Question 1

Which of the following statements correctly describes a characteristic feature of a mixture in basic chemistry?

Accepted Answer

Introduction / Context: In introductory chemistry, students learn to distinguish between elements, compounds and mixtures. Mixtures are particularly important because many everyday materials, such as air, seawater and soil, are mixtures rather than pure substances. Understanding what makes a mixture different from a compound helps in topics like separation techniques and solution chemistry. This question asks which statement correctly describes a key characteristic of mixtures, focusing on composition and the way components are combined. Given Data / Assumptions: Mixtures consist of two or more substances present together. Components of a mixture are not chemically combined but physically combined. Mixtures can be homogeneous (uniform) or heterogeneous (non uniform). Compounds have fixed formulas and are chemically combined, unlike mixtures. Concept / Approach: A mixture is defined as a physical combination of two or more substances where each substance retains its own chemical identity. The composition of a mixture can vary, meaning the components are not present in a fixed ratio. Mixtures can be homogeneous, such as salt dissolved in water, or heterogeneous, such as sand in water. In contrast, a compound is a pure substance with a fixed formula and definite ratio of atoms chemically bonded together. The approach is to identify the option that recognises variability of composition, physical combination and the possibility of both homogeneous and heterogeneous forms. Step-by-Step Solution: Step 1: Recall that mixtures are formed when two or more substances are physically combined without forming new chemical bonds between them. Step 2: Understand that mixtures can have variable composition; the proportions of components can change without altering the nature of the mixture. Step 3: Recognise that mixtures can be homogeneous, like air or sugar solution, where composition is uniform, or heterogeneous, like oil and water, where different regions have different compositions. Step 4: Compounds, unlike mixtures, have fixed formulas and the atoms are chemically combined in definite ratios. Step 5: Therefore, the statement that a mixture can be homogeneous or heterogeneous and that the components are physically combined is the correct description. Verification / Alternative check: Examples from daily life support this definition. Air is a homogeneous mixture of gases, with nitrogen and oxygen present in varying proportions depending on location and conditions. Seawater is a mixture of water and dissolved salts with variable salinity. Granite is a heterogeneous mixture of minerals that can be seen as different coloured grains. In all these cases, the components can be separated by physical methods, and their relative amounts are not fixed by a simple chemical formula. Textbooks consistently define mixtures in this way, confirming the correctness of the chosen statement. Why Other Options Are Wrong: Option A describes a pure substance with a fixed chemical formula, which is characteristic of a compound, not a mixture. Option B states that components are chemically combined in a fixed ratio, again describing a compound like water or sodium chloride. Option C claims that components are always in one specific ratio only, which contradicts the defining feature that mixture composition can vary. Option E suggests that a mixture always consists of only one type of atom, which is actually the definition of an element, not a mixture. Only option D correctly states that a mixture can be homogeneous or heterogeneous and that its components are physically combined rather than chemically bonded in a fixed ratio. Common Pitfalls: A frequent mistake is to confuse mixtures with compounds because both can look uniform at first glance. For example, a sugar solution looks uniform, but it is still a mixture since sugar and water can be separated by evaporation. Another pitfall is assuming that any uniform sample must be a pure substance; in reality, many homogeneous samples are mixtures. To avoid confusion, ask whether the sample could be separated by physical means and whether its composition can vary. If the answer is yes, it is likely a mixture rather than a pure compound. Final Answer: A correct characteristic of a mixture is that It can be homogeneous or heterogeneous and its components are physically combined without a fixed chemical formula.

Question 2

When ionic bonds are formed between metals and nonmetals, metallic atoms generally tend to do what with their outer shell electrons?

Accepted Answer

Introduction / Context: Ionic bonding is one of the basic types of chemical bonding taught in school level chemistry. It typically occurs between metals and nonmetals. Metals tend to form positively charged ions, while nonmetals tend to form negatively charged ions. Understanding what happens to the valence electrons of a metallic atom during ionic bond formation is essential for predicting formulas of ionic compounds, such as NaCl or MgO, and for understanding electrical conductivity in molten salts. Given Data / Assumptions: We are considering ionic bonds, not covalent or metallic bonds. The bonding is between metallic atoms and nonmetal atoms. Valence electrons are the outer shell electrons involved in bonding. Metals usually have low ionisation energies compared with nonmetals. Concept / Approach: Metals have relatively few electrons in their outermost shell and low ionisation energies, so they can easily lose these electrons to achieve a stable noble gas configuration. Nonmetals, on the other hand, have higher electronegativities and tend to gain electrons. In ionic bonding, electrons are transferred from metal atoms to nonmetal atoms. The metal becomes a positively charged cation, and the nonmetal becomes a negatively charged anion. The electrostatic attraction between these oppositely charged ions forms the ionic bond. The correct description of metallic atom behaviour is therefore that they lose electrons and form cations. Step-by-Step Solution: Step 1: Recall that metals, especially those in groups 1 and 2, have one or two electrons in their outermost shell and seek stability by losing these electrons. Step 2: When a metal atom loses one or more valence electrons, it becomes a positively charged ion, called a cation, because it now has more protons than electrons. Step 3: Nonmetals, which have higher electronegativities, accept these electrons and become negatively charged anions. Step 4: The electrostatic attraction between the positively charged metal cations and negatively charged nonmetal anions forms an ionic bond. Step 5: Therefore, the correct statement is that metallic atoms lose electrons to become positively charged cations when ionic bonds are formed. Verification / Alternative check: Examples support this description. Sodium, a group 1 metal, loses one electron to form Na⁺ in sodium chloride. Magnesium, a group 2 metal, loses two electrons to form Mg²⁺ in magnesium oxide. In each case, the metal cation pairs with a nonmetal anion such as Cl⁻ or O²⁻ to form an ionic crystal. This pattern is repeated in many ionic compounds like KBr, CaCl2 and Al2O3. In all these examples, metals clearly lose electrons and form positive ions, confirming the general rule. Why Other Options Are Wrong: Option A describes sharing electrons equally, which is characteristic of nonpolar covalent bonding, not ionic bonding. Option B suggests that metals gain electrons and become anions, which is opposite to their normal behaviour; nonmetals are the ones that gain electrons. Option D claims that metals neither gain nor lose electrons, which contradicts the fundamental definition of ionic bonding as involving electron transfer. Option E suggests that metallic atoms break apart into neutral atoms with no charge, which does not describe bonding at all. Only option C correctly states that metals lose electrons to become positively charged cations when ionic bonds form. Common Pitfalls: Students sometimes confuse ionic bonds with covalent bonds and mistakenly think that metals share electrons in the same way nonmetals do. Another pitfall is to mix up cations and anions, forgetting which is positive and which is negative. To remember, note that cation starts with the same letter as cat, which might be thought of as a positive pet, and anion contains the letter n as in negative. Also remember that metallic behaviour is characterised by electron loss, which explains why metals conduct electricity and form positive ions in electrolysis. Final Answer: When ionic bonds are formed, metallic atoms generally Lose electrons to become positively charged cations, which then attract nonmetal anions to form ionic compounds.

Question 3

With respect to solubility in water, which type of molecules generally dissolve best in water: polar molecules or nonpolar molecules?

