Isotopes: Atoms of an element with different numbers of neutrons.
Average atomic mass: Weighted average of isotope mass and relative abundance (frequency).
Moles
Ideal Gas Law: PV = nRT where:
P = Pressure
V = Volume
n = number of moles
R = Ideal gas constant
T = Temperature
Avogadro’s number: 6.022 \times 10^{23}
At STP (1 atm, 273K), 1 mole occupies 22.4 L.
Molarity (M) = moles of solute / L of solution
Percent composition: (mass of each element/compound / total molar mass of the substance) * 100%
Empirical formula: Simplest whole number ratio of atoms in a compound.
Molecular formula: Actual formula for the substance.
Energy
Electron potential energy increases with distance from the nucleus.
Electron energy is quantized; electrons can only exist at specific energy levels at specific intervals.
Coulomb’s Law: F = \frac{kq1q2}{r^2}, where:
F = Electrostatic force
k = Coulomb's constant
q1 and q2 = magnitudes of the charges
r = distance between the charges
Atoms absorb energy in the form of electromagnetic radiation as electrons jump to higher energy levels. When electrons drop to lower energy levels, atoms give off energy.
Photoelectron Spectroscopy (PES)
Energy is measured in electronvolts (eV).
Incoming radiation energy = Binding energy + Kinetic energy of the ejected electron
Electrons further from the nucleus require less energy to eject and thus will move faster.
Photoelectron Spectrum
Each section of peaks represents a different energy level (1, 2, 3, etc.).
Subshells within each energy level (shape of space electron can be found in orbiting nucleus) are represented by the peaks (1s, 2s, 2p, etc.).
Notation: s(2) - first subshell, p(6) - second subshell
The height of peaks determines the number of electrons in the subshell (e.g., the peak of the p subshell in energy level 2 will be 3x that of the s subshell).
Electron Configuration
spdf notation - shorthand with noble gas configuration first.
Configuration rules
Aufbau principle: Electrons fill the lowest energy subshells available first.
Pauli exclusion principle: Two electrons in the same orbital cannot have the same spin.
Common charges: Zn+2, Ag+1, Al+3, Cd+2; most other transition metal charges vary.
Periodic Trends
Electrons are more attracted if they are closer to the nucleus, or if there are more protons.
Electrons are repelled by other electrons - if there are electrons between the valence electrons and the nucleus, the valence electrons will be less attracted (shielding).
Completed shells are very stable; completed subshells are also stable; atoms will add/subtract valence electrons to complete their shell.
Trends (INCREASING):
Atomic radius: down and to the left
Ionization energy: up and to the right
Electronegativity: up and to the right
Definitions:
Ionization energy: Energy required to remove an electron from an atom.
Electronegativity: How strongly the nucleus of an atom attracts electrons of other atoms in a bond.
Electron affinity: Energy change that occurs when an electron is added to an atom in the gas state (usually exothermic - energy is released).
Molecular and Ionic Compound Structure and Properties
Bonds
Atoms are more stable with full valence shells.
Ionic Bonds
A cation gives up electrons completely.
Electrostatic attractions in a lattice structure.
Between metals and nonmetals (salts).
Coulomb’s law: Greater charge leads to greater bond/lattice energy (higher melting point).
If both ions have equal charges, the smaller radius will have greater coulombic attraction.
Ionic solid: Electrons do not move around the lattice; ionic solids are poor conductors of electricity.
Ionic liquids conduct electricity because ions are free to move around, though electrons are still localized around particular atoms.
Metallic Bonds
Sea of electrons model: A positively charged core is stationary while valence electrons are very mobile.
Metals bond to form alloys
Interstitial alloy: with metals of different radii
Substitutional alloy: with metals of similar radii
Molecular Covalent Bonds
Two atoms share electrons—both atoms achieve complete outer shells.
Single bond: 1 sigma bond - order 1, longest length, least energy
Double bond: 1 sigma and 1 pi bond - order 2, intermediate length, intermediate energy
Triple bond: 1 sigma and 2 pi bonds - order 3, shortest length, greatest bond energy
Bond forms when potential energy is at a minimal level.
Too close: Potential energy is too high due to repulsive forces.
Too far: Potential energy is near 0 because attractive forces are very weak.
Minimal PE occurs when repulsive and attractive forces are balanced.
