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4. Types of chemical reactions and solution stoichiometry

Water, the common solvent

aqueous solutions: solutions in which water is the solvent (dissolving medium)

water can dissolve many different substances

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H₂O molecules

O—H bonds are covalent bonds (electron sharing between oxygen and hydrogen atoms

electrons not shared equally

oxygen has greater attraction for electrons than hydrogen

oxygen atoms gain slight excess of negative charge; hydrogen atoms become slightly positive

polar molecule: unequal charge distribution

δ (delta) indicates partial charge (+ or -)

water dissolves ionic substances

hydration: when “positive ends” of water molecules are attracted to negatively-charged anions and “negative ends” are attracted to positively-charged cations

hydration of ions tends to cause salt to “fall apart” or dissolve into the water

strong forces among the ions are replaced by strong water-ion interactions

when ionic substances (salts) dissolve in water, they break up into the individual cations and aniosn

e.g. ammonium nitrate dissolves in water:

NH₄NO₃ —(H₂Ol)→ NH₄⁺ (aq) + NO₃⁻ (aq)

solubility of ionic substances in water differs greatly

differences typically depend on:

relative attractions of the ions for each other (holding the solids together)

attractions of the ions for water molecules (which cause the solid to disperse in water)

when an ionic compound does dissolve in water, the ions become hydrated and are dispersed

water also dissolves many nonionic substances

e.g. ethanol (C₂H₅OH)

contains a polar O—H bond like in water; very compatible

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many substances do not dissolve in water

e.g. animal fat: fat molecules are nonpolar and do not interact effectively with polar water molecules

“like dissolves like”: polar and ionic substances are expected to be more soluble in water than nonpolar substances

Strong and weak electrolytes

solute: substance dissolved in liquid to form a solution

solvent: dissolving medium in a solution

electrical conductivity: ability of a solution to conduct an electric current

electrolyte: substance that when dissolved in water produces a solution that can conduct electricity

solutions with strong electrolytes can conduct a current very efficiently

solutions with weak electrolytes conduct only a small current

solutions with nonelectrolytes permit no current to flow

basis for conductivity properties of solutions

first identified by Svante Arrhenius (1859-1927)

then a Swedish graduate student in physics

came to believe that the conductivity of solutions arose from the presence of ions

initially scorned by majority of scientific establishment

late 1890s: atoms found to contain charged particles; ionic theory became widely accepted

extent to which a solution can conduct an electric current depends directly on the number of ions present

Strong electrolytes

substances that are completely ionized when they are dissolved in water

soluble salts: contain array of cations and ions that separate and become hydrated when the salt dissolves

strong acids: strong electrolytes that dissociate (ionize) completely in aqueous solution

Arrhenius: found that when HCl, HNO₃, and H₂SO₄ were dissolved in water, they behaved as strong electrolytes

acid: substance that produces H⁺ ions (protons) when it is dissolved in water

ionization of an acid: HA (aq) + H₂O (l) → H₃O⁺ (aq) + A⁻ (aq)

hydrochloric acid, nitric acid, sulfuric acid are aqueous and should be written with (aq) (HCl (aq), HNO₃ (aq), H₂SO₄ (aq)) although they often appear without them

a strong acid is one that dissolves completely into its ions; virtually no HCl molecules exist in aqueous solutions

sulfuric acid: formula indicates that it can produce two H⁺ ions per molecule, but only the first H⁺ ion is completely dissociated

second can be pulled off under certain conditions

aqueous solution of H₂SO₄ contains mostly H⁺ ions and HSO₄⁻ ions

strong bases: soluble ionic compounds contain the hydroxide ion (OH⁻)

when dissolved in water, cations and OH⁻ ions separate and move independently

most common basic solutions: when sodium hydroxide (NaOH) or potassium hydroxide (KOH) is dissolved in water to produce ions

Weak electrolytes

substances that exhibit a small degree of ionization in water

produce relatively few ions when dissolved in water

most common: weak acids and weak bases

weak acids: any acid that dissociates (ionizes) only to a slight extent in aqueous solutions

formulas

often written with acidic hydrogen atom(s) (any that will produce H⁺ ions in solution) listed first

any nonacidic hydrogens written later

example: HC₂H₃O₂ (acetic acid) indicates one acidic and three nonacidic hydrogen atoms

disassociation reaction for acetic acid in water: HC₂H₃O₂ (aq) + H₂O (l) ⇌ H₃O⁺ (aq) + C₂H₃O₂⁻ (aq)

unlike strong acids, very small numbers of molecules (~1% for acetic acid) dissociate in aqueous solutions at typical concentrations

weak bases: resulting solution is a weak electrolyte

most common weak base: ammonia (NH₃)

ammonia dissolved in water: NH₃ (aq) + H₂O (l) ⇌ NH₄⁺ (aq) + OH⁻ (aq)

basic because OH⁻ ions produced

Nonelectrolytes

substances that dissolve in water but do not produce any ions

e.g. ethanol: entire C₂H₅OH molecules dispersed into water

molecules do not break up into ions

resulting solution does not conduct electric current

e.g. table sugar/sucrose (C₁₂H₂₂O₁₁): very soluble in water but produces no ions when dissolved

