Chapter 3 & 13 IB Chemistry HL

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there is a greater attraction between nucleus and outer energy shell, decreasing radius ● electron shielding decreases Atomic radii decrease O Electron shielding: occurs when outer electrons are shielded from the attraction of the nucleus by inner electrons (known as shielding electrons) Atomic radii increase Ionic radius Isoelectronic ions: species have the same electron configuration.BUT different numbers of protons, as they are different elements. (eg. Na+, Mg2+, Al³+/ N3³-, 0²-, F-) Cations: ● Cations atom there are less electrons than protons so electrostatic force between nucleus and valence electron increases so radius decreases the radius of cation is smaller than the parent atom Na 160 lonic radius (pm) Anions: Anions atom CI 100 there are more electrons than protons so the electrostatic force between nucleus and valence electron decreases so radius increases (increase in nuclear shielding) radius is larger than the parent atom 300- ion 200- 102 100- Down a group: radius increases because energy shell gets further from nucleus and higher electron shielding Across a period: 0 Nat ion CI ● decreases because nuclear charge increases 181 1 T T 11 12 13 14 15 16 Na Mg² AP sit pr Atomic number コー 52 Cr 7 18 Ar from Si4+ tp P³- there is a large increase because negatively-charged ions are larger than their parent atoms and positive ions First lonisation energy ● a measure of the attraction between the positively charged nucleus and the negatively charged outer valence electrons Down a group: FIE decreases because the outer shell is further away from nucleus so it takes less energy to remove the valence electron Across a period: FIE increases the electrons are filling the same energy level and also an increase in nuclear charge (lower radius), increasing the amount of energy it takes to remove valence electron Exceptions: ● decrease in ionisation energy from beryllium to boron O Beryllium 1s² 2s² and Boron 1s² 2s²2p¹ O Electrons in p orbitals are of higher energy and further from the nucleus than electrons in s orbitals, therefore they require less energy to remove decrease in ionisation energy from magnesium to aluminium O Magnesium 1s² 2s² 2p 3s² and Aluminium 1s² 2s² 2p 3s² 3p¹ O Electrons in p orbitals are of higher energy and further from the nucleus than electrons in s orbitals, therefore they require less energy to remove decrease in ionisation energy from nitrogen to oxygen O Nitrogen 1s² 2s²2p³ and Oxygen 1s² 2s²2p4 ● ● ● Electronegativity ● Electronegativity: the attraction of an atom for a bonding pair of electrons. it is the tendency an atom will attract a shared pair of electrons when covalently bonded to another atom as atomic radius decreases the electrostatic force between nucleus and valence electron increases SO electronegativity increases down a group it decreases (because of electron shielding) across periods it increases (because nuclear charge increases) ● O electron is removed from a doubly occupied p-orbital which requires less energy to remove than a singly filled p-orbital decrease in ionisation energy from phosphorus to sulfur o Phosphorus 1s²2s²2p63s²3p³ and Sulfur 1s²2s²2p 3s²3p4 O electron is removed from a doubly occupied p-orbital which requires less energy to remove than a singly filled p-orbital ● ● Electron Affinity First electron affinity: energy released when one mol of electrons is added to one mole of gaseous atom (exothermic) X (g) + e → X¯ (g) First electron affinity decreases down a group O ● MPt (k) Melting Points in Period 3 2000- ● depends on structure and attractive forces 1000- ● ● ● Electronegativity increases This added electron has a weaker attraction to the nucleus and therefore releases less energy when added First electron affinity values generally increase across a period metals have low electron affinities and non-metals have higher electron affinities O metals need to lose electrons to be stable whereas non-metals gain electrons for stability. Hence, non-metals release more energy ● tul.….... Si P S Na Mg Al Cl Na, Mg, Al have metallic bonding so their melting point increases as number of electrons increase and their intermolecular forces are high Si is a metalloid and giant macromolecular covalent structure so its melting point is highest P, S, CI are non-metals and simple molecular structures so it has weak intermolecular forces so their melting points are low Ar is a noble gas and monoatomic molecules with weak forces of attraction so lowest melting point Metal oxides and non-metal oxides Na, Mg oxides are basic Al is amphoteric - both basic or acidic ● Si, P, S, CI are acidic Group 1 - Alkali Metals Properties: 1 valence electron ● they are so reactive they are kept under liquid parrafin so it cannot react with oxygen and water vapour • reducing agents because they lose valence electrons (oxidised) ● reactivity increases down the group O as number of shells increase, the attraction between nucleus and valence electron decreases so it is easier to remove an electron metal+water --> metal hydroxide + H₂ (alkali