Chemistry Regents

Question 1 · HS-PS1-1

Two atoms in the same group of the periodic table are most likely to have:

(A) The same mass number
(B) The same number of valence electrons
(C) The same number of neutrons
(D) The same atomic radius

(B). Same group = same valence electrons = similar chemistry. Atomic radius increases down a group. Mass number and neutron count are unrelated to group position.

Question 2 · HS-PS1-1

The ground-state electron configuration of phosphorus (atomic number 15) is:

(A) 2-8-5
(B) 2-7-6
(C) 2-13
(D) 2-8-4-1

(A) 2-8-5. Fill shells inside out: 2 + 8 + 5 = 15. (B) skips shell capacity → excited state. (C) violates shell 2's max of 8. (D) totals 15 but uses shell 4 unnecessarily. The 2025 tables don't print configurations, so derive them.

Question 3 · HS-PS1-3

Why do ionic compounds have higher melting points than most molecular covalent compounds?

(A) Ionic compounds contain more atoms
(B) Electrostatic attractions between ions are stronger than IMFs between molecules
(C) Ionic compounds are always solids at room temperature
(D) Molecular compounds have weaker covalent bonds

(B). Bulk properties from electrical forces between particles. Ionic = strong ion-ion attractions in a lattice. Molecular = only weaker IMFs hold molecules together. (D) confuses IMFs with covalent bonds within molecules.

Question 4 · HS-PS1-3

HF has an unusually high boiling point for its molar mass. Best explanation:

(A) HF molecules are heavier than expected
(B) HF forms ionic bonds with itself
(C) HF molecules form hydrogen bonds between molecules
(D) HF molecules are nonpolar

(C). H bonded to F (also N or O) creates hydrogen bonding, the strongest IMF. Stronger IMFs require more energy to break during boiling. Water shows the same effect.

Question 5 · HS-PS1-1

Across Period 3 from sodium (Na) to chlorine (Cl), atomic radius:

(A) Increases, because more electrons are added
(B) Decreases, because nuclear charge increases while shells stay the same
(C) Stays the same
(D) Increases, because atomic mass increases

(B). Across a period: protons add but no new shell. Same outer electrons pulled tighter by stronger nuclear charge → atoms shrink. Tested constantly.

Question 6 · HS-PS1-3

Which substance will NOT conduct electricity when dissolved in water?

(A) KCl
(B) NaOH
(C) C₆H₁₂O₆ (glucose)
(D) HCl

(C) Glucose. Molecular covalent compound dissolves as whole, uncharged molecules. No ions → no charge carriers → no conductivity. KCl and NaOH are ionic. HCl is a strong acid that ionizes in water.

Question 7 · HS-PS1-9

A gas occupies 4.0 L at 300 K and 1.0 atm. What volume at 600 K and 2.0 atm?

(A) 1.0 L
(B) 2.0 L
(C) 4.0 L
(D) 8.0 L

(C) 4.0 L. P₁V₁/T₁ = P₂V₂/T₂. (1.0)(4.0)/300 = (2.0)(V₂)/600 → V₂ = 4.0 L. Conceptually: doubling T would double V; doubling P would halve V. They cancel.

Question 8 · HS-PS1-9

Which sample of ideal gas occupies the largest volume at STP?

(A) 1 mole of helium
(B) 1 mole of nitrogen gas
(C) 1 mole of oxygen gas
(D) All occupy the same volume

(D). At STP, 1 mole of any ideal gas = 22.4 L. Identity doesn't matter under the ideal-gas assumption. On Reference Tables Page 3.

Question 9 · HS-PS1-10

A solubility curve shows 80 g of KNO₃ dissolves in 100 g water at 50°C. A student dissolves 60 g of KNO₃ in 100 g water at 50°C. The solution is:

(A) Saturated
(B) Unsaturated
(C) Supersaturated
(D) Cannot be determined

(B) Unsaturated. 60 g < 80 g maximum → more could still dissolve. Saturated = exactly 80 g. Supersaturated = more than 80 g (unstable).

Question 10 · HS-PS1-10

How does temperature typically affect the solubility of a gas in water?

