C8 ACIDS AND ALKALIS Flashcards
pH scale
A scale running from 0-14 that measures how acid or alkaline a solution is
Acid
A solution with a pH less than 7
Alkali
A substance with pH greater than 7
Neutral
A substance with a pH equal to 7
Indicator
A substance that changes colour depending on the pH
Common indicators
Litmus:red in acid, blue in alkali
Methyl orange: red in acid, orange alkali
Phenolphthalein: colourless in acid, pink in alkali
Universal indicator
A mixture of serval indicators that is red in strong acids, green when neutral and purple in strong alkali
Acids and ions
Acids dissolve in water to produce an excess of hydrogen ions H+
Alkalis and ions
Alkalis dissolve in water to produce an excess of hydroxide ions OH-
Hydrochloric acid
Formula: HCL
hydrogen ions formed: 1
Anion formed: chloride CL-
nitric acid
formula : HNO3
hydrogen ions formed: 1
anion formed: nitrate, NO3
sulfuric acid
formula: H2SO4
hydrogen ions formed: 2
anion formed: sulfate, SO42-
ions and pH
the higher the hydrogen ion concentration the lower the pH the higher the hydroxide ion concentration , the higher the pH
concentrated solution
a solution with a large amount of solute dissolved in a given volume
dilute solution
a solution with a small amount of solute dissolved in a given volume
pH and hydrogen ion concentration
every step down the pH scale is a ten-fold increase in hydrogen ion concentration and vice versa
pH 3 to 1 = 100 times increase
ph 4 to 7 = 1000 times decrease
dissociation
when an acid dissolves in water, it splits up into positive hydrogen ions and negative anions
strong acids
acids that dissociate fully when dissolved in water, every single molecule splits up
weak acids
acids that do not fully dissociate when dissolved in water, only some molecules split up
acid examples
strong : hydrochloric, sulfuric
weak : ethanoic
properties of strong acids
strong acids react more quickly than weak acids because there are more hydrogen ions available for reactions
base
a substance that neutralises an acid to form a salt and water
salt
a compound formed from the metal cation of a base and the non-metal anion of an alkali
naming salts
two part names . first part = the metal from the base , second part = the anion from the acid
acids and their anions
sulfuric acid - sulfate
nitric acid - nitrate
hydrochloric acid - chloride
reaction of metal oxides with acid
metal oxide + acid - salt + water
e.g magnesium oxide + hydrochloric acid - magnesium chloride
preparing soluble salts
gently warm a beaker of acid
add a spatula of metal oxide and stir until dissolved
repeat until it no longer dissolves
filter to remove excess oxide
allow water to evaporate to produce pure crystals
bases and alkalis
a base is a substrate that neutralises an acid to form a salt and water. an alkali is a base that is soluble in water
common alkalis
sodium hydroxide , NaOH
potassium hydroxide , KOH
calcium hydroxide , Ca(OH)2
reaction of alkalis with acids
acid + alkali - salt + water
eg.
sodium hydroxide nitric acid - sodium nitrate + water
NaOH(aq) + HNO3(aq) - NaNO3(aq) + H2O(L)
balancing equations
use a tally chart to keep track of the number of atoms on each side.
change the coefficients (the big numbers) to add more of things that are missing
DO NOT TOUCH THE LITTLE NUMBERS
acid and alkali ions
acids produced hydrogens ions H+, alkalis produced hydroxide ions , OH-
ions and neutralisation
the H+ ion and OH- ion react together to form H2O (water)
producing a salt by neutralisation
the salt is produced from the ions left over once the H+ and OH- ions have reacted together
burette
a tall glass tube with 0.1cm3 markings on it and a tap at the bottom used for accurately adding variable amounts of liquid
pipette
a piece of glassware used to very accurately measure a fixed amount of liquid
titration
a method used to find out exactly how much acid is needed to neutralise an alkali
titration metod
add alkali to beaker with a pipette
add an alkali to the beaker
gradually add acid from a burette
note how much has been added at the point of neutralisation
titration indicators
use indicators with a sharp colour change - such as phenolphthalein - rather than a gradual one such as universal
reaction of acid with metals
metal + acid - salt + hydrogen
metal and acid observations
bubbles of hydrogen
metal dissloves
warms up
ionic equation
a chemical equation that shows changes to the ions in a reaction
ionic equation for magnesium and acid
mg+2H+ -> MG2+ +H2
Spectator ion
an ion that does not change during a chemical reaction
half equations
an equation that shows what happens to just one of the ions during chemical reaction. two half equations combine to give the overall ionic equation
half equation examples
mg -> MG2+ +2e
2H+ +2e- - H2
combine to give:
MG + 2H+ - MG2+ +H2
reaction of metal carbonates with acid
carbonate + acid -> salt + water + carbon dioxide
carbonate and acid observations
bubbles of CO2 gas
solid carbonate dissolves
carbonate and acid ionic equation
2H+ +CO3’2- - H20 +CO2
soluble
when a substance can be dissolved by a liquid
insoluble
when a substance cannot be dissolved by a liquid
soluble in water
all common sodium potassium and ammonium salts all nitrates most chlorides most sulfates
insoluble in water
sliver and lead chlorides
lead, barium and calcium sulfates
most carbonates
most hydroxides
precipitate
a solid (insoluble) product formed by mixing two solutions. turns the solution cloudy
precipitate reaction
a reaction that produces a solid precipitate by mixing two solutions
predicting precipitate
when mixing two solutions, swap the names of the salts around to find the possible products. if one is insoluble a precipitate forms
predicting equations
AB + YX - AX + YB
eg
sodium chloride + sliver nitrate - silver chloride + sodium
precipitation ionic equations
only include the ions that make that solid precipitate
to prepare insoluble salts
mix your two solutions
filter the mixture
wash the residue by pouring distilled water through the filter
leave somewhere warm to dry
