Academic Chemistry Spring Semester Review 2024-2025

Formula Sheet

• The exam will have 80 multiple choice questions worth 1.25 points each.
• A formula sheet and STAAR periodic table/formula chart will be provided.
• Formula Sheet includes:
  • Percent yield: $\% \text{ yield} = \frac{\text{experimental yield}}{\text{theoretical yield}} \times 100$
  • Pressure conversion: $1 \text{ atm} = 760 \text{ mm Hg} = 101.3 \text{ kPa}$
  • Standard pressure: $1 \text{ atm}$
  • Standard temperature: $273 \text{ K}$
  • Ideal gas law: $PV = nRT$, where $R = 0.0821 \frac{\text{L} \cdot \text{atm}}{\text{K} \cdot \text{mol}}$
  • Boyle's Law: $P<em>1V</em>1 = P<em>2V</em>2$
  • Charles's Law: $\frac{V<em>1}{T</em>1} = \frac{V<em>2}{T</em>2}$
  • Gay-Lussac's Law: $\frac{P<em>1}{T</em>1} = \frac{P<em>2}{T</em>2}$
  • Combined Gas Law: $\frac{P<em>1V</em>1}{T<em>1} = \frac{P</em>2V<em>2}{T</em>2}$
  • Molarity: $M = \frac{\text{moles solute}}{\text{L solution}}$
  • Dilution: $M<em>1V</em>1 = M<em>2V</em>2$
  • pH: $pH = -\log[H^+]$
  • Ion product of water: $[H^+] \times [OH^-] = 1.0 \times 10^{-14}$
  • pOH: $pOH = -\log[OH^-]$
  • Heat transfer: $q = mC\Delta T$
  • Enthalpy of reaction: $\Delta H<em>{\text{reaction}} = \Delta H</em>f \text{ products} - \Delta H_f \text{ reactants}$
  • Alternate Combined Gas Law Form: $\frac{V<em>1T</em>2}{V<em>2T</em>1}=1$
  • Alternate Combined Gas Law Form: $\frac{P<em>1T</em>2}{P<em>2T</em>1}=1$
  • Alternate Combined Gas Law Form: $\frac{P<em>1V</em>1T<em>2}{P</em>2V<em>2T</em>1}=1$
  • Hydrogen ion concentration: $[H^+] = 10^{-pH}$
  • pH and pOH relationship: $pH + pOH = 14$
  • Hydroxide ion concentration: $[OH^-] = 10^{-pOH}$

Chemical Reactions

• Reactants are on the left side of the reaction, and products are on the right side.
• A coefficient is a number in front of a chemical formula used to balance the equation. It indicates the number of moles of that substance involved in the reaction.
• A catalyst speeds up a chemical reaction without being consumed in the reaction. It is indicated above the arrow in a chemical equation.
• The activity series is used in single replacement reactions to determine if a reaction will occur. A reaction will occur if the element that is trying to replace another element is higher on the activity series.
• A double replacement reaction involves the exchange of positive ions between two compounds. Example: $\text{AB + CD} \rightarrow \text{AD + CB}$
• A single replacement reaction involves one element replacing another element in a compound. Example: $\text{A + BC} \rightarrow \text{B + AC}$
• A synthesis reaction involves two or more reactants combining to form a single product. Example: $\text{A + B} \rightarrow \text{AB}$
• A decomposition reaction involves a single reactant breaking down into two or more products. Example: $\text{AB} \rightarrow \text{A + B}$
• A combustion reaction involves the rapid reaction between a substance with an oxidant, usually oxygen, to produce heat and light. It often involves a hydrocarbon reacting with oxygen to produce carbon dioxide and water. Example: $\text{CxHy + O2} \rightarrow \text{CO2 + H2O}$
• The seven diatomic elements are hydrogen (H2), nitrogen (N2), oxygen (O2), fluorine (F2), chlorine (Cl2), bromine (Br2), and iodine (I2).
• The symbol for aqueous is (aq), gas is (g), solid is (s), and liquid is (l).
• Balanced equation for the synthesis reaction between magnesium and sulfur: $Mg(s) + S(s) \rightarrow MgS(s)$
• Balanced equation for the decomposition reaction of aluminum oxide: $2Al<em>2O</em>3(s) \rightarrow 4Al(s) + 3O_2(g)$
• Balanced equation for the single replacement reaction between potassium and copper (II) chloride: $2K(s) + CuCl_2(aq) \rightarrow 2KCl(aq) + Cu(s)$
• Balanced equation for the double replacement reaction between lead (II) iodide and sodium nitrate: $PbI<em>2(aq) + 2NaNO</em>3(aq) \rightarrow Pb(NO<em>3)</em>2(aq) + 2NaI(aq)$
• Balanced equation for the combustion reaction of propane (C3H8): $C<em>3H</em>8(g) + 5O<em>2(g) \rightarrow 3CO</em>2(g) + 4H_2O(g)$
• Balancing reactions:
  • $\underline{2} AlBr<em>3 + \underline{3} Cl</em>2 \rightarrow \underline{2} AlCl<em>3 + \underline{3} Br</em>2$
  • $\underline{1} P<em>4O</em>{10} + \underline{12} KOH \rightarrow \underline{4} K<em>3PO</em>4 + \underline{6} H_2O$
  • $\underline{1} Ca(OH)<em>2 + \underline{2} HNO</em>3 \rightarrow \underline{1} Ca(NO<em>3)</em>2 + \underline{2} H_2O$

