2) Inorganic Chemistry Flashcards
Group 1 metals
Alkali metals - form alkaline solutions when they react with water
-lithium, sodium, potassium, rubidium, caesium, francium
-first column of periodic table
-share similar characteristic chemical properties - all have one electron in outermost shell
-compounds are all ionic
Group 1 properties
-soft metals
-low density
-low melting point
-very reactive - only need to lose one electron to become stable
Group 1 metals - reactive
-react quickly with oxygen –> oxides
-react with water
-NO TOUCHING - enough sweat on skin to give reaction - produces heat and very corrosive
-stored under oil
-Rb + Cs are too reactive - stored in sealed glass tube so no oxygen
Group 1 metals - reaction with water
metal + water –> metal hydroxide + hydrogen
2M + 2H20 –> 2MOH + H2
-form alkalis in water
Lithium + water
-relatively slow reaction
-doesn’t melt
-fizzing seen, heard
–> lithium hydroxide + hydrogen
Sodium + water
-large amounts of heat released, causes sodium to melt
-hydrogen released catches fire, causes ball of sodium to dash across the surface
–> sodium hydroxide + hydrogen
Potassium + water
-reacts more violently than Na
-enough heat released so hydrogen burns with lilac coloured flame
-melts into a shiny ball that dashes around the surface
Group 1 metals - reaction with oxygen
-form metal oxides
-alkali metals tarnish when exposed to air, dull coating
Group 1 metals - Physical trends
As you go down the group
-mp/ bp decreases
-density increases
-reactivity increases
Group 1 metals - reactivity in terms of electrons
-as atoms get bigger the outer electron is further from the nucleus - weaker forces of attraction between them
-less strongly attracted by the nucleus
-less energy required to overcome force of attraction
-lost more easily
Group 1 metals - halogens
-react with halogens to form compounds
-MX (LiCl, KBr)
Group 7 metals
-fluorine, chlorine, bromine, iodine, astatine
-similar reactions - all have 7 outer shell electrons
-non-metallic elements that are poisonous
-diatomic molecules - F2, Cl2
-react with metals and non metals to form compounds
Group 7 metals - physical properties
As you go down the group
-mp/ bp increases - as relative molecular mass increases - intermolecular forces increases
-colour increases
-reactivity decreases
Fluorine - appearance (room temp), characteristics, colour in solution
-yellow gas
-very reactive, poisonous
Chlorine - appearance (room temp), characteristics, colour in solution
-pale yellow-green gas
-reactive, poisonous, dense
-pale green
Bromine - appearance (room temp), characteristics, colour in solution
-red brown liquid
-dense red-brown volatile liquid
-orange
Iodine - appearance (room temp), characteristics, colour in solution
-purple-black solid
-shimmery, crystalline solid, sublimes to form a purple vapour
-dark brown
Halogen + non metal
-form simple molecular covalent structures
-e.g. halogen + hydrogen –> hydrogen halides
-fluorine is the most reactive (reacts with hydrogen at low temperatures in absence of light)
Halogen displacement reactions
-more reactive halogen displaces a less reactive halogen from an aqueous solution of its halide
Chlorine, bromine, iodine - displacement reaction
Chlorine > bromine > iodine
Potassium bromide + chlorine
-chlorine displaces bromide ions
-yellow - orange colour of bromine seen
Potassium iodide + chlorine
-chlorine displaces iodide ions
-brown colour of iodine seen
Potassium iodide + bromine
-bromine displaces iodide ions
-brown colour of iodine is seen
Group 7 metals - reactivity in terms of electrons
-decreases
-if a shell is closer to the nucleus, it is more attracted to the nucleus - stronger tendency to form a 1- ion
-moving down the group, forces of attraction between the nucleus and outermost shell decreases
-harder for atoms to gain electrons
Composition of air
Oxygen - 21%
Nitrogen - 78%
Argon - 0.9%
Carbon dioxide - 0.04
Finding percentage of oxygen - using metals
1.Place wet iron filings inside the end of a burette
2. Using a clamp, stand the burette vertically over a trough of water.
3. Record the starting height of the water in the burette.
4. Leave for a few weeks then record the final height of the water in the burette.
5. Calculate the change in height of the water in the burette. This is the volume of oxygen that was originally in the burette.
