Flashcards · 4CH1 · Topic 1

Principles of chemistry flashcards

90 cards on Topic 1, subtopics 1a to 1i. Pick a subtopic or study them all.

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1a States of matter

Particles in a solid: arrangement, movement and energy: Close together in a regular arrangement. They vibrate about fixed positions. Lowest energy of the three states.
Particles in a liquid: arrangement, movement and energy: Close together in an irregular arrangement. They move around each other. More energy than in a solid.
Particles in a gas: arrangement, movement and energy: Far apart in a random arrangement. They move fast in all directions. Highest energy of the three states.
Name the changes of state: solid → liquid, liquid → gas, gas → liquid, liquid → solid, solid → gas: Melting; boiling (or evaporating); condensing; freezing; sublimation.
Evaporation compared with boiling: Evaporation: only at the surface, below the boiling point. Boiling: throughout the liquid, at the boiling point.
Why does the temperature of a pure substance stay constant while it melts or boils?: The energy supplied is used to overcome the forces between the particles, not to make them move faster.
Diffusion: The random movement of particles from a high concentration to a low concentration until they are evenly spread.
What makes diffusion faster?: A gas rather than a liquid, a higher temperature, and lighter particles. Into a vacuum it is almost instant.
What does repeatedly diluting potassium manganate(VII) solution show?: The colour gets paler but stays visible for many dilutions, so a small amount of solid contains a huge number of tiny particles.
1 cm³ of dye solution is made up to 100 cm³, then this is repeated once. What fraction of the original concentration remains?: Each step dilutes 100 times: 1/100 × 1/100 = 1/10 000.
Saturated solution: A solution in which no more solute can dissolve at that temperature.
Solubility (Paper 2 only): The mass of solute that dissolves in 100 g of solvent at a stated temperature to make a saturated solution. Unit: g per 100 g of solvent.
4.5 g of a solid saturates 15 g of water at 30 °C. Calculate its solubility. (Paper 2 only): 4.5 ÷ 15 × 100 = 30 g per 100 g of water at 30 °C. Divide by the mass of water, not of the solution.
Solubility 64 g/100 g water at 70 °C, 24 g/100 g at 15 °C. Saturated solution in 200 g of water cooled from 70 °C to 15 °C: mass of crystals? (Paper 2 only): (64 − 24) × 200 ÷ 100 = 80 g

1b Elements, compounds and mixtures

Element, compound and mixture: Element: one type of atom. Compound: two or more elements chemically combined, in fixed proportions. Mixture: two or more substances not chemically combined; proportions can vary.
Is chlorine gas, Cl₂, an element or a compound?: An element. It has two atoms per molecule, but both are the same type of atom.
How do melting points show whether a substance is pure?: A pure substance melts at a sharp, fixed temperature (pure water: 0 °C). A mixture melts over a range. Compare with the data-book value.
Choose the separation method: an insoluble solid from a liquid; a soluble solid from its solution; pure solvent from a solution: Filtration; crystallisation; simple distillation.
How do you separate two liquids that mix and have different boiling points?: Fractional distillation. The liquid with the lower boiling point distils first.
Crystallisation: how do you know the solution is saturated?: Dip in a glass rod: when crystals form on it as it cools, the solution is saturated. Then leave to cool, filter, wash with a little cold water and dry.
Chromatography: why is the baseline in pencil, and the solvent below it?: Pencil does not dissolve and run like ink. Solvent below the baseline so the spots do not dissolve into the solvent.
Rf value: Rf = distance moved by substance ÷ distance moved by solvent, both from the baseline. No units; less than 1.
A spot moves 2.1 cm and the solvent front 7.0 cm. Calculate Rf.: 2.1 ÷ 7.0 = 0.30
What does a chromatogram tell you about purity and identity?: One spot: pure. Several spots: a mixture. A spot level with a reference: the same substance (same solvent and paper).

1c Atomic structure

Atom: The smallest particle of an element that keeps the chemical properties of that element.
Molecule: Two or more atoms chemically (covalently) bonded together. It can be an element (O₂) or a compound (H₂O).
Relative mass, relative charge and position of a proton, a neutron and an electron: Proton: 1, +1, nucleus. Neutron: 1, 0, nucleus. Electron: about 1/1840 (negligible), −1, shells around the nucleus.
Atomic number vs mass number: Atomic number: protons only. Mass number: protons + neutrons.
Isotopes: Atoms of the same element with the same atomic number but different mass numbers (different numbers of neutrons).
Relative atomic mass, Ar: The weighted mean mass of the atoms of an element compared with one-twelfth of the mass of an atom of carbon-12.
Mass number vs Ar: Mass number: a whole number for one isotope. Ar: a weighted mean for the element, usually not a whole number.
Protons, neutrons and electrons in ³²₁₆S²⁻: 16 protons, 32 − 16 = 16 neutrons, 16 + 2 = 18 electrons.
Neon-20 and neon-22 occur in the ratio 9 : 1. Calculate Ar.: (20 × 9 + 22 × 1) ÷ (9 + 1) = 202 ÷ 10 = 20.2. Divide by the total of the ratio, not 100.
Atom vs ion of the same isotope: Same numbers of protons and neutrons; only the number of electrons differs.

