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4. Types of chemical reactions and solution stoichiometry

Solutions

solution: solute and solvent

solute: what is dissolved

solvent: what does the dissolving

looks like the solvent

mixtures

heterogenous: two (or more) things mixed together that look like two (or more) things

homogenous: two (or more) things mixed together that look like one thing

examples

solid solute and liquid solvent: salt water

liquid solute and liquid solvent: juice

gas solute and liquid solvent: soda

solid solute and solid solvent: alloy

gas solute and gas solvent: air

gas solute and solid solvent: foam

liquid solute and solid solvent: fillings

Electrolytes

electrolyte: solution that conducts electricity

strong electrolyte: ionic compound that is completely dissolved; conducts electricity well

weak electrolyte: compound that is partially dissolved; conducts electricity poorly

nonelectrolyte: covalent compound; does not conduct electricity

Molarity and dilutions

way to measure concentration of a solution

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example: calculate the molarity of 11.5 g of NaOH in enough water to make 1.50 L of solution

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example: convert above to only Na

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dilution: when you make a solution less concentrated (add more solvent)

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M: molarity

V: volume

1: at start

2: after dilution

example: what volume of 16 M sulfuric acid must be used to make 1.5 L of a 0.10 M solution?

(16 M)(V₁) = (0.10 M)(1.5 L)

V₁ = 9.4 × 10⁻³ L

Creating solutions

get a volumetric flask

add solute to flask

add half the amount of solvent

put on the stopper

mix until the solute dissolves

fill with solvent up to the fill line

mix again

Spectroscopy

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spectroscopy: an experiment that uses lightwaves to determine concentration

absorbance: measurement of how much light gets stopped/absorbed

Beer-Lambert Law

A = abC

A: absorbance

C: concentration

ab: constant

standard: a bunch of known values that help you determine an unknown

Beer-Lambert Law uses standards to make a graph

graphing

A = abC; y = mx + b

A = y

C = x

m = ab

0 = b

y-axis: absorbance

x-axis: concentration (M)

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Saturation

saturated: when a solvent can no longer dissolve a solute

unsaturated: solvent can dissolve more

increasing solvent dissolving

change the temperature of the solvent (mostly increase in temperature)

add more solvent

mix the solvent

increase the surface area of the solute

Types of reactions

Precipitation

precipitation reaction: two aqueous solutions are mixed; precipitate (solid) forms

solubility rules (order matters)

sodium (Na⁺), potassium (K⁺), ammonium (NH₄⁺) compounds are always soluble

nitrates (NO₃⁻), acetates, chlorates are always soluble

most chlorides are soluble

except silver, mercury (I), lead

lead (II) chloride is soluble in hot water

most sulfates are soluble

except those of barium, strontium, lead, calcium, mercury

most carbonates, phosphates, silicates are insoluble

except those of sodium, potassium, ammonium

most sulfides are insoluble

except those of calcium, strontium, sodium, potassium, ammonium

example: Na₂SO₄(aq) + Pb(NO₃)₂(aq)

reactants are always aqueous

double replacement reaction: switch

check charges

NaNO₃(aq) + PbSO₄(s)

example: KNO₃ (aq) + BaCl₂ (aq)

KCl + BaNO₃

KCl + Ba(NO₃)₂

KCl(aq) + Ba(NO₃)₂(aq)

no precipitate

net ionic equation: separate every aqueous solution into ions and remove what is the same on both sides or just keep the ions from the solid

example: what is the precipitate and net ionic equation of NaCl(aq) + AgNO₃(aq)

write out entire equation: NaCl(aq) + AgNO₃(aq) → NaNO₃(aq) + AgCl(s)

for transition metals, find charge from reactants

AgCl is the precipitate

ionic equation: Na⁺(aq) + Cl⁻(aq) + Ag⁺(aq) + NO₃⁻(aq) → Na⁺(aq) + NO₃⁻(aq) + AgCl(s)

net ionic equation: Ag⁺(aq) + Cl⁻(aq) → AgCl(s)

spectator ions: ions that did not participate in the reaction; not present in a net ionic equation

Acids and bases

acid

base

usually has an H in the formula

usually has an OH in the formula

pH less than 7

pH greater than 7

sour

bitter

can burn you

gives protons

accepts protons

e.g. HCl, vinegar, citric acid

e.g. soap, bleach, NaOH

There are no rows in this table

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Arrhenius’s definition

acid: always makes an H⁺ when put into water

base: always makes an OH⁻ when put into water

Brønsted-Lowry definition

acid: donates proton (H⁺)

base: accepts proton (H⁺)

example: HCl + H₂O

acid (HCl) donates H to base (H₂O)

H₃O + Cl⁻

example: NaOH + H₂O

acid (H₂O) donates H to base (NaOH)

OH⁻ + Na⁺ + H₂O

amphoteric: can be an acid or a base

e.g. water

conjugate acid: base becomes an acid

conjugate base: acid becomes a base

example: NH₃ + H₂O

acid (H₂O) donates H to base (NH₃)

NH₄⁺ + OH⁻

NH⁺ is a conjugate acid

OH⁻ is a conjugate base

example: HBr + H₂O

H₃O⁺ + Br⁻

H₃O⁺ is a conjugate acid

Br⁻ is a conjugate base

strong acid: acid that breaks apart completely

HCl

HBr

H₂SO₄

HNO₃

HClO₄

strong base: base that breaks apart completely

if it has OH in the formula

neutralization reaction: acid + base → water + salt (ionic compound)

e.g. HCl + NaOH → H₂O + NaCl

e.g. HBr + LiOH → H₂O + LiBr

Oxidation-reduction

oxidation-reduction reaction: reaction that contains oxygen and/or a transfer of electrons

e.g. rusting, combustion, cellular respiration, photosynthesis, bleaching, batteries

OIL RIG

oxidation: involves loss of electrons

reduction: involves gaining electrons

oxidation numbers: keep track of electrons in a oxidation-reduction reaction

element → 0

monoatomic ion → charge of ion

fluorine → -1

oxygen → -2

peroxide (O₂²⁻) → -1 (only for O—O bonds)

hydrogen → +1

example: CO₂; oxidation number for carbon

oxygen is always -2

2 oxygens → -4

oxidation number for carbon must be +4 to have neutral charge

example: NO₃⁻; oxidation number for nitrogen

oxygens have -6 charge

nitrogen must have +5 charge to have 1- charge

example: 2Na + Cl₂ → 2NaCl; what is oxidized and what is reduced

charges: 2(0) + 0 → 2(+1)(-1)

Na → lost an electron → oxidized

Cl → gained an electron → reduced

example: balance Cu(s) + Ag⁺(aq) → Ag(s) + Cu²⁺(aq)

determine the half-reactions

Cu(s) → Cu²⁺(aq)

Ag⁺(aq) → Ag(s)

add electrons to balance the charges

Cu(s) → Cu²⁺(aq) + 2e⁻

Ag⁺(aq) + e⁻ → Ag(s)

multiply the reaction(s) to make the electrons equal

Cu(s) → Cu²⁺(aq) + 2e⁻

2Ag⁺(aq) + 2e⁻ → 2Ag(s)

combine reactions and cancel

Cu(s) + 2e⁻ + 2Ag⁺(aq) → Cu²⁺(aq) + 2e⁻ + 2Ag(s)

Cu(s) + 2Ag⁺(aq) → Cu²⁺(aq) + 2Ag(s)

Solutions

Electrolytes

Molarity and dilutions

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