GSCE Chemistry Flashcards
describe two advantages of instuctmentla methods of anaylsius
more accurate, gives accuratie readings and measures accurately/gives good results, no ambiguity, set to internaational standards
very sensiitive, can derect evne the smalles tamount/quanitties of substance/can asnslydd very small amounts of substances
faster/can carry out analysis all the time/speed, instructmenets can carry out quickly, robotic arms
mass spectrum
peak furthest to the right (the molecular ion peak) is equal to the Mr value of something - tells us the mass of the whole molecule
molecules form unique fragmentation patterns. computer has a database of these patterns and can tell the chemist exactly what the chemical is
might be more peaks as molecule gets placed in machine and is deflected, it can break/fragment- other peaks represent a fragment of a molecule
explain why fractions are cracked
smaller- higher demand, less supply
larger- less demand, higher supply
cracking turns larger into smaller
explain why the cracking of decane makes a mixture of products
any carbon carbon bond can be broken in the molecule and it isn’t specific
flame test: yellow-orange flame
sodium
sulfate ion
SO 4 ^(2-)
the concentration of hydrogen ions decreases as pH go from 1 to 2 by
a factor of 10
pH decrease is
From pH 14 to pH 1
pH increase is from pH 1 to pH 14
a student is testing sodium carbonate solution
she adds barium chloride followed by excess dilute hydrochloric acid
which of these observations would not be seen
white precipitate formed when dilute acid is added
ammonia chemical test
turns moist red litmus paper blue
chemical tests are used to identify gases, anions and cations
Leila has an unknown solution
She thinks that the solution contains copper (II) ions and bromide ions.
Describe the chemical tests she does to confirm the presence of these two ions in the solution
aqueous sodium hydroxide - blue precipitate - copper ions
aqueous silver nitrate followed by dilute nitric acid - cream precipitate - bromide ions
Solid
Very close
Regular arrangement
Energy of particles low
Fixed volume+shape
Cannot flow, only vibrate in fixed position
Cannot be easily compressed (particles are very close together with no space to move into)
Liquid
Close
Randomly arranged
Slide past and move around eachother
Can flow and takes shape of their container
Greater energy than solid particles
Fixed volume, no fixed shape
Cannot be easily compressed (particles close together with no space to move into)
Gas
Far apart
Randomly arranged
Move quickly in all directions
Highest energy
No fixed shape+volume
Can flow, completely fills their container
Can be compressed (particles are far apart with space to move into)
Change of state
energy transferred to substance: particles gain energy, overcome forces of attraction between particles
energy transferred from substance to environment: particles lose energy, forces of attraction form between them
Some overcome during melting, remaining forces overcome during evaporating, although some weak forces still remain between particles in gas state
Gas to solid - deposition/desublimation
Solid to gas- sublimation
Difference between physical and chemical changes
Chemical- require chemical reaction + chemical difference from reactants to products( particles differently joined and arranged). Normally irreversible/not easily. New substances made. Colour change, precipitate, gas, door change, temperature change.
E.g combustion, methane burning in air=CO2+H2O
Physical- change - requires energy and involves change in state (arrangement, movement, distance) but no change to particles themselves. Easily reversed, no new substances made. Change of state. E.g ice melt to water
Chemical and physical properties
Physical properties- can be measured or observed without changing substance’s chemical composition. E.g density, hardness, colour, melting/boiling
Chemical- ability of a substance to undergo a specific chemical change and change into a different substance. E.g flammability, corrosion/oxidation
Limitations of particle model (theory used to explain physical properties of solids,liquids, gases, describes arrangement and movement of particles in substance)
Does not show space between particles, forces of attraction between particles, particle size.