Accepted Answer

Introduction / Context: Solubility is a key concept in chemistry and biology. Whether a substance dissolves in water or not depends on the nature of both the solute and the solvent. Water is often called the universal solvent because it dissolves many substances, but it does not dissolve everything. The general rule like dissolves like states that polar solvents tend to dissolve polar solutes, while nonpolar solvents dissolve nonpolar solutes. This question asks which type of molecules, polar or nonpolar, generally dissolve best in water and why. Given Data / Assumptions: Water is a highly polar solvent, with a bent molecular shape and a significant dipole moment. Polar molecules have uneven charge distribution and can interact with water through dipole dipole or hydrogen bonding interactions. Nonpolar molecules have little or no permanent charge separation. Ionic compounds dissociate into positive and negative ions in water. Concept / Approach: The essential idea is that polar solvents stabilise charged or polar solutes through electrostatic interactions. Water molecules surround ions or polar solute molecules, forming hydration shells and lowering the energy of the system. As a result, many polar compounds and ionic salts are very soluble in water. Nonpolar molecules, such as many hydrocarbons, do not interact strongly with water and tend to separate, leading to poor solubility. The approach is to recall that water is polar and therefore favours dissolution of polar and ionic substances rather than nonpolar ones. Step-by-Step Solution: Step 1: Recognise that water has a bent shape and a strong dipole, with partial negative charge near the oxygen and partial positive charges near the hydrogens. Step 2: Polar molecules, such as alcohols and sugars, also have regions of partial positive and negative charge, allowing them to form hydrogen bonds or dipole interactions with water. Step 3: Ionic compounds, like sodium chloride, dissociate into ions (Na⁺ and Cl⁻), which are stabilised by water molecules that orient around them according to charge. Step 4: Nonpolar molecules, such as oils and many hydrocarbons, lack strong charge separation and cannot form favourable interactions with water, so they tend to be insoluble and may form separate layers. Step 5: Therefore, the type of molecules that generally dissolve best in water are polar molecules and ionic compounds, not nonpolar molecules. Verification / Alternative check: Everyday examples support this rule. Table salt (ionic) and sugar (polar covalent) both dissolve readily in water, forming clear solutions. Ethanol, a polar molecule, is completely miscible with water. In contrast, vegetable oil, petrol and many waxes, which are nonpolar, do not dissolve in water and often float on the surface. Biology texts also explain that water dissolves many polar nutrients and ions in blood and cytoplasm, while nonpolar lipids aggregate into membranes or fat droplets. These observations confirm that polar and ionic species are the ones that dissolve best in water. Why Other Options Are Wrong: Option B incorrectly states that water is nonpolar and therefore dissolves nonpolar molecules, which contradicts the well known polar nature of water. Option C claims that both polar and nonpolar molecules dissolve equally well, which is clearly false; oils and hydrocarbons are largely insoluble in water. Option D suggests that water dissolves only metals, which is not correct; many metals are not very soluble in water, while nonmetallic ionic salts dissolve easily. Option E states that only large nonpolar molecules dissolve, but in fact large nonpolar molecules tend to be even less soluble in water than smaller ones. Only option A correctly states that polar molecules and ionic compounds are most soluble in polar water. Common Pitfalls: A common mistake is to think that because water is so important in biology, it must dissolve everything equally well. Another pitfall is to confuse solubility of ionic solids with metallic substances, assuming that all solids behave alike. To avoid these errors, remember the like dissolves like guideline and focus on polarity. Water dissolves ionic and polar substances well by stabilising charges, but it does not dissolve nonpolar substances such as oils, which instead form separate phases or micelles when surfactants are present. Final Answer: The type of molecules that generally dissolve best in water are Primarily polar molecules and ionic compounds, because water is a polar solvent that stabilises charges and dipoles effectively.

Question 4

Which of the following biological macromolecules is a polymer formed by linking many smaller subunits together?

Accepted Answer

Introduction / Context: In biochemistry, many important molecules are polymers, meaning they are made of many repeating smaller units called monomers. Recognising which biomolecules are polymers and which are monomers is essential for understanding the structure and function of proteins, nucleic acids and polysaccharides. This question asks you to identify the polymer among several biological molecules, some of which are small building blocks and others which are large chains made from these building blocks. Given Data / Assumptions: Glycerol is a three carbon alcohol used in forming triglycerides. An amino acid is a monomer used to build proteins. Nucleic acids, such as DNA and RNA, are long chains of nucleotides. Fatty acids are long chain carboxylic acids used to build lipids. Monosaccharide sugars are small carbohydrate units. Concept / Approach: A polymer is a large molecule made by joining many similar or identical monomer units through covalent bonds. Nucleic acids are classic examples: DNA and RNA consist of long chains of nucleotide monomers linked by phosphodiester bonds. In contrast, glycerol, amino acids, fatty acids and monosaccharides listed here are monomers or small molecules that can be combined to form larger structures, but are not themselves polymers. The approach is to identify which option describes a macromolecule built from many smaller repeating units rather than a single small unit. Step-by-Step Solution: Step 1: Consider glycerol, which has three carbon atoms with hydroxyl groups and acts as a backbone in triglycerides, but a single glycerol molecule is not a polymer. Step 2: An amino acid is a small molecule consisting of an amino group, carboxyl group and side chain; proteins are polymers made from many amino acids, but one amino acid alone is not a polymer. Step 3: Nucleic acids, including DNA and RNA, are built from many nucleotide monomers linked together, forming long chains with sequence information. Step 4: A single fatty acid molecule is a monomeric unit that can form part of lipids but is not itself a polymer. Step 5: A monosaccharide is a single sugar unit; polysaccharides such as starch or cellulose are polymers, but one monosaccharide molecule is not. Step 6: Therefore, the polymer in the list is a nucleic acid, which is made of many nucleotide subunits. Verification / Alternative check: Biology textbooks classify nucleic acids, proteins and polysaccharides as major classes of biological polymers, with lipids often treated separately. DNA is described as a long polymer composed of nucleotide monomers arranged in a specific sequence, responsible for genetic information. In contrast, amino acids are described as monomers of proteins, and monosaccharides as monomers of polysaccharides. Glycerol and fatty acids are building blocks of triglycerides and phospholipids, which are not classic polymers in the same sense. This classification confirms nucleic acids as the correct answer. Why Other Options Are Wrong: Option A, glycerol, is a small three carbon molecule and not a chain of repeating units. Option B, an amino acid, is a monomer for proteins, not a polymer itself. Option D, a fatty acid, is a single long chain carboxylic acid, again a monomeric unit. Option E, a simple monosaccharide sugar, is a single unit of carbohydrate, whereas polymers such as starch and cellulose are made of many monosaccharides. Only option C, a nucleic acid such as DNA or RNA, fits the definition of a polymer formed by linking many smaller units together. Common Pitfalls: Students sometimes confuse the building blocks with the larger structures they form, assuming that any large molecule is automatically a polymer. Another pitfall is to overlook nucleic acids and focus only on proteins and carbohydrates when thinking about polymers. To avoid confusion, remember that polymer always implies a repeated chain of similar units. Nucleic acids are polymers of nucleotides, proteins are polymers of amino acids and polysaccharides are polymers of monosaccharides. Lipids are built from smaller components but are not usually considered true polymers in the same way. Final Answer: The biological macromolecule that is a polymer in this list is A nucleic acid such as DNA or RNA, which is made of many nucleotide monomers linked in a chain.

Question 5

Which of the following formulas represents a molecule that is overall nonpolar due to its symmetric shape and cancellation of bond dipoles?

Accepted Answer

Introduction / Context: Molecular polarity depends on both bond polarity and molecular geometry. Even if all individual bonds are polar, a molecule can be nonpolar if the shape is highly symmetric and the bond dipoles cancel each other out. Polarity affects properties such as boiling point, solubility and intermolecular forces. This question asks which of the given molecular formulas corresponds to a nonpolar molecule, taking into account both bond polarity and overall shape. Given Data / Assumptions: H2O is known to be a bent molecule with polar O–H bonds. CF4 is a tetrahedral molecule with four C–F bonds. HCl is a diatomic molecule with a bond between hydrogen and chlorine. H2S is a bent molecule similar in shape to water. NH3 is trigonal pyramidal with a lone pair on nitrogen. Concept / Approach: Bond polarity arises from differences in electronegativity between bonded atoms, creating bond dipoles. Whether the molecule is polar depends on the vector sum of these dipoles. In highly symmetric shapes, such as tetrahedral with identical bonds or linear molecules with equal dipoles in opposite directions, the dipoles can cancel, leading to a nonpolar molecule. In bent or pyramidal shapes, the dipoles usually do not cancel, and the molecule is polar. The approach is to consider each molecule's shape and bond polarity, then identify the one where all bond dipoles cancel due to symmetry. Step-by-Step Solution: Step 1: H2O has two polar O–H bonds and a bent geometry, so the dipoles do not cancel; water is polar. Step 2: CF4 has four polar C–F bonds arranged symmetrically in a tetrahedral shape around the central carbon atom. Step 3: Because the four identical C–F bond dipoles are symmetrically oriented, their vector sum is zero, making CF4 nonpolar overall despite each bond being polar. Step 4: HCl is a simple diatomic molecule with a significant electronegativity difference between hydrogen and chlorine, so the molecule has a dipole and is polar. Step 5: H2S is bent, similar to H2O, and has polar H–S bonds, resulting in a net dipole and a polar molecule. Step 6: NH3 is trigonal pyramidal with a lone pair on nitrogen, making it polar due to an uneven distribution of charge. Step 7: Therefore, CF4 is the only molecule in the list that is overall nonpolar. Verification / Alternative check: Molecular polarity data from reference books confirm that water, hydrogen chloride, hydrogen sulfide and ammonia all have measurable dipole moments, indicating they are polar. Carbon tetrafluoride, CF4, despite having highly polar C–F bonds, has a very small or negligible dipole moment because the tetrahedral arrangement cancels the bond dipoles. CF4 is often used as a textbook example of a nonpolar molecule with polar bonds. Compare CF4 with CH4, which is also tetrahedral and nonpolar, reinforcing the idea that symmetry can cancel dipoles. Why Other Options Are Wrong: Option A, H2O, is polar because its bent shape prevents cancellation of O–H bond dipoles. Option C, HCl, is a polar diatomic molecule with a clear separation of charge between H and Cl. Option D, H2S, has a bent geometry and polar bonds, making it polar. Option E, NH3, with a trigonal pyramidal shape and lone pair, is polar and interacts strongly with water through hydrogen bonding. Only option B, CF4, combines a symmetric tetrahedral shape with identical polar bonds in such a way that all bond dipoles cancel and the molecule is overall nonpolar. Common Pitfalls: Students sometimes assume that if all bonds are polar, the molecule must be polar, forgetting the role of geometry. Another pitfall is to think that tetrahedral molecules are always polar because of their three dimensional shape. To avoid these errors, always consider both bond polarity and molecular geometry. Ask whether the polar bonds are arranged symmetrically around the central atom. If they are, as in CF4 or CO2, the molecule can be nonpolar even though individual bonds are polar. Final Answer: The formula that represents a nonpolar molecule is CF4, a tetrahedral molecule with four identical C–F bonds whose bond dipoles cancel due to symmetry.

