Medium MCAT General Chemistry Practice Questions

Mastering General Chemistry is a cornerstone of achieving a high score on the Chemical and Physical Foundations of Biological Systems section of the MCAT. These Medium MCAT General Chemistry Practice Questions are designed to bridge the gap between basic memorization and the complex application required on exam day. By engaging with these problems, you will refine your understanding of stoichiometry, thermodynamics, and kinetics, ensuring you are prepared for the rigors of medical school admissions.

Concept Explanation

MCAT General Chemistry focuses on the physical principles that govern chemical reactions, ranging from atomic structure and periodic trends to equilibrium and electrochemistry. Unlike introductory college courses that may emphasize rote calculation, the MCAT requires students to apply these concepts to biological systems and laboratory scenarios. Key areas of focus include the behavior of gases, the laws of thermodynamics, and the intricacies of acid-base chemistry. Understanding these topics through retrieval practice for medical students is essential for long-term retention. To excel, you must be comfortable with mental math, unit conversions, and interpreting complex data sets under time pressure.

A significant portion of the exam involves high-yield topics such as:

• Stoichiometry: Calculating yields, limiting reagents, and molarity.
• Thermodynamics: Understanding enthalpy ($\Delta H$), entropy ($\Delta S$), and Gibbs free energy ($\Delta G$).
• Kinetics: Determining rate laws and the effect of catalysts.
• Electronic Structure: Predicting electron configurations and periodic properties.

Solved Examples

Reviewing worked problems helps solidify your logic and identifies common pitfalls in calculation or reasoning.

1. Stoichiometry and Limiting Reagents:

   If 10.0 g of magnesium reacts with 10.0 g of oxygen to form magnesium oxide ($2Mg + O_2 \rightarrow 2MgO$), which is the limiting reagent?

   1. Find the moles of Mg: $\frac{10.0 \text{ g}}{24.3 \text{ g/mol}} \approx 0.41 \text{ mol}$.
   2. Find the moles of $O_2$: $\frac{10.0 \text{ g}}{32.0 \text{ g/mol}} \approx 0.31 \text{ mol}$.
   3. According to the balanced equation, 2 moles of Mg react with 1 mole of $O_2$.
   4. To react all 0.31 mol of $O_2$, you would need $0.31 \times 2 = 0.62 \text{ mol}$ of Mg.
   5. Since we only have 0.41 mol of Mg, Mg is the limiting reagent.
2. Gibbs Free Energy Calculation:

   Calculate $\Delta G$ for a reaction at $300 \text{ K}$ where $\Delta H = -150 \text{ kJ/mol}$ and $\Delta S = -0.2 \text{ kJ/mol}\cdot \text{K}$.

   1. Use the formula: $$\Delta G = \Delta H - T\Delta S$$
   2. Substitute the values: $$\Delta G = -150 \text{ kJ/mol} - (300 \text{ K} \times -0.2 \text{ kJ/mol}\cdot \text{K})$$
   3. Calculate the product: $300 \times -0.2 = -60$.
   4. Solve: $-150 - (-60) = -90 \text{ kJ/mol}$. Since $\Delta G < 0$, the reaction is spontaneous.
3. pH of a Weak Acid:

   Determine the pH of a $0.1 \text{ M}$ solution of acetic acid ($K_a = 1.8 \times 10^{-5}$).

   1. Set up the equilibrium expression: $$K_a = \frac{[H^+][A^-]}{[HA]}$$
   2. Assume $[H^+] = [A^-] = x$, so $1.8 \times 10^{-5} = \frac{x^2}{0.1}$.
   3. Solve for $x^2$: $x^2 = 1.8 \times 10^{-6}$.
   4. Take the square root: $x \approx 1.3 \times 10^{-3} \text{ M}$.
   5. Calculate pH: $-\log(1.3 \times 10^{-3}) \approx 3 - \log(1.3) \approx 2.89$.

Practice Questions

Test your knowledge with these Medium MCAT General Chemistry Practice Questions. Ensure you attempt these without a calculator to simulate exam conditions.

1. A sample of gas occupies 2.0 L at 300 K. What will the volume be if the temperature is increased to 450 K at constant pressure?
2. Identify the hybridization of the central atom in $SF_4$.
3. A reaction has a rate law of $\text{Rate} = k[A]^2[B]$. If the concentration of [A] is doubled and [B] is halved, what is the new rate relative to the original?

