Acids, Bases, Salts, Solubility,

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Chapter 6

• Acids, Bases, Salts, Solubility,
• And stuff like that!

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Definitions

• Solubility
• those compounds with low solubility are said to
  be insoluble,
• those compounds with higher solubility are said
  to be soluble

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More Definitions

• saturated solution
• unsaturated solution
• supersaturated solution

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Soluble or Insoluble

• Explain why some substances are soluble and other
  substances are not soluble by giving one example
  of each. Used balanced equations in you
  discussion.
• You may use the solubility rules Use your
  intelligence and understanding if the internet to
  find them! These are observation based no
  explanation needed at this time.

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Solvation

• What happens when substances dissolve? What
  forces are involved? Use water as a solvent for
  specific examples.
• Ionic?
• Covalent?

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Water as a Solvent

• How water dissolves ionic compounds
• water is a
• ions

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Water as a Solvent

• How water dissolves molecular compounds
• nonpolar covalent molecules
• polar covalent molecules dissolve because
• Each individual molecule is

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Electrolytes
Video Link-electrolytes and non-electrolytes Video
Link Weak and strong electrolytes
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Electrolytes

• cations migrate to the negative electrode (the
  cathode)
• anions migrate to the positive electrode (the
  anode)
• the movement of ions constitutes an electric
  current
• electrolyte
• nonelectrolyte
• strong electrolyte
• weak electrolyte

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Arrhenius Acids and Bases

• In 1884, Svante Arrhenius proposed these
  definitions
• acid a substance that produces H3O ions aqueous
  solution
• base a substance that produces OH- ions in
  aqueous solution

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Arrhenius Acids and Bases

• when HCl, for example, dissolves in water, its
  reacts with water to give hydronium ion and
  chloride ion
• we use curved arrows to show the change in
  position of electron pairs during this reaction

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Arrhenius Acids and Bases

• With bases, the situation is slightly different
• many bases are metal hydroxides such as KOH,
  NaOH, Mg(OH)2, and Ca(OH)2
• these compounds are ionic solids and when they
  dissolve in water, their ions merely separate
• other bases are not hydroxides these bases
  produce OH- by reacting with water molecules

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Arrhenius Acids and Bases

• we use curved arrows to show the transfer of a
  proton from water to ammonia

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Acid and Base Strength

• Strong acid one that reacts completely or almost
  completely with water to form H3O ions
• Strong base one that reacts completely or almost
  completely with water to form OH- ions
• here are the six most common strong acids and the
  four most common strong bases

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Acid and Base Strength

• Weak acid a substance that dissociates only
  partially in water to produce H3O ions
• acetic acid, for example, is a weak acid in
  water, only 4 out every 1000 molecules are
  converted to acetate ions
• Weak base a substance that dissociates only
  partially in water to produce OH- ions
• ammonia, for example, is a weak base

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Brønsted-Lowry Acids Bases

• Acid a proton donor
• Base a proton acceptor
• Acid-base reaction a proton transfer reaction
• Conjugate acid-base pair any pair of molecules
  or ions that can be interconverted by transfer of
  a proton

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Brønsted-Lowry Acids Bases

• Brønsted-Lowry definitions do not require water
  as a reactant

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Brønsted-Lowry Acids Bases

• we can use curved arrows to show the transfer of
  a proton from acetic acid to ammonia

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Brønsted-Lowry Acids Bases

• Note the following about the conjugate acid-base
  pairs in the table
• 1. an acid can be positively charged, neutral, or
  negatively charged examples of each type are
  H3O, H2CO3, and H2PO4-
• 2. a base can be negatively charged or neutral
  examples are OH-, Cl-, and NH3
• 3. acids are classified a monoprotic, diprotic,
  or triprotic depending on the number of protons
  each may give up examples are HCl, H2CO3, and
  H3PO4

