Unit 1: Inorganic 1 and Physical 1 Flashcards

The sample is vaporised, and high-energy electrons are fired at it from an electron gun. This knocks off one outer electron per particle to form 1⁺ gaseous ions: X(g) + e⁻ → X⁺(g) + 2e⁻.

\(t = d \sqrt{\frac{m}{2KE}}\), where \(t\) is flight time in seconds, \(d\) is tube length in metres, \(m\) is mass of a single ion in kg (\(m = \frac{A_r \times 10^{-3}}{L}\)), and \(KE\) is kinetic energy in Joules.

\(m = \frac{\text{Isotopic mass in g mol}^{-1} \times 10^{-3}\text{ kg g}^{-1}}{\text{Avogadro constant } (6.022 \times 10^{23}\text{ mol}^{-1})}\).

Cr (Z = 24): [Ar] 3d⁵ 4s¹ (half-filled 3d sub-shell stability). Cu (Z = 29): [Ar] 3d¹⁰ 4s¹ (fully filled 3d sub-shell stability).

PV = nRT. P = Pressure in Pascals (Pa); V = Volume in cubic metres (m³); n = Amount in moles (mol); R = Gas constant (8.31 J K⁻¹ mol⁻¹); T = Temperature in Kelvin (K = deg C + 273).

Empirical formula is the simplest whole-number ratio of atoms of each element in a compound. Molecular formula gives the actual number of atoms of each element in a molecule.

A shared pair of electrons in which both electrons are donated by one of the bonding atoms (the donor atom has a lone pair; the acceptor atom has an empty orbital).

1. Electron pairs around a central atom repel each other. 2. Electron pairs adopt a geometry that minimises repulsion and maximises separation. 3. Lone pair - lone pair repulsion > lone pair - bonding pair repulsion > bonding pair - bonding pair repulsion.

Shape: Trigonal planar. Bond angle: 120 degrees. Explanation: Boron has 3 bonding pairs and 0 lone pairs; the three electron pairs experience equal repulsion and move as far apart as possible.

CF₄ has a symmetrical tetrahedral shape. The four equal C-F bond dipoles point in opposing directions and cancel each other out completely, resulting in no overall molecular dipole.

Both molecules are non-polar and experience only London dispersion (Van der Waals) forces. CF₄ has 42 electrons compared to 10 in CH₄, giving CF₄ stronger temporary induced dipole-dipole attractions requiring more energy to overcome.

HF molecules form strong intermolecular hydrogen bonds due to the high electronegativity of fluorine and lone pairs on F. Hydrogen bonds are significantly stronger than Van der Waals forces.

The enthalpy change when one mole of a substance is burned completely in oxygen under standard conditions (100 kPa, 298 K) with all reactants and products in their standard states.

The enthalpy change when one mole of a compound is formed from its constituent elements in their standard states under standard conditions (100 kPa, 298 K).

q = m * c * ΔT. q = heat energy absorbed/released (J); m = mass of heated substance (g); c = specific heat capacity (4.18 J g⁻¹ K⁻¹ for water); ΔT = temperature change (K or deg C).

1. Significant heat loss to surroundings and the calorimeter. 2. Incomplete combustion of the fuel producing CO or carbon soot instead of CO₂.

The enthalpy change for a chemical reaction is independent of the route taken, provided the initial and final states are the same.

Oxidation is the loss of electrons. Reduction is the gain of electrons (OIL RIG).

MnO₄⁻ + 8H⁺ + 5e⁻ → Mn²⁺ + 4H₂O.

MnO₄⁻ + 8H⁺ + 5Fe²⁺ → Mn²⁺ + 5Fe³⁺ + 4H₂O.

General increase. Nuclear charge increases (more protons) while outer electrons remain in the same shell with similar shielding, increasing electrostatic attraction and pulling outer electrons closer.

