IB Past Paper - Topic 3 Periodic Table M27 HL PDF

Topic 3 – Periodic Table HL S3.1 The periodic table Name………………………………………………………………. S3.1.1 and S3.1.2 The periodic table - Introduction Q1 The periodic table is arranged in order of increasing atomic number. What information about the structure of an atom does the atomic number giv...

Topic 3 – Periodic Table HL S3.1 The periodic table Name………………………………………………………………. S3.1.1 and S3.1.2 The periodic table - Introduction Q1 The periodic table is arranged in order of increasing atomic number. What information about the structure of an atom does the atomic number give you? …………………….. …………………………………………………………………………………………………… Q2 Using coloured pencils, colour and label the s, p, d and f blocks Q3 What is the relationship between the period number and the number of principle (or main) electron energy levels (n)?...................................................................................................... Q4 On the same periodic table, use a pen/pencil to show the distinction between metals, metalloids and non-metals Q5 In which group are the following elements found: a) magnesium: ……. b) carbon: ……. c) iron: …….. Q6 State the name of the following goups a) Group 1: …………….. b) Group 17: …………………. c) Group 18:……………… Q7 What is the name given to the group of elements found between groups 2 and 13? ………………………………………………………………………………………………………. 1 Q8 Complete the following table Element No. of principle No. of valance energy levels electrons Li 3 5 O Be Q9 Write the electronic configuration of the following: a) s-block elements i) Na …………………………. ii) Ca ………………………………. b) p-block elements i) Si …………………………... ii) Se ……………………………….. c) d-block elements ii) V …………………………... ii) Cu ………………………………. 2 S3.1.3 Atomic radii atomar radius The atomic radius of an atom is defined by half the distance between two nuclei of bonded atoms. Q1 Use the data booklet to help you fill in the missing values of the atomic radii in the periodic table below. (All the values are in 10-12 m) 32 H He 130 84 71 64 Li Be B C N O F Ne 140 109 100 Na Mg Al Si P S Cl Ar 200 K Ca Q2 On the periodic table above, label each arrow as “increasing” or “decreasing” to show the general trend in atomic radii across the periods and down the groups. Q3 Explain why a) the atomic radius decreases as you go along a period from left to right...................................................................................................................................................................................................................................................................... b) the atomic radius increases as you go down a group...................................................................................................................................................................................................................................................................... 3 S3.1.3 Ionic radii - cations (positive ions) Q1 a) What is the atomic radius, in pm, of a sodium atom? ……………… b) What is the ionic radius, in pm, of a sodium ion?........................ c) Explain why the ionic radius of a sodium ion is less than the atomic radius of a sodium atom.......................................................................................................................................................................................................................................................................................... + + electron + Q2 a) What is the ionic radius, in pm, of a sodium ion?........................ b) What is the ionic radius, in pm, of a potassium ion?........................ c) Explain why the ionic radius of potassium is greater than the ionic radius + of sodium................................................................................................................................................................................................................................................................................... Q3 a) What is the ionic radius, in pm, of a sodium ion, Na+?........................ b) What is the ionic radius, in pm, of a magnesium ion, Mg2+?........................ c) What is the ionic radius, in pm, of an aluminium ion, Al3+?........................ d) How many occupied electron shells are in each cation? …… e) Explain why crossing a period from sodium to aluminium the ionic radii of the ions decreases................................................................................................................................................................................................................................................................................... 