6. To calculate the percentage by volume of oxygen in air, divide the change in the burette reading by the original volume of air in the burette and multiply by 100
-water level will rise as iron filings react with oxygen in the burette. Water rises to replace oxygen that has reacted
Finding percentage of oxygen - using non metals
- Place phosphorus on an evaporating dish, float dish in trough of water
- Ignite phosphorus, place bell jar into trough, cover dish
- Record starting height of water level in the bell jar
- Leave apparatus until phosphurus is extinguished
- Measure final water level
- Final - initial = volume of oxygen originally in the jar
- (Change in water level/ original volume) x 100
-combustion of phosphorus uses up oxygen, water rises to replace volume of oxygen used up
Combustion
-burning
-reactions involve a chemical change in which oxygen reacts with elements or compounds to produce oxides
-gives out heat - exothermic
Combustion reaction - magnesium
2Mg + O2 –> 2MgO
-intense white flame
-white powder produced (magnesium oxide)
Combustion - Sulfur reaction
S + O2 –> SO2
-blue flame
-colourless
-poisonous gas produced
Thermal decomposition
-reactions where a substance breaks down due to the action of heat
-one such reaction - thermal decomposition of metal carbonates
Thermal decomposition of metal carbonates
-carbonates of metals from the lower half of the reactivity series tend to decompose on heating
-metal carbonate –> metal oxide + carbon dioxide
-e.g copper (II) carbonate occurs readily on heating
-Copper(II) carbonate - green powder
-slowly darkens as black copper(II) oxide is produced
-CuCO3 –> CuO + CO2
The greenhouse effect
-When shortwave radiation from the sun strikes the Earth’s surface it is absorbed and re-emitted from the surface of the Earth as infrared radiation
-Much of the radiation is trapped inside the Earth’s atmosphere by greenhouse gases which can absorb and store the energy
-Carbon dioxide, methane and water vapour are gases that have this effect
Carbon dioxide - greenhouse effect
-traps extra heat
-enhanced greenhouse effect
-sources: burning wood, fossil fuels, respiration
Reactivity series
Please Stop Lying, Calling Me A Crazy Zebra, Instead Try Learning How Copper Saves Gold
-Potassium
-Sodium
-Lithium
-Calcium
-Magnesium
-Aluminum
-Carbon
-Zinc
-Iron
-Tin
-Lead
-Hydrogen
-Copper
-Silver
-Gold
Reactivity series - + water
metal + water –> metal hydroxide + hydrogen
-potassium: reacts violently
-Sodium: reacts quickly
-Calcium: reacts less quickly
-Mg, Fe, Zn + cold water - slow
Reactivity series - + dilute sulfuric/ hydrochloric acids
-only metals ABOVE hydrogen will react with dilute acids
-more reactive the metal, more vigorous the reaction will be
metal + acid –> salt + hydrogen
Reactivity series - metal displacement reactions (zinc + copper oxide)
-more reactive metal will displace a less reactive metal from its compounds
e.g. zinc + copper oxide –> zinc oxide + copper
Reactivity series - displacement reactions between metals and aqueous solutions
-more reactive metal slowly disappears from the solution, displacing the less reactive metal
-e.g. Mg + CuSO4 –> MgSO4 + Cu
-blue colour of CuSO 4 solution fades as colourless magnesium sulfate solution is formed
-copper coats surface of the magnesium, forms solid metal which falls to the bottom of the beaker
Rusting
Chemical reaction between iron, water, oxygen to form hydrated iron (III) oxide
-redox process
-occurs faster in salty water due to the presence of sodium chloride
Rusting - Barrier methods
-coating iron with barriers that prevent the iron from coming into contact with water and oxygen
-if coatings are washed away/ scratched - iron will rust
-galvinising
-barrier method
-sacrificial protection
Galvanizing
Iron coated with a layer of zinc
-zinc is more reactive than iron, reacts with oxygen/ water more readily than iron
Sacrificial protection
-Zinc, Magnesium, Aluminium (more reactive than Fe) attached to metal hulls of ships to prevent iron/ steel from rusting
Barrier method
-coating iron with oil, grease, plastic, metal below it on reactivity series
Oxidation
-reaction where substance gains oxygen
-loss of electrons
Reduction
-reaction where substance loses oxygen
-gain of electrons
Redox
Reactions where both oxidation and reduction take place
Reducing agent
-is usually a metal or a negative ion loses (donates) electrons to another element or ion (reducing the other species)
