NY Regents Chemistry Exam Flashcards

Q: How many flashcards are in this set?

This set includes 50 original cards covering atomic structure and periodic trends, bonding, the mole and stoichiometry, gases and solutions, kinetics and equilibrium, acids/bases/salts, redox and electrochemistry, organic chemistry, nuclear chemistry, energy, and reference-table skills.

Q: Are reference tables provided on the NY Regents Chemistry exam?

Yes. The Reference Tables for Physical Science: Chemistry, 2025 Edition, are supplied during the exam. They hold constants, formulas, solubility data, activity series, and properties, so memorize how to read each table rather than memorizing every value.

Question 1

What does the atomic number of an element tell you, and why can't it change without changing the element?

Accepted Answer

The atomic number equals the number of protons in the nucleus. It defines the element's identity: change the proton count and you have a different element. In a neutral atom it also equals the electron count.

Question 2

Isotopes vs. ions: how do they differ from a neutral atom?

Accepted Answer

Isotopes have the same protons but different neutrons (different mass number, same element). Ions have a different number of electrons than protons, giving a net charge. Neither changes the element's identity.

Question 3

How do you calculate the number of neutrons in an atom?

Accepted Answer

Neutrons = mass number - atomic number. The mass number counts protons plus neutrons, so subtracting the protons (atomic number) leaves the neutrons. Example: carbon-14 has 14 - 6 = 8 neutrons.

Question 4

Why is the atomic mass on the periodic table usually not a whole number?

Accepted Answer

It is a weighted average of the masses of all naturally occurring isotopes, weighted by their relative abundance. The average leans toward whichever isotope is more common, so it rarely equals any single isotope's mass.

Question 5

Ground state vs. excited state of an electron: what is the difference and what happens between them?

Accepted Answer

Ground state is the lowest-energy electron arrangement. When an atom absorbs energy, an electron jumps to a higher level (excited state). When it falls back, it emits a photon of light with a specific energy, producing a bright-line spectrum.

Question 6

Why does each element produce a unique bright-line (emission) spectrum?

Accepted Answer

Each element has a unique set of energy levels, so the energy gaps between levels are unique. Electrons dropping between those gaps emit photons of specific wavelengths, creating a 'fingerprint' pattern that identifies the element.

Question 7

Where are valence electrons, and why do they matter most in chemistry?

Accepted Answer

Valence electrons are in the outermost principal energy level. They are the electrons involved in bonding, so they largely determine an element's chemical behavior and how many bonds it tends to form.

Question 8

What does a group (vertical column) on the periodic table tell you about reactivity?

Accepted Answer

Elements in the same group have the same number of valence electrons, so they show similar chemical properties. Group 1 (alkali metals) and Group 17 (halogens) are highly reactive; Group 18 (noble gases) have full valence shells and are generally unreactive.

Question 9

How does atomic radius change across a period and down a group?

Accepted Answer

Across a period (left to right) radius decreases: protons are added to the same shell, pulling electrons in tighter. Down a group radius increases: each row adds a new principal energy level farther from the nucleus.

Question 10

What is electronegativity, and how does it trend on the periodic table?

Accepted Answer

Electronegativity is an atom's pull on a shared pair of bonding electrons. It increases across a period (toward fluorine) and decreases down a group. Fluorine is the most electronegative element; metals are low.

Question 11

What is first ionization energy, and why does it generally increase across a period?

Accepted Answer

Ionization energy is the energy needed to remove the most loosely held electron from a gaseous atom. It increases across a period because greater nuclear charge holds electrons more tightly, making them harder to remove.

Question 12

How do properties shift as you move from metals to nonmetals across the periodic table?

Accepted Answer

Metals (left side) lose electrons easily, conduct heat/electricity, and are malleable. Nonmetals (right side) gain or share electrons and are poor conductors. Metalloids along the staircase show intermediate behavior.

Question 13

Ionic vs. covalent bonding: what is transferred or shared, and what kinds of elements form each?

Accepted Answer

Ionic bonds form when a metal transfers electrons to a nonmetal, creating oppositely charged ions held by attraction. Covalent bonds form when two nonmetals share electron pairs. Large electronegativity differences favor ionic bonding.

Question 14

What holds a metallic solid together, and why are metals good conductors?

Accepted Answer

Metallic bonding is positive metal ions surrounded by a 'sea' of mobile, delocalized valence electrons. Those free-moving electrons let metals conduct electricity and heat and let the metal bend without shattering (malleability).

Question 15

What makes a molecule polar, and why can a molecule have polar bonds yet be nonpolar overall?

Accepted Answer

A polar bond has unequal electron sharing (electronegativity difference). A molecule is polar if those bond dipoles don't cancel. In symmetric molecules like CO2, equal dipoles point opposite ways and cancel, making the molecule nonpolar.

Question 16

What are intermolecular forces, and how do they differ from the bonds inside a molecule?