4 S3.1.3 Ionic radii - anions (negative ions) Q1 a) What is the atomic radius, in pm, of a chlorine atom? ……………… b) What is the ionic radius, in pm, of a chloride ion?........................ c) Explain why the atomic radius of a chloride atom is smaller than the ionic radius of a chloride ion..................................................................................................................................................................................................................................................................... Q2 a) What is the ionic radius, in pm, of a chloride ion?........................ b) What is the ionic radius, in pm, of a fluoride ion?........................ c) Explain why the ionic radius of a chloride ion is larger than the ionic radius of a fluoride ion..................................................................................................................................................................................................................................................................... Q3 a) What is the ionic radius, in pm, of a nitride ion, N3-?........................ b) What is the ionic radius, in pm, of an oxide ion, O2-?........................ c) What is the ionic radius, in pm, of a fluoride ion, F-?........................ d) How many occupied electron shells are in each anion? …… e) Explain why, crossing a period from nitride to fluoride, the ionic radii of the ions with decreases.............................................................................................................................................................................................................................................................................................................................................................................................................................. 5 S3.1.3 Ionisation Energy The first ionisation energy of an element is the energy required to remove one mole of electrons from one mole of atoms in the gaseous state. The unit is kJ mol -1. First ionisation energies can be found on p10 of your data booklet. Ionisation energies can be represented by chemical equations: eg Mg (g) à Mg+(g) + e– ΔH =+ 738 kJ mol -1 Q1 Write equations representing the first ionisation energy for a) sodium ΔH =.............................. b) sulphur ΔH =.............................. c) hydrogen ΔH =.............................. Q2 Use the data booklet to help you fill in the missing values of the first ionisation energies in the periodic table on below. (All the values are in kJ mol–1) H He 1312 Li Be B C N O F Ne 900 1086 1681 Na Mg Al Si P S Cl Ar 578 1012 1520 K Ca 590 Q3 On the periodic table above, label each arrow as “increasing” or “decreasing” to show the general trend in first ionisation energies across the periods and down the groups. 6 Q4 a) Explain why the first ionisation energy of the elements generally increases as you go across a period. Use the terms “number of protons”, “nuclear charge”, “occupied electron shells”, “attraction to the nucleus”, “remove an outer electron”...................................................................................................................................................................................................................................................................................................................................................................................................................................................................................................................................................................................................................................................................................................................................................................................................................................................... b) Explain why the first ionisation energy of the elements decreases as you go down a group. Use the terms “occupied electron shells”, “further from the nucleus”, “attraction to the nucleus”, “remove outer electron”...................................................................................................................................................................................................................................................................................................................................................................................................................................................................................................................................................................................................................................................................................................................................................................................................................................................... 7 S3.1.3 Electron affinity Electron affinity is the energy change when an electron is added to an isolated atom in the gaseous state ie X(g) + e– à X-(g) For example, the formation of an oxide ion could be considered as a 2 step process. O(g) + e– à O-(g) ΔH = -141 kJ mol -1 O-(g) + e– à O2-(g) ΔH = +753 kJ mol -1 The first step is exothermic and therefore ΔH is negative. Energy is released as there is an attractive force between the positive nucleus and the added electron. The second step is endothermic and is shown by a positive change in enthalpy. Q1 Suggest a reason why the second step is endothermic ie requires energy to take place ……………………………………………………………………………………………………. ……………………………………………………………………………………………………. ……………………………………………………………………………………………………. ……………………………………………………………………………………………………. 