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Brønsted-Lowry Acids Bases

• carbonic acid, for example can give up one proton
  to become bicarbonate ion, and then the second
  proton to become carbonate ion
• 4. several molecules and ions appear in both the
  acid and conjugate base columns that is, each
  can function as either an acid or a base

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Brønsted-Lowry Acids Bases

• the HCO3- ion, for example, can give up a proton
  to become CO32-, or it can accept a proton to
  become H2CO3
• a substance that can act as either an acid or a
  base is said to be amphiprotic
• the most important amphiprotic substance in Table
  8.2 is H2O it can accept a proton to become
  H3O, or lose a proton to become OH-
• 5. a substance cannot be a Brønsted-Lowry acid
  unless it contains a hydrogen atom, but not all
  hydrogen atoms in most compounds can be given up
• acetic acid, for example, gives up only one proton

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Brønsted-Lowry Acids Bases

• 6. there is an inverse relationship between the
  strength of an acid and the strength of its
  conjugate base
• the stronger the acid, the weaker its conjugate
  base
• HI, for example, is the strongest acid in Table
  8.2, and its conjugate base, I-, is the weakest
  base in the table
• CH3COOH (acetic acid) is a stronger acid that
  H2CO3 (carbonic acid) conversely, CH3COO-
  (acetate ion) is a weaker base that HCO3-
  (bicarbonate ion)

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6.4. DISSOCIATION OF ACIDS AND BASES IN WATER
Table 6.1. Dissociation of Acids
dissociated Formula Name Common uses in 1 M
solution Strength H2SO4 Sulfuric Industrial
chemical 100 Strong HNO3 Nitric Industrial
chemical 100 Strong H3PO4 Phosphoric Fertilizer,
food 8 Moderately additive weak H3C6
H5O7 Citric Fruit drinks 3 Weak CH3CO2H
Acetic Foods, industry 0.4 Weak HClO
Hypochlorous Disinfectant 0.02 Weak HCN
Hydrocyanic Very poisonous 0.002 Very weak
industrial chemical electroplating
waste H3BO3 Boric acid Antiseptic,
ceramics 0.002 Very weak
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Acid-Base Equilibria

• we know that HCl is a strong acid, which means
  that the position of this equilibrium lies very
  far to the right
• in contrast, acetic acid is a weak acid, and the
  position of its equilibrium lies very far to the
  left
• but what if the base is not water? How can we
  determine which are the major species present?

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Acid-Base Equilibria

• To predict the position of an acid-base
  equilibrium such as this, we do the following
• identify the two acids in the equilibrium one on
  the left and one on the right
• using the information in Table 10.1, determine
  which is the stronger acid and which is the
  weaker acid
• also determine which is the stronger base and
  which is the weaker base remember that the
  stronger acid gives the weaker conjugate base,
  and the weaker acid gives the stronger conjugate
  base
• the stronger acid reacts with the stronger base
  to give the weaker acid and weaker base
  equilibrium lies on the side of the weaker acid
  and weaker base

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Acid-Base Equilibria

• identify the two acids and bases, and their
  relative strengths
• the position of this equilibrium lies to the right

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Acid-Base Equilibria

• Example predict the position of equilibrium in
  this acid-base reaction

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Acid-Base Equilibria

• Example predict the position of equilibrium in
  this acid-base reaction
• Solution the position of this equilibrium lies
  to the right

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Acid Ionization Constants

• when a weak acid, HA, dissolves in water
• the equilibrium constant, Keq, for this
  ionization is
• because water is the solvent and its
  concentration changes very little when we add HA
  to it, we treat H2O as a constant equal to 1000
  g/L or 55.5 mol/L
• we combine the two constants to give a new
  constant, which we call an acid ionization
  constant, Ka