-is itself oxidised
Oxidising agent
-is normally a non-metal or positive ion cause oxidation reactions to take place
-gains electrons from other atoms or ions (is itself reduced)
Practical: investigating reactions of dilute acids with metals
- add dilute hydrochloric acid to each of three test tubes
- Add magnesium ribbon to the first, iron filings to second, zinc turnings to third, observe
- lit splint test for any gases given off
- Record
- Repeat experiment with dilute sulfuric acid
Magnesium + dilute hydrochloric acid/dilute sulfuric acid
Dilute hydrochloric acid:
-Dissolves quickly
-gets hot
-hydrogen gas given off
-colourless solution left
Mg + 2HCl –> MgCl2 + H2
Dilute sulfuric acid:
-rapid bubbling
-hydrogen gas given off
-metal dissolves
Mg + H2SO4 –> MgSO4 + H2
Zinc + dilute hydrochloric acid/ dilute sulfuric acid
Dilute hydrochloric acid:
-bubbles given off
-metal slowly dissolves
Zn + 2HCl –> ZnCl2 + H2
dilute sulfuric acid:
-metal dissolves forming a colourless solution
-gas given off slowly
Zn + H2SO4 –> ZnSO4 + H2
Iron + dilute hydrochloric acid/ dilute sulphuric acid
Dilute Hydrochloric acid:
-very slow bubbling
Fe + H2SO4 –> FeSO4 + H2
Dilute sulphuric acid:
-slow reaction
-small bubbles seen
-Fe + H2SO4 –> FeSO4 + H2
Metal ore definition
A rock that contains enough of the metal to make it worthwhile extracting
Minerals definition
Compounds of metals and other elements
Sources of metals
Majority of metals on Earth are found combined with other elements in the crust
Some found uncombined - native elements
Native elements
Metals that exist naturally as the uncombined elements
-tend to be low on the reactivity series
-e.g. gold, platinum, silver, copper
Methods of extraction
-carbon extraction (roasting w carbon) - metals below carbon
-electrolysis - metals above carbon
Methods of extraction (metals above carbon) - iron oxide
Heat iron oxide in a blast furnace. Carbon and carbon monoxide are reducing agents
Fe2O3 (s) + 3C (s) –> 2Fe (l) + 3CO (g)
-iron reduces as it loses oxygen
-form iron metal
-carbon oxidised as it gains oxygen
-form carbon monoxide
Methods of extraction (metals below carbon) - aluminum oxide
Aluminum oxide dissolved in a molten salt (cryolite) - lower melting point of Al2O3
Al 3+ + 3e - → Al (cathode)
2O2- –> O2 + 4e- (anode)
Alloy definition
A mixutre of a metal with other metals or with carbon
e.g. brass (copper and zinc), bronze (copper and tin), steel (iron and carbon)
Why are alloys harder than pure metals
-metals in alloys are different sizes, so lattice arrangement is slightly disrupted
-more difficult for layers of ions to slide over one another
Properties and uses of aluminum
-low density, good conductor of electricity, resists corrosion
-planes, electricity cables, pots
-aluminum not very strong, aluminum alloys used for strength
Properties and uses of copper
-good conductor of electricity, unreactive, ductile, malleable, antimicrobial properties
-electrical wires, pots, pans, water pipes, surfaces in hospitals
Three types of steel
Alloy of carbon
-mild
-high-carbon
-stainless steel
Properties and uses of mild steel
-iron and up to 0.25% carbon
-nails, car bodies, ship building, girders
-soft, malleable
Properties and uses of high-carbon steel
-iron and 0.6-1.2% carbon
-cutting tools, masonry nails
-hard
Stainless steel
-Carbon and chromium/ nickel
-cutlery, sinks
-strong, resistant to corrosion
Litmus indicator colours
Acid: red
Alkali: blue
Phenolphthalein indicator colours
Acids: colourless
Alkali: pink
Methyl orange indicator colours
Acids: red
Alkali: yellow
The pH scale
<7 - acid
>7 - alkaline
-lower the pH - more acidic
-higher the pH - more alkaline
7 - neutral
Universal indicator
-wide range indicator, only approximate
-red 1(acidic) –> yellow –> green (7) –> blue –> purple 14 (alkaline)
Acids and water
-form positively charged hydrogen ions
-presence of H+ ions is what makes a solution acidic
Alkaline and water
-when alkalis are added to water, form negative hydroxide ions (OH–)
-presence of OH- ions is what makes an aqueous solution an alkali
Neutralisation reaction
-occurs when an acid reacts with an alkali
-H+ ions react with OH- to produce water
acid + alkali/ base –> salt + water
-hydrochloric acid produces chlorides
-sulfuric acid produces sulfate
-nitric acid produces nitrates
–metal oxides/ metal hydroxides act as bases
Acid alkali titrations