8 S3.1.3 Electronegativity The electronegativity of an atom is defined as the ability of that atom to attract a bonding pair of electrons towards itself in a covalent bond. Large differences in electronegativities between covalently bonded atoms leads to polar covalent bonds (see Topic 4 – Bonding) Q1 Use the data booklet to help you fill in the missing values of the electronegativities in the periodic table on below. H He 2.2 Li Be B C N O F Ne 1.6 2.6 4.0 Na Mg Al Si P S Cl Ar 0.9 1.9 2.2 K Ca 0.8 Q2 On the periodic table above, label each arrow as “increasing” or “decreasing” to show the general trend in electronegativities across the periods and down the groups. Q3 It is difficult to explain the trends in electronegativities, however, there is a general rule. a) State the relationship between the electronegativity of an element and its size:........................................................................................................................... b) Explain this relationship....................................................................................................................................................................................................................................................... 9 S3.1.4 The chemical properties of the Group 1 elements Reactions with water Your teacher may demonstrate the reaction of some alkali metals. Q1 Describe the similarities and differences in the reactions between lithium and sodium with water similarities differences All alkali metals react exothermically with water to form hydrogen gas and hydroxide ions. The general equation is: alkali + water à metal + hydrogen metal hydroxide Q2 Write a balanced formula equations for potassium reacting with water Q3 Write the order of reactivity of the alkali metals with water in the space below starting with the most reactive.............................................................................................................................................. Q4 Use the concept of ionisation energy to explain the order of reactivity of the alkali metals............................................................................................................................................................................................................................................................................................................................................................................................................................................. 10 S3.1.4 The chemical properties of the Group 17 elements The halogens are reactive non-metals. When a reaction takes place between a halogen and a metal halide a displacement reaction can take place. This can be shown using chemical formula equations. The rule for whether a chemical reaction takes place is that a more reactive halogen will displace a less reactive halogen from a metal halide Cl2 (aq) + 2KBr (aq) à 2KCl (aq) + Br2 (aq) Cl2 (aq) + 2NaI (aq) à 2NaCl (aq) + I2 (aq) Br2 (aq) + 2CsI (aq) à 2CsBr (aq) + I2 (aq) I2 (aq) + 2KBr (aq) à no displacement reaction Cl2 (aq) + 2NaF (aq) à no displacement reaction Br2 (aq) + 2CsCl (aq) à no displacement reaction Q1 Considering the reactions above write the order of reactivity of the halogens from most to least reactive. S3.1.4 Trends in metallic and non-metallic character of the elements Q2 Ionisation energy is a good indicator of metallic character. The easier an atom can lose an electron to form a cation the more metallic character it has. What is the trend in metallic character going down group 1. ………………………………………………………………………………………………………… Q3 Electron affinity is a good indicator of non-metallic character. The easier an atom can gain an electron to form an anion the more non-metallic character it has. What is the trend in non-metallic character going down group 17. ………………………………………………………………………………………………………… 11 S3.1.5 Oxides across the 3rd period Q1 Look at the table below showing the oxides of the elements in the third row. Use the information below the table to complete the table. Formula of No argon Na2O MgO Al2O3 SiO2 P4O10 SO2 Cl2O7 oxide oxides Nature of bonding & - structure Nature of acidic - oxide Nature of bonding - Metal oxides tend to consist of ionic lattices. - Silicon dioxide consists of a giant covalent network. - Non-metallic oxides generally consist of covalent molecules. Nature of oxides (basic, amphoteric or acidic) - If a group 1 or group 2 oxides dissolve in water they form a basic solution - Aluminium oxide reacts with both acids and bases. These oxides are called amphoteric. - Non-metal oxides usually form acidic solutions when they dissolve in water. Q2 Write balanced formula equations for: a) sodium oxide (Na2O) reacting with water to form sodium hydroxide (NaOH) b) magnesium