Presents particles as solid, spherical, i elastic spheres
John Dalton theory
John dalton 1803-
Studied pressure of gases - Gases consists of tiny particles in constant motion
atoms of same element are identical, atoms of different elements are different, atoms cannot be created/destroyed/broken down into anything smaller, atoms rearrange during chemical reactions to make new substances
J.J Thomson 1897
Discovered electrons
Used a cathode-ray tube to conduct an experiment which showed that there are small particles inside atoms
Disproved Dalton that atoms cannot be broken down into anything smaller.
Plum pudding model
Spherical mass of positive charge with electrons scattered inside
Democritus
Greek philosopher Democritus 450 BC
All matter was made up of tiny,Indivisible,solid atoms
Ernest Rutherford
The Geiger-marsden experiment (1909-1911)
Aimed beams of positive charged alpha particles at thin gold foil sheet. Some passed, deflected, bounced back
Disproved plum pudding model (they should’ve all gone straight through according to it)
Nuclear model - atoms have tiny central positive charged nucleus with most of the mass, rest of atom is empty space, electrons orbit the nucleus
Niels Bohr
Using mathematical ideas, improved Rutherford’s model
Bohr’s model 1913
electrons orbit in electron shells/energy levels around nucleus
Why does the atomic model change over time
New discoveries and experiments made
Experiment results disprove old model
new model explains the new evidence
Structure of the atom
Positively charged nucleus which contains subatomic particles: positive protons and neutral neutrons
Surrounded by negatively charged electrons orbiting in electron shells
Nucleus contains most of the mass
The radius of the nucleus is much smaller than the radius of atom
Sizes and relative sizes of atom and molecule
Atom= 1x10^-10 m = 0.1nm Diameter/radius/bond length size ——————————— Small molecule 1x10^-9 m
Protons, electrons and neutrons
Relative mass=
proton:1, neutron:1, electron:0.0005
Relative charge=
proton:+1, neutron:0, electron:-1
Define atomic number and mass number
Atomic number (at bottom)=number of protons Mass number (at top): total number of protons and neutrons
Define ion and isotope
Isotopes- atoms of same element with same number of protons but different number of neutron
Ion- electrically charged particles (can be positive or negative) when atoms lose or gain electrons
Mixture, compound, element
Mixture=two or more different substances not chemically joined together
Element- only one type of atoms with same atomic number
Compound- two or more elements chemically joined together
Purity of a substance
Chemistry: Consists only of one element/compound
Everyday language: substance that has nothing added to it (natural state and unadulterated)
melting point- distinguish pure from impure substances
Pure- sharp specific melting point
Impure/mixtures - melt over a range of temperatures
What the relative formula mass (Mr) of a compound is
Relative formula mass= sun of relative atomic masses
Many useful materials are formulations of mixtures
A formulation = mixture that has been designed as a useful product where each chemical in it has a particular purpose
Mixed in carefully measured quantities to ensure product has required properties
E.g medicines, perfumes, paints
Filtration process
Insoluble solid from a liquid
Beaker with mixture, beaker with funnel and filter paper
Pour mixture in filter funnel, liquid drips through (filtrate) but solid particles caught in filter paper (residue)
E.g sand from water
Filter paper
Has tiny pores in it which are large enough to let simple molecules, smaller liquid molecules and dissolved ions through but not large enough for undissolved solid particles
Crystallisation process
Solid crystals from solution
Solution placed in evaporating basin and heated with a Bunsen burner/electric heater. Solvent evaporates, solid crystals begin to form. When all water evaporated, leave the solid crystals to air-dry.
E.g copper surface crystals from solution
To obtain large, regularly shaped crystals
Put solution in evaporating basin and warm it by placing it over boiling water bath.
Stop heating before all the solvent has evaporated
Wait for remaining solution to cool
Pour excess liquid away/filter
Drug crystals using warm oven/air-dry
Simple distillation
Separate solvent from solution
Works cuz solute has higher boiling point than solvent.
Solution heated with Bunsen burner/electric heater, solvent vapour evaporates rising up, cools in condenser and condenses dripping into beaker.
