General Chemistry I: Practice Midterm Guidance (Chapters 1–3)

Q1. What is the Kelvin temperature of 45.3 °F? Show all calculations.

Background

Topic: Temperature Conversions

This question tests your ability to convert temperatures between Fahrenheit, Celsius, and Kelvin scales using the appropriate formulas.

Key Terms and Formulas:

• Fahrenheit (°F)

• Celsius (°C)

• Kelvin (K)

Formulas:

• $C = \frac{5}{9}(F - 32)$

• $K = C + 273.15$

Step-by-Step Guidance

1. Start by converting 45.3 °F to Celsius using the formula above.

2. Plug the value of 45.3 into the formula: $C = \frac{5}{9}(45.3 - 32)$

3. Calculate the value inside the parentheses first: $45.3 - 32$

4. Multiply the result by $\frac{5}{9}$ to get the temperature in Celsius.

5. Once you have the Celsius value, add 273.15 to convert to Kelvin: $K = C + 273.15$

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Q2. Perform the following calculation to the correct number of significant figures with proper rounding: [(1.9 x 10-2) ÷ (2.333 x 10-5)] + 37745.0

Background

Topic: Significant Figures and Scientific Notation

This question tests your ability to perform calculations with numbers in scientific notation and apply the rules for significant figures and rounding.

Key Terms and Formulas:

• Significant figures: The digits in a measurement that are known with certainty plus one that is estimated.

• Rules for multiplication/division: The result should have as many significant figures as the value with the fewest significant figures.

• Rules for addition/subtraction: The result should have as many decimal places as the value with the fewest decimal places.

Step-by-Step Guidance

1. Divide $1.9 \times 10^{-2}$ by $2.333 \times 10^{-5}$ using the rules for scientific notation.

2. Calculate the numerical part: $\frac{1.9}{2.333}$

3. Calculate the exponent part: $10^{-2} \div 10^{-5} = 10^{-2 - (-5)} = 10^{3}$

4. Multiply the results from steps 2 and 3 to get the answer to the division.

5. Add this result to 37745.0, making sure to apply the correct significant figures and rounding rules for addition.

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Q3. An iceberg has a volume of 7655 cubic feet. What is the mass of the ice in Mg? Density of water = 0.917 g/mL. 2.54 cm = 1 in.

Background

Topic: Unit Conversions and Density Calculations

This question tests your ability to convert between different units of volume and mass, and to use density as a conversion factor.

Key Terms and Formulas:

• Density ($d$): $d = \frac{\text{mass}}{\text{volume}}$

• 1 ft = 12 in; 1 in = 2.54 cm; 1 cm3 = 1 mL

• 1 Mg (megagram) = 1,000,000 g

Step-by-Step Guidance

1. Convert the volume from cubic feet to cubic inches: $7655 \text{ ft}^3 \times (12 \text{ in}/1 \text{ ft})^3$

2. Convert cubic inches to cubic centimeters: $\text{in}^3 \times (2.54 \text{ cm}/1 \text{ in})^3$

3. Since 1 cm3 = 1 mL, the volume in cm3 is equal to the volume in mL.

4. Use the density to find the mass in grams: $\text{mass (g)} = \text{volume (mL)} \times 0.917 \text{ g/mL}$

5. Convert the mass from grams to megagrams: $\text{mass (Mg)} = \text{mass (g)} \div 1,000,000$

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Q4. What is the name and symbol of the base SI unit of volume? Explain.

Background

Topic: SI Units

This question tests your knowledge of the International System of Units (SI) and the base unit for volume.

Key Terms:

• SI (Système International d'Unités): The standard system of units used in science.

• Base unit: The fundamental unit for a physical quantity in the SI system.

Step-by-Step Guidance

1. Recall the seven base SI units and identify which one is used for volume.

2. Note that the SI base unit for volume is derived from the base unit for length.

3. State the name and symbol for the SI unit of volume.

4. Explain why this unit is used for volume and how it relates to the unit for length.

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Q5. Water and solid sodium hydroxide are mixed together by stirring. After stirring for 15 minutes all of the solid sodium hydroxide dissolved in the water to form sodium ions and hydroxide ions. There was also a 10°C increase in temperature. Comment on the accuracy of the following statement and give an explanation: "After mixing for 15 minutes we have a heterogeneous pure mixture that contains only elements and this is an example of a chemical change that resulted in lower kinetic energy."

Background

Topic: Chemical and Physical Changes, Mixtures, and Energy

This question tests your understanding of mixtures, chemical changes, and energy changes during reactions.

Key Terms:

• Heterogeneous mixture: A mixture with non-uniform composition.

• Pure substance: A material with a constant composition.