oxide (MgO) reacting with water form magnesium hydroxide (Mg(OH)2) c) carbon dioxide (CO2) reacting with water to form carbonic acid (H2CO3) d) sulphur dioxide (SO2) reaction with water to form sulphurous acid (H2SO3) 12 S3.1.5 Implications of the acidic nature of non-metallic oxides Ocean acidification and acid rain are major global environmental concerns and involve the reaction of non-metal oxides with water. Complete the table below, balancing the chemical equations where appropriate. Some of the information has been completed for you Global Source of non-metal oxide and equation Name of acid and formation equation Global Implications Concern Source Formula equation name of acid Formula equation ocean Burning of fossil CH4 + O2 à CO2 + H2O Carbonic acid CO2 + H2O à H2CO3 acidification fuels eg methane Burning of coal containing S + O2 à SO2 Sulphurous acid SO2 + H2O à H2SO3 sulphur acid rain Petrol car N2 + O2 à NO2 Nitric acid NO2 + H2O à H2NO3 engines S3.1.6 Oxidation states (Oxidation numbers) Oxidation states are assigned to an atom to show the number of electrons transferred in forming a chemical bond. It is the charge the atom would have if the compound were composed of ions. Oxidation states always have the + or – symbol first. To assign oxidation states the following rules apply: 1 For simple ionic compounds the oxidation state is equal to the charge on the ion. The charge of the ion is determined by the position of the element in the periodic table. Eg MgO : Mg2+ à +2 AlCl3 : Al3+ à +3 O2- à -2 Cl- à -1 2 Treat covalent compounds as if they were ionic, where the most electronegative element forms a negative ion. Eg H2 O : H à +1 CO2 : C à +4 NH3 : N à -3 O à -2 O à -2 H à +1 3 The sum of all the oxidation states in a compound equals zero. Eg H2O : [2 x (+1)] + [1 x (-2)] = 0 Al2O3 : [2 x (+3)] + [3 x (-2)] = 0 4 The sum of the oxidation states in a polyatomic ion is equal to its charge. Eg NH4+ : [1 x (-3)] + [4 x (+1)] = +1 MnO4- : [1 x (+7)] + [4 x (-2)] = -1 5 All elements have an oxidation state of zero as atoms have no overall charge and the electrons are distributed equally between the atoms of the element. 6 Oxygen when combined almost always has an oxidation state of -2 [exception include peroxides (-1)] and oxygen fluoride compounds (+2)] 7 Hydrogen when combined almost always has an oxidation state of +1 [exceptions include metal hydrides (-1)] Q1 Complete the table below, assigning oxidation states to the metal and non-metal in each case. compound Oxidation state of metal Oxidation state of non-metal NaCl BaO MgCl2 Na2S Atoms of the same element can have different oxidation numbers in different compounds. Q2 What is the oxidation state of sulphur in the following covalent compounds: a) H2 S à b) S à c) SCl2 à d) SO2 à One property of transition metals is that they can have different oxidation states. To show this roman numeral are used to distinguish between the different oxidation states. Q3 What is the oxidation state of the transition metal in the following compounds a) iron (III) chloride à b) copper (I) oxide à c) manganese (IV) oxide à d) titanium (II) chloride à Q4 Deduce the oxidation state of the bolded atom in the following polyatomic ions a) MnO4- à Mn: ………… b) NO3- à N: ………….. c) CO32- à C: ………….. d) Cr2O72- à Cr: …………… Q5 a) What is the oxidation state of oxygen in hydrogen peroxide, H2O2? ………. b) What is the oxidation state of hydrogen in sodium hydride, NaH? ……… 1  HIGHER LEVEL MATERIAL S3.1.7 Discontinuities of the ionisation trends across a period The diagram below (from kognity) shows the first ionisation energies of the first 3 periods of the periodic table. Q1 a) Describe the general trend in first ionisation energies as you go across a period. (review) …………………………………………………………………………………………………… …………………………………………………………………………………………………… b) Explain the general trend in first ionisation energies as you go across a period. (review) ………………………………………………………………………………………………………… ………………………………………………………………………………………………………… ………………………………………………………………………………………………………… ………………………………………………………………………………………………………… Q2 The are two anomalies in this trend in both period 2 and period 3. a) Between which two elements can the first anomaly be seen in period 2? …………………………………………………………………………………….. b) Between which two elements can the second anomaly be seen in period 2? …………………………………………………………………………………….. 2 These anomalies can be explained if the concept of electron sub-levels is considered. Q3 a) State the electron configuration of i) beryllium : ………… ii) boron : ………… b) State the first ionization energy, in kJ mol-1, of i) beryllium : ………… ii) boron : ………… c) Use the electronic configurations of beryllium and boron to explain why boron has a lower first ionisation energy than beryllium. Use the term “remove electron”.............................................................................................................................................................................................................................................................................. Q4 a) Draw the orbital box notation of i) nitrogen ii) oxygen b) State the first ionization energy, in kJ mol-1, of i) nitrogen : ………… ii) oxygen : ………… b) Use the orbital box notation of oxygen and nitrogen to explain why oxygen has a lower first ionisation energy than nitrogen. Use the term “remove electron”.............................................................................................................................................................................................................................................................................. 