Network Covalent Bonds: Lattice of covalent bonds—poor conductors, high melting and boiling points.
Conductivity
Conductivity of different substances in different phases:
Ionic Compounds: No (s), Yes (aq), Yes (l), No (g)
Molecular Covalent Compounds: No (s), No (aq), No (l), No (g)
Network Covalent Compounds: No (s), N/A (aq), No (l), No (g)
Metallic Compounds: Yes (s), N/A (aq), Yes (l), No (g)
Lewis Dot Structures
Used to represent the bonding between atoms in a molecule.
Resonance: For bond order calculations, average together all possible orders of a specific bond.
BORON (B) is stable with 6 electrons—only one that does not need a full octet.
Expanded octets: Any atom of an element from n=3 or greater (those with a d subshell) can have [8, 12] valence electrons on the center atom.
Noble gases form bonds by filling an empty d orbital with electrons.
Formal charge: Number of valence electrons minus assigned electrons (1 electron for each line “shared” bond)—0 for neutral molecules
Molecular Geometry (VSEPR)
Double and triple bonds have more repulsive strength than single bonds—occupy more space.
Lone electron pairs have more repulsive strength than bonding pairs, so molecules with lone pairs will have slightly reduced angles between terminal atoms.
Hybridization: determined by how many atoms are attached (sp, sp2, sp3, sp3d, etc.).
Intermolecular Forces and Properties
Polarity
Covalent bond where electrons are unequally shared - polar covalent.
Dipoles are caused by polar covalent bonds—a pair of opposite electric charges separated by some distance, like partial charges on atoms in a polar covalent bond.
If 2 identical atoms bond (ex. Cl-Cl), the electrons are equally shared, creating a nonpolar covalent bond with no dipole.
Bonds can be polar; so can molecules, depending on the molecular geometry (and polarity of bonds - secondary).
In polar molecules, more electronegative atoms will gain a negative partial charge.
Usually, the central atom will be positive—exception is hydrogen (terminal), which is usually positive since it has less electronegativity.
Intermolecular Forces (IMF)
Forces between molecules in a covalently bonded substance—need to be broken apart for covalent substances to change phases.
Changing phase: ionic substances break bonds between individual ions; covalent substances keep bonds inside a molecule in place but break bonds between molecules.
Dipole-Dipole Forces
Polar molecules—the positive end of one molecule is attracted to the negative end of another molecule.
Relatively weak overall—melt and boil at low temps.
Hydrogen Bonds
A special type of dipole-dipole attraction where the positively charged hydrogen end of a molecule is attracted to the negatively charged end of another molecule containing an extremely electronegative element (F, O, N).
Much stronger than normal dipole-dipole forces since a hydrogen atom “sharing”/giving up its lone electron to a bond is left with no shielding.
Higher melting/boiling points than substances held together only by other types of IMF.
London Dispersion Forces (LDF)
All molecules—very weak attractions due to random motion of electrons on atoms within molecules (instantaneous polarity).
Molecules with more electrons experience greater LDF (more random motion).
Higher molar mass usually means greater LDF (as mass increases, electrons increase for the molecule to remain electrically neutral).
IMF Strength
Ionic substances are generally solids at room temp—melting them requires lattice bonds to be broken—necessary energy determined by Coulombic attraction.
Covalent substances (liquids at room temp) boil when IMF are broken; for molecules of similar size, from strongest to weakest: hydrogen bonds, permanent dipoles, LDF (temporary dipoles—greater for larger molecules).
Melting/boiling points of covalent substances are LOWER than for ionic substances.
Bonding/Phases
Substances with weak IMF (LDF) tend to be gases at room temp (N2); substances with strong IMF (hydrogen bonds) tend to be liquids at room temp (H2O).
Ionic substances do not experience IMF—since ionic bonds are stronger than IMF, ionic substances are usually solids at room temp.
Vapor Pressure
Molecules in a liquid are in constant motion—if they hit the surface of the liquid with enough kinetic energy, they can escape the IMF holding them to other molecules and transition into the gas phase.
Vaporization (NOT boiling)—no outside energy is added.
Temperature and vapor pressure are directly proportional.
At the same temp, vapor pressure is dependent on the strength of IMF (stronger IMF, lower vapor pressure).
Solution Separation
Solutes and solvents - like dissolves like
Paper Chromatography
A piece of filter paper with a substance on the bottom is dipped in water.
