Chemistry Chapter on Chemical Bonds
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Questions and Answers

What is the correct number of valence electrons for an element with the electron configuration of 2, 8, 5?

  • 15
  • 8
  • 2
  • 5 (correct)

A polar covalent bond involves the equal sharing of electrons between two atoms.

False (B)

What type of bond is formed due to electrostatic attraction between charged ions?

ionic

An alpha particle is equivalent to a(n) _____ nucleus.

<p>helium</p> Signup and view all the answers

Match the following molecules with their shapes:

<p>CCl4 = Tetrahedral N2 = Linear NF3 = Trigonal Pyramidal H2O2 = Angular BH3 = Trigonal Planar</p> Signup and view all the answers

Which of the following best describes a covalent bond?

<p>Electrons are shared between atoms (A)</p> Signup and view all the answers

In a polar covalent bond, electrons are shared equally between two atoms.

<p>False (B)</p> Signup and view all the answers

What type of bond is formed between a metal and a nonmetal, typically resulting in ions?

<p>ionic</p> Signup and view all the answers

In a Lewis structure, a single line between two atoms represents a ______ bond.

<p>covalent</p> Signup and view all the answers

Match the following formulas with their corresponding names:

<p>PbS = lead(II) sulfide SnSO3 = tin(II) sulfite (note: 6H2O is a hydrate, usually ignored in matching) SF6 = sulfur hexafluoride Na2O2 = sodium peroxide</p> Signup and view all the answers

What is the name of the 3D shape of nitrogen trifluoride (NF3)?

<p>Trigonal pyramidal (A)</p> Signup and view all the answers

The balanced reaction for FeCl2 + KMnO4 + HCl -> FeCl3 + KCl + MnCl2 + H2O is 10FeCl2 + 2KMnO4 + 24HCl ->10FeCl3 + 2KCl + 2MnCl2 + 12H2O

<p>True (A)</p> Signup and view all the answers

Using electronegativity, what type of bond is present in BrCl?

<p>polar covalent</p> Signup and view all the answers

What is the average atomic mass of element X given the following isotopic data: Isotope 302X (12.64%, 302.04 u), 304X (18.23%, 304.12 u), 306X (69.13%, 305.03 u)?

<p>304.96 u (A)</p> Signup and view all the answers

Element X has two isotopes, X-56 (56.0 u) and X-59 (59.0 u). If the atomic mass of X is 57.3 u, what is the percent abundance of the isotope X-56?

<p>56.67%</p> Signup and view all the answers

Transition metals are located in Group 1 of the periodic table.

<p>False (B)</p> Signup and view all the answers

Calcium has ______ valence electrons.

<p>2</p> Signup and view all the answers

Which of the following ions is most likely to be stable?

<p>Mg2+ (C)</p> Signup and view all the answers

Which of the following electron configurations represents an alkaline earth metal?

<p>2, 8, 8, 2 (C)</p> Signup and view all the answers

Which of the following pairs illustrates that the first atom has a larger atomic radius than the second?

<p>Na, Li (C)</p> Signup and view all the answers

Match the following elements with their ionization equations:

<p>Calcium = $Ca \rightarrow Ca^{2+} + 2e^-$ Phosphorus = $P \rightarrow P^{3+} + 3e^-$</p> Signup and view all the answers

What happens to a crystal added to a supersaturated solution?

<p>The crystal will cause more solute to precipitate out of the solution. (B)</p> Signup and view all the answers

In the reaction Cu2O(s) + Cu2S(s) -> 4Cu(s) + SO2(g), copper(I) sulfide is the limiting reagent if 250 kg of copper(I) oxide is reacted with 129 kg of copper(II) sulfide.

<p>True (A)</p> Signup and view all the answers

What is the concentration in mol/L of a solution made by dissolving 25 g of sodium chloride in 100 g of water, given that the density of the solution is 1.15 g/mL?

<p>4.26 mol/L</p> Signup and view all the answers

To make 8.00 L of a 1.50 mol/L solution from concentrated phosphoric acid which has a concentration of 18.0 mol/L, you need a volume of ______ L of the concentrated acid.