3 S3.1.8 Transition elements Transition elements can be defined as elements that contain an incomplete d-sublevel of electrons in one or more of their ions. Q1 How many electrons does a filled d-sublevel contain? ……… Q2 a) What is the electronic configuration of Cu2+? …………………………………………… b) Considering the definition of transition element, is copper a transition element? …….. Q3 a) What is the electronic configuration of Zn2+? …………………………………………… b) The only ion of zinc is Zn2+. Why is zinc not considered to be a true transition metal? ………………………………………………………………………………………………… …………………………………………………………………………………………………. Q3 State the electron configuration of the following ions using orbital box notation. ion 3d a) V3+ b) Fe3+ c) Ni3+ As well as the usual properties of all metals, transition elements have special properties: - variable oxidation states - high melting points - magnetic properties - good catalytic properties - form coloured compounds when the metal complexes. 4 S3.1.8 and S3.1.9 Variable oxidation states One property of a transition metal is they can form compounds with various oxidation states. All the first-row transition metals can exist in the +2 oxidation state. Q1 Give 2 examples of transition metal ions in the +2 oxidation state ………….. Some transition metals can exist in the +3 oxidation state Q2 a) What is the charge of the iron ion in FeCl3? ……… b) What is the oxidation state of Fe in FeCl3? ……… c) What is the oxidation state of Cr in CrCl3? ……… (Chromium can also exist in the +6 oxidation state eg in Cr2O7 2-.) Q3 a) What is the oxidation state of Cu in Cu2O ……… b) What is the oxidation state of Mn in MnO2? ……… (Manganese can also exist in the +7 oxidation state eg in MnO4-.) Q4 Consider the following table and graph. a) Why are the first and second ionisation energies of calcium and titanium are very similar? …………………………………………………………………………………… b) Suggest a reason why the third ionisation energy of calcium is much higher than titanium. ………………………………………………………………………………………………. ………………………………………………………………………………………………. ………………………………………………………………………………………………. c) Titanium can exist as 2+, 3+ or 4+ ions as the corresponding successive ionisation energies are relatively close to each other. State the oxidation state of titanium in TiO2 …………………………………………….. 5 S3.1.8 High melting points of the transition elements Below is a bar graph showing the change in melting points of the elements from atomic number 19 to 30. Q1 Shade the bars of the transition elements. Q2 Name the transition element with a) The highest melting point ……………………………. b) The lowest melting point ……………………………. 3.1.8 The magnetic properties of transition metals and their compounds The magnetic properties of the transition elements are due to the presence of unpaired electrons in the d atomic orbitals of the atom. Iron, cobalt and nickel are permanently magnetic, whilst the others are weakly and temporary magnetic. Q3 Use the electronic configuration of zinc to explain why zinc cannot be magnetic. ………………………………………………………………………………………………………… ………………………………………………………………………………………………………… ………………………………………………………………………………………………………… 6 S3.1.8 Catalytic properties of the transition elements A catalyst is a substance, which increases the rate of a reaction without being chemically changed itself. Catalysts also allow reactions to be carried out at lower temperatures. Example 1 – Decomposition of hydrogen peroxide Hydrogen peroxide (H2O2) is a liquid which decomposes to form water and oxygen. Q1 Write a balanced formula equation for the reaction in the space below: Manganese dioxide is an excellent catalyst for this reaction Q2 What is the formula for manganese dioxide? ……….. Example 2 – Nickel in the conversion of alkenes to alkanes Ethene, C2H4, can be converted to ethane, C2H6, by adding an element across the double bond. Q3 Deduce the added element and write a balanced formula equation below: Nickel is the simple catalyst used in this process. 