(A) Solubility increases with temperature
(B) Solubility decreases with temperature
(C) Independent of temperature
(D) First increases, then decreases

(B). Gas solubility decreases as T rises (gas molecules escape water more easily when hot). Opposite of most solid solutes. Why warm soda goes flat faster.

Question 11 · HS-PS1-8

In ²³⁰₉₀Th → ²²⁶₈₈Ra + X, particle X is:

(A) ⁰₋₁e (beta)
(B) ⁴₂He (alpha)
(C) ¹₀n (neutron)
(D) ⁰₊₁e (positron)

(B) ⁴₂He. Mass: 230 = 226 + 4. Atomic: 90 = 88 + 2. Mass 4, atomic 2 = alpha particle.

Question 12 · HS-PS1-8

Isotope has half-life of 10 days. Starting with 80 g, how much remains after 30 days?

(A) 40 g
(B) 20 g
(C) 10 g
(D) 5 g

(C) 10 g. 30 days = 3 half-lives. 80 → 40 → 20 → 10. Each half-life cuts the sample in half.

Question 13 · HS-PS1-3

Which molecule is polar?

(A) CO₂ (linear)
(B) CH₄ (tetrahedral)
(C) H₂O (bent)
(D) N₂

(C) H₂O. Bent shape → polar O-H bonds don't cancel. CO₂ is linear with polar bonds pointing opposite (cancel → nonpolar). CH₄ symmetric (nonpolar). N₂ identical atoms (nonpolar).

Question 14 · HS-PS1-3

Which property is best explained by metallic bonding?

(A) Brittleness of NaCl
(B) Conductivity of copper wire
(C) Low melting point of methane
(D) Hardness of diamond

(B). Delocalized electrons move freely → carries current. (A) is ionic. (C) is IMFs. (D) is network covalent.

Question 15 · HS-PS1-3

According to Coulomb's law (qualitative), force between charges increases when:

(A) Charges become larger or distance becomes smaller
(B) Charges become smaller or distance becomes larger
(C) Only masses change
(D) Particles move faster

(A). F = q₁q₂/d². Larger charges → larger numerator. Smaller distance squared → smaller denominator. Either increases F. Why small, highly-charged ions (Mg²⁺) form especially strong attractions.

Question 16 · HS-PS2-6

Graphite is soft (lubricant); diamond is the hardest natural substance. Both are pure carbon. Best explanation:

(A) Diamond contains more carbon atoms per gram
(B) Graphite has weaker covalent bonds
(C) Graphite has layered sheets with weak forces between; diamond is a 3D network of strong covalent bonds
(D) Diamond is heavier

(C). Same atoms, different particle structure → different bulk properties. Central HS-PS2-6 idea. Graphite sheets slide → lubricant. Diamond 3D network → no slip plane → hard.

Question 17 · HS-PS1-1

Ground-state configuration 2-8-7. Most likely ion charge:

(A) +1
(B) -1
(C) +7
(D) -2

(B) -1. 2+8+7=17 → chlorine. 7 valence electrons → gains 1 → Cl⁻. Losing 7 to form Cl⁷⁺ is energetically prohibitive.

Question 18 · HS-PS1-10

0.50 mol NaCl in 2.0 L solution. Molarity is:

(A) 0.25 M
(B) 0.50 M
(C) 1.0 M
(D) 2.0 M

(A) 0.25 M. n = MV → M = n/V = 0.50/2.0 = 0.25 M. Reference Tables Page 3.

Question 19 · HS-PS1-8

In fission, a heavy nucleus splits and releases:

(A) Small chemical energy from broken bonds
(B) Large nuclear energy from small mass loss
(C) Heat from IMFs
(D) Light from electron transitions

(B). E = mc². Tiny mass difference → enormous energy because c² is huge. Chemical bonds, IMFs, electron transitions all involve much smaller energies.

Question 20 · HS-PS1-1

Two elements with most similar chemical behavior are most likely:

(A) Lithium and beryllium (Period 2 neighbors)
(B) Carbon and silicon (same group)
(C) Sodium and chlorine (Period 3 ends)
(D) Helium and lithium (consecutive)

(B). Same group → same valence electrons → same chemistry. C and Si both Group 14, both 4 valence electrons, both form 4 covalent bonds.