Balancing and Classifying Chemical Reactions

• Silver nitrate reacts with potassium sulfate to form silver sulfate and potassium nitrate:
  • $2AgNO<em>3(aq) + K</em>2SO<em>4(aq) \rightarrow Ag</em>2SO<em>4(s) + 2KNO</em>3(aq)$
  • Type of reaction: Double Replacement
• Propane (C3H8) burns in oxygen to produce carbon dioxide and water.
  • $C<em>3H</em>8(g) + 5O<em>2(g) \rightarrow 3CO</em>2(g) + 4H_2O(g)$
  • Type of reaction: Combustion
• Sodium reacts with chlorine to produce sodium chloride.
  • $2Na(s) + Cl_2(g) \rightarrow 2NaCl(s)$
  • Type of reaction: Synthesis
• Hydrogen peroxide (H2O2) decomposes to produce water and oxygen.
  • $2H<em>2O</em>2(aq) \rightarrow 2H<em>2O(l) + O</em>2(g)$
  • Type of reaction: Decomposition

Stoichiometry

• Use the following reaction for questions #1-2: $N<em>2 + 3 H</em>2 \rightarrow 2 NH_3$
  • If you start the reaction with 10.0 moles of each reactant, hydrogen is the limiting reactant.
  • If the theoretical yield was 22.5 g of $NH<em>3$, but only 19.8 g of $NH</em>3$ were experimentally produced, the percent yield is:
  • $\% \text{ yield} = \frac{19.8 \text{ g}}{22.5 \text{ g}} \times 100 = 88.0 \%$
• Stoichiometry problems all have one common step: Convert grams to moles.
• In the reaction $4 Fe + 3 O<em>2 \rightarrow 2 Fe</em>2O_3$, the mole ratio of iron to iron (III) oxide is 4:2 or 2:1.
• In the reaction $2 Al<em>2O</em>3 \rightarrow 4 Al + 3 O_2$, the mole ratio of aluminum oxide to oxygen is 2:3.
• Consider the reaction $N<em>2 + 3 H</em>2 \rightarrow 2 NH_3$. To completely react 8.0 mol nitrogen gas, 24 mol of hydrogen gas are required.
• For the reaction $C + 2 H<em>2 \rightarrow CH</em>4$, 2.3 mol of carbon are required to produce 2.3 mol of methane, $CH_4$.
• For the reaction $2 H<em>2 + O</em>2 \rightarrow 2 H_2O$, 4.00 L of hydrogen gas at STP will produce 0.179 mol of water.
• For the reaction $2 Na + Cl_2 \rightarrow 2 NaCl$, 92.0 g of sodium are required to react completely with 2.00 mol of chlorine gas.
• For the reaction $CH<em>4 + 2 O</em>2 \rightarrow CO<em>2 + 2 H</em>2O$, 349 L of $CO<em>2$ (at STP) are produced from the combustion of 25.0 g of methane ($CH</em>4$).
• For the reaction $Cl<em>2 + 2 KBr \rightarrow 2 KCl + Br</em>2$, 67.2 g of bromine can be produced from 100. g of potassium bromide.
• For the reaction $2 Mg + O<em>2 \rightarrow 2 MgO$, 0.43 g of $O</em>2$ is used when $1.6 \times 10^{22}$ atoms of Mg are used.
• The limiting reactant is the reactant that is completely consumed in a chemical reaction and determines the amount of product that can be formed. The excess reactant is the reactant that is present in a greater amount than necessary to react completely with the limiting reactant.
• If nitrogen is the limiting reactant in the reaction: $N<em>2 + 3 H</em>2 \rightarrow 2 NH_3$, hydrogen will be present when the reaction stops.
• Consider the reaction $2 H<em>2 + O</em>2 \rightarrow 2 H_2O$. If you start the reaction with 16.0 g hydrogen gas and 68.0 g oxygen gas:
  • Oxygen is the limiting reactant.
  • 76.6 g of water can be produced (theoretical yield).
• Find the percent yield of water for the above reaction if, in an experiment, a student produced 70.0 g of water:
  • $\% \text{ yield} = \frac{70.0 \text{ g}}{76.6 \text{ g}} \times 100 = 91.4 \%$\n

