Unit 4

4.1 Introduction for Reactions, 4.2 Net Ionic Equations, 4.3 Representations of Reactions, 4.4 Physical and Chemical Changes, 4.5 Stoichiometry, 4.6 Introduction to Titration, 4.7 Types of Chemical Reactions, 4.8 Introduction to Acid-Base Reactions, 4.9 Oxidation-Reduction (Redox) Reactions

chemical change

changes chemical composition; bonds must break/form

physical change

change form/appearance/phase; IMFs break/form or dissolution

signs of a chemical reaction

production of gas, production of precipitate, production of heat/light, unexpected color change, mass change (if exposed to air)

general reaction equation

reactants yield products

why do we balance equations

law of conservation of mass

what do you balance first

MINOH, for combustion CHO

net ionic equation

shows what actually reacts; no spectator ions

spectator ions

ions that do not participate in the reaction; stay aqueous before and after reaction

what happens to H₂CO₃ in reactions

weak acid and unstable, decomposes into H₂O and CO₂

what happens to weak acids/bases in water

do not fully dissociate, do not separate into ions in net ionic equations

types of reactions

synthesis, decomposition, combustion, single replacement, double replacement, acid-base, redox

synthesis

multiple reactants, one product

decomposition

one reactant, multiple products

combustion

adding O₂ to a hydrocarbon, making CO₂ and H₂O

single replacement

element replaces another, A + BC → AC + B

double replacement

element replaces another, AB + CD → AD + CB

stoichiometry basics

always convert into moles, then use mole ratios to get the desired product

moles to mass

multiply by molar mass

moles to atoms/molecules/particles

multiply by 6.022 × 10²³

moles to liters at STP

multiply by 22.4

other helpful equations for stoich

PV = nRT, M = moles/liters

titration

determines concentration of an unknown solution by adding standardized solution with known concentration

titrant

what you titrate with, solution with known concentration

analyte

what you are analyzing, solution with unknown concentration

equivalence point

quantitative measure of when the reaction is completed, reactants will neutralize at certain pH

titration formula

nM₁V₁ = nM₂V₂

end point

qualitative measure of the completion of the reaction, when the indicator changes color

which is more accurate, equivalence or end point

equivalence

titration curve

shows pH of analyte solution vs. volume of titrant added

where is equivalence point on titration curve

in the middle of the steep increase

acid-base titration

equivalence point when [H+] = [OH-]

redox titration

equivalence point when number of electrons lost = number of electrons gained

precipitation reaction

type of double replacement that produces an insoluble compound (solid)

acid-base reaction

type of double replacement; H+ from acid reacts with OH- from base, produces salt and H₂O

arrhenius acid

releases H+ or H₃O+ in solution

strong acids

fully dissociate in water; HCl, HNO₃, HI, HClO₄, HClO₃

weak acids

weakly dissociate in water, reversible

arrhenius base

releases OH- in solution, limited definition

strong bases

fully dissociate in water; group I and II metals + OH-

weak bases

weakly dissociate in water, reversible

bronsted-lowry definition

focuses on the transfer of protons

bronsted-lowry acid

donates H+

bronsted-lowry base

accepts H+

amphoteric compound

substance that can act as both a base and acid

conjugate acid-base pairs

acid → conjugate base

base → conjugate acid

redox reaction

electrons transferred, change in charge

oxidation

electrons lost

reduction

electrons gained

oxidation state

shows hypothetical charge of species

general rule for oxidation states

equals charge of atom/ion

elemental/diatomic oxidation state

0

atomic ion oxidation state

charge on atom

hydrogen oxidation state

+1 when bonded to nonmetal, -1 when bonded to metal

oxygen oxidation state

-1 in peroxide, -2 anywhere else

fluorine oxidation state

-1

neutral compound oxidation state

0

polyatomic ion oxidation state

sum of all oxidation numbers = charge of ion

how to balance redox reactions

separate into half reactions and balance the number of electrons transferred

half reactions

separate oxidation and reduction reaction & figure out how many electrons used to oxidize/reduce

where are electrons in oxidation

in product

where are electrons in reduction

in reactant

how to balance electrons in half reactions

use LCM (e.g. if oxidation used 3 electrons but reduction only used 1, multiply reduction by 3)