Elements Of Life Flashcards
1
Q
Nucleus of the atom?
A
Most of the mass of the atom is concentrated in the nucleus.
The diameter of the nucleus is smaller compared to the whole atom.
The nucleus is where all the neutrons and protons are found.
2
Q
Masses + Charge Of Protons, Neutrons + Electrons?
A
Proton - 1 (Mass), +1 (Charge),
Neutron - 1 (Mass), 0 (Charge),
Electron - 1/2000 (Mass), -1 (Charge).
3
Q
Nuclear symbols?
A
This is basically what’s on the periodic table.
(But on a chemistry exam question, it may give you an example of a nuclear symbol and your expected to write it the same way as them).
- Mass number,
- Atomic number,
- Element number.
4
Q
Neutral, negative and positive atoms?
A
For neutral atoms, which have no overall charge, the amount of protons should be the same as the electrons.
Negative ions have more electrons than protons.
Positive ions have more protons than electrons.
E.g. Mg2+ is a positive ion because it has lost 2 electrons and therefore has 2 more protons than electrons.
5
Q
How has the current model of the atom developed?
A
- Ancient Greeks through that all matter was made from invisible particles.
- At the start of the 19th century, John Dalton described atoms as solid spheres, and said that different spheres made up different elements.
- Scientists did more experiments and our current model began to emerge our current model.
In 1897, J J Thomson did a series of experiments that concluded that atoms weren’t solid and invisible.
- He measured the charge and made and showed that an atom must contain smaller, negatively charged particles (electrons). He called electrons ‘corpuscles’.
- The idea of atoms being solid spheres. The new model was known as the ‘plum pudding model’ - positively charged sphere with negative electrons embedded everywhere in it.
6
Q
How was the plum pudding model proved to be wrong?
A
Ernest Rutherford and students - Hans Geiger and Ernest Marsden - 1909.
Conducted Geiger-Marsden experiment. They fired alpha particles (positively charged) at an extremely thin sheet of gold. It was expected that most of the particles would be deflected very slightly by the positive ‘pudding’ that made up most of the atom.
What actually happened? Most of the alpha particles passed straight through the gold atoms, and a very small number were deflected backward (more than 90 degrees).
This showed the plum pudding model couldn’t be right.
So Rutherford came up with a model that could explain this new evidence- the nuclear model of the atom.
Rutherfords model (the nuclear model) stated there is one positively charged nucleus and a cloud of freely orbiting negative electrons surrounded by empty space.
7
Q
How was Rutherfords nuclear model proved wrong?
A
Henry Moseley discovered that the charge of the nucleus increased from one element to another in units of one.
This led Rutherford to investigate the nucleus further he discovered protons in the nucleus. The charges of the nuclei of different atoms could then be explained - atoms of different elements have different number of protons.
There was still one problem with the model - the nuclei of atoms were heavier than they would be if they just contained protons - this lead to the discovery of neutrons by James Chadwick.
8
Q
The Bohr model?
A
Scientists realised that if the electrons were in a ‘cloud’, then they would spiral down into the nucleus, causing the atom to collapse.
Niels Bohr proposed a new model of the atom with four basic principles:
- Electrons can only exist in fixed orbits, or shells,
- Each shell has a fixed energy,
- When an electron moves between shells, electromagnetic radiation,
- Because the energy of shields is fixed, the radiation will have a fixed frequency.
The frequencies of radiation emitted and absorbed by atoms were already known from experiments. The Bohr model fitted these observations.
The Bohr model also explained why some elements (noble gases) are inert. It said that the shells of an atom can only hold fixed numbers of electrons, and that elements reactivity is due to electrons.
When an atom has full shells of electrons, it is stable and does not react.
9
Q
What model of atomic structure do we use today?
A
The Bohr model is widely used to describe atoms because it’s simple and explains may experiments, like bonding and ionisation energy trends.
The most accurate model we have today involves complicated quantum mechanics. Basically, we never know what direction an electron is going in or where it is. But, we can say how likely it is to be at a particular point.
Electrons can act as waves as well as particles.
This model might be more accurate, but it’s harder to get your head around and visualise. It does explain some observations that can’t be accounted for by the Bohr model. We use whatever model is most useful.
10
Q
Difference between empirical and molecular formula?
A
Empirical formula - gives the smallest whole number ratio of a time in a compound.
Molecular formula - gives to actual number of atom (ratio) in a molecule.