Question 6

In molecular geometry and polarity, which description best explains the type of molecule CF4 (carbon tetrafluoride) is in terms of overall charge distribution?

Accepted Answer

Introduction / Context: Carbon tetrafluoride, CF4, is a classic example used in chemistry to test understanding of molecular shape and polarity. Even though each C F bond is polar due to the electronegativity difference between carbon and fluorine, the overall molecule can still be nonpolar depending on its shape. This question checks whether you can correctly relate the tetrahedral geometry of CF4 to its overall charge distribution and classify it as polar or nonpolar. Given Data / Assumptions: - The molecule under discussion is CF4, carbon tetrafluoride. - The central atom is carbon, bonded to four fluorine atoms. - The C F bonds are individually polar because fluorine is more electronegative than carbon. - We assume ideal tetrahedral geometry with identical bond lengths and bond angles. Concept / Approach: Molecular polarity depends on both bond polarity and molecular geometry. A polar bond arises when two atoms have different electronegativities, causing an uneven sharing of electrons. However, if these polar bonds are arranged symmetrically in three dimensional space, the individual bond dipoles can cancel out, producing a nonpolar molecule overall. CF4 has a tetrahedral shape with four identical C F bonds directed to the corners of a regular tetrahedron. The symmetry of this arrangement causes the dipoles to cancel, leaving no net dipole moment. Therefore, CF4 is classified as a nonpolar molecule with a symmetrical distribution of charge. Step-by-Step Solution: Step 1: Identify the shape of CF4. Carbon in CF4 is surrounded by four bonding pairs and no lone pairs, giving a tetrahedral geometry. Step 2: Recognise that each C F bond is polar because fluorine is significantly more electronegative than carbon. Step 3: Visualise the molecule in three dimensions: the four polar bonds are evenly spaced, pointing toward the vertices of a tetrahedron. Step 4: Consider the bond dipoles. Because of the symmetrical arrangement, each bond dipole is balanced by others, so their vector sum is zero. Step 5: Conclude that there is no overall molecular dipole moment, meaning CF4 is nonpolar with a symmetric charge distribution. Verification / Alternative check: Another way to check is to compare CF4 with molecules of similar geometry. Methane CH4 is also tetrahedral and nonpolar because all bonds and positions are identical. Similarly, CF4 has four identical C F bonds in a symmetric arrangement. In contrast, molecules such as CH3Cl, where one substituent is different, become polar because symmetry is broken. Experimental measurements of dipole moment confirm that CF4 has a dipole moment of essentially zero, supporting the conclusion that it is nonpolar and symmetric in its charge distribution. Why Other Options Are Wrong: Option B describes a polar molecule with an asymmetrical distribution of charge. This would apply to molecules like CH3Cl or H2O, not to perfectly symmetric CF4. Option C suggests a nonpolar molecule with an asymmetrical distribution of charge, which is contradictory because an asymmetrical charge distribution usually leads to polarity. Option D describes a polar molecule with symmetrical charge distribution, which is again inconsistent, because true symmetry tends to cancel dipoles and remove net polarity. Common Pitfalls: Students often assume that if all bonds are polar, then the whole molecule must be polar. This is not always true. Another common mistake is to think only in two dimensions and overlook the three dimensional tetrahedral structure. Forgetting about molecular geometry can lead to incorrect classification of CF4 as polar. Always combine bond polarity with the molecule three dimensional shape when deciding whether a molecule is polar or nonpolar. Final Answer: The correct answer is: A nonpolar molecule with a symmetrical distribution of charge around the central atom.

Question 7

In everyday chemistry of food, is the statement Formation of curd from milk is a chemical reaction scientifically correct or incorrect?

Accepted Answer

Introduction / Context: The transformation of milk into curd is a familiar process in households and an excellent real life example of a chemical change in food chemistry. Understanding whether this change is physical or chemical helps students connect classroom concepts to daily life. A chemical change produces new substances with different properties, while a physical change only alters form or state. This question asks whether the common statement about curd formation being a chemical reaction is scientifically correct. Given Data / Assumptions: - The starting material is milk, which contains water, fats, proteins, and lactose. - Curd is obtained from milk by adding a small amount of existing curd or lactic acid bacteria and keeping the mixture warm. - Lactic acid bacteria act on lactose sugar present in milk. - We compare this change with typical physical changes such as melting or boiling. Concept / Approach: In a chemical change, the original substances are converted into new substances with different chemical compositions and properties. During curd formation, lactic acid bacteria ferment lactose sugar into lactic acid. Lactic acid decreases the pH of the mixture and causes milk proteins such as casein to coagulate. The resulting curd has a different taste, texture, and acidity compared to milk. These differences come from the formation of new chemical substances, mainly lactic acid and coagulated protein networks, which shows that a chemical reaction has taken place rather than a simple physical change. Step-by-Step Solution: Step 1: Identify the main components of milk, especially lactose sugar and casein protein. Step 2: Recognise that lactic acid bacteria present in curd or starter culture break down lactose through fermentation. Step 3: During fermentation, lactose is converted into lactic acid, which is a new chemical substance not present in the same amount in fresh milk. Step 4: The lactic acid lowers the pH, causing casein proteins to denature and coagulate, forming a semi solid curd structure. Step 5: Note that milk cannot be easily recovered from curd by simple physical means like cooling or heating, which is typical of irreversible chemical changes. Step 6: Conclude that formation of curd from milk is indeed a chemical reaction, so the statement is correct. Verification / Alternative check: To verify, compare curd formation with a purely physical change such as freezing water into ice. Frozen water can be melted back to liquid water without any change in composition. In contrast, curd retains a sour taste and solidified texture even if warmed or stirred, demonstrating that its composition has changed. Food science texts classify fermentation as a biochemical process involving enzymes and microbes that produce new molecules. This confirms that curd formation from milk is a chemical change based on fermentation and protein coagulation. Why Other Options Are Wrong: Option B claims that curd formation is only a physical change, which is incorrect because new substances like lactic acid are formed. Option C suggests that the change is purely like a change of state such as melting, but curd formation is not simply solidification or liquefaction and involves complex biochemical reactions. Option D states that curd formation is reversible and not a chemical reaction, which is wrong because you cannot easily get back the original milk and the process clearly involves fermentation. Common Pitfalls: Students sometimes judge changes only by appearance. Since both milk and curd are dairy products and look somewhat similar, learners may think no new substances are formed. Another common mistake is to assume that any solidification is a physical change, forgetting that proteins can chemically denature and coagulate. Always check whether the process produces new substances, changes taste and smell, and is hard to reverse. Such clues strongly indicate a chemical change, as in the conversion of milk to curd. Final Answer: The correct answer is: Correct – formation of curd from milk is a chemical change involving new substances.

