DAT Inorganic Chemistry (Natural Sciences) Practice Answers

Answers and Explanations

1. B: Density is defined as mass per unit volume. The mass of the gold is given as 521 grams. The volume of this can be calculated from the increase in volume when added to the water:
77.0 ml – 50 ml = 27 ml
521 g/27 ml = 19.296 g/ml


2. A: The concentration of the ampicillin suspension is 250 mg/5.0 ml, which is 0.05 grams per ml. The needed volume of solution can be calculated as follows:
0.125 grams are needed
0.125 g/0.05 g/ml = (0.125 gallon) / (0.05 gallon/ml) = 2.5 ml

3. A: Self-explanatory; see the periodic table.

4. A: The mass number of an element is defined as the number of protons (atomic number) plus the number of neutrons in the nucleus. The atomic number of potassium is 19, so there are 19 protons. If the atom of potassium has 20 neutrons, the mass number would be 20 + 19, or 39.

5. E: The atomic number of aluminum is 13. This means that the number of protons is 13. Therefore, in a neutral aluminum atom, there are 13 electrons. The first energy level of an atom can hold two electrons. The second and third can each hold eight electrons. In the case of aluminum, this means that the second energy level would be filled, leaving three electrons to occupy the third energy level (13 – 2 – 8 = 3).

6. D: Elements in Group IIA have two electrons in the highest energy level. For example, beryllium (Be), with an atomic number of 4, has two electrons in the first energy level and two in the second, leaving that shell only partially filled with room for six additional electrons. Similarly, magnesium (Mg) with an atomic number of 12 has the occupied energy levels containing two, eight, and two electrons. Again, the highest energy level is unfilled by six electrons. Therefore, to form an ion, the Group IIA elements lose two electrons, rather than gain six, to achieve a net charge of 2+.

7. E: Electronegativity refers to the ability of an atom to attract bonding electrons. Nonpolarity, polarity, and ionic character refer to the charge distributions on atoms and molecules.

8. A: To determine the number of moles of K2SO4 in 15.0 grams, the molecular weight of the substance must be calculated. The molecular weight (MW) is the sum of the atomic weights (AW) of the atoms in the molecule, as shown in the following table:

atom AW # of atoms/molecule contribution to MW
K 39.1 2 78.2
S 32.1 1 32.1
O 16.0 4 64.0
total 174.3

Having determined that the MW of K2SO4 is 174.3 g/mole. We can calculate the number of moles of the compound in 15.0 grams as follows:
15.0 g /174.3 g mole-1 = 15.0 g /174.3 gallon mole-1 = 15.0/174.3 moles = 0.0861 moles

9. C: Ionic bonds generally involve metals and nonmetals and involve one element losing its valence electrons to the other. In general, elements that are close together on the periodic table are more likely to share valence electrons and form covalent bonds.

10. E: In an exothermic reaction, heat is generated when the reactants give rise to product. In such reactions, the reactants have a higher energy level. The energy released, or the heat of reaction, corresponds to the difference in energy level between the reactants and the reaction products.

11. E: In pure water there is an equal concentration, 10-7 M, of hydronium ions and hydroxyl ions. That is, [H30+] = [OH-] = 10-7 M

The ionization product constant of water is the product of the hydronium and hydroxyl ion concentrations. That is, for neutral water
10-7 M x 10-7 M = 10-14 M

Thus, the product of the hydronium and hydroxyl ion concentrations will always be 10-14 M. In this case, the hydroxyl ion concentration is 10-12 M. Thus, we can calculate
[H30+] x 10-12 M = 10-14 M
[H30+] = 10-14 M/10-12 M = 10-2M

12. B: The balanced chemical equation given in this question shows that one mole of silver nitrate (AgNO3) in the presence of sodium chloride (NaCl) will yield the same amount, one mole, of silver chloride (AgCl). We can calculate:
MW of AgNO3 = (107.9) + 14.01 + (3 x 16.0) = 169.91
MW of AgCl = (107.9) + (35.5) = 143.4
100 g of AgNO3 = 100 g/169.91 g per mole = 0.589 moles
0.589 mole x 143.4 g per mole = 84.4 grams

13. D: When a chemical reaction occurs, bonds are generally broken and formed. For the reaction to begin, bonds must be broken in the reactants before any rearrangement of chemical bonds can occur. The reaction energy is defined as the energy needed to break the bonds between the atoms in the reactants.

14. C: Nuclear fission generally involves the bombardment of a fissible isotope, such as 235U, with neutrons. Each atom that is split by a neutron also liberates at least one additional neutron. Thus, each nuclear fission event yields at least two neutrons, which can each cause more fission events, propagating a chain reaction.

15. B: Self-explanatory.

16. B: Kilograms and grams are units of mass, while centimeters and meters are units of length.

17. D: While there is an element, Einsteinium (symbol Es), named after Albert Einstein, he was not a chemist.

18. B: Magnesium and strontium are both in group IIA of the periodic table. They are alkaline earth metals, which have similar chemical properties.

19. B: While chemical equations must obey all laws of thermodynamics, the stoichiometry of chemical equations is constrained by conservation of mass. A balanced chemical equation has the same number of atoms, or mass, of each element on both sides of an equation.

20. B: At equilibrium, chemical reactions that convert product back to starting materials take place at exactly the same rate as the reactions that convert starting materials to products. Hence, no net change in the system can be observed. A system may appear to be at equilibrium if no net change is being observed, but this does not preclude the occurrence of net changes at a rate too slow to be readily observed. Chemical equilibrium is affected by temperature changes, and will adjust to the particular prevailing temperature. The position of a chemical equilibrium is not determined by temperature, but the rate at which equilibrium is achieved is temperature dependent.

DAT Inorganic Chemistry (Natural Sciences) Practice Questions