First and Second Semester Chemistry Final Exam Review Flashcards

Qualitative and Quantitative Observations in Chemistry

Observations are fundamental to the study of chemistry and can be categorized into two primary types: qualitative and quantitative. Qualitative observations rely on the human senses to describe the properties of a substance without using numbers. Examples of qualitative observations include identifying the color of an object, such as when a metal or compound is described as being red, green, or having a specific hue. Other qualitative observations involve the sense of touch, such as noting that a reaction vessel feels warm to the hand, or the sense of smell. Descriptions like a metal being shiny (luster) or malleable (able to be hammered into thin sheets) are also qualitative in nature.

Quantitative observations, in contrast, are characterized by the use of numerical data and measurements to describe a substance or phenomenon. These observations often require scientific instruments. Examples of quantitative observations include measuring the mass of a metal sample as being exactly $10\,g$, determining the volume of a liquid in a container to be $10\,mL$, or counting a specific number of objects, such as observing exactly 2 dogs. Temperature readings, such as noting that the outside temperature is $30\,^\circ C$, are also quantitative observations as they provide a numerical value and specific units.

Laboratory Equipment and Measurement Techniques

A critical skill in the chemistry laboratory is the correct identification and use of various pieces of equipment. Students should be able to recognize and name items including graduated cylinders, beakers, test tube tongs, crucible tongs, scoopulas, and balances. For measuring physical properties with precision, specific tools are preferred. The balance is identified as the best piece of equipment for measuring the mass of an object, with the standard unit for mass being grams $(g)$. For measuring the volume of liquids, the graduated cylinder is the most appropriate and accurate tool, with measurements typically recorded in milliliters $(mL)$.

When reading the volume from liquid measuring devices like a graduated cylinder or beaker, it is essential to look at the meniscus, which is the curve seen at the top of the liquid level. The measurement should be taken at the bottom of the meniscus. For optimal precision, one should always estimate or "guess" one additional digit beyond the smallest marked division on the instrument. For example, if the smallest division is $1\,mL$, the volume should be estimated to the nearest $0.1\,mL$. Specific examples of volume readings include $7.95\,mL$ (where the smallest division is $0.1\,mL$ and the last digit is estimated), $34.5\,mL$ (with an estimated tenths place), $24\,mL$ (measured to the ones place), and $43.0\,mL$ (where the smallest division is $10\,mL$ and the units place is estimated).

Scientific Notation and Significant Figures

Scientific notation is a method for expressing very large or very small numbers in a more manageable format, typically as a number between 1 and 10 multiplied by a power of ten. Examples include rewriting $0.003760$ as $3.760 \times 10^{-3}$ and $79850$ as $7.985 \times 10^4$. Significant figures are used to indicate the precision of a measurement. The number of significant figures is determined by set rules: non-zero digits are always significant, zeros between non-zero digits are significant, leading zeros are never significant, and trailing zeros are significant only if there is a decimal point. For instance, the number $0.003760$ contains 4 significant figures (the leading zeros are placeholders, but the trailing zero after the decimal is significant), while $79850$ is often interpreted as having 4 significant figures, depending on the context of the precision of the trailing zero.

Calculations involving significant figures must follow specific rules to maintain proper precision in the final result. In multiplication and division, the answer should have the same number of significant figures as the measurement with the fewest significant figures. For example, multiplying $101.60$ (5 significant figures) by $2.500$ (4 significant figures) gives a raw result of 254, which must be adjusted to 4 significant figures as $254.0$ or $2.540 \times 10^2$. In addition and subtraction, the result is rounded to the same number of decimal places as the measurement with the fewest decimal places. For instance, adding $0.08477$ to $1.50 \times 10^{-4}$ (which is $0.000150$) would be rounded based on the least precise decimal place. Another example provided is $0.08477 \div 150$, which gives $12.7155$; however, if $150$ has three significant figures, the final result should be rounded to three significant figures, resulting in $0.000565$ or $5.65 \times 10^{-4}$.

Classifying Matter: Substances and Mixtures

Matter can be classified into several categories based on its composition. A pure substance can be either an element or a compound. Elements consist of only one kind of atom, such as Lithium $(Li)$, Nitrogen $(N)$, Helium $(He_2)$, or Hydrogen $(H_2)$. Compounds are formed when two or more different kinds of atoms are chemically bonded together in fixed proportions, with examples including Magnesium Chloride $(MgCl_2)$, Water $(H_2O)$, and Glucose $(C_6H_{12}O_6)$. Mixtures consist of two or more different substances physically combined in the same container but not chemically bonded. Mixtures can be further classified; for example, a mixture of elements like $(He + He_2)$ or $(2Li + Si)$, a mixture of compounds such as $(2H_2O + MgCl_2)$, or a mixture of both elements and compounds.