• Chemical change: A process where substances are transformed into different substances.

• Kinetic energy: The energy of motion, often related to temperature.

Step-by-Step Guidance

1. Analyze whether the resulting solution is heterogeneous or homogeneous.

2. Determine if the mixture is pure or a mixture of substances.

3. Assess whether the mixture contains only elements or compounds/ions.

4. Evaluate if dissolving sodium hydroxide in water is a chemical or physical change.

5. Consider the temperature increase and what it implies about kinetic energy.

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Q6. Two different compounds containing osmium and oxygen have the following masses of oxygen per 1.500 g of osmium: 2.52 dg and 25.2 cg. Are these results consistent with the law of multiple proportions? Explain. If these results are not consistent with the law of multiple proportions then explain what if any law the results are consistent with.

Background

Topic: Laws of Chemical Combination

This question tests your understanding of the law of multiple proportions and how to apply it to experimental data.

Key Terms and Formulas:

• Law of Multiple Proportions: When two elements form more than one compound, the masses of one element that combine with a fixed mass of the other are in ratios of small whole numbers.

• 1 dg = 0.1 g; 1 cg = 0.01 g

Step-by-Step Guidance

1. Convert both masses of oxygen to the same unit (grams).

2. Calculate the ratio of the two masses of oxygen that combine with the same mass of osmium.

3. Simplify the ratio to see if it is a ratio of small whole numbers.

4. Determine if this ratio supports the law of multiple proportions.

5. If not, consider which law (e.g., law of definite proportions) might apply.

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Q7. Which of the statement(s) of Dalton’s atomic theory are consistent with and which are inconsistent with the existence of boron-11 and carbon-12? Explain.

Background

Topic: Dalton's Atomic Theory and Isotopes

This question tests your understanding of Dalton's atomic theory and how the discovery of isotopes relates to it.

Key Terms:

• Dalton's atomic theory: Early model of the atom with several postulates.

• Isotopes: Atoms of the same element with different numbers of neutrons.

Step-by-Step Guidance

1. Recall the main postulates of Dalton's atomic theory.

2. Identify which postulate(s) are challenged by the existence of isotopes like boron-11 and carbon-12.

3. Explain why the existence of isotopes is consistent or inconsistent with each relevant postulate.

4. Provide reasoning for your choices.

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Q8. What is the ratio of the oxidation states of elemental manganese to the oxidation state of manganese in Mn2O7?

Background

Topic: Oxidation States

This question tests your ability to assign oxidation states to elements in compounds and compare them.

Key Terms and Formulas:

• Oxidation state: The hypothetical charge an atom would have if all bonds were ionic.

• Elemental form: Oxidation state is always zero.

• Sum of oxidation states in a neutral compound is zero.

Step-by-Step Guidance

1. Recall that the oxidation state of an element in its elemental form is zero.

2. Assign the oxidation state of oxygen in Mn2O7 (usually -2).

3. Set up an equation for the sum of oxidation states in Mn2O7 to equal zero.

4. Solve for the oxidation state of manganese in Mn2O7.

5. Express the ratio of the oxidation state of elemental manganese to that in Mn2O7.

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Q9. How many protons, neutrons and electrons are there in iodine-127?

Background

Topic: Atomic Structure

This question tests your ability to determine the number of subatomic particles in a neutral atom given its isotope notation.

Key Terms:

• Atomic number (Z): Number of protons

• Mass number (A): Number of protons + neutrons

• Electrons: Equal to protons in a neutral atom

Step-by-Step Guidance

1. Find the atomic number of iodine (I) on the periodic table.

2. The mass number is given as 127.

3. Number of protons = atomic number.

4. Number of neutrons = mass number - atomic number.

5. Number of electrons = number of protons (for a neutral atom).

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Q10. How many protons, neutrons and electrons are there in each of the following species? Are any of these species isotopes of each other? Explain. (a) 40Ar (b) 32S2- (c) 70Ga3+

Background

Topic: Isotopes and Ions

This question tests your ability to determine the number of subatomic particles in various isotopes and ions, and to identify isotopes.

Key Terms:

• Isotopes: Atoms of the same element with different mass numbers.

• Ions: Atoms with a net charge due to loss or gain of electrons.

• Protons: Atomic number

• Neutrons: Mass number - atomic number

• Electrons: Protons minus charge (for cations), protons plus charge (for anions)

Step-by-Step Guidance

1. For each species, find the atomic number (Z) from the periodic table.

2. Determine the number of protons (equals Z).

3. Calculate the number of neutrons: mass number - atomic number.

4. Calculate the number of electrons, adjusting for the charge.

5. Compare the species to see if any are isotopes (same element, different mass number).

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Q11. Give the names of two species and their family(s) that meet the following criteria: representative elements that form ions with a charge of 2+.