More polar components of the substance travel further up the filter paper with the polar water.
The distance a substance travels up the paper is measured by the retention/retardation factor: R_f = \frac{\text{distance traveled by solute}}{\text{distance traveled by solvent front}}
Stronger attraction - larger R_f
Column Chromatography
A column is packed with a stationary substance.
A separable solution (analyte) is injected, adhering to a stationary phase.
Another solution (eluent - liquid/gas) is injected into the column.
More attracted analyte molecules will move through faster and leave the column first.
Distillation
Takes advantage of different boiling points of substances by boiling a mixture at an intermediate point.
Vapor is collected, cooled, and condensed back to a liquid separate from the leftover liquid.
Kinetic Molecular Theory
Kinetic energy of a single gas molecule: KE = \frac{1}{2}mv^2
Average kinetic energy of a gas depends on the temperature (directly proportional), not the identity of the gas (different gases will have the same KE at the same temp).
Ideal gases have insignificant volume of molecules, no forces of attraction between molecules, and are in constant motion without losing KE.
Deviations occur at low temperatures or high pressures (gas molecules are packed too tightly together):
The volume of gas molecules becomes significant (less free space for molecules to move around than predicted).
Gas molecules attract one another and stick together (real pressure is smaller than predicted pressure).
Maxwell-Boltzmann Diagrams
Higher temp -> greater KE -> greater range of velocity.
Smaller masses, greater velocities to have the same KE.
Effusion
The rate at which a gas escapes from a container through microscopic holes.
High to low pressure.
Greater speed, greater temp, greater rate of effusion.
If at the same temp, the gas with the lower molar mass will effuse first.
Equations
Ideal Gas Equation: PV = nRT, where R = 0.0821
Combined Gas Law: \frac{P1V1}{T1} = \frac{P2V2}{T2}
Dalton’s Law: P{\text{total}} = Pa + Pb + Pc + \dots
Partial Pressure: Pa = P{\text{total}} \times \frac{\text{moles of gas A}}{\text{total moles of gas}}
Density: D = \frac{m}{V}
From the ideal gas law: \text{Molar mass} = \frac{DRT}{P}
Electromagnetic Spectrum
E = hv, where
E = energy change
h = Planck’s constant (6.626 \times 10^{-34})
v = frequency
C = \lambda v, where
C = speed of light (2.998 \times 10^8)
v = frequency
\lambda = wavelength
Beer’s Law: A = abc, where
A = absorbance
a = molar absorptivity (constant depending on the solution)
b = path length of light through the solution (constant)
c = concentration of the solution
Colorimetry - direct relationship b/w concentration and absorbance
Chemical Reactions
Types of Reactions
Synthesis: Everything combines to form one compound.
Decomposition: One compound + heat is split into multiple elements/compounds.
Acid-Base reaction: Acid + base -> water + salt
Oxidation-Reduction (Redox) reaction: Changes the oxidation state of some species.
Combustion: Substance with H and C + O2 -> CO2 + H2O
Precipitation: Aqueous solutions -> insoluble salt (+ more aq sometimes)
Can be written as net ionic - Those free ions not in net ionic are spectator ions
Solubility Rules
Alkali metal cations or ammonium (NH_4^+) cations are ALWAYS soluble.
Compounds with a nitrate (NO_3^−) anion are ALWAYS soluble.
Use the law of conservation of mass (if x g of CO2 is produced, find g of C which will be the starting amount).
Gravimetric analysis
When asked to determine the identity of a certain compound, find g of component produced, then use mass percent \frac{\text{g found}}{\text{total sample mass}}
Compare to mass percent of options: \frac{\text{molar mass of component}}{\text{molar mass of entire compound}}
Oxidation States
Neutral atoms not bonded to other atoms have an oxidation state of 0.
Monoatomic ions have an oxidation state equal to the charge on that ion (e.g., Zn2+ will be +2).
Oxygen is -2 (EXCEPTION: in hydrogen peroxide, H2O2, O is -1).
Hydrogen is +1 with nonmetals, -1 with metals.
In the absence of oxygen, the most electronegative element in a compound will take an oxidation state equal to its usual charge (e.g., F is -1 in CF4).
IF none of the above rules apply, determine the oxidation state by adding up all elements’ oxidation states to 0/charge on ion.