Analyzing chemical formulas can reveal the number of molecules and atoms present. For instance, in the formula for $3\,H_2SO_4$, there are 3 total molecules and a total of 21 atoms (3 sulfur atoms, 6 hydrogen atoms, and 12 oxygen atoms), and it is classified as a compound. A mixture of compounds like $(MgCl_2 + CsBr)$ contains 2 total molecules and 5 total atoms. The combination of $3\,H_3PO_4 + 2\,CO_2$ results in 5 total molecules and 28 total atoms, representing a mixture of compounds. Finally, a mixture of elements such as $(2\,Na + N_2 + O_2)$ consists of 4 molecules or discrete units and 6 total atoms.

Physical and Chemical Changes

Physical changes involve a transformation in the form or state of a substance without altering its chemical identity. Examples of physical changes include phase changes such as melting (solid to liquid), boiling (liquid to gas), freezing (liquid to solid), heating up, and condensing (gas to liquid). Other examples of physical changes are dissolving a substance, like making a mixture of sugar and water (e.g., Kool-Aid), breaking or crumbling a material, and dampening or wetting a substance. These changes are characterized by the substance remaining chemically the same throughout the process.

Chemical changes, or chemical reactions, result in the formation of one or more new substances with different chemical properties. There are five primary indications that a chemical change has occurred: the substance is flammable and catches fire, a gas is produced (often observed as bubbles or a new odor), a significant color change occurs, a precipitate is formed (a solid that emerges from the mixing of two liquids), or light is given off during the process. Examples clearly distinguished include a snowball melting (physical change), a window breaking (physical change), a Bunsen burner burning (chemical change), and sugar dissolving (physical change).

Factors Affecting Reaction Rates and Catalysts

The rate at which a chemical reaction occurs can be influenced by several factors, many of which can be explained by the collision theory, which states that reactions occur when particles collide with sufficient energy and correct orientation. Increasing the concentration of reactants leads to more frequent collisions, thereby speeding up the reaction. Similarly, raising the temperature of the reaction mixture increases the kinetic energy of the particles, leading to more frequent and more energetic collisions, thus increasing the reaction rate. Decreasing the size of reactant particles (which increases their surface area) also accelerates the reaction. Conversely, decreasing temperature, increasing the size of reactant particles (decreasing surface area), or lowering the concentration of reactants will generally slow down a chemical reaction.

A catalyst is a substance that can be added to a chemical reaction to increase the reaction rate. The key role of a catalyst is that it speeds up the reaction without being consumed itself in the process. It does not increase the temperature or change the concentration of the reactants, but rather provides an alternative pathway with lower activation energy for the reaction to occur.

States of Matter and Phase Changes

Matter typically exists in three primary states: solid, liquid, and gas, each with distinct physical characteristics related to atomic and molecular arrangement and motion. In a solid state, particles are slowed down and tightly packed together, often in a regular arrangement, where they primarily vibrate in fixed positions. Liquids have particles that are at a medium distance from each other, allowing them to move and slide past one another. The gas state is characterized by particles being farthest apart, moving very quickly, and rarely bumping into each other due to the large distances between them.

Phase changes between these states occur at specific temperatures and involve the absorption or release of energy. The heating curve of a substance like water illustrates these changes. Segment A shows the solid state (e.g., ice). Segment B represents the plateau where melting (solid to liquid) or freezing (liquid to solid) occurs; the temperature remains constant at the melting/freezing point ($0\,^\circ C$ for water). Segment C displays the liquid state. Segment D shows another plateau where boiling or vaporization (liquid to gas) and condensation (gas to liquid) take place; the temperature stays constant at the boiling point ($100\,^\circ C$ for water). Finally, segment E represents the gas state. The segments on the graph showing a substance changing state are the horizontal plateaus, segments B and D.

Thermochemistry: Specific Heat and Phase Change Energy

Specific heat ($SH$) is defined as the amount of energy required to raise the temperature of $1\,g$ of a substance by $1\,^\circ C$, with common units being $J/g\,^\circ C$ or $cal/g\,^\circ C$. Substances with a low specific heat capacity undergo rapid changes in temperature when heat is added or removed, while those with a high specific heat require more energy to change temperature or to maintain a particular temperature. For example, if copper has a specific heat of $0.385\,J/g\,^\circ C$ and iron has a specific heat of $0.449\,J/g\,^\circ C$, copper will reach a higher temperature if the same amount of heat energy is added to equal masses of both metals because it has a lower specific heat.