Composition of solutions

to perform stoichiometric calculations when two solutions are mixed, you must know:

the nature of the reaction (depends on the exact forms the chemicals take when dissolved)

the amounts of chemicals present in the solutions (usually expressed as concentrations)

most commonly used expression of concentration: molarity (M)

moles of solute per volume of solution in liters

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Dilution

dilution: adding water to achieve the molarity desired for a particular solution

common acids are purchased as concentrated solutions and diluted as needed

calculating

determining how much water must be added to an amount of stock solution to achieve a solution of the desired concentration

moles of solute after dilution = moles of solute before dilution

steps

determine the number of moles of substance in the final solution by multiplying the volume by the molarity

use a volume with a known molarity and known amount of substance to solve for the volume

example: 500. mL of 1.00 M acetic acid (HC₂H₃O₂) from what volume of a 17.4-M stock solution of acetic acid?

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to make 500 mL of a 1.00-M acetic acid solution, take 28.7 mL of 17.4 M acetic acid and dilute to a total volume of 500 mL with distilled water

dilution procedure

pipet: device for accurately measuring and transferring a given volume of solution

volumetric (transfer) pipets: come in specific sizes (e.g. 5 mL, 10 mL, 25 mL)

measuring pipets: used to measure volumes for which a volumetric pipet is not available

volumetric flask

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M: molarity

V: volume

Types of chemical reactions

millions of possible chemical reactions

grouped into classes

precipitation reactions

acid-base reactions

oxidation-reduction reactions

Precipitation reactions

precipitation reaction (double displacement reaction): when two solutions are mixed and a precipitate forms (insoluble substance/solid)

predicting identity of precipitate

example: K₂CrO₄(aq) added to Ba(NO₃)₂(aq) form yellow solid

think about what products are possible

in virtually every case, when a solid containing ions dissolves in water, the ions separate

think about the nature of each reactant solution

Ba(NO₃)₂(aq)

barium nitrate (white solid) dissolved in water

contains Ba²⁺ and NO₃⁻ ions (not Ba(NO₃)₂ units)

K₂CrO₄(aq)

potassium chromate (solid) dissolved in water

contains K⁺ and CrO₄²⁻ ions

two ways to represent the mixing of Ba(NO₃)₂(aq) and K₂CrO₄(aq)

K₂CrO₄(aq) + Ba(NO₃)₂(aq) → products

2K⁺(aq) + CrO₄²⁻(aq) + Ba²⁺(aq) + 2NO₃⁻(aq) → products

solution contains:

K⁺

CrO₄²⁻

Ba²⁺

NO₃⁻

know that:

when ions form a solid compound, the compound must have a zero net charge

products contain both anions and cations

K⁺ and Ba²⁺ could not combine

CrO₄²⁻ and NO₃⁻ could not combine

most ionic materials contain only two types of ions: one type of cation and one type of anion

determine possibilities for the solid

possible combinations of a given cation and anion from the list of ions in the solution:

K₂CrO₄

KNO₃

BaCrO₄

Ba(NO₃)₂

K₂CrO₄ and Ba(NO₃)₂ were the reactants, so the solid could be KNO₃ or BaCrO₄

need more facts

K⁺ ion and NO₃⁻ ion are both colorless, so KNO₃ would be white, not yellow

CrO₄²⁻ ion is yellow (K₂CrO₄ is yellow) so the solid is most likely BaCrO₄

K⁺ and NO₃⁻ were left dissolved in the solution

K₂CrO₄(aq) + Ba(NO₃)₂(aq) → BaCrO₄(s) + 2KNO₃(aq)

as long as water is present, the KNO₃ remains dissolved as separate ions

example: AgNO₃(aq) and KCl(aq) form white solid

Ag⁺, NO₃⁻ + K⁺, Cl⁻ → Ag⁺, NO₃⁻, K⁺, Cl⁻ → white solid

possible compounds: AgNO₃, KCl, AgCl, KNO₃

cannot be AgNO₃ or KCl, so either AgCl or KNO₃

KNO₃ is quite soluble in water, so solid KNO₃ will not form; solid must be AgCl

AgNO₃(aq) + KCl(aq) → AgCl(s) + KNO₃(aq)

doing chemistry requires both understanding ideas and remembering key information