E.g. a molecule has the empirical formula C4H3O2. It’s molecular mass is 166. What is the molecular formula?
1. Work out Mr. Mr C4H3O2 = 83.
2. Molecular mass (so mass of atoms actually used in reaction) is 166. So do 166/83 = 2
3. Times everything in empirical formula by 2.
C8H6O4.
11
Q
Relative Atomic Mass?
A
Ar.
The average relative isotopic mass of naturally occurring isotopes of an element, whilst taking into account their abundances. Masses of atoms are compared to Carbon-12.
Not usually a whole number because it is an average.
E.g. 35Cl and 37Cl - isotopic masses are 35 and 37. Relative atomic mass is 35.5.
Ar values have no units.
12
Q
Calculation Ar?
A
Working out Ar:
You can do this normally by adding all the relative atomic mass values of the atoms. E.g. H2O = (1x2) + 16 = 18.
You can also do this using a mass spectrum diagram.
Times all the peak numbers (relative abundance) by the x-axis number that they align with. Then add them all.
Then divide this number by the total sum of all the peaks (which should be the total relative abundance).
Then times this all by 100.
13
Q
Relative Masses Can Be Measured Using?
A
Using a mass spectrometer.
Mass spectrometers can tell you: relative atomic mass, relative molecular mass, relative isotopic abundance, molecular structure, and your horoscope.
14
Q
How Does A Mass Spectrometer Work?
A
Steps Of How It Works:
1. Vaporisation - the sample is turned into a gas (vaporised) using an electrical heater.
- Ionisation - the gas particles are bombarded with high-energy electrons to ionise them. Electrons are knocked off the particles, leaving positive ions.
- Acceleration - the positive ions are accelerated by an electrical field.
- Detection - the time taken for the positive ions to reach the detector is measured. This depends on an ions mass and charge - light, highly charged ions will reach the detector first, whilst heavier ions with a smaller charge will take longer. For each sample analysed, a mass spectrum is produced.
15
Q
How To Read A Mass Spectrum?
A
Y-axis - gives the abundance of ions, often as a percentage. For an element, the height of each peak gives the relative isotopic abundance. E.g. 35Cl has abundance of 75.5%.
If the sample is an element, each line on the graph will represent a different isotope of the element.
The X-axis units are given as a mass/charge ratio. Since the charge of ions is mostly +1, you can often assume the x-axis as simply the relative mass.
16
Q
How Can Ar Be Worked Out From Mass Spectrum?
A
If relative abundance is given as a percentage:
1. For each peak, read the % relative isotopic abundance from the y-axis and the relative isotopic mass from the x-axis. Multiply these numbers together for each isotope.
- Add up these totals.
- Divide the total by 100 (because percentages were used).
If relative abundance is not given as a percentage:
1. For each peak, read the % relative isotopic abundance from the y-axis and the relative isotopic mass from the x-axis. Multiply these numbers together for each isotope.
- Add up these totals.
- Then, add up all of the relative abundance’s (the peak numbers) and add them together. Divide the total from step 1 by this number.
- (You can also work out relative isotopic abundance by timesing this number by 100 and giving the answer as a percentage).
17
Q
Ions?
A
Ions have different numbers of protons and electrons (this number is usually the same).
Negative Ions - have more electrons than protons.
Positive Ions - have fewer electrons than protons.
18
Q
Molar Mass?
A
The same as relative molecular mass.
19
Q
How to balance equations?
A
Same number of atoms on both sides.
Equations can also use 1/2 to make some equations balance.
20
Q
What Are Isotopes?
A
Atoms of the same element.
They have the same atomic number (number of protons - bottom number on periodic table).
They have different mass numbers (number of neutrons - top number on periodic table).
E.g. chlorine: Cl with 17 protons and 18 neutrons. Isotope of Cl with 17 protons and 20 neutrons.
21
Q
Relative Isotopic Mass?
A
Always a whole number. Relative isotopic mass is the mass of an atom of an isotope on a scale where an atom of Carbon-12 is 12.
E.g. 35Cl and 37Cl - isotopic masses are 35 and 37. Relative atomic mass is 35.5.
22
Q
Ionic equations?
A
Ionic equations only show the reacting ions in the reaction and the products that the reacting products form.
Ionic equations can only involve ions in solutions. So we don’t break down H2O into H+ and O- because H2O doesn’t dissolve in solutions.
- Balance the equation.
