FSc Pre-Engineering Chemistry Flashcards

24 sample cards from the Chemistry deck

Sampled from the end of the deck, so these are different cards from the ones shown on the syllabus page.

1. List the four main types of crystalline solids with an example of each.

   Ionic (e.g. $\ce{NaCl}$), covalent/network (e.g. diamond, $\ce{SiO2}$), molecular (e.g. ice, solid $\ce{CO2}$), and metallic (e.g. copper, iron) crystals.

2. Give the charge, mass, and discoverer of the electron, proton, and neutron.

   Electron: charge $-1.6 \times 10^{-19}\ \text{C}$, mass $9.11 \times 10^{-31}\ \text{kg}$, discovered by J. J. Thomson. Proton: charge $+1.6 \times 10^{-19}\ \text{C}$, mass $1.67 \times 10^{-27}\ \text{kg}$, by Goldstein/Rutherford. Neutron: neutral, mass $\approx$ proton, discovered by James Chadwick.

3. Summarise the conclusions of Rutherford's gold-foil ($\alpha$-scattering) experiment.

   Most $\alpha$-particles passed straight through (atom is mostly empty space); a few were deflected and very few bounced back (a small, dense, positively charged nucleus exists at the centre). Electrons revolve around this nucleus, and most of the atom's mass is concentrated in the nucleus.

4. What were the two main defects of Rutherford's atomic model?

   1) According to classical electromagnetic theory, an orbiting (accelerating) electron should continuously radiate energy and spiral into the nucleus, making the atom unstable. 2) It could not explain the existence of discrete line spectra of atoms.

5. State the key postulates of Bohr's atomic model.

   1) Electrons revolve only in certain fixed circular orbits (stationary states) without radiating energy. 2) Angular momentum is quantized: $mvr = \frac{nh}{2\pi}$. 3) Energy is absorbed or emitted only when an electron jumps between orbits, with $\Delta E = h\nu$.

6. Write the formula for the radius and energy of the $n$th Bohr orbit of hydrogen.

   Radius: $r_n = 0.529 \times n^{2}\ \text{\AA}$. Energy: $$E_n = -\frac{1312}{n^{2}}\ \text{kJ mol}^{-1} = -\frac{13.6}{n^{2}}\ \text{eV}$$ The negative sign indicates the electron is bound to the nucleus.

7. Write Rydberg's equation for the hydrogen spectrum and name the visible series.

   $$\frac{1}{\lambda} = R_H\left(\frac{1}{n_1^{2}} - \frac{1}{n_2^{2}}\right)$$ where $R_H = 1.097 \times 10^{7}\ \text{m}^{-1}$. The Balmer series ($n_1 = 2$) lies in the visible region; Lyman ($n_1=1$) is UV; Paschen, Brackett, Pfund are IR.

8. Name the four quantum numbers and what each describes.

   Principal ($n$): main energy level/shell and size. Azimuthal/angular ($l$): subshell and orbital shape ($s,p,d,f$). Magnetic ($m_l$): orbital orientation in space. Spin ($m_s$): direction of electron spin, $+\frac{1}{2}$ or $-\frac{1}{2}$.

9. What values can the azimuthal quantum number $l$ and magnetic quantum number $m_l$ take?

   $l$ takes values $0$ to $(n-1)$, where $0=s$, $1=p$, $2=d$, $3=f$. For each $l$, $m_l$ takes integer values from $-l$ to $+l$ including $0$, giving $(2l+1)$ orbitals in that subshell.

10. Describe the shapes of $s$ and $p$ orbitals and the maximum electrons each subshell holds.

    An $s$ orbital is spherical; $p$ orbitals are dumbbell-shaped, oriented along the $x$, $y$, $z$ axes. Maximum electrons: $s = 2$, $p = 6$, $d = 10$, $f = 14$. Each orbital holds at most $2$ electrons with opposite spins.

11. State the Aufbau principle, Pauli exclusion principle, and Hund's rule.

    Aufbau: electrons fill orbitals from lowest to highest energy. Pauli exclusion: no two electrons in an atom can have all four quantum numbers identical (an orbital holds at most two electrons of opposite spin). Hund's rule: orbitals of equal energy are each singly occupied before pairing begins, and unpaired electrons have parallel spins.

12. Write the ground-state electronic configuration of chromium ($Z=24$) and copper ($Z=29$).

    Cr: $1s^{2}\,2s^{2}\,2p^{6}\,3s^{2}\,3p^{6}\,3d^{5}\,4s^{1}$. Cu: $1s^{2}\,2s^{2}\,2p^{6}\,3s^{2}\,3p^{6}\,3d^{10}\,4s^{1}$. Both deviate from expected filling because half-filled and fully-filled $d$ subshells give extra stability.

