Trends in the Periodic Table (Chpt. 7)

1. Atomic radius (size) 2. Ionization energy 3. Electronegativity The three properties of elements whose changes across the periodic table are to be investigated are:

1.Trends in Atomic Radii The atomic radius (covalent radius) of an atom is defined as half the distance between the nuclei of two atoms of the same element that are joined together by a single covalent bond e.g. in a molecule of hydrogen it is found that the distance between the two nuclei is nm. Therefore the covalent radius of a hydrogen atom is 0.074/2 = 0.037nm *Note: Noble gases do not form covalent bonds with one another so they have NO atomic radius

The size of an atom depends on the attraction between the positively charged protons and negatively charged electrons in the atom: - Large attraction – positive protons will pull outer electrons closer to the nucleus - leads to smaller atomic radius - Small attraction – electrons will be found further from the nucleus – leads to a larger atomic radius *Note – understanding the trends in atomic radii values:

Screening (shielding) Effect: means that the inner shell or shells of electrons help to shield the outer electrons from the positive charge in the nucleus (Please leave space for diagram)

Atomic radius decreases across the period: - increasing nuclear charge - no increase in screening effect Decrease Atomic radius increases down the group: - new shell - screening effect Increase

1. Trends in Atomic radii: A) Across a period radius decreases Increasing effective nuclear charge: - number of protons increases from left to right across a period therefore greater attraction between nucleus and outer electrons – shells drawn closer to nucleus No increase in screening effect: - same number of shells therefore no increase in screening

B) Down a group radius increases n = 1 n = 2 n = 3 n = 4 Although there is an increase in the number of protons: New Shell – additional electrons are going into a new shell which is further from the nucleus – radius increases Screening Effect – inner electrons screen outer electrons from the nucleus

Which is Bigger??? Na or K ? Na or Mg ? Al or I ?

2.Trends in Ionisation Energy Ionisation energy is a term used to describe the tendency of an atom to lose an electron The First Ionisation Energy of an atom is the minimum amount of energy required to completely remove the most loosely bound electron from a neutral gaseous atom in its ground state 1 st ionisation energy equation for hydrogen and sodium H (g) - e - H + (g) Na (g) - e - Na + (g)

The second ionisation energy of an element refers to the removal of a second electron from the positive ion formed when the first electron is removed e.g. second ionisation energy of sodium : Na + (g) - e - Na 2+ (g) Ionisation energy unit – ‘kilojoules per mole’ Table of first ionisation energies given on pg. 80 in the log tables

Increase Decrease Ionisation Energy increases across the period: - increasing atomic charge - decreasing atomic radius Ionisation Energy decreases down the group: - increasing atomic radius - screening effect

A)Down a group ionisation energy decreases Although there is an increase in the number of protons (nuclear charge): Increasing atomic radius – radius increases therefore number of shells of electrons increases – outermost electrons are moving further away from attractive force of nucleus and it becomes easier to remove an electron from the outer shell. Screening Effect – inner electrons screen outer electrons from the positively charged nucleus – becomes easier to remove outermost electrons and ionisation energy values decrease.

B)Across a period ionisation energy increases Increasing effective nuclear charge – number of protons in nucleus is increasing as move from left to right across a period. As a result, the attraction between the nucleus and the outer electrons is increasing. More energy is now required to remove an electron from the outermost shell Decreasing atomic radius – atomic radius decreases from left to right so outer electrons drawn closer to nucleus. Due to increased attraction between electron in outermost shell and nucleus the ionisation energy values increase.

Which has a higher 1 st ionization energy? Mg or Ca ? Al or S ? Cs or Ba ?

Exceptions to the General Trend Across a Period If plot a graph of ionisation energy Vs atomic number for first 20 elements it is clear that, in any one period, some elements do not follow the smooth increase Li Be B C N O F Ne Atomic Number First Ionisation Energy Ionisation Energies for n = 2 period

Ionisation Energies for n = 3 period First Ionisation Energy Atomic Number Na Mg Al Si P S Cl Ar

In n = 2 period beryllium and nitrogen have higher values than expected In the n = 3 period magnesium and phosphorous have higher values than expected This irregularity can be explained by the fact that any sublevel that is completely filled (Be, Mg) or exactly half filled (N, P) has extra stability Because of this extra stability their ionisation energy values are higher Be: 1s 2, 2s 2 N: 1s 2, 2s 2, 2p 3 Mg: 1s 2, 2s 2, 2p 6, 3s 2 P: 1s 2, 2s 2, 2p 6, 3s 2, 3p 3

*Note: If asked to account fully for trend across 2 nd /3 rd period write out electron configurations of Be and N/ Mg and P