E.g water from salt solution
Fractional distillation
Separate different liquids from mixture of liquids
Works cuz different liquids have different boiling points
Mixture heated. Vapour rise through a column (which is hot at bottom, cold at top). Vapour condenses when reach part of column that is below temperature of their boiling point and flows out of column.
Two ways of obtaining different liquids from the column
- Collect different liquids from different parts of column.Lowest boiling point collected at the top of column. E.g crude oil distillation
- heat mixture to increase temperatures in column. Liquid with lowest boiling point collected first
Chromatography
Separate mixtures of soluble substances
Stationary phrase, Mobile phrase
different dissolved substances attracted to phrases in different proportions- causes them to move at different rates
More soluble/more attracted to mobile phrase= travels up more
Less soluble/more attracted to stationary phrase=travels up less
Paper chromatography
separates different pigments in coloured soluble substance
Stationary phrase=paper
Mobile phase=solvent
Paper lowered into solvent and it spreads up paper
Analytical technique separating compounds by their relative speeds in a solvent as it spreads through paper
Thin-layer chromatography
Separate non-volatile mixtures
Stationary phrase=thin layer of inert substance supported on flat interactive surface e.g glass/plastic/aluminium foil
Mobile phase= solvent e.g silica gel/cellulose/aluminium oxide
UV light to see transparent, done in short time
Interpret chromatograms
Pure substances produce one spot on chromatogram
Impure substances produces two or more spots
Rf values calculations
Rf= distance travelled by substance/distance travelled by solvent
Gas chromatography
Mobile phrase= inert carrier gas e.g nitrogen/helium
Stationary phrase= thin layer of unreactive liquid/solid on a solid support(e.g silica beads), lining walls of long coiled column inside thermostatically controlled oven
Detector measures amount of each substance in a mixture as it leaves column
Different substances travel at different speeds through column and leave at different times (retention times)
GC machine plots graph of detector reading against retention time
Solubility dependent on how soluble a substance is in the gas
main info gathered from chromatogram: number of compounds in mixture (number of peaks), amount of each compound present (height of peak)
Metals and non metals
Metals= elements that lose electrons to form positive ions
Shiny, good electrical/heat conductors, high dense, high melting point, malleable, forms basic oxides
Non-metal= elements that gain electrons to form negative ions
Dull, poor electrical conductor, low dense, low melting point, brittle, forms acidic oxides
Position of an element in modern periodic table
Elements arranged in order of atomic number
Elements with similar chemical properties are in columns (groups)
Elements in same group have same amount of electrons in their outer shell
Group - number of outer shell electrons
Period- number of shells
Ionic compounds
Ionic bonding
Regular, repeating arrangement
giant ionic lattice
Metal and non-metal(Cation and anion)
Strong electrostatic forces of attraction between oppositely charged ions
Many strong bonds
Melted/dissolved to free ions so they can move and curry current
Simple molecules
Weak intermolecular forces between the molecules
Intermolecular forces increase with size of molecules so larger molecules have higher boiling+melting points
Strong covalent bonds between atoms
Giant covalent structures
Many strong covalent bonds
Electrostatic attraction between a shared pair of electron
Giant covalent lattice
Polymers
covalent
Very large molecules
Relatively strong Intermolecular forces between polymer molecules
Chain of atoms
Metals
Metallic bonding
Giant metal lattice
Strong electrostatic force attraction between positive metal ions and sea of delocalised electrons(which are free to move through whole structure)
When force applied, Layers of metal ions can slide over each other while still being attracted to the sea of delocalised electrons