Question 8

In basic stoichiometry, what is the gram formula mass (molar mass) of ammonium phosphate, (NH4)3PO4, expressed in grams per mole?

Accepted Answer

Introduction / Context: Calculating gram formula mass, also called molar mass, is a fundamental skill in chemistry that underlies stoichiometric calculations. The gram formula mass of a compound tells you the mass of one mole of its formula units. Ammonium phosphate, written as (NH4)3PO4, is a common example used in problems related to fertilisers and inorganic salts. This question tests whether you can correctly interpret the formula and sum the atomic masses to obtain the correct gram formula mass in grams per mole. Given Data / Assumptions: - The chemical formula is (NH4)3PO4, ammonium phosphate. - Approximate atomic masses used are: N about 14 g per mol, H about 1 g per mol, P about 31 g per mol, and O about 16 g per mol. - There are three ammonium ions NH4 plus and one phosphate group PO4 3 minus in the formula unit. - We assume standard rounding of atomic masses to two decimal places where needed. Concept / Approach: To find the gram formula mass of a compound, multiply the atomic mass of each element by the number of atoms of that element in the formula and then add all contributions. For (NH4)3PO4, we first count how many nitrogen, hydrogen, phosphorus, and oxygen atoms are present. Then we use approximate atomic masses for each element. The total gives the molar mass of one mole of ammonium phosphate formula units, which we can compare with the options given to select the closest value. Step-by-Step Solution: Step 1: Expand the formula (NH4)3PO4. The subscript 3 applies to NH4, so there are 3 nitrogen atoms and 12 hydrogen atoms from the ammonium ions. Step 2: Count atoms of each element. There are 3 N atoms, 12 H atoms, 1 P atom, and 4 O atoms in one formula unit of (NH4)3PO4. Step 3: Use approximate atomic masses: nitrogen about 14.01 g per mol, hydrogen about 1.008 g per mol, phosphorus about 30.97 g per mol, and oxygen about 16.00 g per mol. Step 4: Calculate the total mass contribution of each element: N: 3 * 14.01 ≈ 42.03 g per mol; H: 12 * 1.008 ≈ 12.10 g per mol; P: 1 * 30.97 ≈ 30.97 g per mol; O: 4 * 16.00 = 64.00 g per mol. Step 5: Add all contributions. Total mass ≈ 42.03 + 12.10 + 30.97 + 64.00 = 149.10 g per mol (rounded, about 149.08 g per mol). Step 6: Compare with options and choose 149.08 g per mol as the closest and correct gram formula mass. Verification / Alternative check: As a quick check, you can round atomic masses more roughly: N ≈ 14, H ≈ 1, P ≈ 31, O ≈ 16. Then the mass is 3 * 14 + 12 * 1 + 31 + 4 * 16 = 42 + 12 + 31 + 64 = 149 g per mol. This back of the envelope calculation agrees closely with the more precise value and confirms that a result near 149 g per mol is correct. Since 149.08 g per mol is the only option near 149, it must be the correct answer. Why Other Options Are Wrong: Option A (139.54 g per mol) is too low and would imply missing several oxygen or nitrogen atoms in the calculation. Option B (140.70 g per mol) is also significantly below the correct value and does not match the sum of the individual atomic contributions. Option D (153.24 g per mol) is higher than the calculated value and would correspond to counting extra atoms or using incorrect atomic masses. Common Pitfalls: Students sometimes forget to apply the outer subscript 3 to both N and H in NH4, leading to incorrect counts of nitrogen and hydrogen atoms. Another common mistake is to misread the formula as NH4 3PO4 without recognising the grouping, which changes the stoichiometry. Using inaccurate or rounded atomic masses without checking the result can also introduce large errors. Always carefully count each atom in the formula, multiply by the proper atomic mass, and then sum the contributions to get the correct gram formula mass. Final Answer: The correct answer is: 149.08 g per mol.

Question 9

In basic chemical reaction terminology, what are reactants in a chemical reaction?

Accepted Answer

Introduction / Context: Learning the basic language of chemistry is essential for understanding equations, laboratory procedures, and exam questions. Two of the most important terms are reactants and products. Reactants appear on the left side of a chemical equation and are transformed during the reaction, while products appear on the right side and are formed as a result. This question asks you to select the best definition of reactants in a chemical reaction. Given Data / Assumptions: - We are dealing with ordinary chemical reactions, represented by balanced chemical equations. - Reactants are written on the left side of the equation. - Products are written on the right side of the equation. - Catalysts, if present, are not consumed and are not considered reactants or products in the usual sense. Concept / Approach: A reactant is any substance that takes part in a chemical reaction and is consumed or changed during the process. In symbolic form, a reaction is written as Reactants -> Products, where the arrow symbol means yields or gives. Reactants provide the atoms that are rearranged to form new substances. They may be elements or compounds, and they are often listed with plus signs between them when more than one substance is involved. Understanding this definition helps you correctly interpret reaction equations and identify what is being used up in a reaction. Step-by-Step Solution: Step 1: Recall that in a chemical equation such as H2 + O2 -> H2O, the substances before the arrow are H2 and O2, and these are called reactants. Step 2: Note that these reactants undergo a chemical change: hydrogen and oxygen molecules react to form water molecules. Step 3: Recognise that the substances after the arrow, in this example H2O, are called products because they are produced by the reaction. Step 4: Compare this understanding with the answer choices. The definition that reactants are starting substances that undergo chemical change fits the usual description. Step 5: Confirm that catalysts and solvents, although present, are not usually treated as reactants if they are not consumed during the reaction. Verification / Alternative check: To verify, look at another reaction such as Zn + 2HCl -> ZnCl2 + H2. Here, zinc and hydrochloric acid are on the left side and are consumed as the reaction proceeds, so they are reactants. Zinc chloride and hydrogen gas appear on the right and are formed, so they are products. In almost every textbook example, reactants are defined and illustrated in this way. This consistent usage confirms that reactants are indeed the starting substances that undergo change during the reaction. Why Other Options Are Wrong: Option B describes products, not reactants, because products are the final substances formed after the reaction. Option C refers to catalysts, which may influence the rate of reaction but are not consumed and therefore are not reactants in the usual sense. Option D incorrectly claims that only the solvent is a reactant. While solvents participate physically, they may not be chemically changed, and many reactions occur in the absence of a solvent. Common Pitfalls: Learners sometimes mix up reactants and products because they focus only on the direction of the arrow instead of the position in the equation. Another common mistake is to think that anything present in the beaker is a reactant, including catalysts and solvents. To avoid confusion, always remember the general pattern: substances before the arrow are reactants, and substances after the arrow are products. Reactants are consumed or changed, whereas catalysts are not used up and products are formed. Final Answer: The correct answer is: The starting substances that undergo chemical change during the reaction.

Question 10

In molecular orbital theory, what is the bond order of the oxygen molecule O2 in its ground state?

