Trends In the periodic table Trends in Atomic Radii • The atomic radius of an atom id defined as half the distance between the nuclei of 2 atoms of the same element joined together by a single covalent bond Look at the atomic radii in the following table – notice any patterns? • In general atomic radii decrease across the period and increase down the group • It is important when studying any trends in the periodic tables to remember the electrostatic attraction between the positively charged protons and negatively charged electrons, if the attraction is large the protons will pull the outer electrons nearer to the nucleus giving a smaller atomic radius 2 Reasons for the increase in atomic radius down a group 1. The additional electrons are going into a new energy level thus the outer electrons are becoming further away from the nucleus giving a larger atomic radius 2. Screening effect of inner electrons : even though there are more protons and electrons as you move down a group, the inner shells of electrons help to “shield” the outer electrons from the charge of the nucleus allowing atomic radii to be bigger 2 Reasons Atomic Radii decrease from left to right across a period 1. Increasing Nuclear Charge: The number of protons in the nucleus increase from left to right across any one period. This has a greater attractive force on the outer electrons drawing the energy levels closer to the nucleus and decreasing the atomic radius 2. No increase in screening effect: Even though each element across a period has extra electrons they are being added to the same energy level so there is no increase in screening effect, thus there is nothing to “shield” the extra attractive force from the nucleus Electronegativity What is it? • Electronegativity is the power of an atom to attract electrons to itself in a covalent bond Electronegativity Pauling’s electronegativity scale • The higher the value, the more electronegative the element • Fluorine is the most electronegative element • It has an electronegativity value of 4.0 Electronegativity Pauling’s electronegativity scale H 2.1 He - Li 1.0 Be 1.5 B 2.0 C 2.5 N 3.0 O 3.5 F 4.0 Ne - Na 0.9 Mg 1.2 Al 1.5 Si 1.8 P 2.1 S 2.5 Cl 3.0 Ar - Trends in Electronegativity • In general Electronegativity values decrease down a group and increase across a period 2 Reasons Electronegativity Values Decrease down a group 1. The atomic radius increases down a group so the outermost electrons are becoming further away from the influence of the nucleus, this means there is a smaller force of attraction between the nucleus and shared pair of electrons 2. The screening effect of inner electrons “shields” the outermost electrons and also the shared pair of electrons in a bond from the influence of the nucleus. As you move down a group there is increased screening effect 2 Reasons Electronegativity Increases across a period 1. Increasing nuclear charge (due to more protons in the nucleus)as you move across a period means the attraction between the nucleus and outermost electrons becomes greater 2. Decreasing atomic radius across a period means the electrons in the outer most level are closer to the nucleus and thus there is a greater attraction between the nucleus and these electrons Ionisation energy • Some elements such a sodium and potassium lose their outermost electron very easily and thus are very reactive • Other elements such as gold and silver have very little tendency to lose their electrons and thus are very unreactive Ionisation Energy The first ionisation energy of an element is the minimum energy required to remove the most loosely bound electron from an isolated atom of that element in its gaseous state. Unit 7, 16 • The second ionisation energy of an element would refer to the energy needed for the removal of a second electron from the positive ion • Eg. Sodium • First Ionisation Energy Na Na+ + e• Second Ionisation Energy Na+ Na2+ + e- Trends in Ionisation Energy • In General Ionisation Energy increases across a period and decreases down group 2 Reasons Ionisation Energy decrease down the groups in the periodic table 1. Increasing atomic radius As the atomic radius gets larger down a group the outermost electrons are further away from the nucleus and are less influenced by its attractive force thus it is easier to remove the outermost electrons ie. Less energy is needed 2. Screening Effect of Inner Electrons increases down a period meaning the outermost electrons are not as influenced by the nucleus and thus require less energy to be removed 2 Reasons Ionisation energy Increases across the periods in the periodic table 1. Increasing nuclear charge as the number of protons in the nucleus increases the attraction between the nucleus and outermost electrons is increasing meaning more energy is required to remove one of the electrons from the outermost shell 2. Decreasing atomic radius As the radius of an atom decreases the outermost electrons get nearer the nucleus and are more influenced by its attractive force thus more energy is required to remove the outermost electron Graph of the first Twenty Ionisation Energies Unit 7, 21 Explaining the Graph 1. The maximum values are for the noble gases. Reason: Their atoms are very stable because of their electronic configuration [full outer (sub) level], so it is difficult to remove an electron. 