<p>0.667</p> Signup and view all the answers

Match the following salts with their correct balanced dissociation equation:

<p>KCl(s) = K+(aq) + Cl-(aq) Ba(NO3)2(s) = Ba2+(aq) + 2NO3-(aq) (NH4)2SO3(s) = 2NH4+(aq) + SO32-(aq)</p> Signup and view all the answers

What mass of lead is present in a 60,000 L pool with a lead concentration of 4.8 ppm (parts per million)?

<p>0.288 kg (C)</p> Signup and view all the answers

What is the temperature required to create a saturated solution of KNO3 using 50 g of the salt and 50 g of water?

<p>Approximately 78 degrees Celsius</p> Signup and view all the answers

Which of the following correctly represents the net ionic equation for the reaction Pb(NO3)2(aq) + 2KI(aq)?

<p>Pb2+(aq) + 2I-(aq) -&gt; PbI2(s) (A)</p> Signup and view all the answers

Which of the following is the correct balanced neutralization reaction for H3PO4 (aq) + NaOH (aq)?

<p>H3PO4 (aq) + 3NaOH (aq) -&gt; Na3PO4 (aq) + 3H2O (l) (D)</p> Signup and view all the answers

According to the Arrhenius definition, an acid is a substance that donates a proton (H+) in water.

<p>True (A)</p> Signup and view all the answers

What is the hydronium concentration of a hydrobromic acid solution with a pH of 0.7?

<p>$10^{-0.7}$ mol/L or approximately 0.2 mol/L</p> Signup and view all the answers

A strong acid, such as hydrochloric acid, completely dissociates into ions in water, whereas a weak acid, such as acetic acid, only ______ dissociates.

<p>partially</p> Signup and view all the answers

Match the following gas laws with their correct mathematical expression:

<p>Boyle's Law = $P_1V_1 = P_2V_2$ Charles's Law = $V_1/T_1 = V_2/T_2$ Combined Gas Law = $P_1V_1/T_1 = P_2V_2/T_2$</p> Signup and view all the answers

What is the pH of a 0.00005 mol/L HNO3 solution?

<p>4.3 (A)</p> Signup and view all the answers

At Standard Temperature and Pressure (STP), one mole of any gas occupies a volume of 24.4 L.

<p>False (B)</p> Signup and view all the answers

If 1.00 L of nitrogen gas is reacted with 3.0 L of fluorine gas, how many litres of nitrogen trifluoride gas will be produced?

<p>2.0 L</p> Signup and view all the answers

Which of the following reactions is an example of a decomposition reaction?

<p>2H2O -&gt; 2H2 + O2 (A)</p> Signup and view all the answers

In the reaction C4H10 + 15/2 O2 -> 5CO2 + 5H2O, the products are carbon dioxide and water.

<p>True (A)</p> Signup and view all the answers

What is the empirical formula for a compound with the molecular formula C6H6O2?

<p>C3H3O</p> Signup and view all the answers

A solution where more solute is dissolved than is normally possible at a given temperature is said to be ______.

<p>supersaturated</p> Signup and view all the answers

Match the following ionic compounds with corresponding separated ions:

<p>KCl = K+ + Cl- Ba(NO3)2 = Ba2+ + 2NO3- (NH4)2SO3 = 2NH4+ + SO32-</p> Signup and view all the answers

In the reaction Pb(NO3)2 (aq) + 2KI (aq) -> PbI2 (s) + 2KNO3 (aq), which ion(s) are spectator ions?

<p>K+ and NO3- (A)</p> Signup and view all the answers

What is the pH of a solution with a hydronium ion concentration of 1.30 x 10^-4 M?

<p>3.89</p> Signup and view all the answers

Strong acids do not dissociate completely in solution.

<p>False (B)</p> Signup and view all the answers

Flashcards

Average Atomic Mass

The average mass of all the naturally occurring isotopes of an element, taking into account their relative abundances.

Isotopes

Atoms of the same element that have the same number of protons but a different number of neutrons.

Valence Electrons

The outermost electrons in an atom that are involved in chemical bonding.

Ions

Atoms or molecules that have gained or lost electrons, resulting in a net positive or negative charge.

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Ionization Energy

The energy required to remove one electron from a neutral atom in its gaseous state.

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Electron Affinity

The tendency of an atom to gain an electron.

