Chapter 3 Notes: Periodic Trends Chapter 3.2: Elements show trends in their physical and chemical properties across periods and down groups. Chapter 3 Notes: Periodic Trends

Important terms for this section Periodicity Atomic radii Ionic radii Electronegativity Electron affinity Ionization energy Melting points Effective nuclear charge Chemical properties Noble gases Alkali metals Halogens Period 3 oxides

Effective Nuclear Charge The attraction “experienced” by the outer electrons from the positively charged nucleus Shielding occurs moving down groups from inner electrons ~ same ENC More p+ + e- means higher ENC moving across period

Effective Nuclear Charge

ENC continued: Transition metals Nuclear charge increases from Scandium to copper (e- added to inner 3d sub-level) For the most part, 3d electrons shield 4s electrons from increasing nuclear charge The 4s electrons of the 1st row of transition metals feels only slightly increasing ENC as atomic number increases This means the atomic radii only gradually increases as well Increase in nuclear charge is ~equal to the inc in shielding effect  ENC remains about constant

ENC of first 3 rows

Atomic radius – Values in Table 8 of booklet Measured: half distance between neighboring nuclei Since ENC increases across period, radius decreases Since more e- repulsion moving down (add an E level), radius increases

Ionic radius – 5 important points Cations are smaller than parent atoms Anions are larger than parent atoms Ionic radii decrease from groups 1-14 for cations Increase ENC, results in decrease of size Ionic radii decrease from groups 14-17 for anions (for same reason as above) Positive ions are smaller than negative ions of similar electron config due to the loss of E level Ionic radii increase down a group b/c E level increases

Ionic radii trends

Ionization energies Energy required to remove 1 mol of e- from 1 mol of gaseous atoms in their ground state Measure of attraction between nucleus and its outer electrons IE generally increases across period Increase of ENC causes inc of attraction and so difficult to remove IE decreases down group E- removed is furthest from nucleus ENC inc, but more repulsion of inner e- with outer e- (shielding) Differences in trend shows evidence of sub-levels

Ionization energy

Electron affinity – Values in Table 8 of booklet Energy change when 1 mol of e- is added to 1 mol of gaseous atoms to form 1 mol of gaseous ions As e- is added to pos. nucleus, usu. exothermic Releases energy b/c the attraction is favorable Additional e- added is endothermic due to repulsive nature of the already added e- Thought of as the “negative” of 1st IE values Closely related to electronegativity but not same ENC increases, electron affinity increases

Electron affinity

Electronegativity – Values in Section 8 of booklet The ability of an atom to attract electrons in a covalent bond So attraction of + nucleus with bonding e- on another atom(s) General same trend as with IE Increases across period Inc of ENC, so increase attraction to e- Decreases down group Greater distance, so decrease attraction to e-

Electronegativity

Metals vs Non-metals Metals have lower ionization energies and electronegativities than non-metals The ability of metals to have mobile electrons (which is why they can conduct) also demonstrates this property Electrons that do this are called delocalized “Sea of electrons” is the way metals bond (ch 4)

Melting Points Moving down group 1, MP decreases b/c the attraction of nucleii with delocalized e- decreases with distance Moving down group 17, intermolecular forces (IMFs), specifically London dispersion forces, increases, so MP increases Across periods MP increases from group 1-14 Group 15-18, drop and decrease to minimum

Chemical Properties: Noble Gases – Group 18 Colourless gases Monatomic Very unreactive All have stable octets He has stable first principle E level Highest IE and almost no electronegativity They are “full” of themselves and do not react with others due to their nobility Note: chemical properties of elements are mainly determined by the number of valence electrons (thus, groups have similar properties)

Chemical Properties: Alkali Metals – Group 1 Silvery color and too reactive to be found in nature (in pure form) Form singly charged cations = M+ Low ionization energies Reactivity increases down group Ability to conduct based upon outer e-

