Periodic Trends

• things that repeat in the periodic table

• atomic/ionic radii, electonegativity, ionization energy → physical properties that explain the chemical properties

Effective Nuclear Charge

• nuclear charge = number of protons = atomic number

• inner electrons “shield” the nuclear charge that comes from the nucleus from the outer electrons

  • the outer electrons do not feel the full attracton of the nuclear charge as they are shielded and repelled by the inner electrons

  • the charge they do experience is the “effective charge” which is less than the full nuclear charge

• outer electrons = effective nuclear charge

  • example: 11 protons 10 inner electrons (Na) → 11-10 = +1 effective nuclear charge

    • outermost shell has only 1 electron so 11 total electrons but 10 inner electrons

• effective nuclear charge increases with atomic number as you go across a period (left to right) as there is no change in the number of inner electrons, as all atoms have a noble gas structure

Atomic radius

• to find size of an atom, measure distance between two nuclei in a bonded form

  • used as a diameter, so cut that measurement in half to find the atomic radius

• increasing nuclear charge causes the atom to decrease in size

• if you go down a group the size increases because the number of occupied electron shells (given by the periodic number) increases

• if you go across a period the size decreases because they go in the same energy level when you go across and the electrostatic forces are stronger

  • higher electron energy level = bigger

  • distribution of the charge / electrostatic forces affects the atomic radius

    • decreased pull = bigger atom, outer electrons go further away from the nucleus

Ionic radius

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• positive ions are smaller, negative ions are bigger

  • positive → more pull distributed among each electron as one or more electron leaves

  • negative → increased electron repulsion between electrons in the outer principal energy level causes the electrons to move further apart and so increases the radius of the outer shell

• ionic radii decreases from Groups 1-14 for positive ions even though they have the same electron configuration

  • increase in nuclear charge with atomic number across the period causes increased attraction between the nucleus and the electrons (pulls the outer shell closer to the nucleus)

• ionic radii decreases from Groups 14-17 for negative ions even though they have the same electron configuration

  • increase in nuclear charge across a period

  • positive ions are smaller than the negative ions, as the former have only two occupied electron principal energy levels and the latter have three

    • discontinuity in the table (1-14, 14-17)

• ionic radii increases down a group as the number of electron energy levels increase

Ionization energy

• measure of attraction between nucleus and outer electrons

• increases across a period

  • increase in effective nuclear charge causes an increase in the attraction which makes the electrons more difficult to remove

• decreases down a group

  • electron removed is from the energy level furthest from the nucleus, and although nuclear charges increases, the effective nuclear charge is about the same because of the shielding effect, but the distance between the outer electrons and the nucleus is increased so the attraction is weaker

• higher energy level / higher energy (p>s) means ionization energy is smaller as the increased energy makes the orbital more unstable and more susceptible to “pulling” an electron out

• gaseous atoms

• removing one electron

• energy input (all positive values)

• endothermic

• A (g) → A+ + e-

• 

Electron affinity

First electron affinity of an element

• energy change when one mole of electrons is added to one mole of gaseous atoms to form one mole of gaseous ions

  • A (g) + e- → A-

• exothermic

  • added electron is attracted to the positvely charged nucleus

Second electron affinity of an element

• defined similarly to first electron affinity

  • example: O-(g) + e- → O2-(g)

    • process is endothermic as the added electron is repelled by the negatively charged oxide ion and energy needs to be available for this to occur

• negative values = exothermic, positive = endothermic

• similar to ionization energy graph, but displaced to the right by one and inverted

  • electron affinity minimum values = Group 17, ionization energy maximum values = Group 18

• Group 17 elements have incomplete outer energy levels and a high effective nuclear charge of approximately +7 so attraction is strongest

• Group 1 metals have lowest effective nuclear charge of approximately +1 and so attraction is weakest

• Group 2 and 5 elements are the maximum because they have ns² electron configurations, so the added electron must be put into a 2p orbital which is further from the nucleus and so experiences reduced electrostatic attraction due to shielding from electrons in the ns orbital

• Group 15 elements have the configuration ns2np3 so the added electron must occupy a p orbital that is already singly occupied (arrow diagram)

  • the attraction between the electron and atom is less than expected as there is increased inner-electron repulsion (exothermic only for nitrogen)