Phosphorus has configuration 3s² 3p³ (half-filled 3p sub-shell, one electron per orbital). Sulfur has 3s² 3p⁴ with two paired electrons in one 3p orbital. Electron-electron repulsion in the paired orbital makes removing the electron easier.

Solubility of Group 2 hydroxides INCREASES down the group (Mg(OH)₂ is sparingly soluble; Ba(OH)₂ is soluble). Solubility of Group 2 sulfates DECREASES down the group (MgSO₄ is soluble; BaSO₄ is insoluble).

Add dilute HCl followed by aqueous BaCl₂; a white precipitate of BaSO₄ forms (Ba²⁺(aq) + SO₄²⁻(aq) → BaSO₄(s)). Dilute HCl removes carbonate (CO₃²⁻) or sulfite (SO₃²⁻) ions which would also form white precipitates with barium.

Barium sulfate is extremely insoluble in water and stomach acid. It cannot be absorbed into the bloodstream or tissues, passing through the digestive tract unchanged.

n = (c * V) / 1000, where c is concentration in mol dm⁻³ and V is volume in cm³.

Titres that agree within 0.10 cm³ of each other. Only concordant titres are used to calculate the mean titre.

Dissolve solid with washings into flask, add deionised water until bottom of meniscus aligns with calibration line at eye level, invert 10-15 times to ensure uniform mixing.

A technique where an excess of known reagent is reacted with a sample, and remaining unreacted reagent is titrated to find the amount reacted. Used for insoluble or slow-reacting solids.

% Uncertainty = [(Uncertainty x 2) / Titre volume delivered] x 100. The factor of 2 accounts for both initial and final burette readings.

Ba²⁺(aq) + SO₄²⁻(aq) → BaSO₄(s).

The weighted average mass of an atom of an element relative to 1/12th of the mass of an atom of carbon-12.

The average mass of a molecule of a compound relative to 1/12th of the mass of an atom of carbon-12.

The sample is vaporised, and high-energy electrons are fired from an electron gun. This knocks off one outer electron: X(g) + e⁻ → X⁺(g) + 2e⁻.

\(t = d \sqrt{\frac{m}{2KE}}\), where \(t\) is flight time in seconds, \(d\) is tube length in metres, \(m\) is single ion mass in kg, and \(KE\) is kinetic energy in Joules.

\(m = \frac{\text{Isotopic mass in g mol}^{-1} \times 10^{-3}}{\text{Avogadro constant } (6.022 \times 10^{23}\text{ mol}^{-1})}\).

1s² 2s² 2p⁶ 3s² 3p⁶.

Phosphorus has 3p³ (half-filled sub-shell, unpaired electrons). Sulfur has 3p⁴ with two paired electrons in one 3p orbital. Spin-pair repulsion makes the electron easier to remove.

A shared pair of electrons in which both electrons are donated by one of the bonding atoms.

Shape: Trigonal planar. Bond angle: 120 degrees. Boron has 3 bonding pairs and 0 lone pairs. The 3 electron pairs repel equally to achieve maximum separation.

CF₄ has a symmetrical tetrahedral shape. The four equal C-F dipoles point in opposing directions and cancel each other out completely, leaving no net molecular dipole.

Both molecules are non-polar and have London dispersion forces. CF₄ has 42 electrons compared to 10 in CH₄, resulting in stronger temporary induced dipoles that require more energy to overcome.

HF forms intermolecular hydrogen bonds due to high electronegativity of fluorine and lone pairs on F. Hydrogen bonds are significantly stronger than London dispersion forces.

Lone pair - lone pair > lone pair - bonding pair > bonding pair - bonding pair. Each lone pair reduces bond angles by approximately 2.5 degrees.

Solubility of Group 2 hydroxides INCREASES down the group (Mg(OH)₂ is sparingly soluble; Ba(OH)₂ is soluble).

Solubility of Group 2 sulfates DECREASES down the group (MgSO₄ is soluble; BaSO₄ is insoluble).