The heat of fusion ($H_f$) is the amount of energy needed to melt or freeze $1\,g$ of a substance at its melting point, with units of $J/g$ or $cal/g$. The heat of vaporization ($H_v$) is the energy required to boil or condense $1\,g$ of a substance at its boiling point, also in $J/g$ or $cal/g$. For water, specific heat is $4.184\,J/g\,^\circ C$, the heat of fusion is $335\,J/g$, and the heat of vaporization is $2260\,J/g$. Calculations for heat ($Q$) involve the formulas $Q = SH \times m \times \Delta t$ for temperature changes and $Q = m \times H_f$ or $Q = m \times H_v$ for phase changes.

Specific heat calculations can be performed as follows: to find the specific heat of an unknown substance where $67.8\,g$ absorbs $478\,J$ and the temperature increases from $23\,^\circ C$ to $45\,^\circ C$ ($\Delta t = 22.0\,^\circ C$), use the equation $478\,J = SH \times 67.8\,g \times 22.0\,^\circ C$, yielding $SH = 0.320\,J/g\,^\circ C$. To find the heat needed to boil $67\,g$ of water at $100\,^\circ C$, use $Q = 67\,g \times 2260\,J/g = 151420\,J$, which rounds to approximately $150,000\,J$ (using 2-3 significant figures). For calculating the energy to raise $5.0\,g$ of water from $35\,^\circ C$ to $75\,^\circ C$ ($\Delta t = 40\,^\circ C$), the calculation is $Q = 4.184\,J/g\,^\circ C \times 5.0\,g \times 40\,^\circ C = 836.8\,J$, rounded to about $840\,J$.

In laboratory simulations, changing variables can significantly impact temperature changes. Increasing the amount of water in a calorimeter while keeping other factors constant will decrease the change in water temperature ($\Delta t$) because more heat is needed to raise a larger mass of water. Conversely, increasing the mass of the metal will increase the $\Delta t$ of the water, as the higher mass metal stores and releases more energy. If a metal with the same mass but a higher specific heat is used, the $\Delta t$ of the water will also increase because the higher specific heat metal holds and releases more heat at a given temperature.

Atomic Structure and Electron Configuration

The atomic structure of an element is defined by its number of protons, neutrons, and electrons. The atomic number of an element, such as 29 for copper, specifically refers to the number of protons in its nucleus and determines the element's identity. The mass number is the sum of the protons and neutrons. The average atomic mass, such as $63.54$ for copper, is the weighted average of the masses of all naturally occurring isotopes of that element. Isotopes are atoms of the same element that have the same number of protons but different numbers of neutrons, resulting in different mass numbers.

Subatomic particles can be identified from isotope notation. For an isotope like $^{231}_{90}Th$, the atomic number is 90 (number of protons), the mass number is 231, and the number of neutrons is found by subtracting the atomic number from the mass number ($231 - 90 = 141$). In a neutral atom, the number of electrons equals the number of protons, so there are 90 electrons. For a charged ion like $^{54}_{26}Fe^{2+}$, there are 26 protons, and the number of neutrons is $54 - 26 = 28$. The $2+$ charge indicates that the atom has lost 2 electrons, resulting in $26 - 2 = 24$ electrons. In a chlorine ion like $Cl^-$ with a mass number of 35 and an atomic number of 17 ($^{35}_{17}Cl^-$), there are 17 protons, 18 neutrons ($35 - 17$), and 18 electrons because the $1-$ charge means it gained one electron.

Changing subatomic particles has specific effects: changing the number of protons changes the element's identity, the mass number (by 1), and adds a positive charge. Changing the number of neutrons changes the mass number by 1, while the element's identity and charge remain the same (forming an isotope). Changing the number of electrons keeps the element and mass number the same but changes the charge, creating an ion.

Electron configurations describe the arrangement of electrons in an atom's orbitals and energy levels. Orbitals are categorized as s, p, d, and f, with each energy level having a specific number of these orbitals (s: 1, p: 3, d: 5) and a maximum number of electrons (s: 2, p: 6, d: 10). For Silicon ($Si$), the configuration is $1s^2 2s^2 2p^6 3s^2 3p^2$, or in noble gas notation, $[Ne] 3s^2 3p^2$. Examples of correct electron configurations include Bismuth ($Bi$): $[Xe] 6s^2 4f^{14} 5d^{10} 6p^3$. Positive ions like $Na^+$, $Mg^{2+}$, or $Al^{3+}$ all have the noble gas electron configuration of $1s^2 2s^2 2p^6$ because they have lost their valence electrons.

