Chem 140A Exam 2 Study Guidance

Q1. Which of the following elements will generally form only one covalent bond?

Background

Topic: Covalent Bonding and Valence Electrons

This question tests your understanding of how many covalent bonds different elements typically form, based on their valence electrons and the octet rule.

Key Terms and Concepts:

• Covalent bond: A chemical bond formed by the sharing of electron pairs between atoms.

• Valence electrons: Electrons in the outermost shell, which determine bonding behavior.

• Octet rule: Atoms tend to form bonds to achieve 8 electrons in their valence shell (or 2 for hydrogen).

Step-by-Step Guidance

1. Recall the number of valence electrons for each element listed (hydrogen, carbon, nitrogen, sulfur, argon).

2. Consider how many electrons each element needs to achieve a full valence shell (octet or duet for hydrogen).

3. Determine how many covalent bonds each element typically forms to reach a stable configuration.

4. Identify which element among the choices generally forms only one covalent bond.

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Q2. How many lone pairs of electrons are needed to complete the aspartic acid structure shown below?

Background

Topic: Lewis Structures and Lone Pairs

This question tests your ability to count lone pairs in a molecular structure, which is important for understanding molecular geometry and reactivity.

Key Terms and Concepts:

• Lone pair: A pair of valence electrons not shared with another atom in a bond.

• Lewis structure: A diagram showing the arrangement of atoms, bonds, and lone pairs in a molecule.

Step-by-Step Guidance

1. Examine the aspartic acid structure and identify all atoms that typically have lone pairs (such as oxygen and nitrogen).

2. Recall the typical number of lone pairs for each atom type (e.g., oxygen usually has 2, nitrogen usually has 1).

3. Count the number of lone pairs on each relevant atom in the structure.

4. Add up the total number of lone pairs needed to complete the structure.

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Q3. Which of the following molecules will not contain a multiple bond in its Lewis structure?

Background

Topic: Lewis Structures and Multiple Bonds

This question tests your ability to recognize when a molecule requires double or triple bonds to satisfy the octet rule.

Key Terms and Concepts:

• Multiple bond: A double or triple bond between two atoms.

• Lewis structure: Shows how atoms are bonded and where lone pairs are located.

Step-by-Step Guidance

1. Write the Lewis structure for each molecule listed.

2. Determine if each molecule requires double or triple bonds to satisfy the octet rule for all atoms.

3. Identify which molecule can be drawn with only single bonds.

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Q4. What is the molecular formula for dinitrogen pentoxide?

Background

Topic: Chemical Nomenclature

This question tests your ability to translate a compound's name into its correct chemical formula using prefixes.

Key Terms and Concepts:

• Prefix "di-": Means 2 atoms.

• Prefix "penta-": Means 5 atoms.

• Molecular formula: Shows the number and type of atoms in a molecule.

Step-by-Step Guidance

1. Break down the name: "dinitrogen" means 2 nitrogen atoms, "pentoxide" means 5 oxygen atoms.

2. Write the formula using the correct subscripts for each element.

3. Compare your formula to the answer choices.

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Q5. What is the molecular shape around the phosphorus atom in PH3?

Background

Topic: VSEPR Theory and Molecular Geometry

This question tests your ability to predict molecular shape using the Valence Shell Electron Pair Repulsion (VSEPR) theory.

Key Terms and Concepts:

• VSEPR theory: Predicts the shape of molecules based on electron pair repulsion.

• Lone pairs: Affect the geometry by repelling bonding pairs.

Step-by-Step Guidance

1. Draw the Lewis structure for PH3.

2. Count the number of bonding pairs and lone pairs around the phosphorus atom.

3. Use VSEPR theory to determine the molecular geometry based on these pairs.

4. Match the geometry to the correct shape from the answer choices.

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Q6. Which atom has the lowest electronegativity?

Background

Topic: Periodic Trends

This question tests your understanding of how electronegativity varies across the periodic table.

Key Terms and Concepts:

• Electronegativity: The ability of an atom to attract electrons in a bond.

• Periodic trend: Electronegativity increases across a period and decreases down a group.

Step-by-Step Guidance

1. Locate each element on the periodic table.

2. Recall the trend: electronegativity increases from left to right and decreases from top to bottom.

3. Identify which element is furthest down and to the left (lowest electronegativity).

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Q7. How many polar bonds does acetaldehyde have?

