THEORY

Question 1

(a) atomic number, mass number, isotope, Avogadro constant (NA), relative isotopic mass, relative atomic mass (A_r), relative formula mass and relative molecular mass (M_r)

Answer 1

• Atomic number (Z): Number of protons in nucleus. Equal to charge on nucleus
• Mass number (A): Number of protons and neutrons in nucleus. A=Z+N
• Isotope: Atoms of same element with different mass number, have the same number of protons
  
  • differences in mass number caused by different numbers of neutrons
• Avogadro constant: the number of species’ in one mole any substance. Contains 6x10^23 formula units
• Relative isotopic mass: masses of isotopes
• Relative atomic mass A_r: weighted mean of masses of all isotope and their abundances, relative to carbon-12
• Relative formula mass/Relative molecular mass: mass of substances formula units compared to the mass of carbon-12

Question 2

(w) (v) transitions between electronic energy levels in atoms:
flame colours of Li⁺, Na⁺, K⁺, Ca²⁺, Ba²⁺, Cu²⁺
• flame tests for cations

Answer 2

Dip loop of nichrome wire in conc HCl
Then dip loop of nichrome into sample of compound
Hold loop over blue Bunsen flame
Observe colour

• Li⁺ = crimson red
• Na⁺ = orange/yellow
• K⁺ = lilac
• Ca²⁺ = brick red
• Ba²⁺ = apple green
• Cu²⁺ = blue green

Question 3

(v) the electromagnetic spectrum in order of increasing frequency and energy and decreasing wavelength: infrared, visible, ultraviolet

Answer 3

• Visible light (RED = Longer wavelength + Lower energy + frequency than BLUE)

Question 4

(e) conventions for representing the distribution of electrons in atomic orbitals; the shapes of s- and p-orbitals
The ‘electrons in boxes’ model

Answer 4

• Sub-shells:
  
  • s-sub-shell contains one s-orbital
  • p-sub-shell contains three p-orbitals
  • d-sub-shell contains five d-orbitals
  • f-sub-shell contains seven f-orbitals
• Each atomic orbital - can hold a maximum of two electrons of opposite spin (also called spin-pairing)
• S-orbital: spherical shape
• P-orbital: dumbell shape
• Electrons in orbitals can be represented using arrows in boxes:
  
  • box represents the atomic orbital
  • arrow represents single electron
    
    • (in different directions to represent opposite spin)
  • in box model: electrons in a subshell fill the orbitals singly before pairing up
• document on desktop

Question 5

(h) fusion reactions: lighter nuclei join to give heavier nuclei (under conditions of high temperature and pressure); this is how certain elements are formed
Nuclear equations are required.

Answer 5

• Two lighter nuclei fuse together
  
  • Under high pressure and high temperature
  • Conditions needed to overcome repulsion of positive nuclei
• Forms a single heavier nuclei
  
  • How new elements with nuclei heavier than hydrogen are formed
• Releases large quantities of energy (is it exothermic?)

Question 6

(o) the relationship between the position of an element in the s- or p-block of the periodic table and the charge on its ion; the names and formulae of NO₃^–, SO₄^2–, CO₃^2–, OH^–, NH₄⁺, HCO₃^–, Cu²⁺, Zn²⁺, Pb²⁺, Fe²⁺, Fe³⁺; formulae and names for compounds formed between these ions and other given anions and cations

Answer 6

• S-block:
  
  • Group 1 elements form +1 ions
  • Group 2 elements form +2 ions
• Group 6 elements form -2 ions
• Group 7 elements form -1 ions

Anions ending in -ate means there is one or more non-metal bonded to an oxygen

• Salts formed from positively charged cations and negatively charged anions
  
  • so total positive charge must equal total negative charge (overall charge = 0, neutral)
    
    • when naming: cation goes first, then anion

Question 7

(w) (ii) transitions between electronic energy levels in atoms:
the features of these spectra (emission and absorption), similarities and differences
Similarities: both are line spectra; lines in same position for a given element; lines become closer at higher frequencies; series of lines representing transitions to or from a particular energy level. Differences: bright/coloured lines on a black background or black lines on coloured/bright background.

Answer 7

• Similarities: both are line spectra; lines in same position for a given element; lines become closer at higher frequencies; series of lines representing transitions to or from a particular energy level.
• Differences: bright/coloured lines on a black background or black lines on coloured/bright background.

