ACIDS, BASES AND SALTS 2B Flashcards
Definition of an acid
Proton donor - releases H+ ions when in aqueous solution
Definition of a strong and weak acid
Strong is fully dissociated in solution
Weak is partially dissociated in solution
What is the link between strength of acids and the concentration of H+ ions?
The weaker the acid is, the lower proportion of the acid is dissociated, and so there is a lower concentration of H+ ions in the solution.
How is a strong acid generally written?
HA(aq) -> H(aq)+ + A(aq)-
How is a weak acid generally written?
HA(aq) EQUILIBRIUM ARROWS H(aq)+ + A(aq)-
Definition of a base
Proton acceptor
Name the types of bases and how they are bases
Metal hydroxides - accepts H+ to become water
Metal oxides - accepts H+ to become OH- (soluble oxide)
- accepts 2H+ to become water
Ammonia/amino compounds - N forms a dative bond with H+ ion
How can water act as a base?
H+ +H2O -> H3O+ (hydronium ion)
How can water act as an acid?
H2O -> OH- + H+
Definition of an alkali
Soluble base that releases OH- ions in water
Definition of a salt
Ionic compounds formed when the H+ ions in acids are replaced by metal ions or ammonium ions
Acid + Metal hydroxide
Salt + Water
Acid + Metal oxide
Salt + Water
Acid + Metal carbonate
Salt + Water + Carbon dioxide
Acid + reactive metal
Salt + Hydrogen
Why do ionic compounds have high melting and boiling points?
A lot of energy is needed to break the strong electrostatic forces of attraction between oppositely charged ions.
Why do ionic compounds conduct electricity when molten or in aqueous solution?
When molten, ions are free to move as mobile charge carriers.
Why do ionic compounds dissolve in polar solvents?
Polar molecules break down the lattice structure and surround the ions in solution. If the compound has larger charges, they are less soluble.
Why are ionic compounds hard crystalline substances?
Made up of ions that are strongly attracted to each other by electrostatic forces
Soluble or insoluble - K
Soluble
Soluble or insoluble - Na
Soluble
Soluble or insoluble - NH4+
Soluble
Soluble or insoluble - Halides
Soluble except when combined with Ag+,Pb2+,Hg+
Soluble or insoluble - Nitrates
Soluble
Soluble or insoluble - Sulfates
Soluble except when combined with Ag+,Pb2+,Ca2+,Sr2+,Ba2+
Soluble or insoluble - Carbonates
Insoluble except when combined with Na, K, NH4+
Soluble or insoluble - Oxides
Insoluble except when combined with Group 1 or NH4+
Soluble or insoluble - Phosphates
Insoluble except when combined with Group 1 or NH4+
Soluble or insoluble - Hydroxides
Insoluble except when combined with Group 1 or NH4+, Ca2+, Mg2+, Sr2+, Ba2+
Soluble or insoluble - Sulfides
Insoluble except when combined with Group 1 or NH4+
Soluble or insoluble - Chromates
Insoluble except when combined with Group 1 or NH4+, Ca2+, Mg2+
Acid + Metal hydrogen carbonate
Salt + Water + Carbon dioxide
Acid + Alkali
Salt + Water
Test for halides (including in isolation)
Add silver nitrate to unknown substance
Cl ion = white precipitate
Br ion = cream precipitate
I ion = yellow precipitate
In isolation add dilute ammonia solution -> precipitate dissolves then it’s Cl ions
Add conc. ammonia solution -> dissolves then it’s Br ions
Doesn’t dissolve in either then it’s I ions
What are the ions that don’t change in a reaction called?
Spectator ions
What must ionic equations balance for?
Atoms and charge
Test for sulfates
Add Barium chloride and if a white precipitate forms then it’s a sulfate
Test for carbonates
Add dilute nitric acid and there should be effervescence.
To prove this is CO2, bubble it through limewater and it will turn cloudy
Test for ammonium ions
Add sodium hydroxide solution and heat. Hold damp red litmus paper over it and it will turn blue if ammonium ions are present