Limitations of dot and cross diagrams, ball and stick models and two/three dimensional models
Chemical formula model- does not show charges, bonds, shape
Dot and cross diagram- shows how atoms bonded and electrons, does not show 3d arrangement
2d diagram- does not show 3d arrangement
3d diagram- no spaces between ions, charges
How Mendeleev’s arrangement was refined into the modern periodic table
Order of increasing relative atomic mass across rows
All elements in column had similar chemical properties
Left gaps for elements not discovered yet
Swapped the order of a few pairs of elements
Vast array of natural and synthetic organic compounds occur due to the
due to the ability of carbon to form families of similar compounds, chains and rings as each carbon atom can form 4 covalent bonds
Diamond
Each carbon is covalently bonded to 4 other carbons
No free electrons
Very hard, high melting point, does not conduct electricity
Rigid tetrahedral structure held together by many strong covalent bonds - hardness
Oil rig drills, cutting tools, tipped glass cutters
Regular lattice arrangement
Graphite
Each carbon is covalently bonded to 3 other carbons
Layers of hexagonal rings with weak forces between the layers
Layers can slide over eachother
Soft and slippery
One electron from each carbon atom is delocalised so electron can carry charge
Fullerene
Fullerenes - Large molecules of carbon with hollow shapes. different forms/allotrope of carbon with structure based hexagonal rings of carbon atoms (although the rings could be 5 or 7 too)
Buckminsterfullerene C60-spheres of carbon atoms which are made of large molecules. Weak intermolecular forces exist between each buckyballs. Slippery. Lower melting points. Drug delivery. Lubricants. Catalysts.
Carbon nanotubes- layer of graphene rolled into a tube shape. High tensile strength so can be stretched without breaking. Strong. Conduct heat:electricity. High length to diameter ratios. Nanotechnology, electronics, specialised materials (e.g reinforce materials in tennis rackets)
Allotrope
Two more different forms of the same element
E.g graphite and diamond are allotropes of carbon
Graphene
single layer of graphite. Very strong. High melting point. Large regular arrangement of carbon atoms joined by covalent bonds. Delocalised electrons free to move through structure. Giant covalent structure.
Drug delivery, electric circuits, composites, solar cells
Bulk properties
Individual atoms do not have the physical properties of the substances that contain them
Bulk properties= properties due to the influence of many atoms/ions/molecules acting together
Nanoparticles
1-100nm
Few hundred atoms
Nanoparticles properties and uses
Large surface area:volume ratio, catalysts
Small size, paints, cosmetics, transparent sunscreens, self-cleaning window panes
Conduct electricity, small electrical circuits for computers
Lubricant coatings, reduce friction, artificial joints and gears
Risks of Nanoparticles
Possible to breathe in or pass into cells
Might catalyse reactions harmful
Toxic substances could bind to Nanoparticles as large S:V ratio
or pass into bloodstream and reach the brain
Risks unknown as Nanoparticles are new (still being created and tested)
Group 1 alkali metals
Soft, relatively low melting points
Relatively low densities - first 3 can float on water
All have similar chemical properties - as 1 electron in outer shell. +water= metal hydroxide+H2. Metal hydroxide= Base that dissolves in water to form alkaline solution
Li- fizzes steadily, slowly becomes smaller until all reacted
Na- melts to form ball, fizzes rapidly, quickly becomes smaller until disappears
K- quickly melts to form a ball, burns violently w sparks and lilac flame, reacts rapidly with small explosion
Rb- melt very quickly, burn very violently, disappears almost instantly with an explosion
Group 7 halogens
Exist as simple molecules. Each molecule contains two halogen atoms joined by single covalent bond. Cl- pale green gas, Br- brown liquid, I- purple black solid
