Trends in Periodic Table

Atomic Radius of an element is defined as half the distance between the nuclei of the two atoms of the same element that are joined together by a single covalent bond Atomic Radius Bond Length

- What happens to the size of the atomic radius as you go down a group - What happens to the size of the atomic radius as you go down a group? - What happens to the size of the atomic radius as you go across a period?

Why does the size of the atomic radius differ? New shells and sublevels increase the atomic radius as you go down a group Screening effect of inner electrons cause the outer electrons to experience less nuclear charge so are not as tightly held, increasing the atomic radius Increasing nuclear charge and no increase in the screening effect from L-R across a period causes a decease in the atomic radius

Trends in Atomic and Ionic Size Metals Nonmetals Group 1 Al 143 50 e Group 13 Group 17 e e 152 186 227 Li Na K 60 Li+ F- 136 F Cl Br 64 99 114 e e 95 Na+ Cl- 181 Al3+ e e 133 K+ Br- 195 Cations are smaller than parent atoms Anions are larger than parent atoms

Why cant you measure the size of a single atom Why cant you measure the size of a single atom? [2] cant determine where outermost electron is going to be / heisenberg’s Uncertainty principle What is atomic radius? [4] half distance / between two adjacent nuclei / of same element / joined by single covalent bond What happens to atomic radius as one goes across period? Why? [3] Gets smaller / increased nuclear charge / same screening effect What happens to atomic radius as you go down a group? Why? [4] Gets bigger / new main energy level (shell) / increased nuclear charge / cancelled out by increased screening effect For which group are there no atomic radius values? Explain? [3] noble gases / don’t bond / cant predict position of outermost electron

Ionisation Energy Minimum energy required to completely remove the most loosely bound electron from a mole of gaseous atoms in their ground state Unit: kJ mol-1 The general symbol is D Hi, and for a first ionisation energy it is D Hi1. The process may be shown by the example of calcium as: Ca (g) = Ca+ + e-

Page 80 - Log Tables Ionisation Energy Increases L-R Ionisation Energy decreases going down a group Increasing nuclear charge/ no increase in screening Decreasing atomic radius Increasing atomic radius Larger screening effect

Across a Period Ionisation energy Increases (i) increasing nuclear charge. Number of protons increases so attraction increases electrons in same main energy level so no increase in screening (ii) decreasing atomic radius atomic radius decreases so electrons are closer to positive nucleus so held tighter.

Down a Group Decreases [gets less] (i) Increasing atomic radius means electron further from attractive force of the nucleus (ii) Screening effect of inner electrons. Positive charge of nucleus is increasing but inner e- shells shield the outer electron from this increased charge.

Exceptions

Explanation Beryllium higher than Boron Nitrogen higher than Oxygen Removing electron from a full s sublevel Having a full s sublevel gives Beryllium extra stability making it harder to remove an electron Nitrogen higher than Oxygen removing electron from a half filled p sub-level Having a half full p sublevel gives Nitrogen extra stability making it harder to remove an electron from it than Oxygen

Graph of ionisation energies Why has Aluminium has a lower first ionisation energy than Magnesium?

Mg 1s2, 2s2, 2p6, 3s2 Al 1s2, 2s2, 2p6, 3s2 3px1 Aluminium ionisation energy lower than Mg Removing electron from Mg is from a full s sublevel which requires more energy than removing electron from Al which is from a partially filled p sublevel Having a full s sublevel gives Mg extra stability making it harder to remove an electron

Exceptions to general trend Look at the general trend. Argon has the highest value. Why? Full sublevel [outer shell] Are any elements higher than general trend? Yes Mg and P Why?

