ETEA Engineering Test 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. State the octet rule.

   Atoms tend to gain, lose, or share electrons to achieve a stable noble-gas configuration of 8 electrons (an octet) in their valence shell.

2. Distinguish between sigma ($\sigma$) and pi ($\pi$) bonds.

   A $\sigma$ bond forms by head-on (axial) overlap of orbitals and is stronger; a $\pi$ bond forms by sideways (lateral) overlap of p-orbitals and is weaker. A double bond = $1\sigma + 1\pi$; triple = $1\sigma + 2\pi$.

3. What hybridization, geometry, and bond angle correspond to $sp^{3}$?

   $sp^{3}$: 4 hybrid orbitals, tetrahedral geometry, bond angle $109.5^{\circ}$ (e.g. $\ce{CH4}$).

4. What hybridization, geometry, and bond angle correspond to $sp^{2}$ and $sp$?

   $sp^{2}$: trigonal planar, $120^{\circ}$ (e.g. $\ce{BF3}$). $sp$: linear, $180^{\circ}$ (e.g. $\ce{BeCl2}$).

5. Why is the bond angle in water ($104.5^{\circ}$) less than the ideal tetrahedral angle?

   Oxygen is $sp^{3}$ hybridized with two lone pairs. Lone pair-lone pair and lone pair-bond pair repulsions are stronger than bond pair-bond pair, compressing the H-O-H angle below $109.5^{\circ}$.

6. Define dipole moment and give its formula and unit.

   Dipole moment is a measure of bond/molecular polarity: $$\mu = q \times d$$ where $q$ is charge and $d$ is distance between charges. Unit: Debye (D).

7. Why is $\ce{CO2}$ nonpolar despite having polar C=O bonds?

   $\ce{CO2}$ is linear and symmetrical, so the two equal and opposite bond dipoles cancel, giving a net dipole moment of zero.

8. Define bond energy and state how it relates to bond order and bond length.

   Bond energy is the energy required to break one mole of a particular bond in the gaseous state. Higher bond order means greater bond energy and shorter bond length.

9. Define enthalpy of reaction ($\Delta H$) and distinguish exothermic from endothermic reactions.

   $\Delta H = H_{\text{products}} - H_{\text{reactants}}$. Exothermic: $\Delta H < 0$ (heat released); Endothermic: $\Delta H > 0$ (heat absorbed).

10. Define the standard enthalpy of formation, $\Delta H_f^{\circ}$.

    The enthalpy change when one mole of a compound is formed from its constituent elements in their standard states under standard conditions ($25\,^{\circ}\text{C}$, $1\ \text{atm}$). For elements in standard state, $\Delta H_f^{\circ} = 0$.

11. State Hess's law of constant heat summation.

    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, not the path taken.

12. How is the enthalpy of a reaction calculated from standard enthalpies of formation?

    $$\Delta H_{rxn}^{\circ} = \sum n\,\Delta H_f^{\circ}(\text{products}) - \sum n\,\Delta H_f^{\circ}(\text{reactants})$$

13. State the law of mass action.

    At constant temperature, the rate of a reaction is proportional to the product of the molar concentrations (active masses) of the reactants, each raised to the power of its coefficient in the balanced equation.

14. Write the equilibrium constant expression $K_c$ for $\ce{aA + bB <=> cC + dD}$.

    $$K_c = \frac{[\ce{C}]^{c}[\ce{D}]^{d}}{[\ce{A}]^{a}[\ce{B}]^{b}}$$ Concentrations are equilibrium molar concentrations.

15. State Le Chatelier's principle.

    If a system at equilibrium is subjected to a change in concentration, temperature, or pressure, the equilibrium shifts in the direction that counteracts (relieves) the imposed change.

16. For an exothermic reaction at equilibrium, which way does the equilibrium shift when temperature is increased?

    It shifts in the backward (reverse) direction, favouring reactants, to absorb the added heat. This decreases the value of $K_c$.

17. Define pH and write its formula.

    pH is the negative logarithm of the hydrogen ion concentration: $$\text{pH} = -\log[\ce{H+}]$$ At $25\,^{\circ}\text{C}$, $\text{pH} + \text{pOH} = 14$.

18. State the ionic product of water ($K_w$) and its value at $25\,^{\circ}\text{C}$.

    $$K_w = [\ce{H+}][\ce{OH-}] = 1.0 \times 10^{-14}\ \text{mol}^2\,\text{dm}^{-6}\ \text{at}\ 25\,^{\circ}\text{C}$$

19. Define the rate of a chemical reaction.

    The change in concentration of a reactant or product per unit time: $$\text{rate} = -\frac{\Delta[\text{reactant}]}{\Delta t} = +\frac{\Delta[\text{product}]}{\Delta t}$$

20. Define the order of a reaction.

    The order is the sum of the powers of the concentration terms in the experimentally determined rate law. For rate $= k[\ce{A}]^{m}[\ce{B}]^{n}$, the overall order $= m + n$.

21. List the main factors that affect the rate of a reaction.

    Concentration (or pressure for gases), temperature, surface area of solids, presence of a catalyst, and nature of the reactants. Light can also affect photochemical reactions.

22. How does a catalyst increase the rate of a reaction?

    A catalyst provides an alternative reaction pathway with a lower activation energy ($E_a$), so more molecules can react, without itself being consumed and without changing $\Delta H$ or equilibrium position.

23. Define oxidation and reduction in terms of electrons and oxidation number.

    Oxidation is loss of electrons (increase in oxidation number); reduction is gain of electrons (decrease in oxidation number). Mnemonic: OIL RIG.

24. Distinguish between an electrolytic cell and a galvanic (voltaic) cell.

    A galvanic cell converts chemical energy to electrical energy via a spontaneous redox reaction ($E^{\circ}_{cell} > 0$). An electrolytic cell uses electrical energy to drive a non-spontaneous reaction. In both, oxidation occurs at the anode and reduction at the cathode.

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