4. What is the formal charge on the nitrogen atom in the nitrate ion ($NO_3^-$)?
5. Calculate the standard cell potential ($E^\circ_{ \text{cell}}$) for a voltaic cell using the following half-reactions: $$Ag^+ + e^- \rightarrow Ag \quad (E^\circ = +0.80 \text{ V})$$ $$Zn^{2+} + 2e^- \rightarrow Zn \quad (E^\circ = -0.76 \text{ V})$$
6. According to the Heisenberg Uncertainty Principle, which two properties of a particle cannot be measured simultaneously with infinite precision?
7. A solution is prepared by dissolving 58.5 g of $NaCl$ in enough water to make 500 mL of solution. What is the molarity?
8. Which periodic trend increases as you move from left to right across a period and decreases as you move down a group?
9. For the reaction $N_2(g) + 3H_2(g) \rightleftharpoons 2NH_3(g)$, how will an increase in pressure affect the equilibrium position?
10. Determine the oxidation state of Chromium in $K_2Cr_2O_7$.

Answers & Explanations

1. 3.0 L: This uses Charles's Law ($\frac{V_1}{T_1} = \frac{V_2}{T_2}$). $$\frac{2.0}{300} = \frac{V_2}{450} \Rightarrow V_2 = \frac{2.0 \times 450}{300} = \frac{900}{300} = 3.0 \text{ L}$$.
2. $sp^3d$: Sulfur has 6 valence electrons. In $SF_4$, it forms 4 bonds and has 1 lone pair. This equals 5 electron domains, which corresponds to $sp^3d$ hybridization and a see-saw molecular geometry.
3. 2 times the original rate: If [A] is doubled, the rate increases by $2^2 = 4$. If [B] is halved, the rate is multiplied by $0.5$. Total change: $4 \times 0.5 = 2$.
4. +1: In the most stable Lewis structure of $NO_3^-$, Nitrogen has 4 bonds and 0 lone pairs. Formal charge = Valence electrons - (Bonds + Non-bonding electrons) = $5 - (4 + 0) = +1$.
5. +1.56 V: The cell potential is calculated as $E^\circ_{ \text{cell}} = E^\circ_{ \text{cathode}} - E^\circ_{ \text{anode}}$. Since the silver reaction has a higher reduction potential, it is the cathode. $0.80 - (-0.76) = 1.56 \text{ V}$.
6. Position and Momentum: This is a fundamental principle of quantum mechanics. Using retrieval practice vs spaced repetition can help you memorize these theoretical foundations of chemistry.
7. 2.0 M: Molar mass of $NaCl \approx 58.5 \text{ g/mol}$. 58.5 g is 1 mole. Volume is 0.5 L. Molarity = $\frac{1 \text{ mol}}{0.5 \text{ L}} = 2.0 \text{ M}$.
8. Electronegativity (or Ionization Energy): Electronegativity increases across a period due to increased nuclear charge and decreases down a group due to increased shielding.
9. Shift to the right: According to Le Chatelier's Principle, increasing pressure favors the side with fewer moles of gas. The reactant side has 4 moles, while the product side has 2 moles.
10. +6: Potassium is +1, Oxygen is -2. $2(+1) + 2(x) + 7(-2) = 0 \Rightarrow 2 + 2x - 14 = 0 \Rightarrow 2x = 12 \Rightarrow x = +6$.

1. Which of the following describes an adiabatic process?

• AConstant temperature
• BNo heat exchange with the surroundings
• CConstant pressure
• DNo change in internal energy

Frequently Asked Questions

What is the difference between molarity and molality?

Molarity is the number of moles of solute per liter of solution, whereas molality is the number of moles of solute per kilogram of solvent. Molarity is temperature-dependent because volume changes with temperature, while molality remains constant.

How do catalysts affect the equilibrium of a reaction?

Catalysts do not change the equilibrium position or the value of the equilibrium constant ($K_{ \neq}$). They only increase the rate at which equilibrium is reached by lowering the activation energy for both the forward and reverse reactions.

What are the units of the rate constant for a second-order reaction?

The units for a second-order rate constant are $M^{-1}s^{-1}$. This ensures that when multiplied by two concentration terms (each in M), the resulting rate has units of $M/s$.

Why does water have a higher boiling point than hydrogen sulfide?

Water exhibits strong intermolecular hydrogen bonding due to the high electronegativity of oxygen. Hydrogen sulfide ($H_2S$) only experiences weaker dipole-dipole interactions because sulfur is less electronegative, requiring less energy to break the bonds.

What is the relationship between $\Delta G^\circ$ and $K_{ \neq}$?

The relationship is defined by the equation $\Delta G^\circ = -RT \ln K_{ \neq}$. If $K_{ \neq} > 1$, $\Delta G^\circ$ is negative, indicating that the reaction is spontaneous under standard conditions. This is a vital concept in retrieval practice vs practice tests during MCAT prep.