Chemical Properties: Alkali Metals – Reaction with H2O Alkali metal + H2O  H2 + metal hydroxide Increasing intensity of reaction moving down the group: Li < Na < K < Rb < Cs Example reaction: 2K(s) + 2H2O(l)  2KOH(aq) + H2(g) The KOH is ionic and will dissociate in water: 2K(s) + 2H2O(l)  2K+(aq) + 2OH-(aq) + H2(g)

Chemical Properties: Halogens – Group 17 Diatomic molecules = X2 Accepts an electron very easily High ENC (~ +7) and so pull on e- in other atoms Reactivity decreases down group F2 and Cl2 are gases; Br2 is a liquid; I2 is a solid at room temp

Chemical Properties: Halogens – Reaction with Group 1 Very reactive with alkali metals to form ionic halides (salts) Complementary reactants: Na+ and Cl- 2Na(s) + Cl2(g)  2NaCl(s) High ENC of Cl2 pulls outer Na e- and electrostatic attraction completes transfer

Chemical Properties: Halogens – Displacement RXNs Which is most reactive? Br or F? Cl or I? What is the order of increasing reactivity? What do you think would happen if you pour brominated water into chlorinated water? What do you think would happen if you pour chlorinated water into brominated water?

Chemical Properties: Halogens – Displacement RXNs When a more reactive halogen is added to a solution of less reactive: The less reactive is pushed out of solution Chlorine is green/yellow Bromine is orange/brown Iodine is violet Halogens are more soluble in non-polar solvents

Chemical Properties: Halogens – Identifying Halides When reacting with silver, halides form insoluble salts Adding a halide solution with silver produces a precipitate Ag+(aq) + X-(aq)  AgX(s)

Chemical Properties: Period 3 Oxides – Bonding Na  Na2O Mg  MgO Al  Al2O3 Si  SiO2 P  P4O10  P4O6 S  SO3  SO2 Cl  Cl2O7  ClO Moving across period: Ionic bonding transitions to covalent: NaAl have giant ionic struc. PCl have molecular covalent Si has giant covalent (network) Oxides only conduct in liquid form (moving of ions) Oxidation number related to group number

Chemical Properties: Period 3 Oxides – Bonding

Chemical Properties: Period 3 Oxides – Acid-Base Character Acid-base character is based upon the bonding and structure Acid- has pH of 6 or less (covalent) Base- has pH of 8 or more (ionic) Amphoteric- can react with acids and bases

Chemical Properties: Period 3 Oxides –Basic Oxides Basic oxides – Na and Mg dissolve in H2O to form alkaline solutions (has hydroxide ions) Na2O(s) + H2O(l)  2NaOH(aq) MgO(s) + H2O(l)  Mg(OH)2(aq) 2NaOH can also be written: 2Na+(aq) + 2OH-(aq) Basic oxides react w/ acids salt and H2O O2-(s) + 2H+(aq)  H2O(l) Li2O(s) + 2HCl(aq)  2LiCl(aq) + H2O(l) MgO(s) + 2HCl(aq)  MgCl2(aq) + H2O(l)

Chemical Properties: Period 3 Oxides – Acidic Oxides Acidic oxides – Si, P, S, Cl dissolve in water to produce acidic solutions Phosphorus(V) oxide reacts with H2O  Phosphoric(V) acid

Chemical Properties: Period 3 Oxides – Acidic Oxides

Chemical Properties: Period 3 Oxides – Amphoteric Oxides Aluminum does not affect pH (essentially insoluble) Will act like a base when reacted with acid Al2O3(s) + 6H+  2Al3+(aq) + 3H2O(l) Al2O3(s) + 3H2SO4(aq)  Al2(SO4)3(aq) + 3H2O(l) Will act like acid when reacted with bases Al2O3(s) + 3H2O(l) + 2OH-(aq)  Al(OH)4-(aq)

Period 3 Oxides Trends

You need to know the following: The equations for Na2O MgO Al2O3 amphoteric property P4O10 SO3