1d The Periodic Table

How is the Periodic Table arranged, and why does argon come before potassium?: In order of increasing atomic number. Argon (18) comes before potassium (19) even though its Ar is larger.
Group vs period: Group: a column (same number of outer electrons). Period: a row (same number of shells).
Where are the metals and non-metals in the Periodic Table?: Metals: the left-hand side and the centre block. Non-metals: the right-hand side, beyond the stepped line. Hydrogen is a non-metal.
Electronic configurations of potassium (19) and calcium (20): K: 2,8,8,1 Ca: 2,8,8,2
Period 2, Group 5: configuration and element: Two shells, 5 outer electrons: 2,5 — nitrogen.
Two tests to classify an element as a metal or a non-metal: Conductivity of the solid: metals conduct (graphite is the non-metal exception). Oxide: metal oxides are basic; non-metal oxides are generally acidic.
Why do elements in the same group have similar chemical properties?: Their atoms have the same number of outer-shell electrons, and reactions involve the outer-shell electrons.
Why are the noble gases unreactive?: They have a full outer shell, a stable arrangement, so they have no tendency to lose, gain or share electrons. (Helium’s full shell has 2, not 8.)

1e Chemical formulae, equations and calculations

State symbols: (s) solid, including a precipitate; (l) liquid, including water; (g) gas; (aq) dissolved in water.
Balance: C₂H₆ + O₂ → CO₂ + H₂O: 2C₂H₆ + 7O₂ → 4CO₂ + 6H₂O
Relative formula mass, Mr: The sum of the relative atomic masses of all the atoms in the formula. No units.
Mr of magnesium nitrate, Mg(NO₃)₂ (Ar: N = 14, O = 16, Mg = 24): 24 + 2 × (14 + 3 × 16) = 148
Equation linking moles, mass and Mr: moles (mol) = mass (g) ÷ Mr
Three steps of a reacting-mass calculation: 1. Mass → moles of the known substance (÷ Mr). 2. Use the mole ratio from the balanced equation. 3. Moles → mass of the wanted substance (× Mr).
Empirical formula and molecular formula: Empirical: the simplest whole-number ratio of atoms of each element. Molecular: the actual number of atoms of each element in one molecule.
A hydrocarbon has 2.4 g C and 0.6 g H, and Mr 30. Empirical and molecular formulae? (Ar: C = 12, H = 1): C 0.20 mol, H 0.60 mol → 1 : 3, CH₃. 30 ÷ 15 = 2, so C₂H₆.
Percentage yield: (actual yield ÷ theoretical yield) × 100. It cannot be more than 100%.
Finding x in a hydrated salt such as Na₂SO₄·xH₂O: Heat to constant mass. Mass of water = hydrated − anhydrous. x = moles of water ÷ moles of anhydrous salt (use moles, not masses).
Concentration in mol/dm³ (Paper 2 only): c = n ÷ V, with V in dm³ (cm³ ÷ 1000).
Molar volume of a gas at rtp (Paper 2 only): 24 dm³ or 24 000 cm³ per mole of any gas.
Volume of 0.0250 mol of hydrogen at rtp, in cm³ (Paper 2 only): 0.0250 × 24 000 = 600 cm³
Finding the formula of magnesium oxide by burning magnesium in a crucible: two key precautions: Lift the lid now and then to let air in, but keep it on to stop MgO smoke escaping. Heat, cool and weigh to constant mass.

1f Ionic bonding

Ion: A charged particle formed when an atom loses or gains electrons.
How do metals and non-metals form ions?: Metals lose outer electrons → positive ions. Non-metals gain electrons → negative ions. Each ion has a full outer shell; the number of protons does not change.
Ionic bond: A strong electrostatic attraction between oppositely charged ions (not between atoms).
Ion charges from the group number: Group 1: 1+. Group 2: 2+. Group 3: 3+. Group 5: 3−. Group 6: 2−. Group 7: 1−.
Formulae and charges of the hydroxide, nitrate, carbonate, sulfate and ammonium ions: OH⁻, NO₃⁻, CO₃²⁻, SO₄²⁻, NH₄⁺
Formula of calcium nitrate (Ca²⁺, NO₃⁻): Ca(NO₃)₂. The LCM of the charges is 2, so one Ca²⁺ balances two NO₃⁻; the bracket shows two nitrate ions.
Electronic configuration of the oxide ion, O²⁻: 2.8 — oxygen (2.6) gains two electrons.
Why do ionic compounds have high melting and boiling points?: There are many strong electrostatic attractions between oppositely charged ions throughout the giant lattice, which need a lot of energy to overcome.
Why does solid potassium bromide not conduct, but its solution does?: In the solid the ions are fixed in the lattice. In solution (or when molten) the ions are free to move and carry charge. It is the ions that move, not electrons.