Background

Topic: Periodic Table and Ion Formation

This question tests your knowledge of the periodic table, specifically which representative elements form 2+ ions and their group/family names.

Key Terms:

• Representative elements: Elements in groups 1, 2, and 13–18.

• 2+ ions: Typically formed by group 2 (alkaline earth metals).

• Family: Group name on the periodic table.

Step-by-Step Guidance

1. Identify which groups on the periodic table contain elements that commonly form 2+ ions.

2. List two elements from these groups.

3. State the family/group name for each element.

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Q12. An element has a Z number of 14 and an A number of 30. Give the name and symbol of a metalloid that you would expect to have the most similarity in its chemical reactivity to this element.

Background

Topic: Periodic Trends and Chemical Reactivity

This question tests your understanding of periodic trends and how elements in the same group have similar chemical properties.

Key Terms:

• Z number: Atomic number

• A number: Mass number

• Metalloid: Element with properties intermediate between metals and nonmetals

Step-by-Step Guidance

1. Identify the element with Z = 14 (atomic number 14) and its group on the periodic table.

2. Find the metalloid in the same group as this element.

3. State the name and symbol of this metalloid.

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Q13. How many zinc atoms are found in 4.55 mmoles of zinc?

Background

Topic: Mole Concept and Avogadro's Number

This question tests your ability to convert between moles and number of atoms using Avogadro's number.

Key Terms and Formulas:

• 1 mole = $6.022 \times 10^{23}$ particles (Avogadro's number)

• mmole: 1 millimole = $1 \times 10^{-3}$ moles

Step-by-Step Guidance

1. Convert 4.55 mmoles to moles: $4.55 \times 10^{-3}$ moles

2. Multiply the number of moles by Avogadro's number to get the number of atoms.

3. Set up the calculation: $\text{atoms} = \text{moles} \times 6.022 \times 10^{23}$

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Q14. How many μg of zinc are found in 6.55 x 1021 zinc atoms?

Background

Topic: Mole Concept, Mass, and Avogadro's Number

This question tests your ability to convert from number of atoms to mass, using Avogadro's number and molar mass, and to convert units of mass.

Key Terms and Formulas:

• Avogadro's number: $6.022 \times 10^{23}$ atoms/mol

• Molar mass of Zn: about 65.38 g/mol

• 1 μg = $1 \times 10^{-6}$ g

Step-by-Step Guidance

1. Calculate the number of moles: $\text{moles} = \frac{6.55 \times 10^{21}}{6.022 \times 10^{23}}$

2. Multiply the number of moles by the molar mass of zinc to get the mass in grams.

3. Convert the mass from grams to micrograms (μg): $1 \text{ g} = 1 \times 10^{6} \mu\text{g}$

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Q15. How many grams of aluminum are found in 2.45 moles of aluminum?

Background

Topic: Moles and Molar Mass

This question tests your ability to convert between moles and grams using the molar mass.

Key Terms and Formulas:

• Molar mass of Al: about 26.98 g/mol

• Mass (g) = moles × molar mass

Step-by-Step Guidance

1. Write the formula: $\text{mass (g)} = \text{moles} \times \text{molar mass}$

2. Plug in the values: 2.45 moles and 26.98 g/mol

3. Set up the multiplication to find the mass in grams.

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Q16. Is sodium sulfate an example of an ionic or a molecular compound? Explain.

Background

Topic: Types of Compounds

This question tests your understanding of the difference between ionic and molecular compounds and how to classify them.

Key Terms:

• Ionic compound: Composed of metal and nonmetal ions.

• Molecular compound: Composed of nonmetals only.

• Sodium sulfate: Na2SO4

Step-by-Step Guidance

1. Identify the elements in sodium sulfate and their types (metal or nonmetal).

2. Determine if the compound is formed from ions or molecules.

3. Explain your reasoning based on the composition.

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Q17. Write the formula for triphosphorus heptachloride.

Background

Topic: Nomenclature of Molecular Compounds

This question tests your ability to write chemical formulas from names using Greek prefixes.

Key Terms:

• Tri-: 3

• Hepta-: 7

• Phosphorus: P

• Chloride: Cl

Step-by-Step Guidance

1. Identify the number of phosphorus and chlorine atoms from the prefixes.

2. Write the chemical symbols with the appropriate subscripts.

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Q18. What is the name of the following: Cu(NO3)2 · 8H2O

Background

Topic: Nomenclature of Ionic Compounds and Hydrates

This question tests your ability to name ionic compounds with polyatomic ions and waters of hydration.

