AP Chemistry 1-212

Strong acids

HCl, HBr, HI, HClO₄, HNO₃, H₂SO₄

H₂ bond energy & bond type

Single covalent bond, ~400+ kJ/mol

F₂, Cl₂, Br₂, I₂ bond energies & bond types

Single covalent bonds; bond energies range from ~150-250 kJ/mol

O₂ bond energy & bond type

Double covalent bond, ~500 kJ/mol

N₂ bond energy & bond type

Triple covalent bond, ~900 kJ/mol

Dissociation reactions for strong acids

H+ + rest

Strong bases in AP Chemistry

All soluble metal hydroxides (e.g., NaOH, KOH, Ba(OH)₂)

Celsius to Kelvin conversion

K = 273 + °C

Percent error formula

%Error = [(Experimental - Correct) ÷ Correct] × 100

Importance of sign in percent error

It shows whether the measurement is above or below the actual value.

Significant digits on the AP exam

Three significant digits (3), for 99% of all questions

Situations requiring attention to significant figures

When performing lab measurements using addition/subtraction or simple single operations like molar mass calculations

Relationship between number of sig figs and accuracy

More sig figs = greater accuracy

Most accurate volume-measuring devices

Volumetric flask (fixed volume), burette (variable volume), pipette (variable volume)

Relatively accurate lab devices

Graduated cylinder, electronic balance

Devices not considered accurate for volume measurements

Beaker, Erlenmeyer flask

Molar mass of H₂

2.0 g/mol

Molar mass of N₂

28 g/mol

Molar mass of O₂

32 g/mol

Molar mass of H₂O

18 g/mol

Molar mass of CO₂

44 g/mol

Molar mass of NaOH

40.0 g/mol

Atomic number of an element: location and meaning

The number above the element symbol on the AP periodic table; number of protons

Mass number

Mass number = number of protons + number of neutrons of an atom of a specific isotope of an element; always whole numbers; not shown on AP periodic table

Relationship of mass number to atomic mass

Mass number is the closest whole number to the element's atomic mass

(average) atomic mass

The weighted average of all the naturally occurring isotopes of an element; in amu or g/mol; listen on periodic table for each element.

Implication of average atomic mass close to a whole number

It means the element likely has one dominant isotope

Mass spectrometer

A tool that sorts isotopes of elements by mass and shows their relative abundance

Using mass spectrometry data in AP Chemistry

Use the relative abundances and mass numbers to calculate average atomic mass; compare with periodic table value

Main similarity between isotopes of an element

All isotopes have the same number of protons (atomic number)

Key differences between isotopes of the same element

They differ in the number of neutrons and their atomic mass

Coulombic forces definition

The attractions between opposite charges (protons and electrons) and repulsions of similar charges (electrons)

Coulomb's law can predict:

the magnitude of the forces between charges.

How does sign and magnitude of Charge affect Coulombic forces?

Like charges → repel; Unlike charges → attract; Greater magnitudes → stronger attraction/repulsion.

How does distance affect Coulombic forces?

Greater distance → weaker forces; Smaller distance → stronger forces.

Avogadro's number

6.02 × 10²³ particles/mol.

Molar mass unit

The unit for molar mass is g/mol.

Mole-mass conversion

Use dimensional analysis to convert between mass and moles: grams of substance × (1 mol / g molar mass) = mol substance.

Limiting reactant

The reactant that determines the extent of the reaction and product amounts, found by setting up an ICE Chart and dividing mole amounts by stoichiometric coefficients.

Empirical Formula: definition

The simplest whole-number ratio of atoms in a compound.

Finding empirical formula steps

1. Find the number of mol of each element. 2. Divide all mol values by the smallest mol amount. 3. If necessary, multiply the resulting ratio to eliminate fractions.

How does the molar mass compare to the empirical molar mass?

The molar mass of a compound will be a whole number multiple of the empirical molar mass.

Mole fraction formula

Soluble cations

Na⁺, K⁺, and NH₄⁺ form only soluble compounds and are spectator ions (not included in net ionic equations).

Soluble anion

NO₃⁻ forms only soluble compounds and is a spectator ion (not in net ionic equations).

Millimole solute / volume in mL

Gives you molar concentration in mol/L (M).

Molecular equation for NaCl and AgNO₃

NaCl(aq) + AgNO₃(aq) → NaNO₃(aq) + AgCl(s).

Complete ionic equation for NaCl and AgNO₃

Na⁺ + Cl⁻ + Ag⁺ + NO₃⁻ → Na⁺ + NO₃⁻ + AgCl(s).

Net ionic equation for NaCl and AgNO₃

Ag⁺ + Cl⁻ → AgCl(s).

Spectator ions

Na⁺, K⁺, NH₄⁺, and NO₃⁻ are not found in net ionic reactions because all compounds containing them are soluble and dissociate completely.