Explanation Magnesium higher Removing electron from a full s sublevel Phosphorous higher removing electron from a half filled p sub-level Having a half full p sublevel gives Phosphorus extra stability making it harder to remove an electron from it

Define first ionisation energy Define first ionisation energy. [5] The energy required / to completely remove / the most loosely bound electron / from a [mole of] neutral gaseous atom [s], / in its [their] ground state Write an equation for the first ionisation of potassium. [2] K = K+ + e- What happens to first ionisation energy across a period? Explain [3] Increases / increasing nuclear charge / decreasing atomic radius What happens to first ionisation energy down a group? Explain [4] Decreases / Increasing atomic radius / increased screening effect / of inner electrons. Be and N have higher values than the general trend; explain why [4] Be filled / S sublevel / N ½ filled / P sublevel

Second and successive Ionisation Energy

Second ionisation energy The energy needed to remove a second electron from each ion in a mole of ions is the second ionisation energy. e.g. for calcium: Ca+(g) = Ca2+(g) + e- DHi2 = + 1150 kJ mol-1 The second ionisation energy is much larger than the first. What are the reasons for this?

Reasons The ionisation energies increase because as each electron is removed from an atom, the remaining ion becomes more positively charged. Removing the next electron away from an increasing positive charge is more difficult and the ionisation energy is even larger.

Further evidence for the existence of Energy Levels Take a Na atom and remove the electrons one by one. Plot a graph of the successive ionisation energies. Use log10 because values so big.

Graph of logarithm (log 10) of ionisation energy of sodium against the number of electrons removed. There are one or more particularly large rises within the set of ionisation energies of each element. Why??

Why? Why are there large increases between the first and second ionisation energies and again between the ninth and tenth ionisation energies? Look at the s,p,d configurations to find out

Trends within groups Alkali Metals: What do you think happens to the reactivity as you go down the group? Reactions with oxygen: 2K + ½O2 = K2O Reactions with water: Na +H2O = NaOH +½H2

Halogens What happens to the reactivity of halogens as you move down the group? Reactivity Decreases

Trends in Electronegativity Is the relative attraction of an atom in a molecule has for the shared pairs of electrons in a covalent bond Page 81 - Log Tables

Electronegativity increases Increasing nuclear charge Decreasing atomic radius Increasing atomic radius Larger screening effect Electronegativity Decreases

Is the relative attraction that an atom in a molecule has for the shared pair of electrons in a single covalent bond between two atoms of the same element. In a bond between two identical atoms the pair of electrons are shared equally chemists have found that in many bonds the pair of electrons are attracted to one of the atoms more than to the other. SCC Science Dept

Hydrogen and Chlorine Electrons attracted to chlorine more than to hydrogen [ bigger , but more +ve nucleus] therefore the electrons spend more time near the chlorine than near the hydrogen this gives the chlorine a slightly negative charge [δ- delta minus] it gives the hydrogen a slightly positive charge [δ+ delta plus] SCC Science Dept

H Cl SCC Science Dept

Linus Pauling measured the electronegativity of each element and put them in a table Noble gases are not in this table because they do not form bonds H Cl +  – SCC Science Dept

Differences and Bond Type Difference 0 to 0.45 [Pure] Covalent Bond Difference > 0.45 but < 1.7 Polar Covalent Bond Difference is = or > 1.7 then Ionic Bond SCC Science Dept

Trends in Electronegativity Across a period Goes up Bigger nuclear charge – Smaller atomic radius distance from nucleus Actually closer as diameter of atom gets smaller – SCC Science Dept

Trends in electronegativity Down a Group Goes down Bigger nuclear charge But! But! But! But! Increased Atomic Radius - electron is much further from nucleus Increased shielding by more inner electron shells SCC Science Dept

Who invented the term? [1] Linus Pauling Define electronegativity [3] measure of the attraction of an atom / for an electron / in a single covalent bond / between atoms of the same element Who invented the term? [1] Linus Pauling What happens to it as you go across a period. Explain.[3] Goes up / bigger nuclear charge / same shell or decreasing atomic radius What happens to it as you go down group? Explain. [3] Goes down / further from nucleus / more shielding Give three classes of bond and electronegativity associated with each [6] 0 - 0.4 / pure covalent / 0.4 to 1.7 / polar covalent / over 1.7 / ionic How can you show that a molecule is polar. [4] Rub pen / pour stream past it / polar bends / non-polar does not bend. What group of elements is not present in the table of electronegativities. Why? [2] Noble Gases / don’t bond Total = 22 SCC Science Dept