Key Terms:

• Cu(NO3)2: Copper(II) nitrate

• Hydrate: Compound with water molecules attached

• 8H2O: Octahydrate

Step-by-Step Guidance

1. Name the cation (Cu2+) and anion (NO3-).

2. Combine the names to get the name of the salt.

3. Add the hydrate prefix for 8 water molecules.

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Q19. Write the formula and name of the compound that forms from the most stable ions that are formed from barium and fluorine.

Background

Topic: Ionic Compounds and Ion Charges

This question tests your ability to determine the charges of ions and write the formula and name of the resulting compound.

Key Terms:

• Barium: Ba2+

• Fluorine: F-

• Formula: Use the criss-cross method to balance charges

Step-by-Step Guidance

1. Determine the charges of the most stable ions for barium and fluorine.

2. Balance the charges to write the correct formula.

3. Name the compound using standard ionic nomenclature.

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Q20. Calculate the molar mass of Ga2(HPO4)3

Background

Topic: Molar Mass Calculations

This question tests your ability to calculate the molar mass of a compound by summing the atomic masses of all atoms in the formula.

Key Terms and Formulas:

• Molar mass: Sum of atomic masses of all atoms in a formula (g/mol)

• Atomic masses: Ga (about 69.72), H (1.01), P (30.97), O (16.00)

Step-by-Step Guidance

1. Identify the number of each type of atom in Ga2(HPO4)3.

2. Multiply the number of each atom by its atomic mass.

3. Add up all the masses to get the total molar mass.

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Q21. What is the mass percent of phosphorus in Ga2(HPO4)3?

Background

Topic: Mass Percent Composition

This question tests your ability to calculate the mass percent of an element in a compound.

Key Terms and Formulas:

• Mass percent: $\frac{\text{mass of element in 1 mole}}{\text{molar mass of compound}} \times 100\%$

Step-by-Step Guidance

1. Calculate the total mass of phosphorus in one mole of Ga2(HPO4)3 (number of P atoms × atomic mass of P).

2. Divide by the molar mass of the compound (from Q20).

3. Multiply by 100% to get the mass percent.

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Q22. How many atoms of oxygen are needed to produce 23.0 mg of Ga2(HPO4)3?

Background

Topic: Stoichiometry and Mole Concept

This question tests your ability to convert mass to moles, use the chemical formula to find the number of atoms, and apply Avogadro's number.

Key Terms and Formulas:

• Molar mass (from Q20)

• Number of O atoms per formula unit

• Avogadro's number: $6.022 \times 10^{23}$

Step-by-Step Guidance

1. Convert 23.0 mg to grams.

2. Calculate the number of moles of Ga2(HPO4)3 using its molar mass.

3. Determine the number of O atoms per formula unit.

4. Multiply the number of moles by Avogadro's number and the number of O atoms per formula unit.

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Q23. What is the simplest or empirical formula for the substance with this analysis: 2.677 g Ba and 3.115 g Br?

Background

Topic: Empirical Formula Determination

This question tests your ability to determine the empirical formula from mass data.

Key Terms and Formulas:

• Empirical formula: Simplest whole-number ratio of atoms in a compound

• Moles = mass / atomic mass

Step-by-Step Guidance

1. Calculate the number of moles of Ba and Br using their atomic masses.

2. Divide each by the smallest number of moles to get the simplest ratio.

3. Write the empirical formula based on this ratio.

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Q24. Problem 3.95 (Principles) or 3.127 (Chemistry) in lecture texts.

Background

Topic: Refer to your textbook for the specific problem. This question is a placeholder for a textbook exercise.

Key Terms:

• Consult your lecture text for the exact question and follow the same step-by-step approach as above.

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Q25. Write a balanced chemical equation for the following: Solid lead(II) sulfide reacts with aqueous hydrobromic acid to form solid lead(II) bromide and dihydrogen monosulfide gas.

Background

Topic: Writing and Balancing Chemical Equations

This question tests your ability to write chemical formulas from names and balance chemical equations.

Key Terms:

• Lead(II) sulfide: PbS

• Hydrobromic acid: HBr (aq)

• Lead(II) bromide: PbBr2

• Dihydrogen monosulfide: H2S (g)

Step-by-Step Guidance

1. Write the formulas for all reactants and products.

2. Set up the unbalanced equation.

3. Balance the equation by adjusting coefficients to conserve atoms of each element.

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Q26. Problem 3.81 (Principles) or 3.107 (Chemistry) in lecture text.

Background

Topic: Refer to your textbook for the specific problem. This question is a placeholder for a textbook exercise.

Key Terms:

• Consult your lecture text for the exact question and follow the same step-by-step approach as above.

Try solving on your own before revealing the answer!