Solubility of salts in water

Solubility of salts in water

most nitrate (NO₃⁻) salts are soluble

most salts containing the alkali metal ions (Li⁺, Na⁺, Cs⁺, Rb⁺) and the ammonium ion (NH₄⁺) are soluble

most chloride, bromide, and iodide salts are soluble (notable exceptions: salts containing Ag⁺, Pb²⁺, Hg₂²⁺)

most sulfate salts are soluble (notable exceptions: BaSO₄, PbSO₄, Hg₂SO₄, CaSO₄)

most hydroxides (e.g. NaOH, KOH) are only slightly soluble; Ba(OH)₂, Sr(OH)₂, Ca(OH)₂ are marginally soluble

most sulfide (S²⁻), carbonate (CO₃²⁻), chromate (CrO₄²⁻), and phosphate (PO₄³⁻) salts are only slightly soluble, except those containing the cations in Rule 2

There are no rows in this table

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when solutions containing ionic substances are mixed, determine the products by thinking in terms of ion interchange

take the cation from one reactant and combine it with the anion of the other reactant

use solubility rules to predict whether each product forms as a solid

focus on the actual components of the solution before any reaction occurs and then figure out how these components will react with each other

Describing reactions in solution

formula equation: overall reaction stoichiometry but not necessarily the actual forms of the reactants and products in solution

e.g. K₂CrO₄(aq) + Ba(NO₃)₂(aq) → BaCrO₄(s) + 2KNO₃(aq)

complete ionic equation: represents as ions all reactants and products that are strong electrolytes

e.g. 2K⁺(aq) + CrO₄²⁻(aq) + Ba²⁺(aq) + 2NO₃⁻(aq) → BaCrO₄(s) + 2K⁺(aq) + 2NO₃⁻(aq)

net ionic equation: only those solution components undergoing a change

spectator ions: ions that do not participate directly in the reaction

e.g. Ba²⁺(aq) + CrO₄²⁻(aq) → BaCrO₄(s)

Stoichiometry of precipitation reactions

the same principles of chemical stoichiometry apply to reactions that take place in solutions

convert all quantities to moles

use coefficients of the balanced equation to assemble mole ratios

determine which reactant is limiting

in solution reactions:

it is sometimes difficult to tell immediately what reaction will occur when two solutions are mixed

always first write the species that are actually present in the solution

to obtain the moles of reactants, use the volume of the solution and its molarity

solving stoichiometry problems for reactions in solution

identify the species present in the combined solution, and determine what reaction occurs

write the balanced net ionic equation for the reaction

calculate the moles of reactants

determine which reactant is limiting

calculate the moles of product(s), as required

convert to grams or other units, as required

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Acid-base reactions

acid: proton donor

base: proton acceptor

predicting acid-base reaction:

example: HCl(aq) and NaOH(aq)

focus on the species present in the mixed solution

contains H⁺, Cl⁻, Na⁺, OH⁻

separated ions because HCl is a strong acid and NaOH is a strong base

what reaction will occur?

will NaCl precipitate? no; soluble in water

Na⁺ and Cl⁻ are spectator ions

water is a nonelectrolyte; large quantities of H⁺ and OH⁻ cannot coexist in the solution

react to form H₂O molecules: H⁺(aq) + OH⁻(aq) → H₂O(l)

example: aqueous solution of acetic acid (HC₂H₃O₂) and aqueous solution of potassium hydroxide (KOH)

species

aqueous solution of acetic acid is weak electrolyte; does not dissociate into ions to any great extent

when solid KOH is dissolved in water, it dissociates completely to produce K⁺ and OH⁻

in the solution formed by mixing aqueous solutions of HC₂H₃O₂ and KOH, before any reaction occurs, the principles species are HC₂H₃O₂, K⁺, OH⁻

reaction

precipitation does not occur between K⁺ and OH⁻ because KOH is soluble

hydroxide ion (proton acceptor) and proton donor

the HC₂H₃O₂ are a source of protons

the OH⁻ ion has suh a strong affinity for protons that it can strip them from the HC₂H₃O₂ molecules