Electronegativity

• attraction/pull an atom has on the electron pairs it shares with another atom in a covalent bond (attraction between nucleus and electrons)

  • covalent bond (shared electron pairs) → not equally shared

• similar to ionization energy

  • both measure attraction between the nucleus and its outer electrons - in this case bonding electrons

  • same trends (period and group)

  • differences

    • ionization energy can be measured directly and are a property of gaseous atoms

    • elements with high electronegativities have the most exothermic electron affinities

      • electron affinity is a property of isolated gaseous atoms

    • electronegativity is a property of an atom in a molecule

      • derived indirectly from experimental bond energy data

• increases across a period

  • increase in nuclear charge resulting in an increased attraction between nucleus and bond electrons

• decreases down a group

  • electrons furthest from the nucleus as radius increases so there is reduced attraction

• most electronegative element(s) are on the top right of the periodic table and the least electronegative element(s) are on the bottom left

• relative scale / no units (0-4.00)

• metals have lower ionization energies and electronegativies than non-metals

Melting points

• decreases down Group 1

  • the elements have metallic structures which are held together by attractive forces between delocalized outer electrons and the positively charged ions (attraction decreases with distance)

    • delocalized electrons - electrons that move about atoms in a metallic structure

  • although they have the same charge down a group, there is a greater volume so the distribution of charge is less as the distance between delocalized electrons are greater

    • bigger ions have weaker metallic bonds

    • easier to melt

• increases down Group 17

  • the elements have molecular structures which are held together by London dispersion forces (increases with number of electrons)

• generally rise across a period and reaches a maximum at group 14 then falls to reach a minimum at group 18

  • in Period 3, the bonding changes from metallic (Na, Mg, Al) to giant covalent (Si) to weak van der Waal’s attraction between simple molecules (P4, S8, Cl2) and single atoms (Ar)

    • all Period 3 elements are solid at room temperature except chlorine and argon

• 

• structure explanation of trends on graph

  • Li, Be, B, Na, Mg, Al → metallic

  • C → giant covalent network

  • Si → giant covalent

  • P, S, Cl, Ar → simple molecules, single atoms

Chemical Properties

• determined by electron configuration of the atom

  • elements in the same group contain similar chemical properties as they have the same amount of valence electrons

• intermolecular forces

  • London Dispersion, dipole-dipole, hydrogen bonding

    • London Dispersion - more electrons = stronger forces = bigger atomic mass

Groups in the Periodic Table

Group 18 (noble gases)

• colourless gases

• monoatomic - exists as single atoms

• very unreactive

  • inability to lose or gain electrons

    • do not form negative ions as the electron would be added to an empty outer energy level shell where they would experience a negligible effective nuclear force

  • complete valence energy levels with 8 electrons (except helium which has a complete principal first energy level with 2 electrons)

    • stable octet

Group 1 (alkali metals)

• physical properties

  • good conductors of electricity and heat

    • due to the mobility of their outer electron

  • low densities

  • grey shiny surfaces when cut with a knife

• chemical properties

  • very reactive

  • forms ionic compounds with non-metals

    • forms single charged ions with the stable octet of the noble gases when they react

  • reactivity increases down the group as the elements with higher atomic number have the lowest ionization energies

Reaction with water

• reacts with water to produce hydrogen and the metal hydroxide

  • lithium floats and reacts slowly (releases hydrogen but keeps its shape)

  • sodium reacts with a vigorous release of hydrogen (heat produced is enough to melt the unreacted metal, which forms a small ball that moves around on the water surface)

  • potassium reacts even more vigorously to produce enough heat to ignite the hydrogen produced (produces a lilac coloured flame and moves excitedly on the water surface)

• metals are called alkali metals because the resulting solution is alkaline owing to the presence of the hydrogen ion formed

• reaction gets more vigorous down the group

  • caesium (lowest ionization energy) forms positive ions most readily

Group 17 (halogens)

• exists as diatomic molecules

• physical properties

  • coloured

  • shows a gradual change from gas (F2, Cl2) to liquid (Br2) and solid (I2 and At2)

• chemical properties

  • very reactive non-metals (reactivity decreases down group)

    • readiness to accept electrons, illustrated by their very exothermic electron affinities

    • nuclei have a high effective charge, and so exert a strong pull on any electron from other atoms which then the extra electron completes the valence shell