2. The minimum values are for the group one metals (alkali metals). Reason: Their atoms have only one electron in their outer level, so it is easily removed (as when this is lost it will have noble gas configuration.) This is why group one are so reactive. 3. In general, ionisation energies increase in moving across a period from the alkali metal to the next noble gas. Reason: 1. Increase in nuclear charge. (greater pull for electrons) 2. Decrease in atomic radius. 4. Ionisation energies gradually decrease in moving down a group. Reason: 1. Increase in atomic radius. 2. Screening effect. (This is where the inner shell or shells of electrons help to shield the outer electrons from the positive charge in the nucleus. Unit 7, 22 Exceptions to Rule to 3 – Across a Period There are two exceptions to this generalisation: (a) Group two elements (e.g. Be, Mg) have abnormally high values. This is because the most loosely bound electron comes from a full s orbital. (e.g. 1s2, 2s2, 2p6, 3s2 in Mg) which is a relatively stable state. When the next element in each case (B, Al) is being ionised, the electron being removed is the single electron in the p – orbital (e.g. 1s2, 2s2, 2p6, 3s2, 3p1 in Al). (b) Group five elements also show abnormally high values (e.g. N and P). The reason here is that the electrons being removed are from exactly half – filled p –orbitals, (e.g. 1s2, 2s2, 2p6, 3s2, 3p3 in P) and the half filled orbitals are the next most stable state after that of completely filled orbitals. Unit 7, 23 Ionisation Energy Trends - Summary • Increase going across a period. – – • Decrease going down a group. – – • Increase in atomic radius. Screening effect. Exceptions, Group 2. – • Increase in nuclear charge. Decrease in atomic radius. Full (outer) sublevel. Exceptions, Group 5. – Half full (outer) sublevel. Be 1s 2 2s 2 Mg 1s 2 2s 2 2p6 3s 2 Ca 1s 2 2s 2 2p6 3s 2 3p6 4s 2 N 1s 2 2s 2 2p3 P 1s 2 2s 2 2p6 3s 2 3p3 Unit 7, 24 P.T.E Trends in IonisationEnergy Energies Ionisation Trends Unit 7, 25 Higher Ionisation Energy Levels for the Third Period Unit 7, 26 Reasons Ionisation energies increase significantly for each subsequent electron being removed from an atom • When an ion is created the remaining electrons are more strongly attracted to the nucleus as ions are slightly smaller than neutral atoms • Whenever you move to a new energy level and try to start removing electrons from here a considerable jump in ionisation energy is seen as the new energy level is closer to the nucleus and is more influenced by its attractive force (the values will be in a similar range when removing electrons from the same energy level) Example 1 : The following table gives the first ionisation energies , in KJ mol1 , of the elements in the second period of the Periodic table. Li Be B C N O F Ne 519 900 799 1090 1400 1310 1680 2080 (i) Explain the factors which account for the trend in ionisation energies across a period. (ii) Explain why the values for boron and oxygen are exceptional. Solution : (i) Increase in nuclear charge. Decrease in atomic radius. (ii) The values for boron and oxygen seem exceptional as the values for the atoms before them have abnormal values, Be due to the fact the electron is being removed from a full s - sublevel and N as the electron is being removed from a half full p - sublevel, both of which have extra stability. Unit 7, 28 Example 2 : Explain why the first ionisation energy of oxygen atoms is greater that that of chlorine atoms. Solution : Chlorine is below Oxygen in the periodic table and ionisation energies decrease as you go downwards due to: 1. Increase in atomic radius 2. Screening effect Example 3 : The first ionisation energy of Sodium is 496 KJ mol1 , and the second ionsiation energy is 4562 KJ mol1. Account for the large difference. Solution : 1st ionisation energy: Na(1s 2 , 2s 2 , 2p6 , 3s1 ) Na (1s 2 , 2s 2 , 2p6 ) e 2nd ionisation energy: Na (1s 2 , 2s 2 , 2p6 ) Na2 (1s 2 , 2s 2 , 2p5 ) e The first ionisation energy is removing an electron from a 3s orbital after which the electronic configuration will be that of a noble gas, with a full outer level. The second ionisation energy is removing an electron from a full p - orbital in a full level which is much closer to the nucleus hence a much higher value. Unit 7, 29 Trends in Groups • Trends in Chemical reactivity of Alkali Metals 1. Increasing reactivity down the group as the outer most electron is further from the nucleus 2. Reaction with oxygen, all alkali metals react with oxygen to form oxides 2K + ½O2 K2 O • Lithium in air will oxidise completely to a white lithium oxide powder in hours, it only takes a few seconds for this to happen with caesium! Reaction with water • All alkali metals react with water to form the hydroxide of the metal and release hydrogen gas Na + H2O NaOH + ½H2 The more reactive the metal the more heat will be produced and the hydrogen will catch fire this is why you see flames ! The reaction between alkali metals and acid is so dangerous it must never be attempted as too much explosive hydrogen is released Trends in the Chemical Reactivity of the Halogens • The halogens are the most electronegative elements in the periodic table, since they have such an attraction for electrons they are not found free in nature, Chlorine gas is made from sodium chloride • As fluorine is the most electronegative it is too reactive to be kept in the school laboratory • Reactivity of the Halogens decreases as you move down the group • NB Look up last two pages of this chapter in your book!