13. Distinguish between ionic and covalent bonds.

    An ionic bond forms by complete transfer of electrons from a metal to a non-metal, creating oppositely charged ions held by electrostatic attraction (e.g. $\ce{NaCl}$). A covalent bond forms by mutual sharing of electron pairs between non-metal atoms (e.g. $\ce{H2}$, $\ce{Cl2}$).

14. Define electronegativity and relate its difference to bond type.

    Electronegativity is the tendency of an atom to attract a shared (bonding) electron pair toward itself. A difference $\geq 1.7$ generally gives a predominantly ionic bond; $0 < \Delta < 1.7$ gives a polar covalent bond; $\Delta = 0$ gives a non-polar covalent bond.

15. Compare the general properties of ionic and covalent compounds.

    Ionic compounds: high melting/boiling points, hard crystalline solids, conduct electricity when molten or in solution, usually soluble in water. Covalent compounds: low melting/boiling points, often gases/liquids/soft solids, generally non-conductors, usually soluble in non-polar solvents.

16. State the main idea of VSEPR theory.

    Valence Shell Electron Pair Repulsion theory: electron pairs (bonding and lone) around a central atom arrange themselves as far apart as possible to minimise repulsion, determining molecular geometry. Repulsion order: lone pair–lone pair > lone pair–bond pair > bond pair–bond pair.

17. Give the geometry and bond angle for 2, 3, and 4 bonding pairs (no lone pairs).

    2 pairs: linear, $180^{\circ}$ (e.g. $\ce{BeCl2}$). 3 pairs: trigonal planar, $120^{\circ}$ (e.g. $\ce{BF3}$). 4 pairs: tetrahedral, $109.5^{\circ}$ (e.g. $\ce{CH4}$).

18. Explain why $\ce{H2O}$ is bent ($104.5^{\circ}$) and $\ce{NH3}$ is pyramidal ($107^{\circ}$) despite tetrahedral electron geometry.

    Both have four electron pairs in tetrahedral arrangement. $\ce{NH3}$ has one lone pair and $\ce{H2O}$ has two lone pairs. Lone-pair repulsion is stronger than bond-pair repulsion, so it compresses the bond angles below $109.5^{\circ}$ — giving $107^{\circ}$ (pyramidal) and $104.5^{\circ}$ (bent) respectively.

19. Define hybridization.

    Hybridization is the mixing of atomic orbitals of slightly different energies (e.g. one $s$ and the appropriate number of $p$/$d$ orbitals) on the same atom to form an equal number of new, equivalent hybrid orbitals with definite shapes and orientations suitable for bonding.

20. Relate $sp$, $sp^{2}$, and $sp^{3}$ hybridization to geometry and bond angle.

    $sp$: linear, $180^{\circ}$ (2 hybrid orbitals, e.g. $\ce{C2H2}$). $sp^{2}$: trigonal planar, $120^{\circ}$ (3 hybrid orbitals, e.g. $\ce{C2H4}$). $sp^{3}$: tetrahedral, $109.5^{\circ}$ (4 hybrid orbitals, e.g. $\ce{CH4}$).

21. Distinguish between a sigma ($\sigma$) bond and a pi ($\pi$) bond.

    A $\sigma$ bond forms by head-on (axial) overlap of orbitals along the internuclear axis and is stronger. A $\pi$ bond forms by sidewise (lateral) overlap of parallel $p$ orbitals and is weaker. A single bond is one $\sigma$; a double bond is one $\sigma$ + one $\pi$; a triple bond is one $\sigma$ + two $\pi$.

22. Define system, surroundings, and the types of thermodynamic systems.

    A system is the part of the universe under study; the surroundings are everything outside it. Types: open (exchanges both matter and energy), closed (exchanges energy only), and isolated (exchanges neither matter nor energy).

23. Define enthalpy and the sign convention for exothermic and endothermic reactions.

    Enthalpy ($H$) is the heat content of a system at constant pressure; the enthalpy change is $\Delta H = H_{\text{products}} - H_{\text{reactants}}$. For an exothermic reaction $\Delta H$ is negative (heat released); for an endothermic reaction $\Delta H$ is positive (heat absorbed).

24. State Hess's law of constant heat summation and one of its uses.

    Hess's law states that the total enthalpy change for a reaction is the same whether it occurs in one step or several steps, depending only on the initial and final states (since $H$ is a state function). It is used to calculate enthalpy changes (e.g. heats of formation) that cannot be measured directly.

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