Add dilute HCl (to remove carbonate or sulfite ions), then add aqueous BaCl₂. A white precipitate of BaSO₄ forms: Ba²⁺(aq) + SO₄²⁻(aq) → BaSO₄(s).

Barium sulfate is extremely insoluble in water and stomach acid. It cannot be absorbed into the bloodstream or body tissues, passing through the digestive tract harmlessly.

TiCl₄ + 2Mg → Ti + 2MgCl₂ (under an inert argon atmosphere at 1000 deg C).

They act as bases to neutralise acidic sulfur dioxide gas produced from burning coal: CaO(s) + SO₂(g) → CaSO₃(s).

Boiling point increases down the group. Molecules have more electrons, leading to stronger London dispersion forces that require more thermal energy to overcome.

Reducing ability INCREASES down the group: F⁻ and Cl⁻ do not reduce H₂SO₄; Br⁻ reduces H₂SO₄ to SO₂; I⁻ reduces H₂SO₄ to SO₂, S, and H₂S.

AgCl (white): dissolves in dilute NH₃. AgBr (cream): insoluble in dilute NH₃, dissolves in concentrated NH₃. AgI (yellow): insoluble in both dilute and concentrated NH₃.

A redox reaction in which the same element is simultaneously oxidised and reduced in the same reaction.

Cl₂ + 2NaOH → NaCl + NaClO + H₂O. Chlorine is reduced to -1 in NaCl and oxidised to +1 in NaClO.

Chlorine reacts with water to produce chloric(I) acid (HClO), which kills pathogenic bacteria and prevents cholera and typhoid epidemics. The health benefits far outweigh the toxic risk.

General increase from Na to Ar. Nuclear charge increases while shielding remains similar because electrons enter the same third shell, increasing attraction to the outer electrons.

In Al, the outer electron is removed from a 3p orbital which is higher in energy and shielded by the 3s² electrons, making it easier to remove than the 3s electron in Mg.

In S, the fourth 3p electron is paired in an orbital. Spin-pair repulsion between the two paired electrons in the same 3p orbital makes that electron easier to remove.

Atomic radius decreases from Na to Cl. Nuclear charge increases while shielding is similar, pulling outer valence electrons closer to the nucleus.

Silicon has a giant covalent macromolecular structure with a 3D network of strong covalent bonds requiring massive thermal energy to break.

Sulfur exists as S₈ molecules with 128 electrons, while phosphorus exists as P₄ with 60 electrons. S₈ has stronger London dispersion forces requiring more thermal energy to overcome.

PV = nRT. P = Pressure (Pa); V = Volume (m³); n = Amount (mol); R = Gas constant (8.31 J K⁻¹ mol⁻¹); T = Temperature (K = deg C + 273).

Atom Economy (%) = [(stoichiometric coefficient x molecular mass of desired product) / Sum of (stoichiometric coefficient x molecular mass of all reactants)] x 100.

Empirical formula is the simplest whole-number ratio of atoms of each element in a compound. Molecular formula is the actual number of atoms of each element in a molecule.

Percentage Yield (%) = (Actual mass of product formed / Theoretical maximum mass of product) x 100.

Zero intermediate rounding. Keep all unrounded numbers in your calculator registers and round only at the final reported answer (usually to 3 significant figures).

Divide by 1,000,000 (or multiply by 10⁻⁶).

Oxidation is the process of electron loss. Reduction is the process of electron gain (OIL RIG).

An oxidising agent is an electron acceptor, so it is itself reduced. A reducing agent is an electron donor, so it is itself oxidised.

Cl₂. It accepts electrons and is reduced to Cl⁻. Br⁻ is the reducing agent.

In a neutral compound they add up to 0. In an ion they add up to the charge on the ion. An uncombined element is 0.

H in NaH is -1 (metal hydride). O in H₂O₂ is -1 (peroxide).

x + 4(-2) = -1, so Mn is +7.