Accepted Answer

Introduction / Context: Bond order is an important concept in chemical bonding that gives a simple measure of the number of chemical bonds between a pair of atoms. For diatomic molecules such as O2, bond order is related to bond strength, bond length, and stability. Understanding the bond order of oxygen is especially important in molecular orbital theory and helps explain why O2 has a double bond and is paramagnetic. This question focuses on the numerical bond order for O2 in its ground state. Given Data / Assumptions: - The molecule considered is O2, a homonuclear diatomic molecule. - We are interested in the ground state, not excited or ionised species. - Bond order is defined using the molecular orbital electron configuration. - We use the standard bond order formula based on bonding and antibonding electrons. Concept / Approach: In molecular orbital theory, electrons in a molecule occupy bonding and antibonding molecular orbitals. Bond order can be calculated using the formula: bond order = (number of electrons in bonding orbitals minus number of electrons in antibonding orbitals) divided by 2. For O2, we fill the molecular orbitals in order of increasing energy, considering the total number of valence electrons from both oxygen atoms. The resulting configuration leads to a bond order of 2, which corresponds to a double bond between the two oxygen atoms. Step-by-Step Solution: Step 1: Each oxygen atom has 6 valence electrons, so O2 has 12 valence electrons in total. Step 2: In the molecular orbital scheme for O2, these valence electrons fill the sigma 2s, sigma 2s star, sigma 2p z, pi 2p x and 2p y, and pi 2p star x and 2p star y orbitals in order. Step 3: Count bonding electrons. The bonding molecular orbitals (sigma 2s, sigma 2p z, pi 2p x, pi 2p y) contain a total of 8 electrons. Step 4: Count antibonding electrons. The antibonding molecular orbitals (sigma 2s star and pi 2p star x and 2p star y) together contain 4 electrons. Step 5: Apply the bond order formula: bond order = (8 minus 4) divided by 2 = 4 divided by 2 = 2. Step 6: Conclude that the O2 molecule has a bond order of 2, consistent with a double bond. Verification / Alternative check: You can verify this result by comparing with simple Lewis structure ideas. The Lewis structure of O2 is often written with a double bond between the two oxygen atoms, each also having two lone pairs. A double bond corresponds to a bond order of 2. Furthermore, experimental data such as bond length and bond energy indicate stronger bonding in O2 compared to a single bond and weaker than a hypothetical triple bond, which is consistent with bond order 2. Paramagnetism of O2 arises from unpaired electrons in antibonding orbitals but does not change the bond order calculation. Why Other Options Are Wrong: Option A (1.5) would correspond to a weaker bond, more typical of some resonance averaged structures such as in ozone, not O2 in its ground state. Option C (2.5) would imply more bonding than a double bond, closer to a bond between two atoms with partial triple bond character, which does not match experimental O2 data. Option D (1) suggests a single bond and would predict a much longer and weaker bond than observed for O2, so it is not correct. Common Pitfalls: A common mistake is to confuse O2 with other oxygen species such as O2 2 minus or O2 plus, which have different bond orders. Another pitfall is to rely only on Lewis structures without understanding that paramagnetism requires molecular orbital theory for complete explanation. Students may also incorrectly count electrons in bonding and antibonding orbitals. Careful counting of valence electrons and systematic filling of molecular orbitals helps avoid such errors. Final Answer: The correct answer is: 2.

Question 11

When metallic atoms combine with nonmetallic atoms in ionic bonding, how do the metallic atoms generally behave in terms of electrons and ion formation?

Accepted Answer

Introduction / Context: Ionic bonding is a central concept in basic chemistry. It usually occurs between metals and nonmetals, where electrons are transferred rather than shared. Metals tend to lose electrons, and nonmetals tend to gain electrons. This electron transfer leads to the formation of oppositely charged ions that attract each other. This question asks what metallic atoms typically do when they combine with nonmetallic atoms in such reactions. Given Data / Assumptions: - We are considering reactions between metals and nonmetals, such as sodium with chlorine. - Metallic atoms are located on the left side of the periodic table. - Nonmetallic atoms are located on the right side of the periodic table. - Ionic bonding involves whole electron transfer from metal to nonmetal. Concept / Approach: Metals have relatively low ionisation energy and often have only one, two, or three valence electrons. It is easier for them to lose these electrons and achieve a noble gas configuration. When a metal atom loses one or more electrons, it becomes a positively charged ion called a cation. Nonmetals, by contrast, have higher electronegativity and tend to gain electrons to fill their valence shell, becoming negatively charged anions. In ionic compounds such as NaCl or CaCl2, the metal provides cations and the nonmetal provides anions. Step-by-Step Solution: Step 1: Recall that metals, such as sodium and magnesium, are typically found on the left side of the periodic table and have just a few valence electrons. Step 2: It is energetically favourable for these metals to lose their valence electrons to reach a stable noble gas configuration. Step 3: When a metal atom loses electrons, it now has more protons than electrons, resulting in a net positive charge and forming a cation. Step 4: Nonmetals like chlorine or oxygen gain the electrons released by metals and form negatively charged anions. Step 5: The electrostatic attraction between positively charged cations and negatively charged anions forms an ionic bond and produces an ionic compound. Step 6: Therefore, the correct description is that metallic atoms lose electrons and form positively charged cations. Verification / Alternative check: Consider the example of sodium chloride. Sodium (Na) has one valence electron and loses it to form Na plus, a cation. Chlorine (Cl) gains this electron to form Cl minus, an anion. The resulting Na plus and Cl minus ions combine to form NaCl. Similar patterns are seen in magnesium oxide, aluminium oxide, and many other ionic compounds. These examples consistently show that metallic atoms lose electrons and form cations when reacting with nonmetals. Why Other Options Are Wrong: Option B states that metals gain electrons and form anions, which describes nonmetal behaviour rather than metal behaviour. Option C combines losing electrons with forming anions, which is contradictory, because losing electrons leads to a positive charge, not a negative one. Option D suggests that metals gain electrons yet somehow form cations, which again conflicts with the rule that gaining electrons leads to negative charge. Common Pitfalls: Students sometimes confuse metals with nonmetals when memorising trends. Another common mistake is to think that gaining electrons always happens because atoms want to fill their shells, without considering that metals can reach stability more easily by losing electrons. To avoid confusion, remember a simple rule: metals lose electrons and become positive, while nonmetals gain electrons and become negative in typical ionic bonding situations. Final Answer: The correct answer is: They lose electrons and form positively charged cations.

Question 12

In basic biochemistry of carbohydrates, which of the following is not classified as a monosaccharide sugar?

Accepted Answer

Introduction / Context: Carbohydrates are an essential class of biomolecules, and they are often divided into monosaccharides, disaccharides, and polysaccharides. Monosaccharides are the simplest carbohydrate units, while disaccharides consist of two monosaccharides linked together. Recognising common examples of each type is important in biology and nutrition. This question asks you to identify which of the listed sugars is not a monosaccharide. Given Data / Assumptions: - The sugars listed are fructose, glucose, galactose, and sucrose. - Fructose, glucose, and galactose are simple sugars found in foods and in metabolism. - Sucrose is common table sugar used in households. - We classify sugars based on the number of simple sugar units in their structure. Concept / Approach: Monosaccharides are single sugar units such as glucose, fructose, and galactose. They cannot be hydrolysed into simpler carbohydrate units. Disaccharides are made of two monosaccharide units linked by a glycosidic bond and can be hydrolysed into those units. Sucrose is composed of one glucose molecule and one fructose molecule bonded together, which makes it a disaccharide rather than a monosaccharide. Therefore, sucrose is the sugar that does not belong in the monosaccharide category. Step-by-Step Solution: Step 1: Recall common monosaccharides: glucose (blood sugar), fructose (fruit sugar), and galactose (part of milk sugar) are all single sugar units. Step 2: Recognise that each of these monosaccharides has the general formula C6H12O6 but differs in structure and arrangement of atoms. Step 3: Consider sucrose, which is ordinary table sugar, usually extracted from sugarcane or sugar beet. Step 4: Understand that sucrose is formed when a glucose unit and a fructose unit join together via a glycosidic bond, making it a disaccharide. Step 5: Since monosaccharides cannot be hydrolysed to simpler sugars but disaccharides can, sucrose does not meet the definition of a monosaccharide. Step 6: Conclude that sucrose is the correct answer because it is not a monosaccharide. Verification / Alternative check: Biology and biochemistry textbooks classify glucose, fructose, and galactose as monosaccharides. Lactose (milk sugar), maltose, and sucrose are consistently listed as disaccharides. Laboratory hydrolysis of sucrose using acid or the enzyme sucrase splits sucrose into glucose and fructose, confirming that sucrose is composed of two monosaccharide units. Since monosaccharides cannot be further hydrolysed into smaller carbohydrate units, this property distinguishes sucrose from the other three options. Why Other Options Are Wrong: Option A (Fructose) is incorrect as a choice because fructose is a monosaccharide found in fruits and honey. Option B (Glucose) is incorrect because glucose is the primary monosaccharide used in cellular respiration and found in blood. Option C (Galactose) is incorrect because galactose is a monosaccharide that combines with glucose to form lactose, the sugar in milk. Common Pitfalls: Students sometimes confuse sucrose with glucose or fructose because all are sweet sugars commonly eaten in food. Another common mistake is to assume that all sweet sugars are monosaccharides. To avoid confusion, remember that the ending ose appears in many sugar names and does not itself indicate whether the sugar is mono, di, or polysaccharide. It is helpful to memorise that glucose, fructose, and galactose are monosaccharides, while sucrose is a disaccharide composed of glucose and fructose. Final Answer: The correct answer is: Sucrose.

Question 13

In periodic table trends, which of the following elements has the lowest electronegativity value among the choices given?