7 Example 3 – Iron and cobalt in biological systems. Heme is found in red blood cells and enables blood to carry around oxygen. Q4 Circle the iron in the heme complex on the right. Vitamin B12 is essential to produce red blood cells. Q5 Circle the cobalt in the Vit B12 complex on the right. Example 4 – Palladium and Platinum in catalytic converter These metals convert the polluting gases of car exhaust into less polluting gases: Q6 Complete the table below: After passing through catalytic Pollutant from exhaust converter unburned hydrocarbons carbon monoxide nitrogen oxides 8 Example 5 – The Contact process The Contact process is a method of reacting sulphur dioxide with oxygen to form sulphur trioxide. (which is then further used to make sulphuric acid) Q7 Write a balanced formula equation for the reaction in the space below: Vanadium (V) oxide is the catalyst used in the Contact process. Q8 What is the formula for vanadium (V) oxide? ……….. Example 6 – Iron in the Haber process The Haber process is a method of converting elemental hydrogen and nitrogen into ammonia (NH3). Q9 Write a balanced formula equation, including state symbols, for the reaction in the space below: Iron is the simple catalyst used in the Haber Process. Q10 What is the economic significance of the catalysts used in the Contact and the Haber process? ………………………………………………………………………………………………………… 9 S3.1.8 Complex ion formation Transition metals ions are relatively small and can be surrounded by small molecules and/or small negative ions. These small molecules and ions are known as ligands. The thing all ligands have in common is that they have lone pairs of non-bonding electrons. Common ligands include: water molecule (H2O) cyanide ion (CN-) chloride ion (Cl-) ammonia molecule (NH3) carbon monoxide (CO) Look at the outer electron diagram of these 4 ligands. Q1 Circle the lone pair of electrons. O x x x C x N xx H H x H Cl x N x x H H The non-bonding pairs of electrons on the ligands can then form bonds with the transition metals to form a complex ion. These bonds are called co-ordination bonds as both electrons come from the same atom. 10  [Pt(Cl)4]2- [Fe(CN)6]3- [Cu(NH3)4]2+ [Ag(NH3)2]+ The complex contains the transition metal and the ligands. This is shown by squared brackets in the formula. The co-ordination number of the complex is the number of co-ordinate bonds there from ligand to metal ion. The shape of the complex can be determined by the co- ordination number. Q1 Complete the table below. An example has been given which is not shown above. Coordination Oxidation state Complex Ligand Shape Number of metal [Ni(CO)4] CO 4 0 tetrahedral [Fe(CN)6]3- octahedral [Cu(NH3)4]2+ tetrahedral 2- [Pt(Cl)4] square planar [Ag(NH3)2]+ linear Q2 Ligands can be replaced by other ligands. For example, + H2 O + NH3 2- 2+ [Cu(Cl)4] à [Cu(H2O)4] à [Cu(NH3)4]2+ A B C Give the name and symbol for the ligand found in complex A? …………………………. 11 Some ligands contain more than one lone pair available to co-ordinate with the transition metal ion. Two examples, the oxalate ion and ethylenediamine, are shown below. You do not need to memorise these structures. Ethylenediamine Oxalate ion C₂H₄(NH₂)₂ C2O42- The lone pairs which bond with the metal ion are shown as dots on the molecules above. Since the molecules have two lone pairs available, they are known as bidentate ligands. Q3 The diagram below shows the structure of the tris(ethylenediamine) cobalt (III) ion, which contains the ethylenediamine ligand and is can be described as a polydentate ligand a) How many ethylenediamine molecules are present in the tris(ethylenediamine) cobalt (III) ion? … b) What is the co-ordination number of this complex? ……. c) What is the oxidation state of the cobalt ion? …… d) What shape is the complex? ………………………………. e) If all the ethylenediamine ligands are replaced with the oxalate ion ligand, and the cobalt kept the same oxidation state, what would the charge of the overall complex become? …………….. 12 Some molecules can act as hexadentate ligands. The most common hexadentate ion is EDTA4- (ethylenediaminetetraacetic acid), which is shown below. When EDTA4- binds with a transition metal it surrounds the metal using the six lone pairs which are available to form co-ordinate bonds. This produces complexes such as the copper (II) EDTA complex Q4 a) What is the co-ordination number of copper (II) EDTA complex above? ……… b) What is the shape of the complex? …………………………………. c) What is the overall charge of the complex? ……………… 13 S3.1.10 Coloured d-block complexes The colour of the d-block complexes can be related to the presence of partially filled d orbitals on the transition metal ion. The colour of some transition metal ions are shown below: Q1 Suggest why Sc3+ and Zn2+ are colourless. ……………………………………………………………………………………………………. S3.1.10 The visible spectrum The table below shows the various wavelengths of the visible colours of the electromagnetic spectrum. Q2 Which colour has the highest energy? ………………………. 14 S3.1.10 Complementary colours The colour of a substance is determined by the colour of light it absorbs. For example, copper sulphate appears blue because it absorbs orange light. Orange and blue are said to be complementary colours. Complementary pairs of colours can be deduced from the colour wheel, which can be found in section 17 of the Data booklet. Q1 Using the colour wheel below, complete the following table. colour of wavelength Complex colour of complex absorbed [Co(H2O)6]3+ red [Fe(CN)6]3- blue [Cu(Cl)4]2- violet This concept can be seen in a diagram below: Q2 a) Which colour is absorbed by the iron complex? ……………………………… b) Which colour is the iron complex? …………………………………………….. 