Background

Topic: Bond Polarity

This question tests your ability to identify polar bonds in a molecule based on differences in electronegativity.

Key Terms and Concepts:

• Polar bond: A covalent bond with unequal sharing of electrons due to a difference in electronegativity.

• Electronegativity difference: If the difference is significant (usually >0.4), the bond is polar.

Step-by-Step Guidance

1. Draw the Lewis structure for acetaldehyde (CH3CHO).

2. Identify all the bonds in the molecule.

3. Determine which bonds are between atoms with different electronegativities (e.g., C-H, C-O, C-C).

4. Count the number of polar bonds present.

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Q8. Which bond has its polarity incorrectly labeled?

Background

Topic: Bond Polarity and Delta Notation

This question tests your understanding of how to assign partial positive (δ+) and partial negative (δ−) charges based on electronegativity.

Key Terms and Concepts:

• δ+ (delta plus): Indicates the less electronegative atom in a polar bond.

• δ− (delta minus): Indicates the more electronegative atom.

Step-by-Step Guidance

1. For each bond, compare the electronegativities of the two atoms.

2. Assign δ+ to the less electronegative atom and δ− to the more electronegative atom.

3. Check which answer choice has the labels reversed or incorrect.

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Q9. What is the measure of the bond angle around the sulfur atom in the structure below?

Background

Topic: Molecular Geometry and Bond Angles

This question tests your ability to predict bond angles based on molecular shape using VSEPR theory.

Key Terms and Concepts:

• Bond angle: The angle between two bonds at an atom.

• VSEPR theory: Used to predict bond angles based on electron pair repulsion.

Step-by-Step Guidance

1. Draw or visualize the structure around the sulfur atom.

2. Determine the number of bonding and lone pairs on sulfur.

3. Use VSEPR theory to predict the ideal bond angle for that geometry.

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Q10. Which molecule or ion has only two resonance structures?

Background

Topic: Resonance Structures

This question tests your ability to recognize molecules or ions that can be represented by exactly two resonance structures.

Key Terms and Concepts:

• Resonance structure: Different valid Lewis structures for the same molecule or ion, differing only in the placement of electrons.

Step-by-Step Guidance

1. Draw the Lewis structures for each molecule or ion listed.

2. Count the number of valid resonance structures for each.

3. Identify which one has exactly two resonance structures.

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Q11. How many valence electrons in total does the OBr- ion have?

Background

Topic: Counting Valence Electrons in Ions

This question tests your ability to count the total number of valence electrons in a polyatomic ion, including the effect of the charge.

Key Terms and Concepts:

• Valence electrons: Electrons in the outermost shell of each atom.

• Polyatomic ion: A charged species composed of two or more atoms.

• Negative charge: Add one electron for each negative charge.

Step-by-Step Guidance

1. Find the number of valence electrons for oxygen and bromine from the periodic table.

2. Add these numbers together.

3. Add one more electron for the negative charge.

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Q12. Which compound is named incorrectly?

Background

Topic: Chemical Nomenclature

This question tests your ability to match chemical formulas with their correct names, especially for covalent and ionic compounds.

Key Terms and Concepts:

• Prefixes: Used for naming covalent compounds (mono-, di-, tri-, tetra-, etc.).

• Ionic compounds: Do not use prefixes.

Step-by-Step Guidance

1. Review the formula and name for each compound listed.

2. Check if the name matches the correct naming convention for that type of compound.

3. Identify any mismatches or incorrect use of prefixes.

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Q13. In order to balance the chemical equation that represents the rusting of an iron nail, what coefficient should be added in front of the iron(III) oxide?

Background

Topic: Balancing Chemical Equations

This question tests your ability to balance a chemical equation by adjusting coefficients to conserve atoms.

Key Terms and Concepts:

• Coefficient: The number placed in front of a compound to indicate the number of molecules or moles.

• Law of Conservation of Mass: Atoms are neither created nor destroyed in a chemical reaction.