Accepted Answer

Introduction / Context: Electronegativity is a key periodic trend that describes how strongly an atom attracts electrons in a chemical bond. It plays an important role in predicting bond polarity, molecular structure, and reactivity. Across periods and down groups, electronegativity follows predictable patterns. This question asks which element from a small set has the lowest electronegativity value. Given Data / Assumptions: - The elements listed are bromine, fluorine, carbon, and lithium. - Fluorine is located in Group 17 (halogens) in Period 2. - Bromine is a halogen in Period 4. - Carbon is a nonmetal in Group 14 and Period 2. - Lithium is an alkali metal in Group 1 and Period 2. Concept / Approach: Electronegativity generally increases across a period from left to right and decreases down a group in the periodic table. Nonmetals, especially halogens, are highly electronegative, while alkali metals have low electronegativity. Fluorine is famously the most electronegative element. On the other hand, lithium, located on the far left of Period 2, has a relatively low electronegativity compared to the others in the list. Using these trend rules, we can identify lithium as the element with the lowest electronegativity among the given options. Step-by-Step Solution: Step 1: Recall the general trend: electronegativity increases from left to right and decreases from top to bottom. Step 2: Recognise that fluorine is the most electronegative element in the periodic table, so it cannot have the lowest electronegativity. Step 3: Bromine is a halogen lower in the same group as fluorine, so it is still quite electronegative, though less than fluorine. Step 4: Carbon is a nonmetal in the upper right part of the table, giving it a moderate to high electronegativity value. Step 5: Lithium is an alkali metal at the far left of Period 2. Alkali metals have low electronegativity because they prefer to lose their single valence electron. Step 6: Compare all four and conclude that lithium has the lowest electronegativity among the options listed. Verification / Alternative check: Approximate electronegativity values on the Pauling scale support this conclusion. Fluorine has a value of about 3.9, bromine about 2.8, carbon about 2.5, and lithium about 1.0. Clearly, lithium has the smallest numerical value, confirming it is the least electronegative element in this group. These values are consistent across standard chemical references and confirm the periodic trend reasoning. Why Other Options Are Wrong: Option A (Bromine) is a halogen with relatively high electronegativity compared to metals and many other nonmetals, so it is not the lowest. Option B (Fluorine) is the most electronegative element in the periodic table, so it is the opposite of the correct answer. Option C (Carbon) has a moderate electronegativity but is still significantly more electronegative than lithium. Common Pitfalls: Students sometimes mix up electronegativity with other properties such as ionisation energy or atomic radius. Another common mistake is to assume that all nonmetals are automatically more electronegative than all metals without checking specific positions in the periodic table. While this is often true, it is always safer to recall the exact trend and, when possible, approximate values. Remembering that alkali metals like lithium have very low electronegativity values helps avoid confusion. Final Answer: The correct answer is: Lithium.

Question 14

In corrosion chemistry, rust forms when iron reacts with which gas in the presence of moisture in air?

Accepted Answer

Introduction / Context: Rusting is a common example of corrosion and is often seen on iron objects exposed to air and moisture. It is a chemical process in which iron is oxidised to form hydrated iron oxides. Understanding which components of air are actually responsible for rusting helps clarify basic concepts in environmental chemistry and everyday science. This question focuses on identifying the main gas that reacts with iron during rust formation. Given Data / Assumptions: - Rust is commonly observed on iron and steel when they are exposed to damp air. - Air is mainly a mixture of nitrogen, oxygen, and small amounts of other gases including carbon dioxide and water vapour. - Rust is chemically a form of hydrated iron oxide. - We consider normal atmospheric conditions, not special industrial environments. Concept / Approach: Rusting is an oxidation reaction in which iron reacts with oxygen in the presence of water to form iron oxides, typically Fe2O3·xH2O. Oxygen is the key oxidising agent that combines with iron. Water acts as a medium that facilitates the flow of electrons and ions in electrochemical corrosion. Nitrogen is largely inert under these conditions, and carbon dioxide plays only a secondary role in some corrosion processes. Therefore, oxygen gas from the air is the main gas responsible for rust formation on iron. Step-by-Step Solution: Step 1: Recall the basic rusting reaction, often simplified as iron plus oxygen plus water gives hydrated iron oxide. Step 2: Recognise that the chemical change involves oxidation of iron atoms to iron ions, which requires oxygen as the electron acceptor. Step 3: Note that water is necessary for rusting but is not the primary gas that reacts; it provides the medium in which ions can move. Step 4: Identify oxygen gas as the component in air that reacts chemically with iron to form iron oxides. Step 5: Compare with nitrogen, which does not usually react with iron at ordinary temperatures and pressures, and carbon dioxide, which plays a minor role only in specific conditions. Step 6: Conclude that oxygen gas from the air is the correct answer. Verification / Alternative check: The chemical equations representing rust formation show iron reacting with oxygen. For example, 4Fe + 3O2 + xH2O gives 2Fe2O3·xH2O. Industrial corrosion studies and school experiments that remove oxygen from water demonstrate that rusting is greatly slowed or prevented when oxygen is absent, even if water is present. On the other hand, removing nitrogen or small amounts of carbon dioxide does not stop rusting. These observations confirm that oxygen is the critical gas in rust formation. Why Other Options Are Wrong: Option A (Water vapour only) is incomplete because water alone without oxygen does not produce rust; oxygen is needed for oxidation. Option C (Nitrogen gas) is incorrect because nitrogen is relatively inert at room temperature and does not significantly react with iron to cause rust. Option D (Carbon dioxide gas) may contribute to some forms of corrosion in the presence of water, but the classic rusting of iron is primarily due to oxygen, not carbon dioxide. Common Pitfalls: Students sometimes overemphasise the role of water and think that moisture alone causes rust. While water is necessary, it is only part of the process; oxygen from the air is equally essential. Another common misconception is to blame carbon dioxide or other minor air components without recognising that rust is an oxidation process requiring oxygen. To correctly identify the cause of rusting, always remember the involvement of iron, oxygen, and water together. Final Answer: The correct answer is: Oxygen gas from the air.

Question 15

In periodic trends, how does the atomic mass of the first 30 elements in the periodic table generally change as the atomic number increases?

Accepted Answer

Introduction / Context: Understanding how atomic mass changes with atomic number is a key part of learning the structure of the periodic table. The first 30 elements, from hydrogen to zinc, provide a good sample for seeing general trends. Atomic mass is the weighted average mass of an element isotopes, measured in atomic mass units. This question asks you to identify the overall pattern in atomic mass for the first 30 elements as their atomic number increases. Given Data / Assumptions: - The first 30 elements range from hydrogen (atomic number 1) to zinc (atomic number 30). - Atomic number counts the number of protons in the nucleus. - Atomic mass depends on the total number of protons and neutrons, and on isotopic abundances. - We are looking for the general trend, not small irregular variations. Concept / Approach: In the periodic table, as atomic number increases, each successive element has an additional proton and typically one or more additional neutrons. As a result, the atomic mass generally increases with atomic number. There can be minor irregularities due to isotopes and binding energy differences, but these do not reverse the overall upward trend. When you examine the list of the first 30 elements, you see that hydrogen has the smallest atomic mass, and heavy elements like zinc have much larger atomic masses. Step-by-Step Solution: Step 1: Consider the starting point. Hydrogen (atomic number 1) has an atomic mass close to 1 atomic mass unit. Step 2: Move to helium (atomic number 2), which has an atomic mass near 4 units, already larger than hydrogen. Step 3: Continue along the sequence: lithium, beryllium, boron, carbon, nitrogen, oxygen and so on. Each new element generally has a larger atomic mass than the previous one. Step 4: By the time you reach elements like calcium, iron, and zinc, their atomic masses are many times greater than that of hydrogen. Step 5: Note that there may be small deviations in how much the mass increases from one element to the next, but there is no systematic decrease in mass as atomic number grows. Step 6: Conclude that the overall pattern for the first 30 elements is that atomic mass generally increases as atomic number increases. Verification / Alternative check: You can verify this trend by looking at any periodic table that lists approximate atomic masses beneath each element symbol. Starting at hydrogen and moving step by step to zinc, you will see the numbers rising from about 1 to values around 65 for zinc. Although isotopic composition sometimes creates slight irregularities in exact values, these never change the basic rule that atomic mass becomes larger as atomic number increases across this range. This confirms that a general increase is the correct description. Why Other Options Are Wrong: Option B claims that atomic mass decreases as atomic number increases, which contradicts observations because heavier elements have more nucleons. Option C suggests that atomic mass stays exactly the same for all first 30 elements, which is clearly false since hydrogen and zinc differ greatly in mass. Option D states that atomic mass changes randomly with no trend, but periodic tables and nuclear physics both show a clear overall increase tied to adding protons and neutrons. Common Pitfalls: Some learners focus too much on minor irregularities, such as cases where the increase in mass from one element to the next is small, and mistakenly think there is no pattern. Others confuse atomic number with atomic mass and think they have no relationship. To avoid confusion, remember that atomic number counts protons and that heavier elements usually also carry more neutrons, so the mass naturally increases as you progress through the periodic table. Final Answer: The correct answer is: Atomic mass generally increases as atomic number increases across the first 30 elements.