15 S3.1.10 Splitting of d-orbitals When a ligand donates a pair of non-bonding electrons to a transition metal, the donated electrons interact with the d-orbitals of the transition metal. This makes the d orbitals split, and the orbitals no longer have the same energy. Eg Titanium (II) ion ↑ ↑ ↑ ↑ ↑ ↑ Free ion 5 d atomic orbitals – all the same energy Complex- ground state Complex- excited state 5 d atomic orbitals – An electron in the lower 2 orbitals are higher in d subshell absorbs energy than the other 3 energy. When white light is shone on the complex, an electron in the lower orbitals can be excited to the higher d orbitals. To do this, the electron must absorb energy. A photon can provide this energy. Thus, the complex absorbs light of a specific wavelength and reflects the complementary colour. The energy difference, and thus the colour of light, between the ground and excited states is determined by how the lone pairs of the ligand and the d atomic orbitals intact. Consider the following table: complex [Mn(H2O)6]3+ [Fe(H2O)6]3+ [Fe(H2O)6]2+ [Cu(H2O)4]2+ [Cu(NH3)4]2+ [Cu(Cl)4]2- colour of pale pink yellow light green light blue deep blue yellow/green complex Q1 From the information given, what three factors seem to affect the colour of a complex? 1. …………………………………… 2. …………………………………… 3. …………………………………… 16 Q2 Consider the following d-orbital energy level diagrams, A and B. Compounds which contain the [Fe(H2O)6]2+ complex are often yellow in colour, whereas compounds containing the [Fe(CN)6]4- complex are more often orange. a) Using the colour wheel in section 15 of the data booklet, deduce which orbital diagram, A or B is most likely to represent [Fe(H2O)6]2+ : ……………….. b) Explain your answer to part a) ………………………………………………………………………………………………………. ………………………………………………………………………………………………………. ………………………………………………………………………………………………………. Q3 A transition metal complex appears purple and absorbed yellow light with a wavelength of approximately 𝜆 = 580 nm. Calculate the following to 3 significant figures a) The frequency of light absorbed: ………………………………………………………………………………………………………. ………………………………………………………………………………………………………. ………………………………………………………………………………………………………. b) The energy difference between the two spit d-orbitals ………………………………………………………………………………………………………. ………………………………………………………………………………………………………. ………………………………………………………………………………………………………. 17 Answers p1 - 2 Q1 Atomic number tells you the number of protons in an atom Q2 ----------------- Q3 the period number is equal to the number of main/principal energy levels Q4 ------------------ Q5 a) 2 b) 14 c) 8 Q6 a) alkali metals b) halogens f) noble gases Q7 Transition metals Q8 Element No. of principle No. of valance energy levels electrons Li 2 1 P 3 5 O 2 6 Be 2 2 Q9 a)i) 1s2 2s2 2p6 3s1 ii) 1s2 2s2 2p6 3s2 3p6 4s2 b)i) 1s2 2s2 2p6 3s2 3p2 ii) 1s2 2s2 2p6 3s2 3p6 4s2 3d10 4p4 c)i) 1s2 2s2 2p6 3s2 3p6 4s2 3d3 ii) 1s2 2s2 2p6 3s2 3p6 4s1 3d10 p3 Q1 self check Q2 atomic radius decreases along period and increases along group Q3 a) As you go across the period left to right, the number of protons in the nucleus increase and the nucleus is becoming more positive. The electrons are therefore more strongly attracted to the nucleus and the atomic radius decreases. b) As you go down a group the number of occupied electron levels increases, therefore the atomic radius increases. 18 p4 Q1 a) 160 pm b) 102 pm c) The sodium ion has fewer occupied energy levels Q2 a) 102 pm b) 138 pm c) Potassium ion has one more occupied electron level Q3 a) 102 pm b) 72 pm c) 54 pm d) 10 electrons e) Although all the ions have the same number of occupied electron levels, the number of protons in the nucleus from sodium to aluminium is increasing and so the nucleus is becoming more positive. The outer electrons are therefore more strongly attracted to the more positive nucleus and the ionic radius is smaller ‘ p5 Q1 a) 100 pm b) 181 pm c) In the chloride ion the number of electrons in the valence shell has increased. These electrons will repel each other increasing the size of the ion. Q2 a) 181 pm b) 133 pm c) The chloride ion has one more occupied electron energy level Q3 a) 146 pm b) 140 pm c) 133 pm d) 10 electrons e) Going across the period, the Ions have the same number of occupied electron levels but an increasing number of protons. The outer electrons are therefore more strongly attracted to the more positive nucleus and the ionic radius is smaller p6-7 Q1 a) Na (g) à Na+(g) + e– ΔH = 496kJ mol -1 b) S (g) à S+(g) + e– ΔH = 1000 kJ mol -1 c) H (g) à H+(g) + e– ΔH = 1312 kJ mol -1 Q2 self check Q3 Ionisation energies increase going across a period and decrease going down a group 19 Q4 a) As you go across a period from left to right the number of protons increases and the nucleus charge increases. However, the number of occupied electron shells remain the same across the period. This means the outer electrons have a stronger attraction to the nucleus, meaning more energy is required to remove an outer electron. b) As you go down a group the number of occupied electron shells increases. This means the outer electrons are further from the nucleus and there