    • reactivity decreases down group as atomic radius increases and attraction for outer electrons decreases

  • forms ionic compounds with metals and covalent compounds with non-metals

Reaction with Group 1 metals

• halogens react with Group 1 metals to form ionic halides (stable octets)

  • example: 2Na(s) + Cl2(g) → 2NaCl(s)

• the electrostatic forces between the oppositely charged ions bonds the ions together

  • the outer electron moves like a harpoon from sodium to chlorine and then the opposite charges of the two ions pull them together

• most vigorous reaction occurs between the elements furthest apart in the Periodic Table; the most reactive alkali metal (at the bottom of Group 1) and most reactive halogen (at the top of Group 17)

Displacement reactions

• the relative reactivity can be seen by placing them in direct competition for an extra electron

  • example: 2KBr(aq) + Cl2(aq) → 2KCl(aq) + Br2(aq)

    • chlorine is more reactive as it displaced bromine

    • stronger attraction for an electron because of smaller atomic radius

• colour changes are used to determine whether or not the reaction has occured

Halides

• halogens form insoluble salts with silver

  • adding a solution containing the halide to a solution containing silver ions produces a precipitate that is useful in identifying the halide ion

    • colour of the precipitate helps you identify the halide

Period 3 Oxides

Bonding of the Period 3 oxides

• the transition from metallic to non-metallic character is illustrated by the bonding of the Period 3 oxides

  • ionic compounds are generally formed between metal and non-metal elements so the oxides of elements Na to Al have giant ionic structures

  • covalent compounds are formed between non-metals, so the oxides of phosphorus, sulfur, and chlorine are molecular covalent

  • oxide of silicon (which is a metalloid) has a giant covalent structure

• ionic character of a compound depends on the difference in electronegativity between its elements

  

  • oxygen has an electronegativity of 3.4, so the ionic character of the oxides decreases from left to right, as the electronegativity values of the Period 3 elements approach this value

  • oxides become more ionic down a group as the electronegativity decreases

• 

  • conductivity of the molten oxides gives an experimental measure of their ionic character, they only conduct electricity in liquid form if the ions are free to move (as shown on table)

  • maximum oxidation number of a Period 3 element is related to the group number

    • +1 for Group 1, +2 for Group 2, +3 for Group 13, +4 for Group 14, etc.

Acid-base character of the Period 3 oxides

• acid-base properties of the oxides are closely linked to their bonding and structure

  • metallic elements are basic, non-metal oxides are acidic

    • aluminium oxide (ionic oxide with some covalent character) shows amphoteric properties

• amphoteric - able to react as both an acid and a base

• 

Basic Oxides

• alkaline solutions (because of the hydroxide ions)

  • Na2O(s) + H2O(l) → 2NaOH(aq)

  • MgO(s) + H2O(l) → Mg(OH)2 (aq)

• basic oxide reacting with an acid to form a salt and water

  • O²-(s) + 2H+(aq) → H2O(l)

  • Li2O(s) + 2HCl(aq) → 2LiCl(aq) + H2O(l)

  • MgO(s) + 2HCl(aq) → MgCl2(aq) + H2O(l)

Acidic Oxides

• non-metallic oxides react readily with water to produce acidic solutions

  • P4O10(s) + 6H2O(l) → 4H3PO4(aq)

  • P4O6(s) + 6H2O(l) → 4H3PO3(aq)

  • SO3(l) + H2O(l) → H2SO4(aq)

  • SO2(g) + H2O(l) → H2SO3(aq)

  • Cl2O7(l) + H2O(l) → 2HClO4(aq)

  • Cl2O(l) + H2O(l) → 2HClO(aq)

• silicon dioxide does not react with water, but reacts with concentrated alkalis to form silicates

  • SiO2(s) + 2OH-(aq) → SiO3²-(aq) + H2O(l)

Amphoteric Oxides

• behaves as a base

  • Al2O3(s) + 6H+ → 2Al3+(aq) + 3H2O(l)

  • Al2O3(s) + 3H2SO4(aq) → Al2(SO4)3(aq) + 3H2O(l)

• behaves as an acid

  • Al2O3(s) + 3H2O(l) + 2OH-(aq) → 2Al(OH)4-(aq)