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Acid Ionization Constants

• Ka for acetic acid, for example is 1.8 x 10-5
• because the acid ionization constants for weak
  acids are numbers with negative exponents, we
  commonly express acid strengths as pKa where
• the value of pKa for acetic acid is 4.75
• values of Ka and pKa for some weak acids are
  given in Table 10.2
• as you study the entries in this table, note the
  inverse relationship between values of Ka and pKa
• the weaker the acid, the smaller its Ka, but the
  larger its pKa

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Properties of Acids Bases

• Neutralization
• acids and bases react with each other in a
  process called neutralization.
• Reaction of acids with metals
• strong acids react with certain metals (called
  active metals) to produce a salt and hydrogen
  gas, H2

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Properties of Acids Bases

• Reaction with metal hydroxides
• reaction of an acid with a metal hydroxide gives
  a salt plus water
• the reaction is more accurately written as
• omitting spectator ions gives this net ionic
  equation

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Properties of Acids Bases

• Reaction with metal oxides
• strong acids react with metal oxides to give
  water plus a salt

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Properties of Acids Bases

• Reaction with carbonates and bicarbonates
• strong acids react with carbonates to give
  carbonic acid, which rapidly decomposes to CO2
  and H2O
• strong acids react similarly with bicarbonates

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Properties of Acids Bases

• Reaction with ammonia and amines
• any acid stronger than NH4 is strong enough to
  react with NH3 to give a salt

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Self-Ionization of Water

• pure water contains a very small number of H3O
  ions and OH- ions formed by proton transfer from
  one water molecule to another
• the equilibrium expression for this reaction is
• we can treat H2O as a constant 55.5 mol/L

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Self-Ionization of Water

• combining these constants gives a new constant
  called the ion product of water, Kw
• in pure water, the value of Kw is 1.0 x 10-14
• this means that in pure water

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Self-Ionization of Water

• the product of H3O and OH- in any aqueous
  solution is equal to 1.0 x 10-14 for solutions as
  well.
• for example, if we add 0.010 mole of HCl to 1
  liter of pure water, it reacts completely with
  water to give 0.010 mole of H3O
• in this solution, H3O is 0.010 or 1.0 x 10-2
• this means that the concentration of hydroxide
  ion is

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pH and pOH

• we commonly express these concentrations as pH,
  where
• pH -log H3O
• we can now state the definitions of acidic and
  basic solutions in terms of pH
• acidic solution one whose pH is less than 7.0
• basic solution one whose pH is greater than 7.0
• neutral solution one whose pH is equal to 7.0

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pH and pOH

• just as pH is a convenient way to designate the
  concentration of H3O, pOH is a convenient way to
  designate the concentration of OH-
• pOH -logOH-
• the ion product of water, Kw, is 1.0 x 10-14
• taking the logarithm of this equation gives
• pH pOH 14
• thus, if we know the pH of an aqueous solution,
  we can easily calculate its pOH

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pH and pOH

• pH of some common materials

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pH of Salt Solutions

• When some salts dissolve in pure water, there is
  no change in pH from that of pure water
• Many salts, however, are acidic or basic and
  cause a change the pH when they dissolve
• We are concerned in this section with basic salts
  and acidic salts

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pH of Salt Solutions

• Basic salt raises the pH
• as an example of a basic salt is sodium acetate
• when this salt dissolves in water, it ionizes
  Na ions do not react with water, but CH3COO-
  ions do
• the position of equilibrium lies to the left
• nevertheless, there are enough OH- ions present
  in 0.10 M sodium acetate to raise the pH to 8.88

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pH of Salt Solutions

• Acidic salt lowers the pH
• an example of an acidic salt is ammonium chloride
• chloride ion does not react with water, but the
  ammonium ion does
• although the position of this equilibrium lies to
  the left, there are enough H3O ions present to
  make the solution acidic

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Acid-Base Titrations

• Titration an analytical procedure in which a
  solute in a solution of known concentration
  reacts with a known stoichiometry with a
  substance whose concentration is to be determined