- Break up the equation into ions. E.g. HNO3 would split into H+ and NO3- (remember that H is always positive). We don’t break down anything that doesn’t dissolve in a solution, e.g. H2O.
- Cancel the ions that are repeated on both left and right side of the equations.
Once an ionic equation has been established, check the charges are balanced, as well as the atoms.
If the charges don’t balance, the equation is wrong.
E.g. ionic equation: H+ + OH- —> H2O
Net charge on left = +1 from H and -1 from OH so the net charge is 0. The net charge on the right side is also 0.
23
Q
Number Of Moles Equation?
A
Moles = Number Of Particles You Have / Avogadro’s Constant
24
Q
How Is Mol Calculated In Concentration?
A
Unit: mol dm-3.
Number Moles = Concentration x Volume (dm3)
If the volume is given in cm3, then we need to divide the volume by 100. This will give the units in dm3.
25
A hydrocarbon is burnt in excess oxygen. 4.4g or carbon dioxide and 1.8g of water is made. What is the empirical formula of the hydrocarbon?
A hydrocarbon has a H and an C.
We use n = mass/Mr.
H2O: 1.8/18 = 0.1
CO2: 4.4/44 = 0.1
Now, because there is 2 hydrogen atoms in H2O, we know that the original compound from which is was created should have 0.2 moles of hydrogen atoms (0.1 x 2).
Because there is 1 carbon atom in CO2, we know that the hydrocarbon must has 0.1 moles of carbon atoms (0.1 x 1).
This creates a ratio. C:H = 0.1 : 0.2.
This tells us the empirical formula should be CH2.
26
Relative Molecular Mass?
Also known as relative formula mass and molar mass.
Average mass of a molecule or formula unit on a scale where carbon-12 is 12.
E.g. C2H6O - (2x12) + (6x1) + 16 = 46
This also means that one mole of C2H6O weights 46g.
27
Aluminium reacts with oxygen to form aluminium oxide, Al2O3. Calculate the number of grams of Al2O3 that could be produced if 2.5g of aluminium and 2.5g of oxygen were allowed to react?
Write balanced equation.
4Al + O2 —> 2Al2O
M/Mr = n.
2. 5 / (Al Mr) = 0.09 n
2. 5 / (O Mr) = 0.07 n
Al2O Mr = 102
Use equation to find ratios.
4Al = 2Al2O
0.09 = 0.045. 0.045 x 102 = 4.59g of Al2O.
The ratio of moles of oxygen to Al2O is 3:2.
0. 07/3 = 0.02. 0.02 x 2 = 0.046.
0. 046 x 102 = 4.76g
This means that Al is the limiting reactant because it produces less grams of Al2O than O.
The amount of grams of Al2O that can be made is 4.59g because the reaction is limited by Al.
28
A compound has a percentage composition of 56.5% potassium, 8.7% carbon and 34.8% oxygen by mass. Calculate the empirical formula?
So imagine the percentages are grams.
K: 56.5/39.1 (mr) = 1.45 moles.
C: 8.7/ 12 = 0.725
O: 34.8/16 = 2.18
Divide each by the smallest.
K: 1.45/0.725 = 2
C: 0.725/0.725 = 1
O: 2.18/0.725 = 3
So the ratio is 2:1:3
Empirical formula: K2CO3
29
What is water of crystallisation?
Water molecules that are incorporated into another compound.
Hydrated compounds contain them and anhydrous compounds don’t.
One mole of a hydrated compound always contains the same Humber of moles of water of crystallisation (always a whole number too).
E.g. hydrated iron (III) chloride has six moles of water for every mole of iron chloride. So the formula is FeCl3 •6H2O
You beat a hydrous compound to make it anhydrous. If you know the mass of a compound before and after heating, you can work out it’s formula.
30
4.6g of alcohol, with molar mass 46g, is burnt in excess oxygen. It produces 8.8g of carbon dioxide and 5.4g of water. Calculate the empirical formula?
H2O: 5.4/18 (mr) = 0.3 moles
CO2: 8.8/44 = 0.2 moles
We can always find the amount of carbon and hydrogen atoms from the products.
There is 1 carbon atom in CO2 , so we know that there must be 0.2 moles of C in the original compound.
There is 2 hydrogen atoms in H2O, so we know that there must be 0.6 moles of hydrogen moles in the original comping (0.3 x 2).
We then need to work out the oxygen.
