The Periodic Table of elements can be organised into different blocks based on the orbitals the outer electrons reside in: An s-block element is one which has an atom with highest energy/outer electron in an s-subshell (orbital). A p-block element is one which has an atom with highest energy/outer electron in a p-subshell. When writing electronic configuration of transition metal ions, remember that the electrons are lost from the 4s subshell first. For example: Fe 1s2 2s2 2p6 3s2 3p6 3d6 4s2 Fe2+ 1s2 2s2 2p6 3s2 3p6 3d6 Copper and chromium have unusual electronic configuration: Cr is not 1s2 2s2 2p6 3s2 3p6 3d4 4s2 but it is 1s2 2s2 2p6 3s2 3p6 3d5 4s1 due to the stability of half filled 3d. Cu is not 1s2 2s2 2p6 3s2 3p6 3d9 4s2 but is 1s2 2s2 2p6 3s2 3p6 3d10 4s1 due to the stability of the full 3d. 4 FACTFILE: GCE CHEMISTRY / 1.2 ATOMIC STRUCTURE Ionisation energy – the evidence for shells and sub-shells Ionisation energy is a measure of the amount of energy needed to remove electrons from atoms. As electrons are negatively charged and protons in the nucleus are positively charged, there will be an attraction between them. The greater the pull of the nucleus, the harder it will be to pull an electron away from an atom. First ionisation energy is the energy required to convert one mole of gaseous atoms into gaseous ions with a single positive charge. For example: Na(g) → Na+(g) + e- Note that state symbols should be given. Second ionisation energy is the energy required to convert one mole of gaseous ions with a single positive charge into gaseous ions with a double positive charge. Na+(g) → Na2+(g) + e- Third ionisation energy is the energy required to convert one mole of gaseous ions with a double positive charge into gaseous ions with a triple positive charge. Na2+(g) → Na3+(g) + e- The trend in first ionisation energies across Periods of the Periodic Table provides evidence for shells and sub-shells, for example Period 3: 1st ionisation energy (kJ mol–1) Na Mg Al Si P S Cl Ar The dip at aluminium, for example, can be explained by considering its electronic configuration; aluminium’s outer electron is in a 3p sub-shell which is further from the nucleus than the outer electron in magnesium which is in the 3s sub-shell. The dip at sulfur is explained by the outer electron configuration, 3p4, which compares with the outer configuration of 3p3 in phosphorus which is half filled and stable. The presence of one fully filled 3p orbital (compared to three half-filled 3p orbitals in phosphorus) increases repulsion between the paired electrons in the 3p orbital in sulfur which makes it easier for an electron to be removed. 5 FACTFILE: GCE CHEMISTRY / 1.2 ATOMIC STRUCTURE Ionisation energy increases across a period due to Increasing nuclear charge Shielding is constant as the electron is being removed from the same shell and so there is greater attraction between the nucleus and the outer electron. Ionisation energy decreases down a group because Atomic radius increases Shielding increases due to increased number of shells and so there is less attraction between the nucleus and the outer electron Considering successive ionisation energies for individual elements allows an element’s group to be identified. Consider the successive ionisation energies of aluminium, 1s2 2s2 2p6 3s2 3p1: The 1st ionisation energy is fairly low because the 3p electron is shielded from the nucleus by all the other electrons. The 2nd and 3rd ionisation energies are significantly higher than the 1st because the 3s electrons are being removed and are not as shielded as the 3p electron. 1st: 578 kJ mol-1, 2nd: 1817 kJ mol-1, 3rd: 2745 kJ mol-1 There is a huge jump to the 4th ionisation energy, since a 2p electron is now being removed. The shielding has reduced as the 4th electron being removed is in a shell closer to the nucleus. 4th: 11578 kJ mol-1, 5th: 14831 kJ mol-1, 6th: 18378 kJ mol-1 With three electrons being removed before a huge jump in the ionisation energy, this suggests the element is in Group III. 6 FACTFILE: GCE CHEMISTRY / 1.2 ATOMIC STRUCTURE Revision Questions 1 Which one of the following represents the first five ionisation energies in kJ mol-1 of an s-block element? 1st 2nd 3rd 4th 5th A 580 1800 2700 11600 14800 B 740 1500 7700 10500 13600 C 1000 2300 3400 4600 7000 D 14800 11600 2700 1800 580 2 Neon has several isotopes. a) Complete the table below. Number of Number of Number of protons electrons neutrons Neon-20 Neon-21 Neon-22 b) The table below gives the abundance of each isotope of neon. Calculate the relative atomic mass of neon to two decimal places. Isotope % abundance Neon-20 90.92 Neon-21 0.26 Neon-22 8.82 c) Name the isotope used as the standard to compare the relative atomic mass of atoms.. 7 FACTFILE: GCE CHEMISTRY / 1.2 ATOMIC STRUCTURE Revision Questions 2 d) Label the sub-shells below and draw the electronic sctucture of neon in the ground state. e) Draw the shape of an s and of a p orbital. s orbital p orbital 3 Which one of the following represents the ground state electronic configuration of a nitrogen atom? A B C D 2p 2s 1s 8 FACTFILE: GCE CHEMISTRY / 1.2 ATOMIC STRUCTURE Revision Questions 4 Aluminium is the most abundant metal in the Earth's crust. It is used in electrical cables and is present in high strength alloys. a) Atoms of aluminum have a mass number of 27. How many neutrons are present in the nucleus of these atoms? b) (i) Write the equation, including state symbols, which represents the first ionisation energy of aluminium. (ii) E  xplain why the first ionisation energy of boron has a larger value than the first ionisation energy of aluminium. (iii) E  xplain why the first ionisation energy of magnesium has a larger value than the first ionisation energy of aluminium.  (iv) Give the ground state electronic configuration of the Al4+ ion. (v) Sketch a graph to show the successive ionisation energies of aluminium. log (ionisation energy) 1 2 3 4 5 6 7 8 9 10 11 12 13 © CCEA 2016 9