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Acid-Base Titrations

• An acid-base titration must meet these
  requirement
• 1. we must know the equation for the reaction so
  that we can determine the stoichiometric ratio of
  reactants to use in our calculations
• 2. the reaction must be rapid and complete
• 3. there must be a clear-cut change in a
  measurable property at the end point (when the
  reagents have combined exactly)
• 4. we must have precise measurements of the
  amount of each reactant

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Acid-Base Titrations

• As an example, let us use 0.108 M H2SO4 to
  determine the concentration of a NaOH solution
• requirement 1 we know the balanced equation
• requirement 2 the reaction between H3O and OH-
  is rapid and complete
• requirement 3 we can use either an acid-base
  indicator or a pH meter to observe the sudden
  change in pH that occurs at the end point of the
  titration
• requirement 4 we use volumetric glassware

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Acid-Base Titrations

• experimental measurements
• doing the calculations

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pH Buffers

• pH buffer a solution that resists change in pH
  when limited amounts of acid or base are added to
  it
• a pH buffer as an acid or base shock absorber
• a pH buffer is common called simply a buffer
• the most common buffers consist of approximately
  equal molar amounts of a weak acid and a salt of
  the conjugate base of the weak acid
• for example, if we dissolve 1.0 mole of acetic
  acid and 1.0 mole of its conjugate base (in the
  form of sodium acetate) in water, we have an
  acetate buffer

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pH Buffers

• How an acetate buffer resists changes in pH
• if we add a strong acid, such as HCl, added H3O
  ions react with acetate ions and are removed from
  solution
• if we add a strong base, such as NaOH, added OH-
  ions react with acetic acid and are removed from
  solution

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pH Buffers

• The effect of a buffer can be quite dramatic
• consider a phosphate buffer prepared by
  dissolving 0.10 mole of NaH2PO4 (a weak acid) and
  0.10 mole of Na2HPO4 (the salt of its conjugate
  base) in enough water to make 1 liter of solution

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pH Buffers

• Buffer pH
• if we mix equal molar amounts of a weak acid and
  a salt of its conjugate base, the pH of the
  solution will be equal to the pKa of the weak
  acid
• if we want a buffer of pH 9.14, for example, we
  can mix equal molar amounts of boric acid
  (H3BO3), pKa 9.14, and sodium dihydrogen borate
  (NaH2BO3), the salt of its conjugate base

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pH Buffers

• Buffer capacity depends both its pH and its
  concentration

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Blood Buffers

• The average pH of human blood is 7.4
• any change larger than 0.10 pH unit in either
  direction can cause illness
• To maintain this pH, the body uses three buffer
  systems
• carbonate buffer H2CO3 and its conjugate base,
  HCO3-
• phosphate buffer H2PO4- and its conjugate base,
  HPO42-
• proteins discussed in Chapter 21

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Henderson-Hasselbalch Eg.

• Henderson-Hasselbalch equation a mathematical
  relationship between
• pH,
• pKa of the weak acid, HA
• concentrations HA, and its conjugate base, A-
• It is derived in the following way
• taking the logarithm of this equation gives

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Henderson-Hasselbalch Eg.

• multiplying through by -1 gives
• -log Ka is by definition pKa, and -log H3O is
  by definition pH making these substitutions
  gives
• rearranging terms gives

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Henderson-Hasselbalch Eg.

• Example what is the pH of a phosphate buffer
  solution containing 1.0 mole of NaH2PO4 and 0.50
  mole of Na2HPO4 dissolved in enough water to make
  1.0 liter of solution

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Henderson-Hasselbalch Eg.