Second and Subsequent Ionisation Energies Evidence for the fact that electrons are arranged in shells of different energies is also provided by studying the values of a number of ionisation energies of any one particular element. Second ionisation higher than the first because removing an electron from: - an ion so there is more positive charges per electron i.e. greater nuclear charge - closer to nucleus as atomic radius of ion is smaller than corresponding atom

Successive Ionisation Energies - Bigger increase if an electron is removed from a half filled sublevel - Much bigger if new shell entered because electron being removed is: a) closer to the nucleus b) in a full sublevel c) has less shielding

Investigation of Successive Ionisation energies of Aluminium First 3 ionisation energies increase steadily as electrons are removed from the 3 rd shell. 4 th is a big jump as second shell is entered 5 th to 11 th get steadily bigger as successive electrons are removed from the second shell 12 th very big as 1 st shell is entered Note: Jumps in ionisation energies are evidence for the existence of the energy levels (shells)

3.Trends in Electronegativity Electronegativity is the relative attraction that an atom in a molecule has for the shared pair of electrons in a covalent bond Concept proposed by Linus Pauling

Increase Decrease Electronegativity increases across the period: - increasing nuclear charge - decreasing atomic radius Electronegativity decreases down the group: - increasing atomic radius - screening effect

A)Down a group electronegativity decreases Even though nuclear charge increases down a group: Increasing atomic radius: - atomic radius increases – outermost electrons are moving further away from attractive force of nucleus. Therefore smaller attraction between the nucleus and the shared pair of electrons Screening Effect: - inner electrons screen outer electrons from the positively charged nucleus – since it is outermost electrons involved in bonding, the attraction of the nucleus for these electrons decreases going down the group i.e. electronegativity decreases.

B)Across a period electronegativity increases Increasing effective nuclear charge – number of protons in nucleus is increasing as move from left to right across a period. As a result, the attraction between the nucleus and the outer electrons is increasing. Therefore the electrons involved in bonding are being more strongly attracted to the nucleus i.e. electronegativity increases Decreasing atomic radius – atomic radius decreases from left to right so outer electrons drawn closer to nucleus. Due to increased attraction between electron in outermost shell and nucleus the electronegativity values increase.

Trends Within Groups The chemical properties of elements are largely determined by the number of electrons in the outermost shell: All elements in group 1 have one electron in outermost shell – all have similar chemical properties All elements in group 7 have seven electrons in their outermost shell – all have similar chemical properties

1. Trends in chemical reactivity of Alkali Metals (Group 1) Very reactive elements Low ionisation energies and electronegativity values tend to loose electrons to form ionic compounds. Reactivity of alkali metals increases down the group (as ionisation energy decreases – more easily outer electron is lost – the more reactive the metal)

a) Reaction of Alkali Metals with oxygen: All alkali metals react with oxygen to form oxides: Potassium + Oxygen Potassium Oxide K + O 2 K 2 O 2½

b) Reaction of Alkali Metals with water: All alkali metals react with water to form the hydroxide of the metal and hydrogen gas is given off: sodium + water sodium+ hydrogen hydroxide Na + H 2 O NaOH + H 2 ½

c) Reaction of Alkali metals with dilute acids: Extremely dangerous reaction so much hydrogen produced that an explosion occurs: Sodium + hydrochloric sodium + hydrogen acid chloride Na+HClNaCl + H 2 ½

2. Trends in chemical reactivity of Halogens (Group 7): elements in group 7 are the most electronegative elements in periodic table. electronegativity values decrease down the group – fluorine being the most electronegative element due to strong attraction for electrons – very reactive elements as they tend to remove electrons from other substances The reactivity of the halogens increases up the group (as one moves up the group the electronegativity values increase)

Halogens as oxidising agents: an oxidising agent is a substance that removes electrons from other substances e.g. If chlorine gas is bubbled through a solution of bromide ions (Br - ) the chlorine takes the electrons from the bromide ions and converts them into bromine: chlorine + bromide ion chloride ion + bromine Cl 2 + Br - Cl - + Br 2 *Note: electronegativity of Cl greater than that of Br ½½

A similar reaction occurs between bromine and iodide ions (I - ) where the bromine (more electronegative) would take electrons from the iodide ions( less electronegative). In each case the more reactive halogen displaces the less reactive halogen from solution

3. Trends in Physical Properties Noble Gases = Inert gases = very little reactivity Even though it is not possible to consider any trends in chemical properties of the noble gases it is possible to study trends in their physical properties

Boiling Points of Noble Gases: - steady increase in boiling points down the group - caused by the increasing atomic radius down the group - as atoms increase in size temporary dipoles are produced due to the larger electron clouds - Van der Waals forces are stronger between bigger atoms - giving rise to higher boiling points A similar trend can be noted in the halogens as you go down the group