Question 8

(b) (i) the concept of amount of substance (moles) and its use to perform calculations involving: masses of substances, empirical and molecular formulae, percentage composition, percentage yields, water of crystallisation

Answer 8

• Mole: Mole is a unit used to measure the amount of substance

Question 9

(d) balanced full and ionic chemical equations, including state symbols​

Answer 9

• Balanced equations show, relative quantities involved in reaction
  
  • Must be equal numbers of each type of atom on both sides of arrow
  • Can only be balanced by putting numbers in front of formulae
• State symbols:
  
  • gas: (g)
  • liquid: (l)
  • solid: (s)
  • aqueous solution: (aq)

Question 10

(i) chemical bonding in terms of electrostatic forces; simple electron ‘dot-and-cross’ diagrams to describe the electron arrangements in ions and covalent and dative covalent bonds
In covalent bonds there is a balance between the repulsive forces between the nuclei and the attractive forces between the nuclei and the electrons.

Answer 10

• Dot and cross diagram: Diagrams used to represent the way that atoms bond together
  
  • Outer shell of one atom represented by dots
  • Outer shell of another atom represented with crosses
• Covalent bonds - occur between non-metals
  
  • Sharing of electrons
  • Both have full outer shells of electrons (because shared electrons count as part of both atoms outer shell)
  • Both positive nuclei are attracted electrostatically to the shared electrons
  • Repulsion between both positive nuclei
    
    • to maintain covalent bond there is a balance between attraction and repulsion

Question 11

(t) the terms: acid, base, alkali, neutralisation; techniques and procedures for making soluble salts by reacting acids and bases and insoluble salts by precipitation reactions
Knowledge of the names and formulae of the mineral acids, HCl, HNO₃ and H₂SO₄ will be expected.
• making salts (including percentage yield)

Answer 11

• Acid: A compound that dissociates in water to produce hydrogen ions (acids are aqueous, HCl_(g) not an acid)
  
  • pH less than 7
  • Common acids:
    
    • HCl: Hydrochloric acid
    • HNO₃: Nitric acid
    • H₂SO₄: Sulfuric acid
• Base: A compound that reacts with an acid (is a proton acceptor), to produce water (and a salt)
  
  • pH more than 7
• Alkali: A base that dissolves in water to produce hydroxide, OH^- ions
• Making ionic salts: (neutralisation reactions)
  
  • acid + alkali –> salt + water
  • acid + base –> salt + water
  • acid + carbonate –> salt + water + carbon dioxide
  • acid + metal –> salt + hydrogen
• Insoluble salts: (made by precipitation reactions)
  
  • barium/calcium/lead/silver sulfates
  • silver/lead halides
  • all metal carbonates
  • metal hydroxides (except Group 1 hydroxides and ammonium hydroxide)
• Soluble salts:
  
  • lithium, sodium, potassium and ammonium salts
  • nitrates
  • chlorides/bromides and iodides except silver halides as well as copper iodide, lead chloride, lead bromide (all 3 = white precipitates) and lead iodide (yellow precipitate)
  • Sulfates except barium/calcium/lead and silver sulfates (form white precipitates)
  • lithium/sodium/potassium/strontium/calcium/barium and ammonium hydroxides

Question 12

(u) the basic nature of the oxides and hydroxides of Group 2 (Mg–Ba)
Description only, including equations, for reactions of Group 2 oxides and hydroxides with water and acids.

Answer 12

• Elements of group 2 are reactive
• Form compounds containing ions with a +2 charge, because they contain 2 electrons in their outer shell, that are lost when they react
• Reactions of Group 2 metals with oxygen:
  
  • when Group 2 metal burnt in oxygen
  • white solid metal oxide forms
  • (M=general G2 metal)
  • 2M_(s) + O_2(g) –> 2MO_(s)
• Reactions of Group 2 metals with water:
  
  • Forms metal hydroxide and hydrogen
  • M_(s) + 2H₂O_(l) –> M(OH)_2(aq) + H_2(g)
  • Group 2 metals do not react with water as vigorously as Group 1 metals
• Group 2 oxides and hydroxides form alkaline solutions in water
  
  • Metal oxide forms a hydroxide in water:
    
    • MO_(s) + H₂O_(l) –> M²⁺_(aq) + 2OH^-_(aq)
  • if solubility of hydroxide is very low like magnesium hydroxide, then it is sparingly soluble
  • Hydroxyl makes solution alkaline
  • non metal oxides make acidic solutions
• Group 2 oxides and hydroxides react with acids to form salts
  
  • because both are bases, they neutralise dilute acids
  • MO_(s) + 2HCl_(aq) –> MCl_2(aq) + H₂O_(l)
  • M(OH)_2(s) + H₂SO_4(aq) –> MSO_4(aq) + 2H₂O_(l)

Question 13

(w) (iii) transitions between electronic energy levels in atoms:
the relationship between the energy emitted or absorbed and the frequency of the line produced in the spectra, ∆E = hν

Answer 13

• When an electron moves to a higher/lower shell it absorbs/emits electromagnetic radiation with a certain frequency (particle theory of light: light is packets of energy called photons)

Question 14

(w) (iv) transitions between electronic energy levels in atoms:
the relationship between frequency, wavelength and the speed of electromagnetic radiation, c = νλ

Answer 14

• Light (form of electromagnetic radiation)
  
  • behaves like a wave with a wavelength (λ) and frequency (ν)
  • speed of light given on the data sheet (3.00 x 10⁸ m/s - when light travelling in a vacuum)

Question 15

(r) charge density of an ion and its relation to the thermal stability of the Group 2 carbonates
Smaller ions with the same charge have higher charge density and thus distort the large carbonate ion, so that the compound decomposes at lower temperature.