OH⁻(aq) + HC₂H₃O₂(aq) → H₂O(l) + C₂H₃O₂⁻(aq)

the hydroxide ion is such a strong base that for purposes of stoichiometric calculations it can be assumed to react completely with any weak acid (that will be encountered)

performing calculations for acid-base reactions

list the species present in the combined solution before any reaction occurs, and decide what reaction will occur

write the balanced net ionic equation for this reaction

calculate the moles of the reactants

reactions in solution: use the volumes of the original solutions and their molarities

determine the limiting reactant where appropriate

calculate the moles of the required reactant or product

convert to grams or volume (of solution), as required

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Acid-base titrations

volumetric analysis: technique for determining the amount of a certain substance by doing a titration

titration: delivery (from a buret) of a measured volume of a solution of known concentration (titrant) into a solution containing the substance being analyzed (analyte)

titrant contains substance that reacts in a known manner with the analyte

equivalence point/stoichiometric point: point in the titration where enough titrant has been added to react exactly with the analyte

indicator: substance added at the beginning of the titration that changes color at (or very near) the equivalence point

endpoint: point where the indicator actually changes color

goal: choose an indicator such that the endpoint occurs exactly at the equivalence point

requirements for a successful titration

the exact reaction between titrant and analyte must be known (and rapid)

the stoichiometric (equivalence) point must be marked accurately

the volume of titrant required to reach the stoichiometric point must be known accurately

acid-base titration: when the analyte is a base or acid, the required titrant is a strong acid or strong base

common indicator: phenolphthalein

colorless in acidic solution

pink in basic solution

endpoint occurs approximately one drop of base beyond the stoichiometric point

in doing solution reactions, first write down all the components in the solution and focus on the chemistry of each one

Oxidation-reduction reactions

oxidation-reduction (redox) reaction: one or more electrons are transferred

most reactions used for energy production are redox reactions

photosynthesis

combustion

Oxidation states

way to keep track of electrons in oxidation-reduction reactions, particularly redox reactions involving covalent substances

electrons are shared by atoms in covalent bonds

obtained by arbitrarily assigning the electrons to particular atoms

two identical atoms: split equally

two different atoms (unequal): assigned completely to atom with stronger attraction for electrons

e.g. H₂O: oxygen has greater attraction

oxygen has excess of two electrons (oxidation state is -2)

each hydrogen has no electrons (oxidation state is +1)

oxidation states (oxidation numbers): the imaginary charges the atoms would have if the shared electrons were divided equally between identical atoms or were assigned to the atom in each bond that has the greater attraction for electrons

the sum of the oxidation states must be zero for an electrically neutral compound

writing charges

ionic (actual): n±

oxidation states (not actual): ±n

noninteger oxidation states can occur because of the arbitrary way electrons are divided

e.g. in Fe₃O₄

each oxygen has usual oxidation state of -2

each iron atom has an oxidation state of +8/3

can be viewed as four O²⁻ ions, two Fe³⁺ ions, one Fe²⁺ ion per formula unit

Assigning oxidation states

Assigning oxidation states

Oxidation state of:

Summary

Examples

an atom in an element is zero

element: 0

Na(s), O₂(g), O₃(g), Hg(l)

a monatomic ion is the same as its charge

monatomic ion: charge of ion

Na⁺, Cl⁻

fluorine is -1 in its compound

fluorine: -1

HF, PF₃

oxygen is usually -2 in its compounds (peroxides: oxygen is -1)

oxygen: -2

H₂O, CO₂

hydrogen is +1 in its covalent compounds

hydrogen: +1

H₂O, HCl, NH₃

There are no rows in this table

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Characteristics of oxidation-reduction reactions

oxidation: increase in oxidation state (loss of electrons)

reduction: decrease in oxidation state (gain of electrons)

mnemonic: OIL RIG (Oxidation Involves Loss; Reduction Involves Gain)

oxidizing agent (electron acceptor): accepts electrons

reducing agent (electron donor): gives electrons

when the oxidizing or reducing agent is named, the whole compound is specified

oxidized

reduced

loses electrons

gains electrons

oxidation state increases

oxidation state decreases

reducing agent

oxidizing agent

There are no rows in this table

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Balancing oxidation-reduction reactions

Oxidation states method

balancing oxidation-reduction reactions by oxidation states

write the unbalanced equation

determine the oxidation states of all atoms in the reactants and products

show electrons gained and lost using “tie lines”

use coefficients to equalize the electrons gained and lost

balance the rest of the equation by inspection

add appropriate states

example: solid copper and silver ions in aqueous solution: Cu(s) + Ag⁺(aq) → Ag(s) + Cu²⁺(aq)

assign oxidation states

by element

Cu: 0

Ag⁺: +1

Ag: 0

Cu²⁺: +2

Ag gains one electron

Cu loses two electrons

balance redox equation

need equal numbers of electrons gained and lost

Cu(s) + 2Ag⁺(aq) → 2Ag(s) + Cu²⁺(aq)

example: H⁺(aq) + Cl⁻(aq) + Sn(s) + NO₃⁻(aq) → SnCl₆²⁻(aq) + NO₂(g) + H₂O(l)

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