Chemistry 1 - atomic structure and the periodic table Flashcards
Isotopes
Isotopes are different forms of the same element, which have:
The same number of protons
Different number of neutrons
Same atomic number but different mass numbers
Relative atomic mass (Ar)= Sum of( isotope abundance * isotope mass number) / sum of abundance’s of all the atoms
Example:
Copper has two stable isotopes. Cu-63 has an abundance of 69.2% and Cu-65 has an abundance of 30.8%. Calculate the relative atomic mass of copper to 1 decimal place.
(69.263) + (30.865) / 69.2 + 30.8
= 4359.6+2002 / 100
= 63.616
=63.6
Paper Chromatography - Practical
1) Draw a line near the bottom of a sheet of filter paper (using pencil as it is insoluble)
2) add a spot of ink to the line and place the sheet in a beaker of solvent e.g. water
3) The solvent used depends on what’s being tested. Some compounds dissolve well in water but sometimes other solvents like ethanol are needed
4) make sure the ink isn’t touching the solvent( as it will dissolve)
5) place a lid on top of the container to stop the solvent evaporating
6) the solvent seeps up the paper carrying the ink with it
7) Each different dye in the ink will separate out . Each will form a spot in a different place ( 1 spot per dye in the ink)
8) if any of the dyes in the ink are insoluble they will stay on the base line ( pencil line)
9) when the solvent has nearly reached the top take the filter paper out and leave it to dry
10) end result is called a chromatogram
Separation Techniques
Filtration - separates insoluble solids from liquids
Evaporation - separate soluble solids from solutions
Crystallisation - separate soluble solids from solutions if the salt doesn’t decompose when it’s heated
Filtration - Filter paper folded into a cone shape - the solid is left in the filter paper
Evaporation -
1) pour the solution into an evaporating dish.
2) slowly heat the the solution.(the solvent will evaporate and the solution will become more concentrated forming crystals)
3) keep heating the evaporating until all you have left is crystals
Crystallisation -
1) pour the solution into an evaporating dish and gently heat the solution. Some of the solvent will evaporate and the solution will get more concentrated.
2) once some of the solvent has evaporated, or when you start to see crystals forming ( the point of crystallisation), remove the dish from heat and leave the solution to cool.
3) the salt should start to form crystals as it becomes insoluble in the cold, highly concentrated solution.
4) filter the crystals out of the solution, and leave them in a warm place to dry.
Separating Rock Salt
Using Filtration and Evaporation/Crystallisation
1) grind the mixture to make sure salt crystals are small ( so will dissolve easily)
2) put the mixture in water and stir ( the salt will dissolve but the sand won’t)
3) filter the mixture. The grains of sand won’t fit through the tiny holes of the filter paper, the salt will as it’s a part of the solution
4) evaporate or crystallise the water from the salt so that it forms dry crystals
Simple Distillation - separate out solutions
E.g separating pure water from sea water
1) simple distillation is used for separating out a liquid from a solution.
2) the solution is heated. The part of the solution that has the lowest boiling point evaporates first.
3) the vapour is then cooled, condenses and is collected
4) the rest of the solution is left behind in the flask
Fractional Distillation - separate a mixture of liquids
E.g crude oil
1) you put your mixture in a flask and stick a fractionating column on top. Then you heat it.
2) The different liquids will all have different boiling points - so they will evaporate at different temperatures.
3) the liquid with the lowest boiling point evaporates first. When the temperature on the thermometer matches the boiling point of this liquid it will reach the top of the column
4) liquids with higher boiling point may also start to evaporate but the column is cooler towards the top so they will condense back down.
5) when the first liquid has been collected you raise the temperature until the next one reaches the top
History of the atom
John Dalton (19th century) - describes atoms as solid spheres made up of different elements
JJ Thompson (1897) - experiments of charge and mass showed that atoms must contain smaller negatively charged particles (electrons). Called Plum Pudding model ( ball of positive charge with electrons stuck in it)
Ernest Rutherford(1909) - conducted alpha particle scattering experiment ( fired positively charged alpha particles at a thin sheet of gold.
Some were deflected more than expected and some were deflected backwards disproving the plum pudding model.
He then invented nuclear model:
Tiny positively charged nucleus at centre ( most of mass concentrated)
A cloud of negative electrons surrounding the nucleus.
Niels Bohr - electrons in a cloud around the nucleus would be attracted to it causing the atom to collapse
Proposed electrons orbit the nucleus in fixed shells at fixed distances from the nucleus
Development of the periodic table
Early 1800s: elements were arranged by atomic mass ( only thing they could measure)
Dmitri Mendeleev(1869): put elements mainly in order of atomic mass but switched the order of the properties meant it should be changed Gaps were left in the table to make sure that elements with similar properties stayed in the same group. He predicted the existence of undiscovered elements.
Transition metals - centre of the periodic table
Typical metals : good conductors of heat and electricity and they are dense, strong and shiny
Transition metals can have more than one ion e.g. Cu^+ , Cu^2+
Group 1 (alkali metals)
Reactive soft metals:
All have one electron in outer shell so they are very reactive
As you go down group 1: Increasing reactivity( outer electron is lost easier as it it further away from the nucleus) Lower melting & boiling points Higher relative atomic mass
Form ionic compounds with non-metals:
Reaction with water:
React vigorously with water to produce hydrogen gas and metal hydroxides ( salts that dissolve in water to produce alkaline solutions)
Reaction with Chlorine:
React vigorously when heated in chlorine gas to form white metal chloride salts.
Reaction with oxygen:
React with oxygen to form a metal oxide
Different oxides form:
Lithium > Lithium oxide Li(2)0
Sodium > Sodium oxide Na(2)0 and Sodium peroxide Na(2)0(2)
Potassium > Potassium peroxide K(2)0(2) and Potassium superoxide (KO(2))
Group 7 - Halogens
Non metals with coloured vapours
Fluorine - very reactive, poisonous yellow gas
Chlorine - fairly reactive, poisonous dense green gas
Bromine - dense, poisonous red-brown volatile liquid
Iodine - dark grey crystalline solid or purple vapour
All exist as pairs of atoms
Trends as you go down: Less reactive ( harder to gain an extra electron as outer shell is further away from nucleus) Have higher melting and boiling points Have higher relative atomic masses
Halogens form molecular compounds with other non-metals by covalent bonding
Form ionic bonds with metals
More reactive halogens will displace less reactive ones
Group 0 - Noble gases
All have full outer shells meaning they don’t react with much at all.
Exist as monoatomic - single atoms
Trends as you go down:
Boiling points increase ( increase number of electrons leading to greater intermolecular forces)