solution) Lithium floats, Sodium floats and fizzes, Potassium ignites a lilac flame ● ● Group 7 - Halogens ● 7 valence electrons (need to gain 1 electron for a full shell) • oxidising agents because they gain an electron (reduced) ● reactivity decreases down the group O as number of shells increase, the attraction between nucleus and valence electron decreases so it is harder to gain an electron displacement reactions happen when a more reactive halogen displaces a less reactive halogen Properties of some you need to know: Chlorine (Cl₂) dense pale-green gas smelly and poisonous - occurs as chloride in the sea - relative atomic mass 35.5 Bromine (Br₂) deep-red liquid with red-brown vapour smelly and poisonous - occurs as bromide in the sea relative atomic mass 80 lodine (1₂) grey solid with purple vapour smelly and poisonous - occurs as iodides and iodates in some rocks and in seaweed relative atomic mass 127 C13.1 Properties of Transition Metals Transition metals: element that has an incomplete d sub-level in one or more of its oxidation states they occupy the d-block Exceptions: O Scandium ion Sc³+ because it has no electrons in its d-orbital O Zinc because it has a full d-orbital in all its ions ● in ionisation, electrons are first lost in the 4s orbital first then 3d ● Properties: ● variable oxidation states ● form complex ions ● form coloured compounds catalysts ● conduct electric/thermal energy • hard/shiny Complex ions ● Complex ions: ligands pair with a central metal ion by forming a co-ordinate covalent bond transitional metals acts as lewis acids as they attract elements with electrons Ligands: neutral or anion compounds with a lone pair of electrons H₂0: #₂0: H₂0 H₂0 •0th₂ •Ott₂ 2+ ● co-ordination number is the number of co-ordinate bonds bonded to the metal ion ● monodentate ligands: form 1 co-ordinate bond (H₂O, NH3, CI, OH, CN-, SCN-) ● bidendate ligands: form 2 co-ordinate bonds (OOC-COO, NH₂CH₂-CH₂NH₂) hexadentate ligands: form 6 co-ordinate bonds (EDTA4-): Magnetism ● transition metals and its complexes that have a lone pair of electrons will be magnetic Ferromagnetism 11111 Paramagnetism 111]10[ ה Diamagnetic 12 12 12 12 12 • [Ar]3d6 ● Iron (II) complex ion magnetism: 5 d-orbitals are split by ligands Explanation ● Permanent magnetism. The lone pair of electrons are aligned parallel to each other Paramagnetism: if ligand is low in the spectrochemical series the split will be small = 4 lone pairs of electrons weak ligand field splitting 1 1 11 1 16 Temporary magnetism. Some electrons are paired and some are lone. The opposite spins of the electrons cancel their charges. The lone electrons creates a small magnetic field. 12 No magnetism. All paired electrons. Diamagnetism: if ligand is high in he spectrochemical series the split will be big = no lone pairs of electrons ● strong ligand field splitting 1 ΔΕΙ 12 ΔΕ, ΔΕΣ ΔΕ Catalytic behaviour Catalyst: increases rate of a reaction without being used up ● Homogeneous catalyst: is in the same physical state as the reactants Catalyst Iron Vanadium (V) oxide Nickel Manganese (IV) oxide Platinum/Palladium Use Haber process: N₂ (g) + 3H₂(g) → 2NH3(g) Contact Process: 2SO₂ (g) + O₂(g) = 2SO3(g) Hydrogenation reactions Hydrogen peroxide decomposition: 2H₂O₂ (aq) → 2H₂O (1) + O₂(g) Catalytic converters in cars C13.2 Colour of complex ions Process: 1. when ligand approaches metal along axes the d-orbital splits into 2 higher energy level and 3 lower energy level sub-levels 2. white light falls on the complex solution 3. 1 electron is promoted to the higher energy level 4. colour corresponding to a wavelength of visible light is absorbed = energy change 5. transmitted light will be the complementary (opposite) colour of the energy absorbed ДЕ IGE 700nm ------ 400nm 14 ● 647nm Exceptions: Red Violet 424nm 1 12 1 Orange Blue 585nm Yelow Green 491nm 575nm Sc3+ has no d-electrons Cu+ and Zn2+ has a complete d-sublevel ● hence, these species will not create a coloured complex (colourless) because no electrons can be promoted from the lower to higher energy sub-level Factors that affect the colour of complex ion: ● oxidation state of the transition metal ● geometry of the complex ion - affects the strength of the repulsion between the ligands and d-orbital electrons nature of transition metal - they have different nuclear charges and attractions ligand bonded Spectrochemical series: I< Br<SCN< Cl<S² <N< F<ONO<OH <SO<NO <C₂07² <0-²<H₂0~NCS <EDTA<NH3~Py<en <bpy~phen <NO₂ <PR3<CH3 <CN ~CO From 1 To H₂0 are weak field ligands From NCS TO CO are strong field ligands complex ions absorb energies corresponding to the electromagnetic spectrum of visible light if ligand changes (eg. from water to ammonia) the wavelength absorbed changes, which changes the colour of the complex ion ● light energy absorbed increases when NH3 is substituted for H2O in a complex ion ● this causes the splitting to increase and wavelength to decrease so transmitted light changes colour ● Factors that affect the split: number of protons - affects levels ● oxidation state - determines number of electrons in d-orbital ligands - affects electron density ● ●