There’s 4.6g of the compound all together, so we do:
0.6 x 1 (H) = 0.6g
0.2 x 12 (C) = 2.4g
4.6 - 0.6 - 2.4 = 1.6g (this is the grams of oxygen).
So then, 1.6/16 (mr) = 0.1 moles
This can then be put as a ratio of C:H:O = 2:6:1
This is empirical formula: C2H6O
30
Heating 3.210g of hydrated MgSO4•XH2O forms 1.567g of anhydrous magnesium sulfate. Find the value of X and write the formula of the hydrated salt.
(You did this by yourself first time).
3. 210 - 1.567 = 1.643g (mass of water lost through heating).
1. 643 / 18 (mr) = 0.09
1. 567 / 120.4 (mr of MgSO4) = 0.013 (we use this mass and not the 3.210 because we’re finding the mass without the water in).
Divide both by smallest.
0. 09/0.013 = 7
0. 013 / 0.013 = 1
Ratio is 7:1
MgSO4•7H20
X = 7
31
Percentage yield equation?
Actual yield / theoretical yield x100%
Theoretical yield = mass of product that should be made in the reaction if no chemicals are lost in the process.
Actual yield = the actual mass of the product (always less). Reasons for this are: sometimes chemicals don’t fully react and some chemicals are lost (e.g. left on filter paper).
32
Ethanol can be oxidised to form ethanal: C2H5OH + (O) —> CH3CHO + H2O. 9.2g of ethanol was reached with an excess of oxidising agent and 2.1g of ethanal was produced. Calculate the percentage yield?
(You did this first time).
(O) just means an oxidising agent.
We know from the question that the actual yield is 2.1g cause that is what has been produced.
We need to work out theoretical yield.
9.2g of C2H5OH.
M / mr = n
9.2g / 46 = 0.2 moles
0.2 x 44 (mr of CH3CHO) = 8.8g so this is theoretical yield.
Use equation.
Percentage yield = actual /theoretical x100%
2.1 / 8.8 x100 = 24%
33
What is a standard solution?
Any solution that you know the concentration of.
How to make one:
E.g. make 250cm3 of a 6.00mol dm-3 solution of sodium chloride.
- Work out how many moles of solute you need. Moles = concentration x volume (/1000 to find dm3)
= 6.00 x 250 = 1.5mol
- Now work out how many grams of solute is needed.
Mass = moles x Mr = 1.5 x 58.5 (mr of sodium chloride).
(Practical bit)
Place beaker on scales and zero scales.
Weigh out this mass of solute.
Add distilled water to beaker until the solute has dissolved.
Tip solution into volumetric flask (should be an appropriate size). Use a funnel to make sure it all goes in.
Rinse the beaker and rod with distilled water and add that to the flask too. This makes sure there’s no solute clinging to beaker or rod.
Now top the flask up with the correct volume (250cm3) with more distilled water. Make sure bottom of the meniscus reaches the line - use a pipette the closer to the line you get).
Use a stopper on the bottle and mix it.
Check the meniscus again.
When making a standard solution, always make sure it s a suitable concentration compared to the solution your titrating it against. If it’s too dilute, you’ll have to add loads but if it’s too concentrated, the results can be inaccurate easy. A good rule: it should be a similar concentration to the solution it’s being titrated against.
34
Flame Tests?
Flame test is a test to detect metal ions.
Compounds of some metals give a characteristic colour when heated. This is the idea behind flame test. Here is how to do a flame test:
1) Dip a nichrome wire loop in concentrated hydraulic acid.
2) Then dip the wire loop into the same sample of the compound.
3) Hold the loop in the clear blue part of the Bunsen flame.
4) Observe the colour change in the flame.
You can use these colour changes to detect and identify different ions.
35
Flame Test Result For Lithium?
Lithium (Li+) turns a crimson colour.
36
Flame Test Result For Sodium?
Sodium (Na+) turns a orange-yellow colour.
37
Flame Test Result For Potassium?
Potassium (K+) turns a lilac.
39
Flame Test Result For Calcium?
Calcium (Ca2+) turns a brick red.
39
Flame Test Result For Barium?
Barium (Ba2+) turns a green.
41
Flame Test Result For Copper?
Copper (Cu2+) turns a blue-green.
41
Test For Hydroxide Metal Ions?
Many metal hydroxides are insoluble and precipitate out of solution when formed. Some of these hydroxides have a characteristic colour. So in this test you:
1) Add a few drops of sodium hydroxide solution to a solution of your mystery compound in the hope of forming an insoluble hydroxide.
