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Home , Ionic radius, Periodic trends, Van der Waals radius

3.2 Mrs. Page – IB SL – 2015-2016 Essential Idea

• Elements show trends in their physical and chemical properties across periods and down groups. Understandings

• Vertical and horizontal trends in the exist for , , , , and . • Trends in metallic and non-metallic behavior are due to the trends above. • Oxides change from basic through amphoteric to acidic across a . Application and Skills

• Prediction and explanation of the metallic and non-metallic behavior of an element based on its position in the periodic table. • Discussion of the similarities and differences in the properties of elements in the same , with reference to alkali (group 1) and (group 17). • Construction of equations to explain the pH changes for reactions of Na2O, MgO, P4O10, and the oxides of and with water. Nature of Science

• Looking for patterns – the position of an element in the periodic table allows scientists to make accurate predictions of its physical and chemical properties. This gives scientists the ability to synthesize new substances based on the expected reactivity of elements. Trends in Electron Configuration • As we have already seen, there is a pattern when looking at the electron configuration and the periodic table. • electrons – electrons in the highest main energy level (outer shells  s and p orbitals) • Valence electrons for the main group elements are related to their group (1 valence for alkali metals, 2 for alkaline earth  7 for halogens and 8 for noble gases) • The period number indicates the number of shells (main energy levels)

Trends in Atomic Radius • Used to describe the size of an  larger radius = larger atom • The electron cloud doesn’t have a definite edge • Get around this by measuring diatomic molecules and measuring the distance between their nuclei  half of this measure is the radius of each atom (this is sometimes called the ) • Trends: • atomic radius increases down group • atomic radius decreases across period Atomic Radius H Li • Increases down a group • More electron shells Na • More electrons so more repulsion K

Rb Atomic Radius • Decreases across period • nuclear charge increases (more protons), more attraction to electrons pulls shells closer to nucleus • (Note: the e- are in same energy level so there is no additional shielding)

Na Mg Al Si P S Cl Ar Trends in Atomic Radius • Note: it is not possible to measure the atomic radius for the noble gases because the do not form covalent bonds (unreactive) • Therefore there is another measure atomic radius – the van der Waals’ radius • The van der Waals’ radius assumes that when two collide, there is little penetration in the electron clouds of each atom, therefore if frozen the atomic radius would be ½ the diameter between the nuclei of the non-bonded atoms • Bond length can be estimated from the atomic radii Effective Nuclear Charge & Shielding • Shielding (S): Core electrons (those in inner shell) repel the outer shell electrons creating a smaller electrostatic force between the outer electrons and the nucleus. • The outer shell electrons are pushed away from the nucleus being held less tightly. • The net charge experienced by an electron is called the effective nuclear charge, Zeff

http://wps.prenhall.com/ Trends in Ionic Radius • Cations form by losing electrons (+ charge) • Metals form cations • Cations of an element have configuration. • Cations are smaller than the atom they come from • Lost electrons therefore the same number of protons are attracting fewer electrons • Less repulsion between electrons due to fewer electrons Trends in Ionic Radius • Anions are form by gaining electrons (- charge) • Non-metals form anions • Anions of an element have noble gas configuration. • Anions are larger than the atom they come from • More electrons means more repulsion between them so increases electron cloud • Number of protons remained the same and more electrons so attractive force weakened Trends in Ionic RadiI -3 B+3 N O-2 F-1 Li+1

Be+2 C+4 First Ionization Energy

• The energy required to remove one electron from each atom in one of gaseous atoms under standard conditions. + - • X(g)  X (g) + e First Ionization Energy • Down groups first ionization energy decreases • Size of atom increases down group so outer electrons are further from nucleus and so less strongly attracted to the nucleus • Increased shielding First Ionization Energy • Across a period first ionization energy increases • Larger nuclear charge w/out adding energy levels attracts electrons more strongly • Atom is smaller so electrons are more strongly held First Ionization Energy •EXCEPTIONSBoron has a lower ionization energy than • Be 1s2 2s2 B 1s2 2s2 2p1 • 2p sub level is higher energy so less energy is needed to remove the electron from B First Ionization Energy •EXCEPTIONSOxygen has a lower ionization energy than nitrogen • N 1s2 2s2 2p3 O 1s2 2s2 2p4 • has two electrons paired in same p orbital making it easier to remove an electron due to the repulsive force between electrons in the same shell. Electron Affinity Eea • The change in energy (in kJ/mole) of a neutral atom (in the gaseous phase) when an electron is added to the atom to form a negative . - - • X(g) + e  X (g) • See data book Section 8 • Negative signs indicate energy is released (exothermic) • More negative = greater attraction for the electron Electron Affinity Down Group 17 • Not same clear trend as others

• General trend is that Eea decreases down a group • Atoms larger so weaker attraction to electrons • In smaller atoms electron repulsion force is greater (less exothermic when electrons are added to smaller atoms) Electron Affinity Across period