C, N, S, P frequently vary oxidation states (low electronegativity).
Redox Reactions
Write full reaction as 2 half-reactions (oxidation and reduction; OIL RIG).
Add H2O to compensate for oxygen on one side.
Add H+ to compensate for H from H2O on other side.
Balance 2 half-reactions to have the same number of electrons and add them together to produce one complete reaction.
ACIDIC: stop here
BASIC: Add OH- to both sides - enough for all H+ on one side to be converted to H2O; then cancel out H2O so it only remains on one side.
Acids and Bases (briefly)
A color change signals the end of a titration (can be redox or acid/base).
Species with the H+ ion are acids, the same species but without H+ is a base - conjugate acid/base pairs.
Water can act as an acid or base - amphoteric.
Kinetics
Rate Law
\text{Rate} = k[A]^x [B]^y [C]^z
Can calculate x, y, z via a table from (\text{concentration factor})^x = (\text{rate factor})
k is only dependent on temperature (always increases with T).
K{eq} = \frac{K1}{K_2}, where
K_1 = rate constant of forward reaction
K_2 = rate constant of the reverse reaction
k calculated by dividing any rate in the table by the concentrations to their respective powers.
Units for rate are M/s, units for conc are M -> calculate units for k from there.
If A + 2B + C -> D; rate of formation of D = rate of disappearance of A and C = 0.5 * rate of disappearance of B.
Orders
Zero-Order
Rate = k
Concentration vs. time has slope -k
First-Order
Rate = k[A]
\ln[A] vs. time has slope -k
\ln[A]t = -kt + \ln[A]0
Second-Order
Rate = k[A]^2
\frac{1}{[A]} vs. time has slope k
\frac{1}{[A]t} = kt + \frac{1}{[A]0}
Half-Life
First-order reactions only have a constant half-life.
t_{1/2} = \frac{\ln(2)}{k} = \frac{0.693}{k}
Collision Theory
Chemical reactions occur because reactants are constantly moving and colliding with one another.
When reactants collide with sufficient energy (activation energy Ea), a reaction occurs.
Gaseous/aqueous: increased concentration increases rate of reaction (more likely to collide).
Stirring increases reaction rate for heterogeneous mixtures (causing heterogeneous mixture to move around increases collisions; insignificant once the mixture becomes homogeneous due to the number of collisions happening due to inherent motion of aq molecules).
Greater temperature increases the reaction rate (a greater fraction of reactant molecules has sufficient energy to exceed the activation energy barrier - vertical line on Maxwell-Boltzmann with multiple temps).
Reactions only occur if reactants collide with the correct orientation to break the right bonds.
Reaction Energy Profile
Reaction Mechanisms
Species that are produced in a mechanism but are also fully consumed and do not appear in the balanced equation are intermediates.
Adding up all mechanism steps and canceling out different species leads to the balanced reaction.
Elementary steps with 2 reactants (even if they are the same) are bimolecular; elementary steps with 1 reactant are unimolecular.
Speed is determined by a slow step (rate determining step).
Consistency is determined by the slow step and those leading up to it.
Make rate for the slow step (e.g., If X + B -> Y, rate = k[X][B]).
Substitute in the rate for X from above equilibrium reaction.
Compare to the actual reaction’s rate equation.
The slow step has the highest activation energy.
Catalysts
Catalysts increase the rate of a chemical reaction without being consumed in the process.
Catalysts do not appear in the balanced equation.
In a reaction mechanism, catalysts enter first, then exit.
Catalysis (reaction with a catalyst)
Surface catalysis: A reaction intermediate is formed.
Enzyme catalysis: A catalyst binds to reactants to reduce activation energy.
Acid-base catalysis: Reactants lose/gain protons to change the reaction rate.
Thermodynamics
Temperature/Heat
Temperature is the average amount of kinetic energy due to molecular motion in a given substance.
Heat is the energy flow between 2 different substances at different temperatures.
First law of thermodynamics: energy can be neither created nor destroyed.
When bonds are formed, energy is released; when bonds are broken, energy is absorbed.
Exothermic: energy transferred from the system to surroundings (delta H is negative).
More energy is released when the product bonds form than is necessary to break reactant bonds.
Endothermic: energy transferred from surroundings into the system (delta H is positive).
