Please enable JavaScript.
Coggle requires JavaScript to display documents.
inorganic chemistry - Coggle Diagram
inorganic chemistry
chemical tests
flame test method
clean a nichrome wire in the Bunsen burner with HCl. Then add some solid metal to the wire and place it back in the Bunsen burner
gases
hydrogen
use a burning splint held at the end of a test tube and you should hear a squeaky pop
oxygen
relights a glowing splint
carbon dioxide
turns limewater cloudy
ammonia
turns damp red litmus paper blue
chlorine
turns damp blue litmus paper red then bleaches it
flame test colours
lithium
red
sodium
yellow
copper
blue green
potassium
lilac
calcium
orange red
carbonate ions
add dilute hydrochloric acid and barium chloride solution to the unknown solution. a thick white precipitate confirms the presence of a sulfate
cations tests
copper II, iron II, iron III
add a few drops of aqueous sodium hydroxide to a solution to a solution of the unknown substance and identify the metal ion by the colour of the precipitate
copper -> blue
iron II -> green
iron III -> brown
ammonium
warm the suspected ammonium salt gently with aqueous sodium hydroxide
if ammonium ions are present a gas is evolved, which is pungent and turns damp red litmus paper blue
halide ions
add dilute nitric acid and then silver nitrate solution to the solution containing the suspected halide and identify by the colour of the precipitate
chloride -> white
bromide -> cream
iodide -> yellow
water
turns anhydrous copper sulfate from white to blue
sulfate ions
add dilute hydrochloric acid and barium chloride solution to the unknown solution and a thick white precipitate confirms the presence of a sulfate
acids, bases and salt preparations
metal oxides, metal hydroxides and ammonia can act as bases, and that alkalis are bases that are soluble in water
how to make a salt from an insoluble reactant
add excess insoluble base to the acid
filter to remove unreacted base
heat the solution so that water evaporates and crystals of the salt remain
solubility
sodium
all soluble
potassium
all soluble
ammonium
all soluble
nitrates
all soluble
chlorides
all soluble except silver and lead
sulfates
all soluble except lead, barium and calcium
carbonates
only sodium, potassium and ammonium are soluble
hydroxides
only sodium, potassium and ammonium are soluble
proton transfer
acids donate protons
bases accept protons
equations
acid + base → salt + water
acid + metal → salt + hydrogen
acid + metal carbonate → salt + water + carbon dioxide
how to make a soluble salt from an acid and an alkali
use a titration to find the exact volume of the alkali that reacts with the acid
mix the exact volumes of the acid and base
warm solution so that water evaporates and crystals of the salt remain
how to make an insoluble salt from two soluble salts
mix solutions of 2 soluble reactants
filter mixture (insoluble salt will remain on filter paper)
wash salt with distilled water
leave salt to dry
acids, alkalis and titrations
neutralisation
A neutralisation reaction is one between an acid and a base
the ionic equation for any alkali-acid neutralisation reaction is: H+(aq) + OH-(aq) -> H20(l)
aqueous acids
acids produce H+ ions in aqueous solutions
alkalis produce OH- ions in aqueous solutions
pH scale
the pH scale (0 - 14) measures the acidity or alkalinity of a solution, and can be measured using universal indicator.
7 is neutral
<7 is acidic
7 is alkaline
universal indicator
add a couple of drops of solution to a piece of universal indicator paper and
observe what colour it goes (compare to pH scale)
indicators
phenophthalein
alkaline = pink
acidic = colourless\
methyl orange
alkaline = yellow
acidic = red
litmus
litmus solution
alkaline = blue
acidic = red
paper
Blue litmus paper goes red in acidic & stays blue in alkaline
Red litmus paper goes blue in alkaline & stays red in acidic
titrations
Wash burette using acid and then water
Fill burette to 100cm3 with acid with the meniscus’ base on the 100cm3 line
Use 25cm3 pipette to add 25cm3 of alkali into a conical flask, drawing alkali into the pipette using a pipette filler
dd a few drops of a suitable indicator to the conical flask (eg: phenolphthalein which is pink when alkaline and colourless when acidic)
Add acid from burette to alkali until end-point is reached (as shown by indicator)
The titre (volume of acid needed to exactly neutralise the acid) is the difference between the first (100cm3) and second readings on the burette)
Repeat the experiment until you get concordant results
reactivity series
definitions
oxidation: gain of oxygen OR loss of electrons
reduction: loss of oxygen OR gain of electrons
redox: a reaction in which both oxidation and reduction occur
oxidising agent: causes another reactant to be oxidised and is reduced itself
reducing agent: causes another reactant to be reduced and is oxidised itself
displacement reactions
You can see if one metal is more reactive than another by using displacement
reactions:
Easily seen when a salt of the less reactive metal is in the solution
More reactive metal gradually disappears as it forms a solution
Less reactive metal coats the surface of the more reactive metal
reactions of metals
A few reactive metals will react with cold water
products are a metal hydroxide (forming and alkaline solution) and hydrogen gas
most meats react with acid
acid + metal → salt + hydrogen
almost all metals react with oxygen
metal + oxygen → metal oxide
Only metal that does not react with any of the above is gold, because it is extremely unreactive
You can therefore deduce the relative reactivity of some metals by seeing if they react with water (i.e. VERY reactive), acid (reactive), and oxygen (not that reactive)
rusting