• General trend is that Eea becomes more exothermic (increases) • Increased nuclear charge & smaller atomic radii  electrons more strongly attracted Electronegativity • A measure of the attraction of an atom in a molecule for the electron pair in the of which it is a part. • In a covalent bond  atoms do not share the pair of bonding electrons equally Electronegativity – Across Period • Increases across period • Increased nuclear charge & no additional shielding Electronegativity – DOWN Group • Decreases down a group • Bonding electrons far from nucleus – less attraction, more shielding Trends in metallic and Non-metallic properties

• Metallic properties decrease across period • Metallic properties increase down a group • Metals lose electrons – get oxidized • Non-metals gain electrons – get reduced Melting Points • Depends on the type of bonding and structure of compound • Down a group – elements bond in a similar way • For metals – melting point decreases down a group (elements held together by delocalized electrons and positive ), attraction decreases with distance • Nonmetals – melting points increase down a group (more in topic 4) • Generally increase across a period until group 14 and then decrease to group 18. (more in topic 4) Group Trends – Alkali Metals • Group 1 Elements: Alkali Metals • Highly reactive – not found in nature (1 ) • Reactivity increases down the group (ionization energies decrease) • Soft silvery metals • Low melting points (decreasing down group) • Form M+ ions • React vigorously with oxygen (tarnish quickly)

• 4M(s) + O2(g)  2M2O(s) Group Trends – Alkali Metals • Group 1 Elements: Alkali Metals • React rapidly with water (stored in oil) • Good conductors • Reactivity with water increases down a group

• 2M(s) + H2O(l)  MOH(aq) + H2(g) • Forms an alkali (strong base), which ionizes completely in aqueous solution

• Ex: 2K(s) + H2O(l)  2KOH(aq) + H2(g) Group Trends – Halogens • Group 17 Elements: Halogens • Nonmetals

• Most diatomic molecules (X2) • Melting points increase down the group • 7 valence electrons  react by gaining an electron or forming covalent bonds • Reactivity decreases down the group (F is most reactive element) • React with alkali metals to form salts

• 2M(s) + X2(g)  2MX(s) Group Trends – Halogens

• Displacement Reactions of Halogens • Reaction of with a solution containing halide ions • K, Cl, Br, and I solutions are colorless KCl(aq) KBr(aq) KI(aq) • Cl2(aq) + 2KBr(aq)  2KCl(aq) + Br2(aq) No Orange Dark red/ • Cl2(aq) + 2KI(aq)  2KCl(aq) + I2(aq) Cl2(aq) reaction solution brown • Br2(aq) + 2KI(aq)  2KBr(aq) + I2(aq) solution No No Dark red Br2(aq) reaction reaction /brown Color change to brown due to solution Color change to orange due to No No No reaction I2(aq) reaction reaction Group Trends – Noble Gases

• Colorless gases • Monoatomic • Least reactive elements – full valence shell (stable octet) • Highest ionization energies Trends in metal & Non-metal Oxides • Oxides  when element combines with oxygen • Oxygen forms on O2- ion

• Ex: Na2O, CaO, Al2O3 • Metal oxides form giant ionic structures • Metal oxides are basic & react with water to form metal

• Ex: Na2O(s) + H2O(l)  2NaOH(aq) Trends in metal & Non-metal Oxides • Non-metal oxides are formed by covalent bonding and form giant covalent structures • Non-metal oxides are acidic and react with water to form acidic solutions

• CO2(g) + H2O(l)  H2CO3(aq) carbonic acid • SO3(l) + H2O(l)  H2SO4(aq) sulfuric acid • SO2(g) + H2O(l)  H2SO3(aq) sulfurous acid • P4O10(s) + H2O(l)  4H3PO4(aq) phosphoric acid Trends in metal & Non-metal Oxides • Non-metal oxides are acidic and react with water to form acidic solutions

• CO2(g) + H2O(l)  H2CO3(aq) carbonic acid • SO3(l) + H2O(l)  H2SO4(aq) sulfuric acid • SO2(g) + H2O(l)  H2SO3(aq) sulfurous acid • P4O10(s) + H2O(l)  4H3PO4(aq) phosphoric acid Interesting Oxides • dioxide does not dissolve in water but is an acidic oxide

• SiO2(s) + 2NaOH(aq)  Na2SiO3(aq) + H2O(l) • Aluminum oxide is amphoteric (reacts as both an acid and a base)

• Al2O3(s) + 2NaOH(aq) + 3H2O(l)  2NaAl(OH)4(aq) • Al2O3(s) + 6HCl(aq)  2AlCl3(aq) + 3H2O(l) Polyatomic Ions to KNOW - p. 86 FORMULA NAME FORMULA NAME 2- PO 3- Phosphite CO3 Carbonate 3 2- PO 3- Phosphate C2O4 Ethanedioate (oxalate) 4 - ClO- Hypochlorite NO2 Nitrite - ClO - Chlorite NO3 Nitrate 2 2- ClO - Chlorate SO3 Sulfite 3 2- ClO - Perchlorate SO4 Sulfate 4 OH- Hydroxide