More energy is required to break reactant bonds than is released when bonds in products form.
Energy Diagrams
Enthalpy
Enthalpy of Formation
Change in energy when one mole of a compound is formed from its component pure elements under standard conditions (25C/298K).
\Delta Hf = (\Delta Hf \text{ for products}) - (\Delta H_f \text{ for reactants})
Multiply delta Hf for each product/reactant by the coefficient.
If delta Hf is negative, energy is released when the compound is formed, so the product is more stable (exothermic).
If delta Hf is positive, energy is absorbed when the compound is formed, so the product is less stable than its constituent elements (endothermic).
Heat of formation is 0 when the pure element is in its standard state (ex. H2(g) or F2(g)).
Bond Energy
\Delta H (J) = (\text{bond energies of reactants}) - (\text{bond energies of products})
Multiply bond energies for each bond by the coefficient.
Hess’s Law
Finding delta H for the overall reaction from knowing delta H for the steps of the reaction.
Flipping the equation flips the sign of delta H.
Multiplying/dividing the equation by a coefficient multiplies/divides delta H by that coefficient.
Adding/subtracting equations adds/subtracts their delta H values.
Enthalpy of Solution
Ionic substances dissolving in water.
1: Breaking of solute bonds
Energy required is equal to the lattice energy (positive delta H since bonds are being broken).
2: Separation of solvent molecules
Water molecules must spread out to make room for the solute ions (requires energy to weaken the IMF between water molecules - positive delta H).
3: Creation of new attractions
Free-floating ions are attracted to the dipoles of water molecules (energy is released - negative delta H).
Hydration energy = step 2 + step 3 energies
Coulombic energy: increases with charge magnitude, decreases as size increases
Enthalpy of solution=step 1 + 2 + 3 energies
Phase Changes
Solid to gas is sublimation, gas to solid is deposition.
When vapor pressure equals the surrounding atmospheric pressure, the liquid boils
Lower atmospheric pressure (high elevation) means a lower boiling point.
Enthalpy of fusion: energy to melt a solid; heat of fusion: heat given off by a substance when it freezes.
Enthalpy of vaporization: energy to turn a liquid into a gas; heat of vaporization: heat given off by a substance condensing.
IMF is stronger for a liquid than a gas, and for a solid than a liquid, and the stronger IMF is more stable; therefore, going from a gas to a liquid or a liquid to a solid releases energy (exothermic).
As heat is added to a substance, the temperature of the substance can increase OR it can change phases, but not both at once.
When a substance is changing phases, the temperature of the substance remains constant.
Calorimetry
Specific heat: the amount of heat required to raise the temperature of one gram of a substance by one degree C/K.
Large specific heat: can absorb much heat without a significant temperature change
Low specific heat: quickly changes temperature
Heat added equation: q = mc\Delta T, where
q = heat added (J or cal)
m = mass
c = specific heat
\Delta T = Change in temp
q1 = q2 for mixtures
Calorimetry: measurement of heat changes during chemical reactions.
Find J from q, find moles from stoichiometry, divide the two to find \Delta H. (\Delta H \text{ measured in J/mol})
Heating Curves
For problems where a solid completely melts or the like, add q from mc\Delta T to (moles) \times (\text{heat of fusion}) for the total heat required for the process to occur.
Equilibrium
Keq
A reaction is at equilibrium when all concentrations stop changing.
The reaction does not stop—the rate of forward and reverse reactions becomes equal.
All concentrations do NOT sum to the initial concentration of reactants.
In reaction 2A -> B, the concentration of A will decrease 2x as much as the concentration of B increases.
Equilibrium Expression/Law of Mass Action
For the reaction aA + bB -> cC + dD: K_{eq} = \frac{([C]^c [D]^d)}{([A]^a [B]^b)}
[A], etc. are molar concentrations/partial pressures at equilibrium.
Products are in the numerator, reactants are in the denominator.
Coefficients in the balanced equation become exponents in the equilibrium expression.
Only gaseous and aqueous species are included in the expression.
Keq has no units.
K > 1 favors the forward reaction; K < 1 favors the reverse reaction.
Different Equilibrium Constants
Kc is for molar concentrations.
Kp is for partial pressures.
Ksp is the solubility product (no denominator because reactants are solids).
Ka is the acid dissociation constant for weak acids.
Kb is the base dissociation constant for weak bases.