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Metallic Character

The measure of an element's ability to lose electrons and form positive ions.

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Electron Affinity

The energy required to add one electron to a neutral atom in its gaseous state.

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Limiting Reactant

The reactant that is completely consumed in a chemical reaction, limiting the amount of product that can be formed.

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Theoretical Yield

The maximum amount of product that can be formed in a chemical reaction, based on the amount of limiting reactant.

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Percentage Yield

The ratio of the actual yield (the amount of product actually obtained) to the theoretical yield, expressed as a percentage.

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Saturated Solution

A solution that contains the maximum amount of solute that can dissolve at a given temperature.

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Unsaturated Solution

A solution that contains less solute than the maximum amount that can dissolve at a given temperature.

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Supersaturated Solution

A solution that contains more solute than the maximum amount that can dissolve at a given temperature. It is unstable and can cause the excess solute to precipitate out.

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Percent by Mass Concentration

The concentration of a solution expressed as the mass of solute per 100 g of solution.

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Molarity (mol/L)

The concentration of a solution expressed as the number of moles of solute per liter of solution.

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Neutralization reaction

A reaction where an acid and a base react to form salt and water, resulting in a neutral solution. The number of H+ ions from the acid equals the number of OH- ions from the base.

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Bronsted-Lowry acid

An acid that donates a proton (H+) to a base in a reaction.

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Bronsted-Lowry base

A base that accepts a proton (H+) from an acid in a reaction.

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Arrhenius acid

A substance that dissolves in water to produce H+ ions, which are responsible for the acidity of the solution.

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Arrhenius base

A substance that dissolves in water to produce OH- ions, which are responsible for the basicity of the solution.

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pH

The measure of the concentration of hydrogen ions (H+) in a solution. It is a logarithmic scale, with lower pH values indicating higher acidity.

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Vapor pressure

The pressure exerted by the vapor of a liquid in equilibrium with its liquid phase at a given temperature.

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Charles's Law

The volume of a gas at constant pressure is directly proportional to its absolute temperature.

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Covalent bond

A covalent bond forms when two atoms share electrons equally, resulting in a molecule with a non-polar character. The atoms involved have similar electronegativity values.

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Polar covalent bond

A polar covalent bond forms when two atoms share electrons unequally because of a difference in electronegativity. The atom with higher electronegativity attracts the shared electrons more strongly, creating a partial negative charge on that atom and a partial positive charge on the other atom.

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Ionic bond

An ionic bond forms when one atom completely transfers electrons to another atom, resulting in the formation of oppositely charged ions that attract each other via electrostatic forces. This occurs between atoms with a large difference in electronegativity values.

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Predicting bond type using electronegativity

The electronegativity difference between two atoms can be used to predict the type of bond between them. If the difference is less than 0.5, the bond is generally considered non-polar covalent. If the difference is between 0.5 and 1.7, the bond is generally considered polar covalent. If the difference is greater than 1.7, the bond is generally considered ionic.

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Lewis structure

A Lewis structure is a diagram that shows the bonding between atoms in a molecule, including the lone pairs of electrons. It helps to visualize the arrangement of electrons in a molecule.

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3D shape of a molecule

The shape of a molecule is determined by the arrangement of the atoms and their bonding electrons. There are several common 3D shapes for molecules, such as linear, trigonal planar, tetrahedral, pyramidal, bent, and octahedral.

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Balancing chemical equations

Balancing a chemical equation ensures that the number of atoms of each element on the reactant side is equal to the number of atoms of each element on the product side. This follows the law of conservation of mass, which states that matter cannot be created or destroyed in a chemical reaction.

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Predicting chemical reactions

A chemical reaction occurs when reactants are chemically transformed into products. Predicting whether a reaction will occur can be done by considering various factors such as the nature of the reactants, the presence of catalysts, and the enthalpy change of the reaction. A word equation is a simple representation of a chemical reaction using the names of the reactants and products.

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Decomposition Reaction

A chemical reaction where a single reactant breaks down into two or more products. The reaction is typically driven by heat or electricity.

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Single Displacement Reaction

A chemical reaction where one element replaces another in a compound. The more reactive element displaces the less reactive element.

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Synthesis Reaction

A chemical reaction where two reactants combine to form a single product.