Accepted Answer

Intermolecular forces are attractions between separate molecules; they are much weaker than the covalent bonds within a molecule. Stronger intermolecular forces (like hydrogen bonding) raise boiling point and lower vapor pressure.

Question 17

Why does water have an unusually high boiling point for such a small molecule?

Accepted Answer

Water molecules form hydrogen bonds, a strong type of intermolecular attraction between H and highly electronegative O. These extra attractions require more energy to overcome, raising water's boiling point well above similar-sized molecules.

Question 18

How can you tell from properties whether a solid is ionic, molecular (covalent), or metallic?

Accepted Answer

Ionic solids have high melting points and conduct only when molten or dissolved. Molecular solids have low melting points and don't conduct. Metallic solids conduct as solids and are malleable. Properties reveal the bonding type.

Question 19

What is a mole, and why is it chemistry's key counting unit?

Accepted Answer

A mole is 6.02 x 10^23 particles (Avogadro's number). Because atoms are too small to count individually, the mole lets chemists relate countable masses in grams to numbers of particles for reactions and calculations.

Question 20

How do you convert between mass and moles, and where do you find the needed values?

Accepted Answer

Moles = mass (g) / molar mass (g/mol), or mass = moles x molar mass. Molar mass is the sum of the atomic masses (from the periodic table/reference tables) of every atom in the formula.

Question 21

What does the law of conservation of mass require when balancing an equation?

Accepted Answer

Atoms are neither created nor destroyed, so each element must have the same number of atoms on both sides. You balance by adjusting coefficients only, never subscripts, since changing subscripts would change the substance.

Question 22

What do the coefficients in a balanced equation actually represent?

Accepted Answer

Coefficients give the mole ratio between reactants and products. For 2 H2 + O2 -> 2 H2O, two moles of hydrogen react with one mole of oxygen to form two moles of water. This ratio drives stoichiometry calculations.

Question 23

What is the difference between an empirical formula and a molecular formula?

Accepted Answer

The empirical formula is the simplest whole-number ratio of atoms (e.g., CH2O). The molecular formula gives the actual atom counts in one molecule (e.g., C6H12O6) and is a whole-number multiple of the empirical formula.

Question 24

How do you find the percent composition of an element in a compound?

Accepted Answer

Percent by mass = (mass of that element in one mole / molar mass of the compound) x 100. Multiply each element's atomic mass by its subscript first, then divide by the total molar mass.

Question 25

How do you name and write formulas for ionic compounds, including transition metals?

Accepted Answer

Combine the positive ion and negative ion so total charge is zero, using subscripts to balance charges. For metals with multiple charges (like iron), a Roman numeral shows the charge: FeCl3 is iron(III) chloride.

Question 26

What do the kinetic molecular theory assumptions say about an 'ideal gas'?

Accepted Answer

Ideal gas particles have negligible volume, are in constant random straight-line motion, have no attractive forces, and undergo elastic collisions. Real gases behave most like ideal gases at high temperature and low pressure.

Question 27

How are pressure, volume, and temperature related for a fixed amount of gas?

Accepted Answer

At constant temperature, pressure and volume are inversely related (Boyle's law). At constant pressure, volume and Kelvin temperature are directly related (Charles's law). Always use the Kelvin scale (K = degrees C + 273) for gas temperature.

Question 28

What happens to temperature during a phase change, and where does the added energy go?

Accepted Answer

Temperature stays constant during a phase change even though heat is added. The energy goes into breaking attractions between particles (potential energy) rather than increasing their motion (kinetic energy/temperature).

Question 29

How does a substance's vapor pressure relate to its boiling point and intermolecular forces?

Accepted Answer

A liquid boils when its vapor pressure equals the surrounding pressure. Substances with weaker intermolecular forces have higher vapor pressure and boil more easily (lower boiling point); stronger forces mean lower vapor pressure.

Question 30

What three factors increase how fast a solute dissolves, and how does temperature affect gas vs. solid solubility?

Accepted Answer

Stirring, heating, and increasing surface area speed up dissolving. Raising temperature usually increases the solubility of solids but decreases the solubility of gases in water (warm soda goes flat).

Question 31

How do molarity, dilution, and concentration relate?

Accepted Answer

Molarity (M) = moles of solute / liters of solution. A solution is concentrated if it has much solute per volume, dilute if little. Adding solvent lowers molarity but does not change the number of moles of solute present.

Question 32

What three things must happen for a chemical reaction to occur (collision theory)?

Accepted Answer

Particles must collide, with enough energy to reach the activation energy, and in the correct orientation. Collisions lacking sufficient energy or proper alignment do not produce a reaction.

Question 33

Name four ways to increase the rate of a reaction and the reason each works.

Accepted Answer

Higher temperature (more energetic collisions), higher concentration/pressure (more frequent collisions), greater surface area (more contact), and a catalyst (lowers activation energy). Each increases effective collisions per second.