Question 16

In the periodic table, which one of the following elements is classified as a transition element (d-block metal)?

Accepted Answer

Introduction / Context: Transition elements, also called transition metals, occupy the central block of the periodic table and are characterised by partially filled d subshells in one or more of their common oxidation states. They often show variable oxidation states, form coloured compounds, and can act as good catalysts. This question asks you to identify which among the given symbols represents a transition element. Given Data / Assumptions: - The options are hydrogen (H), arsenic (As), iron (Fe), and strontium (Sr). - Transition elements are typically found in groups 3 to 12 of the periodic table. - Elements outside the d block are not considered transition elements in the standard sense. Concept / Approach: The definition of a transition element usually requires that the element has an incomplete d subshell in either its neutral atom or in one of its common ions. Elements in the central d block of the periodic table, such as iron, copper, and nickel, meet this criterion. Hydrogen is a nonmetal located at the top of Group 1, arsenic is a metalloid in the p block, and strontium is an alkaline earth metal in Group 2. Iron, with atomic number 26, is a typical d block metal and a classic example of a transition element. Step-by-Step Solution: Step 1: Locate hydrogen (H) in the periodic table. It is placed in Group 1 but is a nonmetal and belongs to the s block, not the d block. Step 2: Locate arsenic (As). It is in Group 15 in the p block and is classified as a metalloid, not a transition metal. Step 3: Locate iron (Fe). It is found in the central region of the periodic table in the d block, typically classified as a transition metal. Step 4: Locate strontium (Sr). It is in Group 2 as an alkaline earth metal, in the s block, not in the transition metal block. Step 5: Check the electron configuration of iron. It has partially filled 3d orbitals in many oxidation states and clearly satisfies the transition element definition. Step 6: Conclude that Fe (iron) is the only transition element among the options. Verification / Alternative check: Textbooks and periodic table charts list iron as a transition metal and often use it as a key example when discussing transition metal properties such as magnetism, formation of coloured compounds, and catalytic activity. Hydrogen is discussed separately as a unique nonmetal, arsenic is grouped with p block elements, and strontium is treated as an alkaline earth metal. This classification consistently supports the conclusion that iron is the transition element in the given list. Why Other Options Are Wrong: Option A (H) is wrong because hydrogen is a nonmetal and does not have d electrons in the ground state. Option B (As) is incorrect because arsenic is a p block metalloid, not a d block transition metal. Option D (Sr) is incorrect because strontium is an alkaline earth metal in Group 2 and belongs to the s block. Common Pitfalls: Some students confuse the term metal with transition metal and may think that any metallic element is a transition metal. Others may misread the periodic table layout and not clearly see the separation between s, p, d, and f blocks. To avoid these mistakes, remember that transition elements are specifically the d block metals in groups 3 to 12, and that classic examples are iron, copper, and nickel, not hydrogen or alkaline earth metals. Final Answer: The correct answer is: Fe (iron).

Question 17

In chemical bonding theory, which statement best describes the electronegativity of an element atom?

Accepted Answer

Introduction / Context: Electronegativity is a central idea in chemical bonding that helps explain bond polarity, dipole moments, and the behaviour of atoms in molecules. Different elements have different tendencies to attract shared electrons in a covalent bond. Understanding the correct definition of electronegativity is essential for predicting which side of a bond will be partially negative or positive. This question asks you to choose the best description of electronegativity from several related statements. Given Data / Assumptions: - We are dealing with atoms that form covalent or polar covalent bonds. - Electronegativity is used to compare different elements, usually on a scale such as the Pauling scale. - The options describe various abilities related to electrons: gaining, losing, attracting, or sharing. - We need the definition that matches standard textbook usage. Concept / Approach: Electronegativity is defined as the tendency or ability of an atom in a chemical bond to attract the bonding electrons toward itself. It is not the same as ionisation energy or electron affinity, although it is related to both. Electronegativity explains why polar covalent bonds form when atoms with different electronegativities share electrons unequally. The correct definition must emphasise attraction for shared electrons in a bonding situation, not just losing or gaining electrons to form ions. Step-by-Step Solution: Step 1: Recall that in a covalent bond, electrons are shared between two atoms, but they may not be shared equally if the atoms have different electronegativities. Step 2: The atom with higher electronegativity pulls the shared electron pair closer, acquiring a partial negative charge, while the other atom becomes partially positive. Step 3: Review the standard definition: electronegativity is the ability of an atom to attract electrons in a chemical bond toward itself. Step 4: Compare this definition with Option B, which states that electronegativity is the ability to attract electrons during bond formation. This matches the textbook description. Step 5: Check that the other options describe different properties, such as losing electrons or simply sharing them, which are not correct definitions of electronegativity. Step 6: Conclude that Option B is the best and correct statement. Verification / Alternative check: Chemistry references and data tables define electronegativity in almost identical words: the tendency of an atom to attract the shared electron pair in a covalent bond. For example, fluorine has a very high electronegativity and strongly attracts electrons, leading to very polar bonds when paired with less electronegative elements. Numerical electronegativity scales are always used in the context of bonding electrons, not free electrons, which confirms that Option B, focused on bond formation, is the correct interpretation. Why Other Options Are Wrong: Option A refers to producing energy while losing an electron, which relates more to ionisation energy or energy change, not to electronegativity. Option C describes the ability to lose electrons and form a cation, which is more closely related to metallic behaviour and low electronegativity, not to electronegativity itself. Option D mentions sharing electrons during covalent bond formation but does not emphasise attraction; equal sharing occurs when electronegativities are similar, not as a definition of electronegativity. Common Pitfalls: Students sometimes confuse electronegativity with electron affinity or ionisation energy because all involve electrons and energy changes. Another common error is to think of electronegativity as a property only of isolated atoms rather than atoms in bonded situations. To avoid these mistakes, always associate electronegativity with the pull on shared electrons in a bond and remember that it is a relative scale comparing different elements. Final Answer: The correct answer is: Electronegativity of an atom is its ability to attract electrons toward itself during bond formation.

Question 18

In the structure of liquid water, what name is given to the attractive force that occurs between neighboring water molecules?