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Double Displacement Reaction

A chemical reaction where two reactants exchange ions to form two new products.

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Combustion Reaction

A chemical reaction where a substance reacts with oxygen, producing heat and light.

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Expansion reaction

A chemical reaction in which there is an increase in volume.

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Contraction reaction

A chemical reaction in which there is a decrease in volume.

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Study Notes

Exam Review Questions (Part 1)

  • Classifying Properties: Silver tarnishing, gold's conductivity, potassium iodide dissolving in water, oxygen supporting combustion, and mercury evaporating are classified as chemical or physical properties. Classify the changes as chemical (C) or physical (P). Examples include twisting copper, burning coal, melting wax, and fermentation.
  • Classifying Substances: Classify each of the following as a pure substance (PS), a solution (S), or a mechanical mixture (MM). Examples include plain Jell-O, neon gas, vinegar, Raisin Bran cereal, calcium carbonate, and air.
  • Smallest Particle in Covalent Compounds: The smallest particle of a covalently-bonded compound is a molecule. Pure substances are always homogeneous.
  • J.J. Thomson's Model: J.J. Thomson discovered the electron while studying cathode rays. Thomson's atomic model describes the atom as a sphere of positive charge with negatively charged electrons embedded within it.
  • Rutherford's Experiment: Rutherford's experiment disproved Thomson's model and introduced the concept of a nucleus. It focused on positive alpha particles and a thin gold foil, which demonstrated a dense, positively charged nucleus in the center of the atom. This nucleus contains most of the atom's mass.
  • Isotopes: Atoms of the same element with different numbers of neutrons are isotopes. Examples provided include 168X, 179Y, 178Z. Isotopes of sodium-23 (2311Na) and phosphorus-31 (3115P) are illustrated by Bohr-Rutherford diagrams.
  • Average Atomic Mass: The average atomic mass of an element is calculated by considering the abundance and mass of each isotope. Using the provided data, the average atomic mass of element X is determined.

Exam Review Questions (Part 2)

  • Isotope Abundance: Element X, composed of isotopes X-56 and X-59, has an atomic mass of 57.3 u. Calculate the % abundance for each.
  • Periodic Table Groups: Label the following groups: metals, nonmetals, transition metals, halogens, noble gases, alkali metals, metalloids, inner transition metals.
  • Valence Electrons: Determine the number of valence electrons for calcium (Ca) and fluorine (F).
  • Stable Ions: Identify the stable ions among Mg2+, K+, B3-, and S2-.
  • Electron Configuration and Elements: Determine the alkaline earth metal, order of increasing atomic radii, atom with the lowest ionization energy, and elements in the same family from given electron configurations.
  • Trends on the Periodic Table: Understand trends in atomic size, ionization energy, and electron affinity, based on occupied electron shells and the nuclear charge.
  • Ionization Equations: Provide ionization equations for calcium and phosphorus.
  • Types of Bonds: Explain the difference between covalent, polar covalent, and ionic bonds.

Exam Review Questions (Part 3)

  • Bond Type Determination: Determine the bond type in the following compounds using electronegativity values. Determine the compound type for BrCl, NI3 and Al2O3, and CS2.
  • Lewis Structures and Shapes: Draw Lewis structures and determine the 3-dimensional shape of compounds like carbon tetrachloride, nitrogen trifluoride, hydrogen peroxide, and boron trihydride.
  • Chemical Formulas and Names: Complete the table of chemical formulas, names, and associated information. (Provides a list of compounds)

Exam Review Questions (Part 4)

  • Balancing Chemical Reactions: Balance the provided chemical reactions.
  • Reaction Prediction: Predict if reactions will occur for given pairs of reactants and complete the balanced equations, or indicate NR (no reaction).
  • Type of Reactions: Complete and balance chemical reactions and classify them as synthesis, decomposition, single displacement, double displacement, combustion, or neutralization).
  • Single Displacement Reactions: Describe single displacement reactions using metals W, X, Y, and Z based on provided observations.