-method of analysing the concentratino of solutions
1. Using a pipette measure out 25cm3 of alkali or acid into a conical flash
2. Add a named indicator –> phenolphthalein (turns pink)
3. Fill burette with acid, note initial volume
4. Add avid slowly until the end point is reached - dropwise near the endpoint - swirl regularly to ensure mixing
5. End point indicated by colour change (pink –> colourless)
6. Note end volume and work out how much has been added
7. Repeat until you have concordant results - within 0.2cm3
Solubility rules
-all nitrates are soluble
-common sodium, potassium, ammonium compounds are solubles
-chlorides are soluble, except lead and silver
-sulfates are soluble, except barium, calcium, lead
-carbonates are insoluble, except for sodium, potassium, ammonium
-hydroxides are insoluble, except sodium, potassium, calcium (calcium hydroxide is slightly soluble)
Acids & bases - proton transfer
Acids - proton donors as they ionize in solution, producing protons H+ ions
-these H+ ions make the aqueous solution acidic
Bases - proton acceptors as they ionize in solution, producing OH- ions which can accept protons
-these OH- ions make the aqueous solution alkaline
Reactions of acids + metal carbonates
-form salt, carbon dioxide, water
-these reactions are easily distinguishable from acid/ metal oxide/ hydroxide reactions due to the presence of effervescence caused by the carbon dioxide gas
Bases vs alkalis
-bases are substances which can neutralise an acid, forming a salt and water
-bases - usually oxides, hydroxides, carbonates of metals
-bases - ammonia solution - produces hydroxide ions
-base which is water soluble - alkali
Prepare a soluble salt - acid + insoluble base
-insoluble reactant is added in excess to ensure that all of the acid has reacted
1. Add insoluble base to acid in a beaker bit by bit until in excess
2. Filter solution using filter paper and a funnel into evaporating dish
3. Heat the solution gently to evaporate off some water
4. Leave solution to cool and crystallise
-e.g. prepare Copper (II) Sulfate
-Hydrated copper (II) sulfate crystals should be bright blue, regularly shaped
Prepare a insoluble salt - from two soluble salts
- Pour two soluble salts together. Precipitate forms
- Filter solution through filter paper in a funnel
- Wash precipitate with deionised water
- Leave solid in an evaporating dish to dry
-e.g. magnesium sulfate and lead nitrate to make Lead (II) Sulfate
Prepare a soluble salt - acid and alkali
- Measure out a set amount of acid into a conical flash using a pipette. Add a few drops of indicator
- Slowly add alkali to the acid, using a burette, until you reach the end point - when acid has been exactly neutralised
- Carry out reaction with exactly the same volumes of alkali and acid but with no indicator
- Solution that remains when the reaction is complete contains only the salt and water
- Slowly evaporate off some of the water and leave solution to crystallise
- Left with a pure, dry salt
Ammonia test
Turns damp red litmus paper blue
-colourless, pungent smell
Carbon dioxide test
Bubble through limewater, turns cloudy
-colourless and odourless
Chlorine test
Bleaches damp blue litmus paper
-pale green, choking smell
Hydrogen test
Hold a lit splint in mouth of test tube, burns with a squeaky pop sound
-colourless and odourless
Oxygen test
Relights a glowing splint
-colourless and odourless
Flame test
-dip a clean wire loop into a solid sample of the compound being tested.
-put the loop into the edge of the blue flame from a Bunsen burner.
-observe and record the flame colour produced
Test for cations
Flame test - Li+, Na+, K+ Ca2+, Cu2+
Red
Yellow
Lilac
Red-orange
Blue-green
Test for cations - Cu2+, Fe2+, Fe3+
Colour of precipitate formed when sodium hydroxide added
-light blue
-green
-red brown
Tests for anions - Sulfate SO4 2-
-acidify with dilute nitric acid
-add aqueous barium nitrate
-white precipitate formed
Test for anions - carbonate CO3 2-
-Add dilute acid
-test the gas released
-turns limewater milky
Test for anions - chloride Cl-, bromide Br-, iodide I-
-acidify with dilute nitric acid
-add aqueous silver nitrate
-white
-cream
-yellow
Test for water - chemical
Anhydrous copper(II) sulfate turns from white to blue on the addition of water
Test for water - physical
Check boiling point - 100C
-any impurities tend to raise boiling point, depress melting point of pure substance