Gas Laws

• According to the kinetic-molecular theory, the particles of a gas are in constant, random motion, are far apart relative to their size, have elastic collisions, and have no attractive or repulsive forces between them.
• To convert 55°C to Kelvin: $55 + 273 = 328 K$
• A barometer measures atmospheric pressure.
• If the temperature of a fixed quantity of gas decreases and the pressure remains unchanged, the volume decreases.
• The air pressure inside the tires of a car increases when the car is driven because the temperature of the air inside the tires increases due to friction.
• At constant temperature, if the volume of a flexible container is increased, the pressure will decrease.
• The volume of a gas is 800.0 mL when the pressure is 2.00 atm. At the same temperature, the pressure at which the volume of the gas is 3.50 L is 0.457 atm.
• The volume of a gas is 2.5 L when the temperature is 8.0ºC. If the temperature is increased to 25.0ºC without changing the pressure, the new volume is 2.7 L.
• A 220. mL volume of gas is measured at 63.0ºC. If the pressure remains unchanged, the volume of the gas at standard temperature is 179 mL.
• On a cold winter morning when the temperature is –11.0ºC, the air pressure in an automobile tire is 1.20 atm. If the volume does not change, the pressure after the tire has warmed to 25.0ºC is 1.36 atm.
• The pressure of a 900. mL sample of gas at 15.0ºC increases from 500. mm Hg to 800. mm Hg. If the volume is unchanged, the new temperature is 188ºC.
• The volume of a sample of oxygen is 700.0 mL when the pressure is 1 atm and the temperature is 37.0ºC. The Kelvin temperature at which the volume is 2.40 L and the pressure is 0.600 atm is 638 K.
• A sample of gas at 25.0ºC has a volume of 9.00 L and exerts a pressure of 740. mm Hg. The number of moles of gas in the sample is 0.358 mol.

Solutions

• Carbon dioxide in water is an example of a gas solute in a liquid solvent.
• Sugar in water is an example of a solid solute in a liquid solvent.
• Factors that increase the rate of dissolving a solid in water include increasing the temperature, stirring, and using smaller crystals.
• If you wanted a solid substance to dissolve slowly, you should use large crystals and not stir.
• If the amount of solute present in a solution at a given temperature is less than the maximum amount that can dissolve at that temperature, the solution is said to be unsaturated.
• If the amount of dissolved solute in a solution at a given temperature is greater than the amount that can permanently remain in solution at that temperature, the solution is said to be supersaturated.
• The rule “like dissolves like” means that polar solvents dissolve polar solutes, and nonpolar solvents dissolve nonpolar solutes.
• Water is the polar solvent called the “universal solvent.”
• A substance that is not soluble in a polar solvent is nonpolar.
• Milk is soluble in water because the molecules in milk are polar.
• As pressure increases, the solubility of gases in liquids increases.
• As temperature increases, the solubility of gases in liquids decreases.
• In most cases, as temperature increases, the solubility of solids in liquids increases.
• The molarity of a solution that contains 1.4 mol $NH_3$ in 3.0 L solution is 0.47 M.
• The molarity of a solution that contains 65.0 g KCl in 3.00 L solution is 0.291 M.
• The number of moles of HCl present in 1.50 L of a 0.880 M HCl solution is 1.32 mol.
• The mass % of a solution composed of 15 g salt in 100 g water is 13%.
• The amount of NaCl dissolved in a 9.0% NaCl solution if the total solution mass is 100. g is 9.00 g.
• You have a solution of 18.0 M HCl and you want to make a dilution so that you have 2000. mL of a 1.00 M solution. You will require 111 mL of the original acid and will have to mix it with 1889 mL of water.