1g Covalent bonding

Covalent bond: A shared pair of electrons between two atoms.
Why does a covalent bond hold the atoms together?: There is an electrostatic attraction between the shared pair of electrons and the nuclei of both atoms.
Shared pairs and lone pairs in H₂O, NH₃ and N₂: H₂O: 2 shared, 2 lone. NH₃: 3 shared, 1 lone. N₂: 3 shared (a triple bond), 2 lone.
How many shared pairs of electrons are in a molecule of ethane, C₂H₆?: 7 — one C–C and six C–H.
Why do simple molecular substances have low melting and boiling points?: Only the weak intermolecular forces between molecules are overcome, which needs little energy. The strong covalent bonds within the molecules do not break.
Why is chlorine a gas at room temperature but iodine a solid?: Both are simple molecular. Iodine has a higher Mr (254 vs 71), so its intermolecular forces are stronger and need more energy to overcome.
Why do giant covalent substances have high melting points?: Melting means breaking many strong covalent bonds throughout the network, which needs a lot of energy.
Diamond vs graphite: bonds per carbon, hardness and conductivity: Diamond: 4 bonds, very hard (rigid 3-D network), does not conduct. Graphite: 3 bonds, soft and slippery (weak forces between layers), conducts (one delocalised electron per carbon).
Why is C60 fullerene soft and a poor conductor?: It is simple molecular: only weak intermolecular forces hold the molecules together, and electrons cannot move from molecule to molecule.
Do covalent compounds conduct electricity?: Not usually. Graphite is the exception to learn.

1h Metallic bonding

Metallic bonding (Paper 2 only): The electrostatic attraction between positive metal ions and delocalised electrons.
What must a 2-D diagram of a metal lattice show? (Paper 2 only): Regular, closely packed positive ions (labelled +, or 2+ for magnesium) with delocalised electrons (e⁻) between them. Both labels are needed.
Why do metals conduct electricity, as solids and liquids? (Paper 2 only): The delocalised electrons are free to move through the structure and carry charge.
Why are metals malleable? (Paper 2 only): The layers of positive ions can slide over each other, and the attraction to the delocalised electrons is kept.
Why do metals usually have high melting points? (Paper 2 only): There are many strong attractions between the positive ions and the delocalised electrons to overcome.

1i Electrolysis

Electrolysis (Paper 2 only): The breakdown of an ionic compound, molten or in solution, by a direct current.
Cathode and anode: charge, and what happens there (Paper 2 only): Cathode: negative; cations arrive; reduction. Anode: positive; anions arrive; oxidation.
What carries the charge in the electrolyte, and in the wires? (Paper 2 only): Electrolyte: ions move. Wires: electrons move.
Aqueous solution, inert electrodes: what forms at the cathode? (Paper 2 only): A metal less reactive than hydrogen (Cu, Ag) is deposited. If the metal is more reactive than hydrogen (K, Na, Ca, Mg…), hydrogen is given off.
Aqueous solution, inert electrodes: what forms at the anode? (Paper 2 only): A halide in concentrated solution gives the halogen. Sulfate or nitrate: oxygen is given off.
Products of electrolysing concentrated sodium chloride solution, and what is left (Paper 2 only): Cathode: hydrogen. Anode: chlorine. Sodium hydroxide is left, so the solution becomes alkaline.
Products of electrolysing copper(II) sulfate solution with inert electrodes (Paper 2 only): Cathode: copper (pink-brown deposit). Anode: oxygen. The blue colour fades and sulfuric acid is left.
Half-equation for oxygen forming at the anode (Paper 2 only): 4OH⁻ → O₂ + 2H₂O + 4e⁻ (oxidation)
Half-equations for electrolysing molten lead(II) bromide (Paper 2 only): Cathode: Pb²⁺ + 2e⁻ → Pb Anode: 2Br⁻ → Br₂ + 2e⁻ Heat in a fume cupboard: bromine is toxic.
Dilute sulfuric acid is electrolysed and 50 cm³ of hydrogen is collected. Volume of oxygen? (Paper 2 only): 25 cm³. Hydrogen : oxygen = 2 : 1 (4e⁻ give 2H₂ and 1O₂).