2x + 7(-2) = -2, so each Cr is +6.

1. Balance the element being oxidised or reduced. 2. Balance O by adding H₂O. 3. Balance H by adding H⁺. 4. Balance charge by adding electrons.

MnO₄⁻ + 8H⁺ + 5e⁻ → Mn²⁺ + 4H₂O

Multiply so the electrons lost equal the electrons gained. Add the half-equations. Cancel the electrons and any H⁺ or H₂O on both sides. Check atoms and charge.

Ar = Σ(isotopic mass × abundance) ÷ Σ(abundance). For 1+ ions, the m/z of each peak is the isotopic mass and the peak height is the abundance.

(20 × 90.9 + 21 × 0.3 + 22 × 8.8) / 100 = 2017.9 / 100 = 20.2

The molecular ion peak is the peak with the highest m/z (ignore any tiny M+1 peak). Its m/z value gives Mr.

Na⁺(g) → Na²⁺(g) + e⁻

Each electron is removed from an increasingly positive ion, so it is more strongly attracted to the nucleus.

The electron is now removed from a new shell, closer to the nucleus. The number of electrons removed before the jump gives the group. Aluminium jumps after the 3rd, so it is in Group 3 (13).

The outer electron is further from the nucleus and more shielded, so it is less strongly attracted, despite the greater nuclear charge. This is evidence for shells.

Dalton (solid spheres), Thomson (plum pudding), Rutherford (nucleus), Bohr (electrons in shells).

0.500 × 6.02 × 10²³ × 3 = 9.03 × 10²³ atoms

Al₂(SO₄)₃ (two Al³⁺ balance three SO₄²⁻) and (NH₄)₂CO₃ (two NH₄⁺ balance one CO₃²⁻).

As an arrow from the atom that donates the electron pair. NH₄⁺: one arrow from N to H. Al₂Cl₆: two arrows from Cl to Al.

Mg²⁺ has a greater charge and is smaller, and two electrons per atom are delocalised (one for Na). So the attraction between the ions and the delocalised electrons is stronger.

A giant ionic lattice: each Na⁺ is surrounded by 6 Cl⁻ and each Cl⁻ by 6 Na⁺. It conducts only when molten or dissolved, because the ions are then free to move.

In graphite each C bonds to 3 others, leaving delocalised electrons between the layers. In diamond each C bonds to 4 others, so there are no delocalised electrons.

It is a molecular lattice of I₂ molecules. Only weak van der Waals forces between the molecules are overcome; the covalent I-I bonds do not break.

Each H₂O hydrogen-bonds to 4 others in an open tetrahedral lattice. The molecules are held further apart than in liquid water.

The energy taken in is used to overcome the forces between the particles, not to raise the temperature. Freezing and condensing release energy.

The enthalpy change to break one mole of a given type of bond in the gas phase, averaged over a range of compounds.

Mean bond enthalpies are averages over a range of compounds, not the actual values in these molecules. (The data-book enthalpy of combustion, -890 kJ mol⁻¹, differs even more because there the water is a liquid.)

Melting point decreases. The ions get larger, so the delocalised electrons are further from the nucleus and the metallic bonding is weaker. (Mg is lower than the trend predicts: different crystal structure.)

Add dilute HCl, then BaCl₂(aq): a white precipitate of BaSO₄. The acid removes carbonate (and sulfite) ions that would also give a white precipitate. H₂SO₄ is not used because it adds sulfate ions.

Al: the electron is in a 3p orbital, higher in energy and shielded by 3s. S: the electron is paired in a 3p orbital, and the repulsion makes it easier to remove.

The atoms get larger, so the bonding pair is further from the nucleus and more shielded.

8HI + H₂SO₄ → 4I₂ + H₂S + 4H₂O. Sulfur goes from +6 to -2.

Benefit: it kills bacteria and prevents disease. Risk: chlorine is toxic and can form chlorinated organic compounds. The health benefits outweigh the risks.