Exam Review Questions (Part 5)

  • Limiting and Excess Reactants: Explain the difference between limiting and excess reactants, actual yield, theoretical yield, empirical formulas, and molecular formulas.
  • Percent Composition: Calculate the percent composition of sodium thiosulfate (Na2S2O3).
  • Empirical Formula Determination: Determine the empirical formula for a compound with given percentage composition of potassium, carbon, and oxygen.
  • Molecular Formula Determination: Calculate the molecular formula of hydroquinone, given its molecular mass and elemental composition.
  • Stoichiometry Calculations: Use the equation 2H2(g) + O2(g) → 2H2O(g) to solve stoichiometry problems involving moles, grams, and volumes.
  • Neutralization Reaction: Determine the mass of sodium carbonate needed to produce a given volume of carbon dioxide according to a provided chemical equation.
  • Gas Law Calculations: Determine the volume of water vapor produced in a combustion reaction, given the mass of reactants and conditions. Calculate limiting reagents. Perform stoichiometry calculations.

Exam Review Questions (Part 6)

  • Solubility Curves: Use solubility curves to determine the solubility of KNO3 at a given temperature and to estimate the temperature required for creating a saturated solution. Calculate the mass of KNO3 that precipitates when a saturated solution cools.
  • Concentration Calculations: Calculate the percentage by mass and molarity of sodium chloride solutions and concentration of acetic acid in solution, given density and mass percent data.
  • Balanced Dissociation Equations: Complete the balanced dissociation equations for given salts (e.g., KCl, Ba(NO3)2, (NH4)2SO3).
  • Precipitation Reactions: Complete balanced equations for precipitation reactions between given pairs of reactants.
  • Net Ionic Equations: Write total and net ionic equations for neutralization and precipitationreactions including sate symbols (s,aq).

Exam Review Questions (Part 7)

  • Bronsted-Lowry Acid-Base Definitions: Write an equation illustrating the autoionization of water and describe acid-bases in terms of Bronsted-Lowry theory. Explain the difference between strong and weak acids with regards to their dissociation.
  • Properties of Acids and Bases: List three properties of acids and three properties of bases.
  • pH Calculations: Calculate the pH of hydrochloric acid and nitric acid solutions.
  • Concentration Calculation: Determine the hydronium concentration of a hydrobromic acid solution from its pH reading.
  • Neutralization Reactions: Calculate the volume of a sodium hydroxide solution needed to neutralize a given volume of hydrobromic acid solution. Include the balanced chemical equation.
  • Titration Calculation: Calculate the concentration of a base required to titrate a given volume of an acid solution. Include the balanced chemical equation.
  • Heating Curves: Sketch and label a heating curve of benzene (mp = 5.5°C; bp = 80.1°C). Interpret relationships in heating curves.
  • Gas Law Calculations: Calculate the volume of a gas at a different temperature via gas law relationships.
  • Gas Collection Over Water: Calculate the number of moles of a gas collected over water, considering the vapour pressure of water. Perform corrections for water pressure.
  • Molar Mass Calculations: Calculate the molar mass (or density) of given quantities of gases, adjusting for various conditions of temperature, pressure, volume, and mass.
  • Stoichiometric Calculations: Solve stoichiometry problems in gas-phase reactions, involving gases like nitrogen gas, fluorine gas, and nitrogen trifluoride and adjust to specified temperature and pressure conditions.

Exam Review Questions (Part 8)

  • Balanced Chemical Equations: Balance chemical equations involving various elements and compounds. Examples are provided on the page.

Exam Review Questions (Page 9)

  • Acid-Base Reactions: Perform and balance various acid-base reactions with solutions (e.g., single displacement, double displacement).
  • Solubility Rules: Refer to solubility rules.
  • Solution Calculations: Calculate various concentrations and volumes in solution reactions.
  • Conversions: Perform conversions between various quantity types (e.g., moles from volume, mass to volume in solution). Include temperature conversions using the appropriate unit conversions.
  • Stoichiometry (various): Carry out stoichiometric calculations in balanced chemical equations, focusing on limiting reagents, theoretical yields, and percentages of yield. Perform calculations related to amounts of substances converted to other substances in reactions.

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Test your understanding of chemical bonds, electron configurations, and molecular shapes with this comprehensive quiz. Explore topics such as polar covalent bonds, ionic bonds, and Lewis structures. This quiz is perfect for students preparing for exams in chemistry.

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