• Example what is the pH of a phosphate buffer
  solution containing 1.0 mole of NaH2PO4 and 0.50
  mole of Na2HPO4 in enough water to make one liter
  of solution
• Solution
• the equilibrium we are dealing with and its pKa
  are
• substituting these values in the H-H equation
  gives

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PREPARATION OF ACIDS Combination of H with
nonmetal H2 Cl2 ? 2HCl Nonmetal with
water Cl2 H2O ? HCl HClO Nonmetal
oxide plus water SO3 H2O ? H2SO4 Evolution
of volatile acid 2NaCl(s) H2SO4(l) ? 2HCl(g)
Na2SO4(s) HCl gas collected in water gives
hydrochloric acid Organic acids, such as acetic
acid, have the carboxylic acid group
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PREPARATION OF BASES Active metal plus water 2K
H2O ? 2K 2OH- H2(g) Metal oxide plus
water CaO H2O ? Ca(OH)2 Substances that
generate OH- in water NH3 H2O ? NH4 OH- Salt
anions that react with water to produce OH-
From Na2CO3 2Na CO32- H2O ? 2Na HCO3-
OH- This reaction is a hydrolysis
reaction Organic bases, particularly amines
(CH3)3N H2O ? (CH3)3NH OH-
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PREPARATION OF SALTS Reaction of acid with base
2NaOH H2SO4 ? 2H2O Na2SO4 (sodium
sulfate) Reaction of metal and nonmetal Ca F2
? CaF2 (calcium fluoride) Metal reacting with
acid Mg H2SO4 ? H2 MgSO4 (magnesium
sulfate) Active metal reacting with base 2Al
6NaOH ? 3H2(g) Na3AlO3 (sodium
aluminate) Addition of a base to a salt to form
another salt and an insoluble base 2KOH MgSO4
? Mg(OH)2(s) K2SO4(aq) Evolution of a volatile
acid leaving a salt 2NaCl(s) H2SO4(l) ?
2HCl(g) Na2SO4(s) Displacement of a metal from
a salt, such as in cementation Fe(s)
CdSO4(aq) ? Cd(s) FeSO4(aq) Specialized
processes, such as the Solvay synthesis of
NaHCO3 NaCl NH3 CO2 H2O ? NaHCO3 NH4Cl
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6.10. ACID SALTS AND BASIC SALTS Acid salts are
salts that contain H and can act as acids
NaHSO4 NaOH ? Na2SO4 H2O Sodium
bicarbonate NaHCO3 Sodium dihydrogen
phosphate, NaH2PO4, used to prepare buffers
Disodium hydrogen phosphate, Na2HPO4, buffers
Potassium hydrogen tartrate, KH4C4H4O6, acid in
baking powder Basic salts contain OH and can
react with H ion Example Calcium
hydroxyapatite source of phosphorus
Ca5OH(PO4)3 Many rock-forming minerals are basic
salts
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6.11. WATER OF HYDRATION Water molecules bound to
other compounds, typically salts Example
Sodium carbonate decahydrate, Na2CO310H2O
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6.12. NAMES OF ACIDS, BASES, AND SALTS Acids H
and a nonmetal Hydro-ic acid Hydrochloric
acid, HCl Oxygen-containing acids H2SO3,
sulfurous acid Table 6.7. Names of Oxyacids of
Chlorine Formula Name Anion name HClO4
perchloric acid perchlorate HClO3 chloric
acid chlorate HClO2 chlorous acid chlorite HClO
hypochlorous acid hypochlorite Bases For ionic
bases containing OH, name of cation followed by
hydroxide NaOH, sodium hydroxide Ca(OH)2,
calcium hydroxide
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Names of Salts Name of cation followed by name of
anion See Table 6.8 for some important ions and
their names Examples Na2SO4, sodium sulfate
KH2PO4, potassium dihydrogen sulfate Ca(ClO)2,
calcium hypochlorite Formulas of Salts Sum of
charge on cations times their subscripts plus sum
of charge on anions times their subscripts must
equal zero
Example Iron(III) sulfate Formula before
adding subscripts Fe(SO4) Cation charge 3
Anion charge -2 2 Fe3 cations gives a
total cation charge of 2 x 3 6 3 SO42- gives
an anion charge of 3 x (-2) -6 Therefore, the
formula is Fe2(SO4)3