Step-by-Step Guidance

1. Write the unbalanced equation: $4 \text{Fe}(s) + 3 \text{O}_2(g) \rightarrow \_\_\_ \text{Fe}_2\text{O}_3(s)$

2. Count the number of Fe and O atoms on both sides.

3. Adjust the coefficient in front of Fe2O3 so that the number of Fe and O atoms is equal on both sides.

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Q14. How many carbon atoms are in 3.85 mol of carbon?

Background

Topic: Avogadro's Number and Mole Concept

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

Key Terms and Formulas:

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

• Conversion formula: $\text{Number of atoms} = \text{moles} \times 6.022 \times 10^{23}$

Step-by-Step Guidance

1. Write down the number of moles given (3.85 mol).

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

3. Set up the calculation, but do not compute the final value yet.

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Q15. What is the formula weight of KCl?

Background

Topic: Formula Weights and Molar Mass

This question tests your ability to calculate the formula weight (molar mass) of an ionic compound using atomic masses.

Key Terms and Formulas:

• Formula weight: The sum of the atomic masses of all atoms in a formula unit.

• Atomic masses: Potassium (K): ~39.10 amu, Chlorine (Cl): ~35.45 amu

• Formula: $\text{Formula weight} = \text{mass of K} + \text{mass of Cl}$

Step-by-Step Guidance

1. Write down the atomic masses for K and Cl.

2. Add the two values together to get the formula weight.

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Q16. What is the mass of 3.81 mol of PH3?

Background

Topic: Moles and Molar Mass

This question tests your ability to convert moles of a compound to grams using its molar mass.

Key Terms and Formulas:

• Molar mass of PH3: P: ~30.97 g/mol, H: ~1.01 g/mol (3 H atoms)

• Formula: $\text{Mass} = \text{moles} \times \text{molar mass}$

Step-by-Step Guidance

1. Calculate the molar mass of PH3 by adding the atomic masses of P and H (3 times).

2. Multiply the molar mass by the number of moles given (3.81 mol).

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Q17. If 8 molecules of H2 react with 4 molecules of O2, how many molecules of water will be produced?

Background

Topic: Stoichiometry and Mole Ratios

This question tests your ability to use balanced chemical equations to determine the amount of product formed from given reactant quantities.

Key Terms and Concepts:

• Balanced equation: $2 \text{H}_2 + \text{O}_2 \rightarrow 2 \text{H}_2\text{O}$

• Mole ratio: The ratio of coefficients in the balanced equation.

Step-by-Step Guidance

1. Write the balanced equation and identify the mole ratio between H2, O2, and H2O.

2. Determine the limiting reactant by comparing the available molecules to the required ratio.

3. Use the limiting reactant to calculate the number of water molecules produced.

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Q18. How many molecules are in 2.40 moles of sucrose (C12H22O11)?

Background

Topic: Mole Concept and Avogadro's Number

This question tests your ability to convert moles of a compound to the number of molecules using Avogadro's number.

Key Terms and Formulas:

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

• Formula: $\text{Number of molecules} = \text{moles} \times 6.022 \times 10^{23}$

Step-by-Step Guidance

1. Write down the number of moles given (2.40 mol).

2. Multiply by Avogadro's number to find the total number of molecules.

3. Set up the calculation, but do not compute the final value yet.

Try solving on your own before revealing the answer!

Q19. Complete the following table using the given balanced equation and the initial quantities of reactants.

Background

Topic: Stoichiometry and Limiting Reactants

This question tests your ability to use a balanced equation to determine how many molecules of product are formed from given amounts of reactants.

Key Terms and Concepts:

• Balanced equation: $\text{N}_2 + 2 \text{O}_2 \rightarrow 2 \text{NO}_2$

• Limiting reactant: The reactant that is completely consumed first, limiting the amount of product formed.

Step-by-Step Guidance

1. Identify the initial quantities of N2 and O2.

2. Use the balanced equation to determine the stoichiometric ratio between reactants and product.

3. Determine which reactant is limiting and calculate the number of NO2 molecules formed.

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Q20. How many grams of AgCl are generated when 85 g of BaCl2 (molar mass = 208.23 g/mol) reacts?

Background

Topic: Stoichiometry and Mass-Mass Calculations

This question tests your ability to use stoichiometry to convert from grams of one reactant to grams of a product.

Key Terms and Formulas:

• Molar mass: The mass of one mole of a substance.