is a weaker attraction to the nucleus, meaning less energy is required to remove the outer electron. p8 Q1 The second step requires energy because you are adding a negative electron to a negative ion. This means you have to overcome the forces of repulsion between the negative ion and the electron, which requires energy. p9 Q1 self check Q2 Electronegativities generally increase going across a period left to right and decrease going down a group. Q3 a) The larger the atom the less electronegative it is. b) The larger the atom,the further the positive nucleus is from the bonding electrons, therefore the lower the electronegativity. p10 Q1 Similarities: the metal floats/moves on the surface; fizzing/effervescence/bubbles; (accept sound is produced) solution gets hot; solution becomes alkaline/basic (seen if an indicator is used) Differences: Sodium - metal melts, faster reaction Lithium - no melting, slower reaction Q2 2K + 2H2O à 2KOH + H2 Q3 Potassium > sodium > lithium Q4 As you go down the group the ionisation energy of the element decrease and less energy is required to remove an eletron. This means the futher down group 1 you go the easier the an ion is formed and the faster the reaction 20 p11 Q1 F2 > Cl2 > Br2 > I2 Q2 As you go down group 1 the metallic character of the element increases. Q3 As you go down group 17 the non-metallic character of the element decreases. p12 Q1 No Formula Na2O MgO Al2O3 SiO2 P4O10 SO3 Cl2O7 argon of oxide oxides Nature of bonding Ionic Ionic Ionic covalent covalent covalent covalent - & lattice lattice lattice network molecule molecule molecule structure Nature of alkaline alkaline amphoteric acidic acidic acidic acidic - oxide Q2 a) Na2O (s) + H2O (l) à 2NaOH (aq) b) MgO (s) + H2O (l) à Mg(OH)2 (aq) c) CO2 (g) + H2O (l) à H2CO3 (aq) d) SO2 (g) + H2O (l) à H2SO3 (aq) p13 21 p15 Q1 Q2 a) -2 b) 0 c) +2 d) +4 Q3 a) +3 b) +1 c) +4 d) +2 Q4 a) +7 b) +5 c) +4 d) +6 Q5 a) -1 b) -1 p16 Q1 a) As you go across a period the ionisation energy increases. b) As you go across a period from left to right the number of protons increases and the nucleus charge increases. However, the number of occupied electron shells remain the same across the period. This means the outer electrons have a stronger attraction to the nucleus, meaning more energy is required to remove an outer electron. Q2 a) Between beryllium and boron b) Between nitrogen and oxygen Q3 a) i) Be: 1s2 2s2 ii) B: 1s2 2s2 2p1 b) i) 900 ii) 801 c) The removed electron from a Be atom comes from the 2s subshell, whilst the removed electron from B comes from the 2p subshell. Since the 2s subshell is closer to the needs it will require more energy to remove the first electron from a boron atom. Q4 a) i) ii) b) i) 1402 ii) 1314 c) The removed electron from an oxygen atom comes from the 2p orbital which contains a pair of electrons. Since these two electrons repel each other, it will take less energy to remove the electron 22 p18 Q1 10 Q2 a) 1s2 2s2 2p6 3s2 3p63d9 b) Yes Q3 a) 1s2 2s2 2p6 3s2 3p63d10 b) No – as zinc (II) ion has a complete d-sublevel. Q2 ion 3d a) V3+ ↑ ↑ b) Fe3+ ↑ ↑ ↑ ↑ ↑ c) Ni3+ ↑↓ ↑↓ ↑ ↑ ↑ p19 Q1 any two appropriately 2+ charged cations Q2 a) 3+ b) +3 c) +3 Q3 a) +1 b) +4 Q4 a) The first and second ionisation energies of calcium and titanium are very similar as, in both cases, the electrons are coming from the 4s energy level. b) The 3rd electron removed in calcium is coming from the 3p sub-shell, whereas the 3rd electron removed from titanium is coming from the 3d sub-shell. The 3p subshell is closer to the nucleus, therefore the energy required to remove an electron is greater. c) +4 (not 4+) p20 Q1 from atomic number 21 – 29 (not 30) Q2 a) Vanadium b) Copper Q3 Zn has an electronic configuration of [Ar] 4s23d10. This means that zinc has no unpaired d- electrons and is thus not magnetic. 23 p21-23 Q1 2H2O2(l) à 2H2O(l) + O2(g) Q2 MnO2 (s) Q3 C2H4(g) + H2(g) à C2H6(g) Q4 In the middle Q5 In the middle Q6 unburned hydrocarbons à carbon dioxide and water carbon monoxide à carbon dioxide nitrogen oxides à nitrogen and oxygen Q7 2SO2(g) + O2(g) à 2SO3(g) Q8 V2O5 Q9 3H2(g) + N2(g) à 2NH3(g) Q10 They speed up the chemical reaction, therefore increasing profit. p25-27 Q1 Coordination Oxidation state Complex Ligand Shape number of metal [Ni(CO)4]3- CO 4 0 tetrahedral 3- - [Fe(CN)6] CN 6 +3 octahedral [Cu(NH3)4]2+ NH3 4 +2 tetrahedral [Cu(Cl)4]2- Cl- 4 +2 square planar [Ag(NH3)2]+ NH3 2 +1 linear Q2 chloride ion; Cl- Q3 a) 3 molecules b) 6 c) +3 d) octahedral e) 3- Q4 a) 6 b) octahedral c) 2- p28 Q1 Neither Sc3+ nor Zn2+ have partially filled d orbitals. Q2 Violet 24 p29 Q1 colour of wavelength Complex colour of complex absorbed [Co(H2O)6]3+ red green [Fe(CN)6]3 orange blue [Cu(Cl)4]2- violet yellow Q2 a) violet b) orange p30-31 Q1 1. The type of metal ion 2. The charge on the metal 3. The type of ligand Q2 a) B b) [Fe(H2O)6]2+ is yellow à complex is absorbing violet light, which has a lower wavelength and thus higher energy. Complex A contains a larger splitting, so is more likely to be yellow in colour. Q3 a) 5.17 x 1014 Hz b) 3.43 x 10-19 J 25