All have similar chemical properties - as all have 7 electrons in outer shell. +metals=salts. F- cold iron wool burns to produce white iron III fluoride, Cl- hot iron wool burns vigorously- orange-brown, Br- hot iron wool burns quickly- red-brown, I-hot iron wool reacts slowly in iodine vapour to produce grey iron iodide. hydrogen+=hydrogen halides (gas at room temp, dissolve in water to produce acidic solutions): F- explodes in cold and dark, Cl- explodes in flame/sunlight, Br-vigorous;
~needs burning hydrogen, I- very slow when heated strongly, forms some hydrogen iodide. At- react very slowly even when heated, little forms
Group 0- noble gases
Exist as single atoms
Low boiling points and densities
Inert
Reactivity of metals
- more- vigorous- easily loses electrons to form cations
- metal+water~>metal hydroxide+hydrogen
- Al=reactive but surface naturally forms thin layer aluminium oxide that keeps water away from metal below)
- Mg slowly reacts first added to water but layer of insoluble MgOH forms, protects metal/stops it from reacting
- but if steam passes over hot Mg, vigorous reaction happens to form MgO+H2
- metal+dilute acid-> salt+hydrogen -metal below hydrogen in RS will not react with dilute acids
Spectator ions
Ions that don’t take part in the reaction
Test for gases
- oxygen: glowing splint relight, as supports combustion
- hydrogen: lighted splint ignited with squeaky pop as H ignites in air
- carbon dioxide: when bubbled through limewater (CaOH solution), it turns cloudy white as CO2 reacts with CaOH solution to form white precipitate
- chlorine: red damp litmus paper bleached as it’s acidic gas that acts as a bleach. If blue litmus, turns red then bleach.
Flame test - metal ions
- Li+ Lithium Red
- Ca2+ Calcium Orange-Red
- Na+ Sodium Yellow
- Cu2+ Copper Green-Blue
- K+ Potassium Lilac
- clean w/ HCl/Nitric acid, rinse w/ deionised water. Dip clean nichrome wire loop into a sample of compound being tested. Put loop to edge of blue flame from Bunsen burner. Observe n record flame colour
Testing for Aqueous Metals Ions
- Dilute NaOH solution + metal ions -> metal hydroxide (but insoluble precipitate)
- Ca2+ and Zn2+ = white
- Fe2+ = green, Fe3+ orange-brown
- Cu2+ = blue
- add excess NaOH to differentiate Zn2+/Ca2+ = CaOH remains unchanged, ZnOH dissolves to colourless solution
Anions (testing for carbonate ions)
- Co3 2+
- add dilute HCl = bubbles given off, caused by carbon dioxide, bubbled through limewater to confirm its carbon dioxide
Anions (testing for sulfate ions)
- barium ions (Ba 2+) + sulfate ions (SO4 2-) -> insoluble white barium sulfate
- add dilute HCl to sample. Add dilute barium chloride solution. White precipitate
- acidified w/HCl to remove any carbonate ions present as it also w barium chloride to produce white precipitate
Halide ions testing
- Add dilute nitric acid to sample. Add dilute silver nitrate solution. Silver ions react with halide ions
- CBI, white cream yellow precipitate
- carbonate ions = white w silver nitrate solution so nitric acid reacts w it to remove them
Identifying ions in unknown salts
- dissolve sample of salt in a little of distilled water if solid salt and not salt solution
- very dilute solutions( of CBI in silver nitrate) makes it difficult to tell whether precipitate is a colour that is too pale
- harmful/toxic if inhaled/swallowed so only dilute solution used
Instrumental methods of analysis
Relies on machines
- extremely fast, quick speed
-sensitive, can detect the substance even in the smallest amounts of sample
-accurate, reliably identify elements and compounds
—————————————-
Very expensive, laboratory glassware is cheaper and more readily available
————
Examples: flame photometer, emission spectroscopy, gas chromatography
Flame photometer
Identify metal ions
Coloured light from vaporised sample is split to produce emission spectrum
Each metal ion produces unique emission spectrum, compare spectrum to known metal ion reference spectra and if match then same
Emission spectroscopy
Analyses spectrum from hot sample, records exact wavelength of light emitted by sample
matches pattern and wavelength to reference data for known elements
E.g identify elements in stars/substances in steel industry
Calibration curve
Graph with readings from machine plotted against known amounts of a substance
Reactivity series
