Topic 5: acids, bases and salts Flashcards
The bronsted lowry definition of acids and bases
The bronsted lowry reaction
an acid is a substance that can donate a proton (H+) to another substance
a base is a substance that can accept a proton (H+) from another substance
the bronsted lowry reaction involves the transfer of a proton from the acid (proton donor) to the base (proton acceptor) in water
- even water can be an acid
bases examples
- include all metal oxides and metal hydroxides, carbonates and alkali
metal oxides and hydroxides examples - Na2O (sodium oxide)
- ZnO (zinc oxide)
- Al(OH)3 (aluminium hydroxide)
alkalis are bases that are soluble in water and produces OH- - NaOH (sodium hydroxide)
- KOH (potassium hydroxide)
- Ca(OH)2 (calcium hydroxide)
Strength of acids and bases
Strong acid
: dissociates fully in solution to give protons (H+)
- approx 100% degree of dissociation
- type of arrow used: a single, irreversible arrow
Strong base
: dissiciates fully in solution to give hydroxide ions (OH-)
- - approx 100% degree of dissociation
- type of arrow used: a single, irreversible arrow
Weak acid
: dissociates partially in solution to give protons (H+)
- <100% degree of dissociation
- type of arrow used: reversible one
Weak base
: dissociates partially in solution to give hydroxide ions (OH-)
- <100% degree of dissociation
- type of arrow used: reversible one
mono, di and tribasic acid
monobasic acid : a substance that produces 1 hydrogen ion per mole of acid
dibasic acid : a substance that produces 2 hydrogen ions per mole of acid
tribasic acid : a substance that produces 3 hydrogen ions per mole of acid
strength vs concentration of an acid
strength of an acid is the measure of the extent of dissociation in an acid in solution
concentration of an acid is the number of moles of undissociated acid per unit volume
*stength of an acid/ base is independant of its concentration
* in order to measure the pH of 2 acids, concentration must be kept at the same
pH of a solution
: the negative logarithm to the base 10 of the concentration of hydrogen ions in solution in mol dm^-3
pH = -log10 (H+)
(H+) = 10 ^-pH
The larger the pH value, the smaller the concentration of H+
pH scale details
- at 25 degrees celsius, the pH ranges from 0 to 14 inclusive
- used to indicate whether a solution is acidic, neutral or alkaline
at the temperature of 25 degrees celsius, - neutral solution: H+ = OH- and pH=7
- acidic solution: H+ > OH- and pH <7
- alkaline solution: H+ < OH- and pH >7
limewater: 12pH
indicators
: dyes/ mixture of dyes which change colour when added to acids or alkakis
A. universal indicator
- a mixture of dyes than can be used to estimate the pH of a solution by tallying the colour of the solution with a universal colour chart
red: acidic
green: neutral
violet: alkaline
B. Indicators used in titration
litmus (not used in titration), 5-8, red, blue
methyl orange, 3-4.4, red, yellow
screened methyl, 3-5, magenta, green
thymol blue, 8-9.6, yellow, blue
phenolpthalein, 8-10, colourless, pink
(indicator name, pH range at which indicator changes colour: end point clr is the middle, clr in acidic, clr in alkaline solution)
properties of acids
-have a sour taste
-pH of less than 7
-turn damp blue litmus paper red, turn universal indicator orane/red
-may be solids (citric acid, C6H8O7), liquids (sulfuric acid H2SO4), or gases (hydrogen chloride gas HCl)
*these must be dissolved in water before they can act as acids
-can conduct electricity (exist as free mobile ions in aq solutions –> can act as charge carriers)
properties of alkalis
-have a bitter taste and feel soapy
-turn damp red litmus paper blue, turn universal indicator blue/ violet
-have pH greater than 7
-can conduct electricity. (excist as free mobile ions in aq solutions and can act as charge carriers)
neutralisation between acids and bases
: refers to the reaction between an acid and a base to form salt and water
acid + BASE –> salt + water
The haber process details
: an industrial method for the manufacture of ammonia (weak base). The reaction is a reversible reaction
- N2 (g) + 3H2 (g) <==> 2NH3 (g) (triangle H= -92 kJ mol ^/1) (- meaning it produces heat)
Common operating conditions
1. high pressure of 250 atm
2. moderate temperature of 450 degrees celsius
3. iron catalyst
4. molar ratio of N2 : H2 is 1:3
Acid base nature of period 3 oxides
changes from basic (metal oxides) to amphoteric to acidic (non metal oxides)
basic oxides
- can react with acids to form salt and water
- metal oxides
- Na2O, MgO
amphoteric oxides