Kw describes the ionization of water (K_w=1 \times 10^{-14}).
Manipulating Keq
Keq for a flipped reaction is the reciprocal of Keq for the initial reaction.
Keq for a reaction multiplied by a coefficient is the initial Keq to the power of the coefficient.
Keq for two reactions added together is their respective initial Keq values multiplied together.
Le Chatelier’s Principle
Increasing the concentration of reactants shifts rxn to favor products (forward) and vice versa.
Increasing pressure increases the partial pressure of all gases in the container and shifts the reaction to the side with fewer gas molecules (moles of gas).
Increasing volume decreases pressure and vice versa.
Adding a non-reacting gas (noble gas) to a non-rigid container causes the volume to increase while not changing the total pressure.
Adding a non-reacting gas to a rigid container would increase the total pressure of the container and not affect the partial pressures of other species—no reaction shift occurs.
Increasing temperature in an endothermic reaction shifts the reaction to favor products (forward); increasing temperature in an exothermic reaction shifts the reaction to favor reactants (reverse).
Treat “heat” as a reactant (endothermic) or product (exothermic) to see shifts like with concentration change.
Diluting aqueous equilibriums shifts the reaction to favor the side with more aqueous species; removing water (evaporation) shifts the reaction to favor the side with less aqueous species.
Shifts caused by concentration/pressure are temporary shifts and do not change Keq; shifts caused by temperature permanently affects Keq and ratio of products to reactants since it adds/removes energy from the system.
Reaction Quotient Q
Q can be calculated at any point with current concentrations/pressures; Keq can only be calculated with equilibrium values.
For the reaction aA + bB -> cC + dD: Q= \frac{([C]^c [D]^d)}{([A]^a [B]^b)}
[A], etc. are initial molar concentrations or partial pressures.
If Q < K, reaction shifts right; if Q > K, reaction shifts left; if Q = K, reaction is at equilibrium.
Solubility
A salt is considered soluble if more than 1g can be dissolved in 100mL of water.
Soluble salts are assumed to dissociate completely in aqueous solutions.
Most solids become more soluble in a liquid as temperature increases.
Solubility Product (Ksp)
For the reaction AaBb(s) ⇄ aAb+(aq) + bBa-(aq): K_{sp} = [A^{b+}]^a [B^{a-}]^b
Molar solubility is determined by subbing x, 2x, 3x, etc. in for concentrations in Ksp expression (x if the coefficient is 1 in the balanced reaction, 2x if the coefficient is 2, etc.).
The molar solubility of a salt is equal to the concentration of any ion that occurs in a 1:1 ration with a salt.
Molar solubility typically increases with temperature since there is more energy available to force water molecules apart to make room for solute ions.
Common Ion Effect
Newly added ions from a separate solution affect the equilibrium of the initial solution if some elements are present in both, even though the newly added ions did not come from the initial compound.
Ex. Adding NaCl to AgCl affects Cl which affects AgCl equilibrium.
Acids and Bases
pH
Formulas
pH = -\log([H^+])
pOH = -\log([OH^-])
pKa= -\log(Ka)
pKb = -\log(Kb)
pKw = -\log(Kw)
[H^+] = [OH^-] => neutral, pH = 7
[H^+] > [OH^-] => acidic, pH < 7
[H^+] < [OH^-] => basic, pH > 7
Increasing pH means decreasing [H^+] (less acidic solution) and vice versa
Strong Acids
Strong acids dissociate completely in water (rxn goes to completion); no equilibrium, equilibrium constant, or dissociation constant.
Important strong acids/bases
No tendency for reverse rxn to occur (-> not ⇄) so the conjugate base of a strong acid is very weak.
The pH of a strong acid solution can be found directly from [H^+] since it dissociates completely.
The best conductors of electricity.
Weak Acids
Weak acid + water causes a small fraction of its molecules to dissociate into H^+ and A^- (conjugate base) ions.
Ka and Kb are measures of the strengths of strong/weak acids—equilibrium constants specific to acids/bases.
Base dissociation constant: K_b = \frac{[HB^+][OH^-]}{[B]}
A greater K_a means a greater extent of dissociation and a stronger acid.
A greater K_b means a stronger base; base is not dissociating but rather accepting a proton (hydrogen ion) from an acid (protonates/ionizes, not dissociates).