Question 34

How does a catalyst speed a reaction without being used up?

Accepted Answer

A catalyst provides an alternate pathway with lower activation energy, so more collisions succeed. It is regenerated by the end, so it does not appear in the overall balanced equation and does not change the products or the energy released.

Question 35

What does it mean for a reversible reaction to be at dynamic equilibrium?

Accepted Answer

Forward and reverse reactions occur at equal rates, so the concentrations of reactants and products stay constant. The reaction has not stopped; both directions continue at the same speed, keeping amounts steady.

Question 36

How does Le Chatelier's principle predict the response to a stress on equilibrium?

Accepted Answer

A system at equilibrium shifts to relieve a stress. Adding a substance shifts away from it; removing one shifts toward it. Higher pressure shifts toward fewer gas moles; changing temperature shifts based on whether heat is a reactant or product.

Question 37

How do Arrhenius acids and bases differ in the ions they release in water?

Accepted Answer

An Arrhenius acid releases H+ (hydronium, H3O+) as its only positive ion in water; an Arrhenius base releases OH- (hydroxide) as its only negative ion. These ions drive acidic and basic behavior.

Question 38

What does the pH scale measure, and what does a one-unit change represent?

Accepted Answer

pH measures hydronium-ion concentration. Below 7 is acidic, 7 is neutral, above 7 is basic. The scale is logarithmic, so each one-unit change is a tenfold change in H+ concentration (pH 3 has 10x the H+ of pH 4).

Question 39

What are the products of an acid-base neutralization reaction?

Accepted Answer

An acid plus a base produces a salt and water. The H+ from the acid combines with OH- from the base to form water, while the remaining ions form an ionic salt. Example: HCl + NaOH -> NaCl + H2O.

Question 40

What is the purpose of a titration, and how do you find the unknown concentration?

Accepted Answer

Titration finds an unknown acid or base concentration by reacting it with a known solution until the endpoint (often shown by an indicator). At neutralization, moles of H+ equal moles of OH-, letting you solve for the unknown.

Question 41

How do you tell oxidation from reduction in a redox reaction?

Accepted Answer

Oxidation is loss of electrons (oxidation number increases); reduction is gain of electrons (oxidation number decreases). Use 'LEO the lion says GER': Lose Electrons = Oxidation, Gain Electrons = Reduction. They always occur together.

Question 42

What role do the anode and cathode play, and how do oxidation/reduction map to them?

Accepted Answer

Oxidation always happens at the anode; reduction always happens at the cathode ('an ox, red cat'). This holds for both voltaic cells and electrolytic cells, though their charge signs differ.

Question 43

How does a voltaic (galvanic) cell differ from an electrolytic cell?

Accepted Answer

A voltaic cell uses a spontaneous redox reaction to produce electrical energy (a battery). An electrolytic cell uses an external power source to drive a nonspontaneous reaction (electroplating). One makes electricity; the other consumes it.

Question 44

How does the activity series predict whether a single-replacement reaction occurs?

Accepted Answer

A more active (more easily oxidized) metal will replace a less active metal or hydrogen from a compound. If the lone element ranks above the one it would replace on the activity series, the reaction proceeds; otherwise no reaction.

Question 45

What distinguishes saturated from unsaturated hydrocarbons?

Accepted Answer

Saturated hydrocarbons (alkanes) have only single C-C bonds and the maximum hydrogen. Unsaturated hydrocarbons contain double bonds (alkenes) or triple bonds (alkynes), so they hold fewer hydrogens and are generally more reactive.

Question 46

What is a functional group, and why does it matter for naming and reactivity?

Accepted Answer

A functional group is a specific atom arrangement (like -OH in alcohols or -COOH in organic acids) that gives a family of organic compounds its characteristic properties and reactions. It determines the compound's class and behavior.

Question 47

What are isomers in organic chemistry?

Accepted Answer

Isomers are compounds with the same molecular formula but different structural arrangements of atoms. Because structure controls properties, isomers can have different physical and chemical properties despite identical formulas.

Question 48

How do fission and fusion differ as nuclear reactions?

Accepted Answer

Fission splits a large heavy nucleus into smaller nuclei, releasing energy (nuclear power plants). Fusion combines small nuclei into a larger one (powering the sun). Both convert a small amount of mass into large energy and change the nucleus.

Question 49

What is half-life, and how do you use it?

Accepted Answer

Half-life is the time for half of a radioactive sample to decay. After each half-life the remaining amount is halved (100% -> 50% -> 25% -> 12.5%). Half-life is constant for a given isotope and is used in radioactive dating.

Question 50

What is the difference between an exothermic and an endothermic reaction in terms of energy and products?

Accepted Answer

Exothermic reactions release energy to the surroundings, so products have lower potential energy than reactants (heat is a product). Endothermic reactions absorb energy, so products have higher potential energy (heat is a reactant).

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