• Stoichiometric ratio: From the balanced equation: $\text{BaCl}_2 + 2 \text{AgNO}_3 \rightarrow 2 \text{AgCl} + \text{Ba(NO}_3)_2$

• Conversion steps: $\text{grams BaCl}_2 \rightarrow \text{moles BaCl}_2 \rightarrow \text{moles AgCl} \rightarrow \text{grams AgCl}$

Step-by-Step Guidance

1. Convert 85 g of BaCl2 to moles using its molar mass.

2. Use the stoichiometric ratio from the balanced equation to find moles of AgCl produced.

3. Multiply the moles of AgCl by its molar mass to find the mass of AgCl produced.

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Q21. Suppose the theoretical yield in a reaction is 10.5 g and the percent yield is 75.5%. What is the actual yield of product obtained?

Background

Topic: Percent Yield Calculations

This question tests your ability to calculate the actual yield from the theoretical yield and percent yield.

Key Terms and Formulas:

• Theoretical yield: The maximum amount of product that can be formed.

• Percent yield: $\text{Percent yield} = \frac{\text{actual yield}}{\text{theoretical yield}} \times 100$

• Rearranged formula: $\text{Actual yield} = \frac{\text{percent yield}}{100} \times \text{theoretical yield}$

Step-by-Step Guidance

1. Write down the theoretical yield (10.5 g) and percent yield (75.5%).

2. Convert the percent yield to a decimal by dividing by 100.

3. Multiply the decimal percent yield by the theoretical yield to find the actual yield.

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Q22. How many moles of sulfur trioxide are formed from 3 moles of sulfur dioxide and 3 moles of oxygen using the given balanced equation?

Background

Topic: Limiting Reactant and Stoichiometry

This question tests your ability to use a balanced equation to determine the amount of product formed from given reactant amounts, considering the limiting reactant.

Key Terms and Concepts:

• Balanced equation: $2 \text{SO}_2 + \text{O}_2 \rightarrow 2 \text{SO}_3$

• Limiting reactant: The reactant that is used up first and limits the amount of product formed.

Step-by-Step Guidance

1. Write the balanced equation and identify the mole ratio between SO2, O2, and SO3.

2. Calculate how many moles of SO3 can be formed from each reactant separately.

3. Determine which reactant is limiting and use it to find the moles of SO3 produced.

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Q23. What is the limiting reactant when 0.480 moles of Sn are mixed with 0.720 moles of HF?

Background

Topic: Limiting Reactant Calculations

This question tests your ability to identify the limiting reactant in a chemical reaction using mole ratios.

Key Terms and Concepts:

• Balanced equation: $\text{Sn} + 2 \text{HF} \rightarrow \text{SnF}_2 + \text{H}_2$

• Limiting reactant: The reactant that is completely consumed first.

Step-by-Step Guidance

1. Write the balanced equation and identify the mole ratio between Sn and HF.

2. Calculate how many moles of HF are needed to react with all the Sn present.

3. Compare the available moles of HF to the amount needed, and determine which reactant is limiting.

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Q24. In the balanced redox reaction: 2 C2H6(g) + 7 O2(g) → 4 CO2(g) + 6 H2O(g), which species is oxidized?

Background

Topic: Redox Reactions and Oxidation States

This question tests your ability to identify which species is oxidized in a redox reaction by assigning oxidation numbers.

Key Terms and Concepts:

• Oxidation: Loss of electrons (increase in oxidation number).

• Reduction: Gain of electrons (decrease in oxidation number).

• Oxidation number: A value assigned to an atom to indicate its degree of oxidation.

Step-by-Step Guidance

1. Assign oxidation numbers to each atom in the reactants and products.

2. Identify which atom increases in oxidation number (is oxidized).

3. Determine which species contains that atom.

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Q25. How many of the following represent oxidation half reactions?

Background

Topic: Redox Reactions and Half-Reactions

This question tests your ability to distinguish between oxidation and reduction half-reactions.

Key Terms and Concepts:

• Oxidation half-reaction: Shows the loss of electrons.

• Reduction half-reaction: Shows the gain of electrons.

Step-by-Step Guidance

1. Examine each half-reaction and determine if electrons are lost (products) or gained (reactants).

2. Count the number of half-reactions that show electrons as products (oxidation).

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