Chem Winter Q

solvent

solution component that is the greatest amount

solute

any component in a solution that isnt the solvent

proof is

defined as 2x the precent of alcohol

80proof= 40%alcohol and 60% water

Strong acids

So I Brought No Clean Cloths

1. H₂SO_{4 (sulfuric acid)}

2. HI (hydroiodic acid)

3. HBr

4. HNO₃

5. HCl

6. HClO₄

Strong Bases

Can't Stand Rich Bad Liars Now

(Ca, Sr, Rb, Ba, Li, Na). 

1. Ca(OH)₂

2. Sr(OH)₂

3. Rb(OH)₂

4. Ba(OH)₂

5. LiOH

6. NaOH

Particles that all acids produce

Hydronium ions

what does a base release when dissolved in water

OH^-

What does a acid release in water

H⁺

Strong electrolytes

• nearly 100% dissociated into ions

• conduct current efficiently

• high electrical conductivity (bc of free moving ions)

• high solubility

Salts (ionic bonds)

strong bases

strong acids

How do you find the limiting reactant

which ever element has the least moles

when asked to find mole of concentration(=molarity) use limiting reactant moles

How to identify base

OH

NH₃

how to identify acid

starts w H

has carboxyl COOH

O.N for F, H, O, group 7, 6, 5

F= -1

H=+1

O=-2

7= -1

6=-2

5=-3

if its a diprotic acid (pure element) its always going to be 0

OIL RIG

oxidation is loss

reduction is gained

Ammonium

NH₄ ⁺

Acetate

CH₃COO^- OR C₂H₃O₂^-

Carbonate

CO₃^-2

Bicarbonate/Hydrogen carbonate

HCO₃^-

Hydroxide

OH^-

Nitrite

NO₂^-

Nitrate

NO₃^-

Chromate

CrO₄^2-

Dichromate

Cr₂O₇^2-

Phosphate

PO₄^3-

Hydrogen phosphate

HPO₄^2-

Dihydrogen Phosphate

H₂PO₄^-

Ammonium

NH₄⁺

Hypochlorite

ClO^-

Chlorite

ClO₂^-

Chlorate

ClO₃^-

perchlorate

ClO₄^-

permanganate

MnO₄^-

Sulfite

SO₃^2-

Hydrogen Sulfite/ Bisulfite

HSO₃^-

Sulfate

SO₄^2-

Hydrogen Sulfate/ bisulfate

HSO₄^-

Cyanide

CN^-

Peroxide

O₂^2-

Properties of Gas

• neither definite shape nor volume

• fills container

• exerts pressurfe on surroiundings

• volume changes w temp and pressure

• mixes completely with other gases

• much less dense than solids and liquids

KMT

1. Particles in Motion: Gases consist of numerous tiny particles (atoms or molecules) in continuous, random, straight-line motion.

2. Negligible Volume: The volume of the gas particles themselves is insignificant compared to the total volume of the container.

3. No Intermolecular Forces Particles exert no attractive or repulsive forces on each other (except during collisions)

4. Elastic Collisions: Collisions between particles and container walls are perfectly elastic, meaning kinetic energy is conserved.

5. Temperature & Kinetic Energy: The average kinetic energy of the particles is directly proportional to the absolute temperature (Kelvin).

Average Kinetic energy of Gas molecues

Parameters Affecting Gases

Pressure

volume

temp

number of moles

effusion

process where gas escapes from a tiny hole into a region of lower pressure

Graham’s Law of Effusion

rate of effusion of gas is inversely proportional to square root of its molar mass

if particles of gas A have higher speeds than B then gas A particles collide more frequently w the walls of the container→ inc effusion

Diffusion

the spread of one substance through another

odors

mean free path

avg distance a particle can travel through air before colliding w another particle

Atmospheric Pressure

Pressure: ratio of force F to surface area A

• P=F/A

Barometer

Instrument used to measure pressure

how do barometers work

• using mercury to measure pressure

• High pressure: when atmospheric pressure inc (heavier air) it pushes harder on the dish, forcing mercury levels up the tube

• low pressure: when pressure dec, theres less downward push, and mercury in the tube falls back down into the reservoir

• the height in mm indicates current air pressure, higher #=higher pressure

Manometer

Using water to measure pressure or any liquid

• u tube contains water, pressure differences pushes liquid and the height differences shows pressure

Precipitation reactions

a reaction where dissolved substances react to form a solid product

• double displacement reaction

Method for writing complete and net ionic equations

1. use solubility rules to deternine physical state

2. separate all aqueous ionic compounds into individual ions to obtain complete ionic equation

3. remove sperctator ions from the complete ionic equation for net ionic equation

Boyle’s Law

• gases are compressible

  • pressure inc as Volume dec

• dec volume inc collesions/area $

Charles’ Law

Avogadro’s Law

Amonton’s Law

combined gas laws

Balancing redox in acidic solution steps

1. find ON to determine oxidation and reduction

2. balance other atoms than O & H

3. balance O by H2O

4. balance H by H+

5. balance charges by adding e-

6. make # of e- = and add half reacdtions

change in oxidation states results from gain/loss if e-

How to determine which gas will effuse the fastest

• lowest molar mass

• grahams law

Atmospheric pressure

Standard points for Ideal Gases

P= 1 atm

T= 273K

V=22.4L

Precipiation Rules

5. insoluable

Reducing agent

the species that gets oxidized (loses electrons).