Both air and water are necessary for iron to rust – i.e. oxidation – gain of oxygen
results in corrosion
prevention of rusting
barrier methods
rust can be prevented by coating iron with barriers that prevent the iron from coming into contact with water and oxygen
However, if the coatings are washed away or scratched, the iron is once again exposed to water and oxygen and will rust
galvanisation/ sacrificial protection
Galvanising is a process where the iron to be protected is coated with a layer of zinc
ZnCO3 is formed when zinc reacts with oxygen and carbon dioxide in the air and protects the iron by the barrier method
If the coating is damaged or scratched, the iron is still protected from rusting because zinc preferentially corrodes as it is higher up the reactivity series than iron
The iron stays protected as it accepts the electrons released by zinc, remaining in the reduced state and thus it does not undergo oxidation
The electrons donated by the zinc react with hydrogen ions in the water producing hydrogen gas:
Zinc therefore reacts with oxygen and water and corrodes instead of the iron
reactivity series
potassium, sodium, lithium, calcium, magnesium, aluminium, carbon, zinc, iron, hydrogen, copper, silver and gold
gases in the atmosphere
climate change
Greenhouse gas effect maintains temperatures on Earth high enough to support life
Greenhouse gases include: water vapour, CO2 & CH4
explanation of the greenhouse gas effect:
Electromagnetic radiation at most wavelengths from the sun passes through the Earth’s atmosphere
The Earth absorbs some radiation and thus warms up (essential for life on Earth). But some heat is radiated from the Earth as infrared radiation.
Some of this IR radiation is absorbed by greenhouse gases in the atmosphere
Atmosphere warms up leading to the greenhouse effect and global warming
global warming is an 'enhanced greenhouse effect'
an increase in average global temperature is a major cause of climate change
thermal decomposition
Metal carbonate –(heat)-> metal oxide + carbon dioxide
combustion
combustion is an example of oxidation
In an oxidation reaction, a substance gains oxygen
Metals and non-metals can take part in these reactions
experiment example with copper
100 cm3 of air passed from side to side over copper that was being heated with a bunsen burner
All oxygen in air will react with copper
It’s a closed system – therefore, no air could get in or out
As it is passed, the volume of air will decrease
Continued until the volume stops decreasing, then record the volume of remaining air
There would be about 79cm3 left, showing that 21cm3 of the original 100cm3 of air was oxygen
percentages of gases in the air
78% nitrogen
21% oxygen
0.96% argon
0.04% carbon dioxide
group 7
physical properties
poisonous non-metals
diatomic gases
7 electrons in their outer shell
appearances, characteristics and colour
fluorine
yellow gas
very reactive
chlorine
pale green gas
reactive and dense gas
bromine
red-brown liquid
orange in solution
dense red-brown volatile liquid
melting points increase as you go down the group due to the atoms size increasing the intermolecular forces, so more energy is required to overcome these forces
predicting properties
halogens react with some metals to form ionic compounds (metal halide salts)
halogens decrease in reactivity moving down the group
rate of reaction is slower for halogens which are further down the group
displacement reactions
a halogen displacement reaction occurs when a more reactive halogen displaces a less reactive halogen in a solution
chlorine displaces both bromine and iodine, whereas bromine only displaces iodine and iodine doesn't displace anything
group 1
properties
they have characteristic properties due to the single electron in the outer shell
all the metals in group 1 react vigorously with water to create an alkaline solution and hydrogen
reactions with water
more bubbles = more vigorous reaction = more reactive alkali metal
lithium
fizzes steadily
sodium
melts into a ball and then fizzes quickly
potassium
gives off sparks and hydrogen burns with a lilac flame
there is an increase in reactivity as you go down the group
they all react with oxygen to make an oxide
electron configurations
down the group makes it easier to lose electrons and form positive metal ions (cations
it is easier to lose electrons due to the increase in electron shells as you go down the group as there is less attraction between the nucleus and the electrons
extraction of metals
most metals are extracted from ores found in the earth's crust and that unreactive metals are often found as the uncombined element
methods of extracton
reduction by carbon
only if the metal is less reactive so that carbon displaces the metal from the ore
electrolysis
used if the metal is more reactive than carbon
expensive due to the large amounts of energy to melt the compounds and to produce the electrical current
uses of metals
aluminium
used for aircrafts and trains
copper
used for electrical wiring and plumbing
iron
hard but too brittle for most uses so normally converted to steel
low carbon steel
0.25% carbon used for car body panels
high carbon steel
2.5% carbon, used for cutting tools
stainless steel
chromium and nickel, used for cutlery and surgical equipment
alloy
Most metals in everyday uses are alloys. Pure copper, gold, iron and aluminium are all too soft for everyday uses and so are mixed with small amounts of similar metals to make them harder for everyday use.
in a pure metal, the ions are all the same size and are in a regular arrangement of layers, meaning that they can slide over each other easily, making them soft. In an alloy, there are different sized ions, which disrupts the regular arrangement
and prevents layers being able to slide over each other so easily.