Answer 15

• Group 2 carbonates thermally decompose (break down when heated)
  
  • form group 2 metal oxide and carbon dioxide
  • MCO_3(s) –> MO_(s) + CO_2(g)
  • More thermally stable group 2 carbonates require more heat to break down
  • Change in thermal stability down the group –> due to charge density of group 2 cations
    
    • charge density = charge relative to volume
    • smaller the group 2 cation the higher the charge density
  • cations with a higher charge density distort/polarise the negatively charged cloud around the carbonate anion by drawing the electrons on the carbonate towards itself
  • lower charge density cations (larger cations) cause less distortion (more stable carbonate ion)
  • greater distortion makes the large carbonate ion less stable and easier to break up when heated

Question 16

(s) the solubility of compounds formed between the following cations and anions: Li⁺, Na⁺, K⁺, Ca²⁺, Ba²⁺, Cu²⁺, Fe²⁺, Fe³⁺, Ag⁺, Pb²⁺, Zn²⁺, Al³⁺, NH^{4 +}, CO₃^2–, SO₄^2–, Cl^–, Br^–, I^–, OH^– , NO₃^–; colours of any precipitates formed; use of these ions as tests e.g. Ba²⁺ as a test for SO₄^2–; a sequence of tests leading to the identification of a salt containing the ions above

Answer 16

*

Question 17

(k) use of the electron pair repulsion principle, based on ‘dot-and-cross’ diagrams, to predict, explain and name the shapes of simple molecules (such as BeCl₂, BF₃, CH₄, NH₃, H₂O and SF₆) and ions (such as NH₄⁺) with up to six outer pairs of electrons (any combination of bonding pairs and lone pairs); assigning bond angles to these structures
No treatment of hybridisation or molecular orbitals is expected but ideas of bond angles being altered by the lone pairs present should be included, for example the bond angles of: CH₄ (109.5°), NH₃ (107°), H₂O (104.5°).

Answer 17

• 3D shape of molecule determined by:
  
  • number of regions of high electron density around central atom
    
    • each of the following count as one region of high electron density includes; single covalent bond, double covalent bond, triple covalent bond and lone pair
  • electron pair repulsion theory
• Electrons = negatively charged, so repel each other so that they arranged as far apart as possible
• SHAPES OF MOLECULES DOC on desktop

Question 18

(l) structures of compounds that have a sodium chloride type lattice
Ionic bonding is the overall attraction in a lattice and is made up of attraction between ions of different charge and repulsion between ions of the same charge.

Answer 18

• Ionic compounds = solid at rtp (room temperatue + pressure)
• giant lattices of ions
  
  • lattice = regular structure
• Consists of repeating positive and negative ions in all three dimensions
  
  • In the lattice –> ions with different charges attract each other, and ions with the same charge repel each other
  • arrange themselves to minimise repulsions and maximise attractions
  • because of regular arrangement of ions, ionic compounds form regularly shaped crystals
• Overall atraction in the lattice = ionic bonding

Question 19

(x) use of data from a mass spectrum to determine relative abundance of isotopes and calculate the relative atomic mass of an element.

Answer 19

• Mass spectrometry measures atomic/molecular mass of different particles in a sample + the relative abundances of different isotopes in an element
  
  • in mass spectrometer, sample atoms/molecules ionised to positvely charged cations and separated according to their mass to charge ratios
  • separated ions are detected and their relative abundances
• Relative abundance = height of each peak
• To calculate relative atomic mass of an element from mass spectrum
  
  • Multiply each relative isotopic abundance by its corresponding relative isotopic mass (m/z value)
  • add the totals
  • divide by 100 if percentage abundance used
    
    • if relative abundance not given as a percentage then divide by total relative abundance instead
    • to calculate percentage relative abundance: (relative abundance/total relative abundance) x 100%

Question 20

(f) (i) the electronic configuration, using sub-shells and atomic orbitals, of:
atoms from hydrogen to krypton

Answer 20

• Chromium: 1s²2s²2p⁶3s²3p⁶3d⁵4s¹
• Copper: 1s²2s²2p⁶3s²3p⁶3d¹⁰4s¹