Accepted Answer

Introduction / Context: Water has many unusual properties, such as a relatively high boiling point and surface tension, which are crucial for life. These properties arise largely from the special interactions between water molecules, not just within a single molecule. Inside each water molecule, the O H bonds are polar covalent, but between different water molecules there is a specific intermolecular attraction known as hydrogen bonding. This question asks you to name the bond or interaction between neighboring water molecules. Given Data / Assumptions: - A water molecule has the formula H2O and bent shape. - The O H bonds inside one water molecule are polar covalent bonds. - Neighboring water molecules attract each other strongly compared to many other small molecules. - We are interested in the interaction between different water molecules, not within a single molecule. Concept / Approach: Hydrogen bonding is a special type of strong dipole dipole interaction that occurs when hydrogen is covalently bonded to a very electronegative atom such as oxygen, nitrogen, or fluorine, and is then attracted to a lone pair on another electronegative atom in a nearby molecule. In water, the oxygen atom is highly electronegative and pulls electron density away from the hydrogen atoms, making them partially positive. These partially positive hydrogens are strongly attracted to lone pairs on the oxygen atoms of neighboring water molecules, creating hydrogen bonds. Step-by-Step Solution: Step 1: Recognise that within a single water molecule, each O H bond is a polar covalent bond because electrons are shared unevenly between oxygen and hydrogen. Step 2: Understand that the water molecule has a bent shape, leading to an overall molecular dipole with a partially negative region near oxygen and partially positive regions near hydrogen. Step 3: Between molecules, the partially positive hydrogen of one water molecule is attracted to the lone pair electrons on the partially negative oxygen of another water molecule. Step 4: This strong, directional attraction is called a hydrogen bond, a specific type of intermolecular force. Step 5: Hydrogen bonds are much weaker than covalent bonds within a molecule but stronger than many other intermolecular forces, which explains water unusual physical properties. Step 6: Therefore, the correct name for the bond between water molecules is a hydrogen bond. Verification / Alternative check: Evidence for hydrogen bonding in water includes its abnormally high boiling point compared to other similar sized molecules like H2S, and its high specific heat capacity. These properties cannot be explained by simple polar interactions alone and require stronger intermolecular forces. Structural studies and theoretical calculations confirm that each water molecule can form hydrogen bonds to several neighbours, producing a network in the liquid and a lattice in ice. These observations strongly support the conclusion that the intermolecular attraction in water is hydrogen bonding. Why Other Options Are Wrong: Option A refers only to general polar interactions and does not capture the specific, stronger hydrogen bond nature of water intermolecular forces. Option B mentions a polar covalent bond between separate water molecules, which is incorrect because covalent bonds are within molecules, not between them. Option D claims that none of the options apply, which is wrong because hydrogen bonding is the standard textbook explanation for water intermolecular attraction. Common Pitfalls: Students often confuse the polar covalent O H bonds inside the water molecule with the hydrogen bonds between molecules. Another common error is to think that hydrogen bonding is the same as covalent bonding or that it requires hydrogen alone without considering the electronegative partner. To avoid these mistakes, remember that hydrogen bonds are intermolecular forces that involve a hydrogen atom already bonded to an electronegative atom and interacting with a lone pair on another electronegative atom. Final Answer: The correct answer is: A hydrogen bond between the hydrogen of one molecule and the oxygen of another.

Question 19

In the general structure of amino acids, which three components are common to all standard amino acids found in proteins?

Accepted Answer

Introduction / Context: Amino acids are the building blocks of proteins and have a characteristic core structure that is repeated in all standard protein forming amino acids. Understanding this core structure is fundamental in biochemistry and molecular biology because it explains how amino acids link together and why they have different properties. This question asks you to identify which three components are common to all amino acids found in proteins. Given Data / Assumptions: - We are considering the standard alpha amino acids that are incorporated into proteins. - Each amino acid has a central alpha carbon atom. - Each amino acid has an amino group and a carboxyl group attached to this central carbon. - A variable side chain, called the R group, is also attached to the central carbon and differs among amino acids. Concept / Approach: The general formula for an alpha amino acid can be written as H2N–CH(R)–COOH, where the central carbon (the alpha carbon) is bonded to four different groups: an amino group (H2N), a carboxyl group (COOH), a hydrogen atom, and a variable side chain R. The amino group and carboxyl group are present in every amino acid, while the R side chain varies and determines the specific identity and properties of each amino acid. Other biological molecules such as nucleotides and lipids contain groups like phosphate, sugars, or glycerol, but these are not part of the basic amino acid structure. Step-by-Step Solution: Step 1: Recall the general structure of an amino acid as H2N–CH(R)–COOH. Step 2: Identify the amino group (H2N) attached to the central alpha carbon; this group gives amino acids their name. Step 3: Identify the carboxyl group (COOH) attached to the same central carbon; this group gives amino acids acidic properties. Step 4: Recognise that the central alpha carbon is bonded to both the amino group and the carboxyl group as well as a hydrogen and an R side chain. Step 5: Understand that the R side chain varies between amino acids (for example, a hydrogen in glycine or a methyl group in alanine), but the presence of some R group attached to the alpha carbon is common to all. Step 6: Conclude that the three core components common to all amino acids are the amino group, the carboxyl group, and the central alpha carbon with its attached variable R side chain. Verification / Alternative check: Standard biochemistry diagrams of amino acids always show the same backbone: an amino group, the central alpha carbon, and a carboxyl group. The only structural difference among amino acids is the nature of the R group attached to the alpha carbon. In contrast, nucleotides (building blocks of DNA and RNA) consist of a sugar, a phosphate group, and a nitrogenous base, while triglycerides (common fats) contain a glycerol backbone and three fatty acid chains. These structures are clearly distinct from amino acids and support the identification of the correct set of components. Why Other Options Are Wrong: Option B describes a nucleotide structure (amino group, phosphate, nitrogenous base) rather than an amino acid structure. Option C combines a carboxyl group with a sugar and a long fatty acid chain, which does not match the standard structure of amino acids and instead resembles parts of other biomolecules. Option D describes a typical triglyceride structure with glycerol and three fatty acid chains, not an amino acid. Common Pitfalls: Students sometimes confuse different classes of biomolecules because they all contain carbon, hydrogen, and oxygen, and often nitrogen or phosphorus. Another common mistake is to forget that the R group, although variable in composition, is structurally present in every amino acid. To avoid confusion, always visualise the amino acid backbone as amino group plus alpha carbon plus carboxyl group, with the R group branching off from the alpha carbon. Final Answer: The correct answer is: An amino group, a carboxyl group, and a central alpha carbon with a variable R side chain.

Question 20

In cell biology, which important polymer is located in the nucleus of eukaryotic cells and carries genetic information?

Accepted Answer

Introduction / Context: The nucleus is often called the control centre of the eukaryotic cell because it houses the cell genetic material. This genetic information is stored in the form of a specific polymer that encodes instructions for building and maintaining the organism. Understanding which polymer carries this information is fundamental in biology, genetics, and biotechnology. This question asks you to identify the key polymer located in the nucleus that stores genetic information. Given Data / Assumptions: - We are considering eukaryotic cells, which contain a membrane bound nucleus. - Genetic information is stored in a long chain polymer made of repeating nucleotide units. - Chromosomes are structures within the nucleus. - Ribosomes and many proteins may be found outside the nucleus, especially in the cytoplasm. Concept / Approach: Deoxyribonucleic acid, or DNA, is the polymer that stores genetic information in almost all living organisms. It is located mainly in the nucleus in eukaryotic cells, packaged into structures called chromosomes. Each chromosome is made of DNA wrapped around proteins called histones. Ribosomes are complexes of RNA and protein responsible for protein synthesis, and they are found in the cytoplasm and on the rough endoplasmic reticulum. While proteins are essential cellular components, they do not act as the primary long term information storage polymer. Step-by-Step Solution: Step 1: Recall that hereditary information is encoded in sequences of nucleotides in a polymer called DNA. Step 2: Recognise that in eukaryotic cells, most DNA is enclosed within the nucleus and arranged into chromosomes. Step 3: Understand that chromosomes are structures composed of DNA and proteins, but they are not themselves a single type of polymer. Step 4: Note that ribosomes are responsible for protein synthesis and consist of ribosomal RNA and protein, mainly located in the cytoplasm and on the rough endoplasmic reticulum. Step 5: Realise that many proteins are found throughout the cell and perform various functions, but they are not the main genetic information storage molecules. Step 6: Conclude that DNA is the important polymer located in the nucleus that carries genetic information. Verification / Alternative check: Classical experiments by scientists such as Avery and Hershey Chase showed that DNA, not protein, is the genetic material responsible for inheritance in bacteria and viruses. In eukaryotic cells, microscopy and staining techniques reveal chromosomes that contain DNA. Molecular biology has further confirmed that genes are specific sequences of nucleotides in DNA, and that replication of DNA passes genetic information to daughter cells. All of this evidence points to DNA as the polymer in the nucleus that carries genetic information. Why Other Options Are Wrong: Option A (Ribosomes) are not primarily located in the nucleus and do not store genetic information; instead, they read messenger RNA to synthesise proteins. Option B (Chromosomes) are structures made of DNA and proteins, but the information storing polymer within them is DNA, so naming chromosomes alone is less precise than naming DNA. Option D (Proteins in the cytoplasm) describes functional molecules but not the main genetic information storage polymer, and many are outside the nucleus. Common Pitfalls: Students sometimes confuse chromosomes and DNA and use the words as if they were identical. While closely related, chromosome refers to a compact structure containing DNA plus proteins, whereas DNA is the specific polymer that carries the code. Another mistake is to think that ribosomes or cytoplasmic proteins hold genetic information because they are essential for cell function. Keeping the distinction clear between information storage (DNA) and functional expression (RNA and proteins) helps avoid confusion. Final Answer: The correct answer is: DNA (deoxyribonucleic acid).

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