- can react with both acids and bases to form salt and water
- Al2O3
acidic oxides
- can react with bases to form salt and water
- non metal oxides
- SiO2, P4O6 and P4O10, SO2 and SO3
basic oxides details
- group 1 and 2 oxides are metal oxides that are basic in nature
reaction with water: usually insoluble in water. however sodium oxide (Na2O) and potassium oxide (K2O) are soluble in water to give an ALKALINE solution
Na2O + H2O –> 2NaOH (sodium hydroxide)
K2O + H2) –> 2KOH (potassium hydroxide)
reaction with acids: react with acids to give salt and water
(eg react with 2HCl to form chlorides)
acidic oxide details
- group 14, 15 and 16 oxides are non metal oxides that are acidic in nature
reaction with water: some can dissolve in water to form acidic solutions
SO2 + H2O –> H2SO3 (SULFUROUS acid)
SO3 + H2O –> H2SO4 (SULFURIC acid)
reaction with bases: react with bases to form salt and water
P4O10 (s) +12 NaOH (aq) –> 4Na3PO4 (aq)(soldium phosphate) + 6H2O (l)
SO2 (g) + 2NaOH (aq) –> Na2SO3 (sodium sulfite) (aq) + H2O (l)
SO3 (g) + 2NaOH (aq) –> Na2SO4 (aq) (soldium sulfate) + H2O (l)
SiO2 (s) + concentrated hot/ molten 2NaOH (l) –350 degrees–> Na2SiO3 (l) (sodium silicate) + H2O (g)
amphoteric oxide details
reaction with water: does not react with water
reaction with acid and bases: react with both acids and bases to form salt and water
eg al2O3, PbO (lead ii oxide), ZnO
Al2O3 + 6HCl –> 2AlCl3 (aluminium chloride) + 3H2O
Al2O3 + 2NaOH + 3H2O –> 2NaAl(OH)4 (sodium aluminate) *
amphoteric oxide details
reaction with water: does not react with water
reaction with acid and bases: react with both acids and bases to form salt and water
eg al2O3, PbO (lead ii oxide), ZnO
Al2O3 + 6HCl –> 2AlCl3 (aluminium chloride) + 3H2O
Al2O3 + 2NaOH + 3H2O –> 2NaAl(OH)4 (sodium aluminate) *
neutral oxide details
do not react with acids or bases or water
carbon monoxide (CO)
water
nitric oxide (NO)
nitrous oxide (N2O)
common atmospheric pollutants
carbon monoxide, nitrogen dioxide, sulfur dioxide
oxide that causes acid rain
sulfur dioxide and nitrogen dioxide
salt details
- ionic compounds
- contain a metal cation and a non metal anion
- are usually formed by the replacement of one or more hydrogen ions of an acid by a metallic ion or an ammonium ion
eg CaO + H2SO4 –> CaSO4 +H2O
Hydrated and anhydrous salts details
- salts form regularly shaped crystals as they are ionic with a regular ionice lattic
- the water bonded chemically within the crystal is known as the water of crystallisation
: salts that contain water of crystallisation are said to be hydrated
: salts that do not contain water of crystallisation are said to be anhydrous
hydrated salts are named like this:
Na2CO3 ·(the presence of woc) 10H2O (10 moles of woc for every mole of Na2CO3)
Hydrated and anhydrous salts details
- salts form regularly shaped crystals as they are ionic with a regular ionice lattic
- the water bonded chemically within the crystal is known as the water of crystallisation
: salts that contain water of crystallisation are said to be hydrated
: salts that do not contain water of crystallisation are said to be anhydrous
hydrated salts are named like this:
Na2CO3 ·(the presence of woc) 10H2O (10 moles of woc for every mole of Na2CO3)
reactivity series
from most to least: (king napolean can manage all zn fe pb)
K
Na
Ca
Mg
Al
Zn
Fe
Pb
——– will react with acids ^
H
——– unreactive metals (below)
Cu
Ag
Au
Reaction of an acid and a metal
metal + acid –> salt + hydrogen gas
- reactive metals above H in the reactivity series + dilute acid
- do not use the extremely reactive (group 1 metals) or unreactive metals
extremely reactive metals
group 1 metals: K, Na react explosively and violently with dilute acids to form salt and hydrogen gas
reaction of an acid and a carbonate/ hydrogen carbonate
carbonate/ hydrogen carbonate (eg NaHCO3) + acid –> salt + water + carbon dioxide gas
reaction of an acid and a base , acid and ammonia
called a neutralisation reaction
- base include metal oxides, metal hydroxides and ammonia
acid + base –> salt + water
acid + ammonia –> salt only
reaction of an alkali and an ammonium salt (HEAT!!)
ammonium salt (NH4+) + base –(heat)–> salt + water + ammonia gas
- when an ammonium salt is warmed in the presense of an alkali
precipitation reaction details
insoluble salts can be formed by a precipitation reaction through the mixing of 2 soluble (aqueous) reagents
- when an aq solution containing the anion of an Xaq salt is mixed with an aq solution containing the cation of that salt, the Xaq salt will precipitate out of the mixture