Solutions: Solubility Curves

• How many grams of $KClO_3$ can be dissolved in 200 g of water at 50°C? Answer should reference solubility curve.
• Is a solution of 10 g $NH_3$ at 90°C: (circle) unsaturated, saturated, or supersaturated? Answer should reference solubility curve.

Acids and Bases

• The two products in a neutralization reaction are a salt and water.
• Write and balance the chemical reaction between nitric acid and calcium hydroxide: $2HNO<em>3 + Ca(OH)</em>2 \rightarrow Ca(NO<em>3)</em>2 + 2H_2O$
• Write and balance the chemical reaction between carbonic acid and potassium hydroxide: $H<em>2CO</em>3 + 2KOH \rightarrow K<em>2CO</em>3 + 2H_2O$
• Determine the formulas, given the following names:
  • phosphoric acid $H<em>3PO</em>4$
  • sulfurous acid $H<em>2SO</em>3$
  • hydrofluoric acid $HF$
  • magnesium hydroxide $Mg(OH)_2$
  • Determine the names, given the following formulas:
  • $NH_3$ ammonia
  • $H<em>3PO</em>3$ phosphorous acid
  • $HCl$ hydrochloric acid
  • $NaOH$ sodium hydroxide
  • $H<em>2CO</em>3$ carbonic acid
• Acids taste sour.
• Litmus paper turns red in an acid and blue in a base.
• Acids react with bases to produce salt and water.
• Bases taste bitter.
• An electrolyte is a substance that conducts electricity when dissolved in water. Acids, bases, and salts form electrolytes when dissolved in water.
• The substances produced when $Ba(OH)<em>2$ neutralizes $H</em>2SO_4$ are barium sulfate and water.
• The formula for finding the pH of a solution if given $[H^+]$ is $pH=-\log[H^+]$.
• The pH of an acidic solution is less than 7. The pOH of an acidic solution is greater than 7.
• The pH of a basic solution is greater than 7. The pOH of a basic solution is less than 7.
• The pH of a solution of HCl with $[H^+] = 1.0 \times 10^{-4} M$ is 4.
• If $[OH^-] = 4.7 \times 10^{-5} M$, the pOH of the solution is 4.33.
• Determine whether the following solutions are acidic, basic, or neutral.
  • $[OH^-] = 1.26 \times 10^{-5} M$ basic
  • pH = 3.3 acidic
  • $[H^+] = [OH^-]$ neutral
  • pOH = 5 basic

Thermochemistry

• Calculate the heat of reaction ($\Delta H$) for the following reaction: $CaO(s) + 2 HCl(g) \rightarrow CaCl<em>2(s) + H</em>2O(g)$. (Value needed from reference data)
• Is the reaction above endothermic or exothermic? (depending on calculated $∆H$)
• Define the following terms:
  • Activation energy: The minimum energy required for a chemical reaction to occur.
  • Change in enthalpy ($\Delta H$): The heat absorbed or released during a chemical reaction at constant pressure.
  • Reactants: The starting materials in a chemical reaction.
  • Products: The substances formed as a result of a chemical reaction.
• Determine whether the following are endothermic or exothermic:
  • $\Delta H<em>{\text{products}} – \Delta H</em>{\text{reactants}} < 0$ exothermic
  • $\Delta H_{\text{reaction}} = – 355 kJ/mol$ exothermic
• Determine whether the following are endothermic or exothermic:
  • Melting: solid to liquid, endothermic
  • Freezing: liquid to solid, exothermic
  • Condensing: gas to liquid, exothermic
  • Boiling: liquid to gas, endothermic
  • Deposition: gas to solid, exothermic
  • Sublimation: solid to gas, endothermic
• How is evaporation similar to and different from boiling?
  • Evaporation occurs at the surface of a liquid at any temperature, while boiling occurs throughout the liquid at a specific temperature (boiling point).
• What is the specific heat of a substance if it takes 250 J of heat energy to raise the temperature of a 20. g sample from 20°C to 45°C? (0.50 J/g°C)
• A 100.0 g sample of a metal with a specific heat of 0.836 J/g°C is heated from 10.0°C to 80.0°C. How much thermal energy in Joules did it absorb? (5852 J)
  *The surroundings lose heat as the reaction occurs. exothermic
  *$2 HCl + 44 kcal \rightarrow H<em>2 + Cl</em>2$ endothermic