Free NY Regents Chemistry Exam Flashcards

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.

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.

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.

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.

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.

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.

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.

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.

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.

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.

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.

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.

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.

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).

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.

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.

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.

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.

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.

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.

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.

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.

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.

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.

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.

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.

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.

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).

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.

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).

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.

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.

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.

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.

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.

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.

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.

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).

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.

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.

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.

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.

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.

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.

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.

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.

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.

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.

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.

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).