Please- Potassium Stop- Sodium Calling- Calcium Me- Magnesium A- Aluminium Careless- (carbon) Zebra- Zinc Instead- Iodine Try- Tin Learning- Lead How- (hydrogen) Copper- Copper Saves- Silver Gold- Gold Platinum- Platinum
Transition metals
- form coloured compounds
- catalysts
- slow react w water/air/acid
- high melting point+density
Theoretical yield
Max possible mass of a product that can be made in chemical reaction
Actual yield
Mass of product actually obtained from chemical reaction
- incomplete reactions
- practical losses (reactants left on apparatus)
- unwanted side effects/competing reactions
Atom economy
Measure of how many reactant atoms form a desired product
-can be improved by finding a use for the other product, which makes it another desired product
PAG titration
- pipette and pipette filler to add 25cm3 of dilute sodium hydroxide solution to clean conical flask
- add few drops of phenolphthalein indicator and put conical flask on white tile
- fill burette with dilute HCl and note starting volume
- slowly add the acid form the burette to the conical flask, swirl to mix
- stop adding acid when end-point (pink to colourless) reached
- note final volume
- repeat steps until gotten concordat titres
How to obtain accurate results from titration
- vertical burette
- take readings from bottom of meniscus
- near the end-point, rinse the inside of flask w distilled water and add acid drop by drop
Why a pipette is use to measure the acid and not a measuring cylinder
Allows same volum of acid to be added each time, helps to make results repeatable
End-point and equivalence point
End-point: when indicator colour first permanently changes from pink to colourless
Equivalence: when chemical reaction in titration mixture ends, indicated by end-point
Conversions
1dm^3=1000cm^3
mol/dm^3 x Mr = g/dm^3
mol dm ^-3 = M= mol/dm^3
Rate of reaction
Measure of how quickly a reactant is used up/ a product is formed
Collision theory - for chemical reaction to happen, particles must collide with enough energy for them to react
Rate of reaction speeds
Slow: rusting, photosynthesis, fermentation
Fast: combustion, explosions, neutralisation, precipitation reactions
What does measuring rates of reaction/mass/volume depend on
Rate - reactants, products,easy to measure changes in them, length
Mass - not suitable for hydrogen/gases with small relative formula mass
Volume- gas syringe, measuring cylinder upside-down burette. cm3/s or cm3/min
Pressure on rate of reaction
Pressure (concentration) increase= reactant particles more crowded= collision frequency between particles increase = rate of reaction increases as more frequency of successful collisions
Lumps/ powders on rate of reaction
Total volume stays same, area of exposed surface increases so surface area: volume ratio increases. More reactant particles exposed at the surface= frequency of successful collisions between reactant particles increase
Catalyst
Substance that increases reaction rate but does not alter the products of the reaction and is unchanged chemically/mass at end of reaction
- allows alternative reaction pathway that has a lower activation energy than uncatalysed reaction
Enzymes
- biological catalyst
- controls reactions in cells
- lowers temp+pressure needed in some industrial reactions
e. g enzymes in yeast used to produce wine by fermenting sugars
PAG - colour change p, rate of reaction
- add 50cm3 of dilute sodium thiosulfate solution to a conical flask
- place conical fast on paper w/ black cross
- use different measuring cylinder to add 10cm3 of dilute HCl to conical flask
- immediately swirl flask to mix its contents, and start stop clock
- measure temp of reaction mixture, clean apparatus
- repeat with different starting temps of sodium thiosulfate solution
- 1000/time for each temp= proportional to reaction rate
Why same person should look at black cross each time
Different ppl may decide that they cannot see the cross at different amounts of cloudiness leading to errors in deciding reaction time
Copper sulfate
Hydrated copper (II) sulfate (blue). The copper ions in its crystal lattice structure are surrounded by water molecules. This water is driven off when it’s heated, leaving white anhydrous copper sulfate. Reversible. CuSO4.SH2O (s)