Nuclear Chemistry and Reactions

Nuclear chemistry involves changes in the nucleus of an atom. Radioactive decay occurs spontaneously and can take several forms. Alpha decay involves the emission of an alpha particle ($\alpha$ or $^4_2He$), as seen in the reaction $^{222}_{86}Rn \rightarrow ^{214}_{84}Po + \alpha$. Beta decay involves the emission of an electron-like beta particle ($\beta$), such as in $^{228}_{90}Th \rightarrow ^{228}_{91}Pa + \beta$. Other nuclear processes include positron emission (e.g., $^{65}_{30}Zn \rightarrow ^{65}_{29}Cu + ^{0}_{+1}e$) and electron capture. Artificial transmutation is the process of bombarding a nucleus to create new elements or isotopes, such as $^{238}_{92}U + ^2_1H \rightarrow ^{239}_{94}Pu + ^1_0n$.

Fission is a nuclear reaction where a large nucleus splits into smaller nuclei, releasing energy and neutrons, a process illustrated by $^{239}_{94}Pu + ^1_0n \rightarrow ^{106}_{44}Ru + ^{131}_{50}Sn + 3^1_0n + \text{energy}$. Fusion is the process where small nuclei combine to form a larger nucleus, releasing significant energy, such as $^2_1H + ^2_1H \rightarrow ^3_2He + ^1_0n + \text{energy}$. Balancing nuclear equations requires that the total sum of mass numbers and the total sum of atomic numbers are equal on both sides of the equation. For example, in the reaction $^{228}_{90}Th \rightarrow ^{4}_{2}He + ^{224}_{88}Ra$, the mass numbers balance ($228 = 4 + 224$) and the atomic numbers balance ($90 = 2 + 88$).

Periodic Table Families, Trends, and Properties

The periodic table is organized into families with similar chemical and physical properties because they share the same number of valence electrons. Key families include the alkali metals (Group 1, 1 valence electron), alkaline earth metals (Group 2, 2 valence electrons), halogens (Group 17, 7 valence electrons), and noble gases (Group 18, 8 valence electrons). For example, Magnesium ($Mg$) is an alkaline earth metal, Rubidium ($Rb$) is an alkali metal, and Iodine ($I$) is a halogen. Metals are located to the left of the metalloid staircase, nonmetals are to the right, and metalloids bridge the two groups. Metals are generally shiny, conductive of electricity, malleable, and reactive with materials like $HCl$ and $CuCl_2$.

Periodic trends include ionization energy and atomic radius. Ionization energy is the energy required to remove one valence electron from an atom. Ionization energy generally increases as you move from left to right across a period and from bottom to top within a group; thus, Helium has the highest ionization energy. For example, the elements sodium, aluminum, sulfur, and argon are in order of increasing ionization energy ($Na < Al < S < Ar$). Atomic radius increases as you move down a group (due to adding energy levels and decreasing attractive force) and decreases as you move across a period (due to more protons in the same energy level increasing attractive force).

Other properties include the formation of acids when nonmetal oxides are added to water, such as with Phosphorus ($P$). Valence charges can also be determined based on an element's group: alkaline earth metals like $Mg$ have a $2+$ charge, alkali metals like $Rb$ have a $1+$ charge, and halogens like $I$ have a $1-$ charge. Specific locations can be identified like Iron ($Fe$) being in period 4 as a transition metal, Chlorine ($Cl$) in period 3 as a halogen, Lithium ($Li$) in period 2 as an alkali metal, and Oxygen ($O$) in period 2 as part of the oxygen family.

Solutions, Solubility, and Chemical Bonding

A solution is a homogeneous mixture of two or more substances. It consists of a solute (the substance that is dissolved, such as sugar) and a solvent (the substance that does the dissolving, such as water). For example, in sugar water, sugar is the solute, water is the solvent, and sugar water is the solution. Solubility is the ability of a solute to dissolve in a particular solvent and is influenced by the polarity of the molecules, often following the rule "like dissolves like." In comparisons of compounds like propane ($C_3H_8$), glycerol, and propylene glycol, propane is a nonpolar covalent molecule and would be the most soluble in a nonpolar solvent like hexane but least soluble in water. In contrast, polar covalent molecules, especially those with multiple hydroxyl ($-OH$) groups like glycerol, will be more soluble in water and less soluble in hexane.