Weak acids and bases in net ionic equations

They are shown as molecules, not ions (undissociated, while the strong acids/bases fully dissociate).

Net ionic equation for strong acid-strong base reaction

H⁺ + OH⁻ → H₂O(l).

Convert torr or mmHg to atm

torr ÷ 760 = atm.

Metric unit for pressure

The pascal (Pa); the kPa is also used, but doesn’t need to be memorized.

Ideal gas law equation

PV = nRT.

Watch out: pressure when gas collected over water

Subtract water vapor pressure from the total pressure.

Standard conditions for STP

1 atm, 0°C.

Molar gas volume at STP

22.4 L/mol.

Using 22.4 L/mol directly

When a question gives STP conditions, allowing you to skip PV = nRT and calculate moles directly.

Formula to calculate molar mass of a gas using density

Molar mass = (Density × R × T) / P (molar mass of a gas is DiRTy over P)

Density of a gas @STP

Density@STP = molar mass / 22.4 L/mol

Molecular speed of gases

Lighter molecules move faster.

Average speed on Maxwell-Boltzmann curve

Just to the right of the peak.

Effect of molecular speed on Maxwell-Boltzmann curve

Faster average speed makes the curve wider and more spread out.

Non-ideal gas conditions

High pressure and low temp.

Non-ideal gas behavior at high pressure

Molecules are closer together, so intermolecular attractions interfere with motion.

Non-ideal gas behavior at low temperatures

Molecules move more slowly, allowing intermolecular forces to dominate.

Source of higher-than-expected pressure in non-ideal gas

Large molecular volume reduces free space for movement.

Lower-than-expected pressure in non-ideal gas

Strong intermolecular attractions cause molecules to stick together.

Partial pressure of a gas in a mixture

Specific heat capacity of water

4.18 J/g°C (pretty high)

Substances with low specific heat capacities

Metals (typically less than one)

Formula for heat in a calorimeter

q = mass × ΔT × specific heat capacity

Limiting factor in accuracy of q = mCΔT for sig figs

Usually temp because it is the difference between two close temperature values.

Formula for calculating ΔH from calorimetry

Calorimeter ΔT+; what is ΔH and endo/exo?

The reaction is exothermic; ΔH is negative.

Calorimeter ΔT-; what is ΔH and endo/exo?

The reaction is endothermic; ΔH is positive.

What is ΔH?

The change in potential energy (enthalpy) of a reaction.

Negative ΔH

Exothermic; potential energy decreases; bond energy increases.

Positive ΔH

Endothermic; potential energy increases; bond energy decreases.

Units of ΔH

kJ/molrxn or kJ per mole of product.

Bonds in an endothermic reaction

Stronger bonds in reactants are broken, weaker bonds in products are formed.

Kinetic and potential energy in an endothermic reaction

KE is absorbed from surroundings; PE increases.

Thermodynamic implications of positive ΔH

Reaction is thermodynamically unfavorable unless driven by an increase in entropy (S) or external energy (e.g., electrolytic cell).

Exothermic Reaction

A reaction where weaker bonds in the reactants are broken, and stronger bonds in the products are formed (bond energy increases).

Kinetic Energy in Exothermic Reaction

KE is released to surroundings; the surroundings to get hotter.

Potential Energy in Exothermic Reaction

Potential energy decreases (ΔH = -).

Thermodynamic Favorability of Exothermic Reaction

Exothermic reactions are thermodynamically favorable.

Heat (KE) in Exothermic Reaction

Heat is a product; ΔH is negative.

Example of Thermodynamically Favorable Reaction

A → B + kinetic energy; exothermic

Heat (KE) in Endothermic Reaction

Heat is a reactant; ΔH is positive.

Example of Thermodynamically Unfavorable Reaction (endothermic)

A + kinetic energy → B

ΔH(f), Enthalpy of Formation

The reaction enthalpy to make 1 mole of a substance from its elements in their most common form at 25°C; most are exothermic.

ΔHf of Elements

0 kJ/mol by definition.

Exception to Standard ΔH(f) of Elements

If the element is not in its most common form at 25°C, its ΔH(f) is not zero.

Calculating ΔHrxn Using ΔHf Values

ΔHrxn = ΣΔHf (products) - ΣΔHf (reactants)

Calculating ΔHrxn Using Bond Energies (BE)

ΔH(rxn) = ΣBE (reactants) - ΣBE (products); the only equation where reactants come first.

Relationship Between Wavelength, Frequency, and Photon Energy

Shorter wavelengths have higher frequencies and greater photon energies.

Unit for Frequency

Hertz (Hz), or 1/second (1/s or s⁻¹).

Calculating Energy of a Photon from Wavelength in nm (2-step method)

Because c is in m/s you must change nm to E-9 m