Electrolyte

solute that produces ions in solution, allows solution to conduct electricity

weak electrolytes

• only partial disassociation

• slightly conductive but poor conductivity

• disassociation is reversible

• weak bases, weak acids, and insoluble compounds

Non-electrolytes

• substance w/no ionization occurs, no conduction of electrical current

• dissolve as neutral molecules, not ions

• covalent bonds not ionic

• sugar, ethanol, urea

Brønsted-Lowry theory

defines an acid as a proton (H⁺) donor and a base as a proton acceptor

H₂O is an acid and base

molecular equation

• reactants written as undisassociated molecules (are whole and not separated)

Overall ionic equation

• distinguishes between molecular and ionic substances, all elements are seperated

net ionic equation

• where spectator ions are removed from ionic equation

Neutralization

reaction where acid reacts w/base profucing salt and H₂O

• salt

  • product of neutralization reaction

  • cation base +anion acid

• HCl + NaOH → NaCl + H₂O

Titration

• analytical method to determine the concentration of a solute in a sample in a sample by reacting it w/ a standard solution

• standard solution: solution w/ known concentrations (titrant)

Saturated Solutions & Supersaturation

• aqueous solubility of most solids inc w/temp

• saturated solution

  • solution that contains the max amount of solute possible at a given temp

• unsaturated solution

  • solution that contains less than the max quantity of solute

• Supersaturated solution

  • holds more solute than normally possible by heating it and cooling it slowly, making it unstable

ideal gas

in an ideal gas the individual particles of the gas don’t interact

First law of thermodynamics

• energy cannot be created/destroyed only change form

• total energy of universe is constant

spontaneous

• once started occurs w/o outside interaction

non-spontaneous

• energy is need for reaction to occur

Second law of Thermodynamics

• entropy of the universe inc in any spontaneous process

colligative properties

• characteristics of a solution that depend on the ratio of the number of solute of the number of solute particles to solvent particles

  • NOT IDENTITY OF PARTICLES

  • changes

    • freezing point depression

    • boiling point elevation

    • osmotic pressure

    • vapor pressure lowering

van’t Hoff factors

• i= how many molecules the element disassociates into

• non electrolytes

  • sugar: glucose, sucrose, fructose

  • alcohol: ethanol, methanol

  • urea

  • protein and polymers

reverse osmosis

• Desalination

  • process that removes most ions from seawater

  • distillation

• reverse osmosis

  • purification where solvent is forced through semipermeable membranes leaving dissolved impurities behind

Vapor pressure

• in a closed container molecules at surface of the liquid are constantly escaping into gas phase (evaporation) & crashing back into the liquid (condensation)

• when rate of evaporation = rate of returning gas → equilibrium

  • pressure exerted by gas molecules at equilibrium is vapor pressure

adding non volatile solute to pure solvent

• when adding a nonvolatile solute (sugar/salt) to a pure solvent (water) vapor pressure drops

• this is because the entire surface of the water WAS pure water and the molecules were ready to jump in the air

  • in the solution the surface spots are now taken up by solute particles → less solvent molecules can jump into the air→ lowers VP

factors that affect vapor pressure

• temp

• suface area

• intermolecular forces

  • stronger forces = higher kinetic E needed to enter gas phase

Volatility

how easy a substance turns into a gas (vapor) at a given temp → bonds hold together

Volatile substances

• have HIGH VP bc molecules are weakly attracted to each other and doesn’t take much energy for them to break bonds into a gas

• characteristics

  • evap quickly at room temp

  • low boiling point

    • alcohols

    • liquid gases

    • sublimation

non volatile substances

• LOW VP

• molecules/ ion are held together by strong IMF and its difficult for them to become gases at room temp

• characteristics

  • doesn’t noticeably evap at room temp

  • usually odorless

  • high boiling point

    • table salt, sugar, oils

ideal solutions

used for raoult’s law

• attraction between different molecules solvent-solute is exactly the same as the attraction between the molecules of solvent-solvent

BUT if molecules repel/attract eachother you get deviations

positive deviation

• molecules repel each other →escaoes easier→ higher pressure than predicted

  • solute-solvent interactions are weaker than solvent-solvent/solute-solute bonds

negative deviation

• molecules stick together strongly → escape less easily→ lower pressure than predicted

  • solute solvent interactions are STRONGER than solvent-solvent/ solute solute interactions

solubility of gases

depends on T&P

• solubility inc as P inc

• solubility dec as temp inc

which is negative/positive deviation

a= negative

b= positive

determining a rate law

Fundamentals of Rate Laws

• Rate law exponents must be determined through experiments; they cannot be figured out just by looking at the balanced chemical equation's

3 types of Reaction Rates

1. Average Rate: The change in concentration of a reactant or product over a specific, measurable time interval (Δ[A]/Δt)

2. Instantaneous Rate: The rate at a specific single point in time, determined graphically as the tangential slope of a concentration vs. time plot

3. Initial Rate: The instantaneous rate at the very start of the reaction (t = 0), immediately after reactants are mixed

Rate Law

equation that defines the relationship between the reactant concentrations and the reaction rate.

Method of determining Initial Rates

• practical technique to determine the rate law by comparing how the initial rate changes when the starting concentration of one reactant is varied while others are held constant

  • If doubling a concentration doubles the rate, the order is 1st order.

  • If doubling a concentration quadruples the rate, the order is 2nd order.

  • If changing the concentration has no effect on the rate, the order is 0th order.

Factors Affecting Reaction Rates

• temp: high temp inc kinetic energy of molecules leading to more frequent and energetic collisions

  • Activation energy: min energy required for a collision to result in a reaction

  • catalyst

  • temperature

  • change in concentration of a reactant

  • physical state of the reactants

pesudo first order

occurs when all reactants except one are in such high concentrations that they stay virtually constant leading the limiting reactant to control the rate