Chemical bonding occurs in two primary types: ionic and covalent. Ionic bonding involves the transfer of electrons between a metal and a nonmetal, resulting in the formation of ions and the creation of compounds like zinc sulfide ($ZnS$), calcium phosphate ($Ca_3(PO_4)_2$), or sodium chlorate ($NaClO_3$). Covalent bonding involves the sharing of electrons between nonmetals, as seen in molecules like phosphorus monofluoride ($PF$) or water ($H_2O$). Nomenclature rules differ for each: ionic compounds are named by the metal then the nonmetal with an "-ide" ending (with Roman numerals for transition metals like Iron(II) nitrate $Fe(NO_3)_2$ or Iron(III) carbonate $Fe_2(CO_3)_3$), while covalent compound names use prefixes to indicate the number of atoms, such as trisulfur pentoxide ($S_3O_5$).

Chemical Formulas and Balancing Reactions

Writing chemical formulas requires balancing the charges of the constituent ions to ensure overall neutrality. Examples include zinc sulfide ($Zn^{2+}$ and $S^{2-}$ giving $ZnS$), hydrogen nitrate ($H^+$ and $NO_3^-$ giving $HNO_3$), phosphorus monofluoride ($PF$), iron(II) nitrate ($Fe^{2+}$ and $NO_3^-$ giving $Fe(NO_3)_2$), and copper(I) carbonate ($Cu^+$ and $CO_3^{2-}$ giving $Cu_2CO_3$). Other examples include sodium chlorate ($NaClO_3$), iron(III) carbonate ($Fe_2(CO_3)_3$), and potassium oxide ($K_2O$). Molar mass calculation involves summing the atomic masses of all atoms in the formula; for instance, the molar mass of iron(III) sulfide ($Fe_2S_3$) is $207.9\,g/mol$ ($2 \times 55.8\,g + 3 \times 32.1\,g$) and magnesium nitrate ($Mg(NO_3)_2$) is $148.3\,g/mol$.

Chemical reactions are balanced using coefficients to ensure a conservation of mass, confirming that the same number of each type of atom exists on both sides of the reaction. Several reaction types exist: double replacement ($DR$), combustion ($OC$), decomposition ($Decomp$), and single replacement ($SR$). Examples include the double replacement reaction $SnCl_2 + (NH_4)_2S \rightarrow SnS + 2\,NH_4Cl$, the combustion of butane $2\,C_4H_{10} + 13\,O_2 \rightarrow 8\,CO_2 + 10\,H_2O$, the double replacement reaction of potassium sulfide and nickel(II) nitrate $K_2S + Ni(NO_3)_2 \rightarrow NiS + 2\,KNO_3$, and the single replacement reaction $2\,Al + 3\,CuCl_2 \rightarrow 2\,AlCl_3 + 3\,Cu$. Other examples include the synthesis and single replacement reactions of chlorine with sodium iodide ($Cl_2 + 2\,NaI \rightarrow 2\,NaCl + I_2$) and cadmium with oxygen ($2\,Cd + O_2 \rightarrow 2\,CdO$), and the reaction between hydrogen sulfide and ammonium carbonate which can produce water and carbon dioxide alongside ammonium sulfide.

Stoichiometry and Moles

Stoichiometry involves the quantitative relationship between reactants and products in a chemical reaction. A mole is a fundamental unit representing Avogadro's number ($6.02 \times 10^{23}$ particles, which can be atoms, molecules, or formula units). Calculations can convert between grams and moles using molar mass, or between moles and number of particles using Avogadro's number. For example, $2500\,g$ of iron(III) oxide is equivalent to approximately $12\,moles$ ($2500 \div 207.9$). The mass of $0.301\,mol$ of barium fluoride is $52.8\,g$ ($0.301 \times 175.3$). Moles of silicon from $2.80 \times 10^{24}$ atoms is $4.65\,mol$. The number of particles in $1.14\,moles$ of $SO_3$ is $6.86 \times 10^{23}$. The number of molecules in $70.3\,g$ of disulfur tetraoxide ($S_2O_4$) is $3.30 \times 10^{23}$. The mass of $3.4 \times 10^{10}$ formula units of potassium sulfate is $9.84 \times 10^{-12}\,g$.

In a reaction like the processing of iron ore $Fe_2O_3 + 3\,H_2 \rightarrow 2\,Fe + 3\,H_2O$, molar ratios from the balanced equation are used for calculations. For instance, $25\,moles$ of $Fe_2O_3$ will produce $50\,moles$ of iron. The molecules of hydrogen needed to make $30.\,moles$ of iron is $2.7 \times 10^{25}$ molecules ($30 \times (3/2) \times 6.02 \times 10^{23}$). Additional stoichiometry includes finding the mass of oxygen produced from $12.0\,moles$ of $NaClO_3$ in the reaction $2\,NaClO_3 \rightarrow 2\,NaCl + 3\,O_2$, which is $576\,g$. The moles of sodium chloride produced from $80.0\,g$ of $NaClO_3$ in the same reaction is $0.751\,mol$. A limiting reactant is the substance in a reaction that is exhausted first, thereby stopping the reaction, such as paper being the limiting reactant when it is burned in an excess of oxygen. Molarity ($M$), a measure of concentration, is defined as moles of solute per liter of solution. For a solution with $4.5\,g$ of $NaOH$ in $210\,mL$, find the moles ($4.5 \div 40.0 = 0.1125\,mol$), then divide by liters to get a concentration of $0.54\,M$.

Acids and Bases: pH, Titration, and Neutralization

Acids and bases are chemical species characterized by their hydrogen ion ($H^+$) and hydroxide ion ($OH^-$) concentrations, respectively. The pH scale measures acidity, with pH values below 7 being acidic (more $H^+$), and values above 7 being basic (more $OH^-$). The concentration of hydrogen ions and hydroxide ions are inversely related; if $[H^+]$ increases, $[OH^-]$ must decrease. A pH change of 2 units (e.g., from 5 to 3) indicates a 100-fold increase in the concentration of hydrogen ions, making the solution more acidic. Neutralization reactions occur when an acid reacts with a base to produce water and a salt, such as $HCl + NaOH \rightarrow H_2O + NaCl$. Diluting an acid with water results in a decrease in the absolute concentration of hydrogen ions.

Titration is a laboratory procedure used to determine the concentration of a solution by reacting it with a solution of known concentration. A burette is often used for precise delivery of the reactant, and the reading, such as $27.81\,mL$, should be recorded at the meniscus. An indicator like phenolphthalein is used to mark the endpoint; if it turns pink and stays pink, the acid has been neutralized and the solution is no longer acidic. For a neutralization involving $Ba(OH)_2$ and $HCl$, the mass of barium hydroxide needed to neutralize $325\,mL$ of $0.510\,M$ $HCl$ is $14.2\,g$. Based on the molarity and volume of potassium hydroxide ($78\,mL$ of $0.10\,M$) required to neutralize an $HCl$ solution ($38.8\,mL$), the molarity of the $HCl$ is found to be $0.201\,M$. When titrating an acid whose concentration has increased, a larger amount of base will be required to reach the neutralization point.

Gas Laws and Particle Behavior

Gas pressure is an emerging property of gas particles resulting from their frequent collisions with the walls of their container and objects within it. Several factors can increase the pressure of a gas, described by the gas laws. Decreasing the volume of a container (while temperature remains constant) leads to an increase in pressure because the particles are confined to a smaller space and collide more frequently. Increasing the temperature of a gas also increases pressure, as the particles gain more kinetic energy, move faster, and strike the walls with more force. Adding more gas molecules to a fixed volume container will also increase the pressure because more particles result in more collisions occurring per unit of time and area. Conversely, if the temperature increases, the speed of the gas particles also increases.

Questions and Discussion

The review materials incorporate several practice questions designed to assess understanding of various topics. These include calculations involving scientific notation and significant figures, such as determining that $0.003760$ has 4 significant figures and resolving calculation precision based on the rules for different operations. Practical applications are tested, like identifying specialized glass equipment (beakers, flasks, and graduated cylinders) and correctly reading volumes from their scales to include an estimated digit. Conceptual understanding is also addressed through multiple-choice questions assessing factors that speed up chemical reactions, the role of catalysts, and the distinguishing characteristics of elements, compounds, and mixtures. Questions regarding atomic structure challenge students to use isotope notation to find subatomic particle counts and identify corresponding electron configurations and names for periodic table groups. Finally, students use provided activity series data to predict if a single replacement reaction like cadmium metal being placed into a chromium (III) bromide solution would occur, determining that it would not occur because cadmium is less reactive than chromium.