2) If you get a coloured insoluble hydroxide, you can then identify the metal ion that was in the compound.
NaOH isn’t the only solution that can be used for precipitation reactions - you can use other hydroxide solutions too.
43
Identify Metal Ion From A Coloured Insoluble Hydroxide?
Silver (Ag+) - brown (precipitate of silver oxide),
Calcium (Ca2+) - white,
Lead (||) (Pb2+) - white,
Copper (Cu2+) - blue,
Iron (||) (Fe2+) - green,
Iron (|||) (Fe3+) - reddish brown,
Zinc (Zn2+) - white at first, redissolved in excess NaOH to form colourless solution,
Aluminium (Al3+) - white at first, redissolved I’m excess NaOH to form a colourless solution.
44
Testing For Carbonates Using Hydrochloric Acid?
Hydrochloric acid can help detect carbonates.
This is a test for carbonate ions (negative ions).
1) Take dilute Hydrochloric acid. Carbonates (CO3,2-) will fizz in hydraulic acid to give off carbon dioxide.
2) Capture the gas given off.
3) Test for carbon dioxide using lime water.
Carbon dioxide turns lime water cloudy. To do this: bubble the gas through a test tube of lime water and watch what happens. If the water goes cloudy, you've identified a carbonate ion.
44
Test For Sulfates With HCL?
You can test for sulphate with hydrochloric acid (HCL) and barium chloride (BaCl).
To identify a sulphate ion:
1) Add dilute HCL, followed by barium chloride solution.
2) If a white precipitate of barium sulphate forms, it means the original compound contains a sulphate.
(The hydrochloric acid is added to get rid of any traces of carbonate ions before you do the test. This would also produce a precipitate, so they’d confuse the results).
45
Litmus Paper Tests For Ammonia Gas?
Ammonia gas (NH3) is alkaline so you can check for it using a damp piece of red litmus paper. If there is ammonia present, the paper will turn blue.
(The litmus paper must be damp so the ammonia gas can dissolve and make the colour change).
46
Litmus Paper Test For Ammonium Ions?
You can also use the damp litmus paper test for when a substance contains ammonium ions (NH4+):
1) Add some sodium hydroxide to your mystery substance in a test tube and gently heat the mixture.
2) Use litmus paper in the gas that is given off. The litmus paper will turn from red to blue like the test for ammonia gas.
2) If there is ammonia given off, this means there are ammonium ions in your mystery substance.
47
Testing For Hydroxides Using Litmus Paper?
Litmus paper is also used to test for hydroxides. Testing for hydroxide ions is just the same as testing to see whether a solution is alkaline.
1) Dip a piece of red litmus paper into the solution.
2) If the hydroxide is present, the paper will turn blue.
(Hydroxides aren't the only alkaline substances so it's best to use as a test to eliminate the presence of other items first.)
48
Test For Halides Using Silver Nitrate Solution?
To test for:
- Chloride ions (Cl-),
- Bromide ions (Br-),
- Iodide ions (I-),
1) Add dilute nitric acid (HNO3),
2) Following this, add silver nitrate solution (AagNO3).
Results:
Chloride - white precipitate of silver chloride,
Bromide - cream precipitate of silver bromide,
Iodide - yellow precipitate of silver iodide.
49
Test For Nitrates With NaOH and Aluminium?
1) To identify nitrate ions, you need to warm the solution of your mystery compound.
2) Add sodium hydroxide solution and aluminium foil to your mystery compound solution.
( The aluminium reduces the nitrate ions to ammonium ions, NH4+ + OH- —> NH3 + H2O).
3) The ammonium ions then reacts with the hydroxide ions to produce ammonia gas and water.
4) So if ammonia is given off, you know your mystery solution must contain nitrate ions. You can test for ammonia gas in the same way as before - with litmus paper.
50
When Testing For Ions, How Do You Watch For False Positives?
If you're testing for sulfate ions, you want to make sure that there are no carbonate ions or sulfite ions first - because they will react with the barium chloride. As well as barium sulphate, barium carbonate and barium sulfite are also insoluble. So you will get a false positive and your results wouldn't mean much at all.
If you're testing for a halide ions, you want to rule out the presence of sulphate ions first. This is because sulphate ions will also produce a precipitate of silver nitrate. A way of getting round this is, is first add dilute acid to your test solutions. This will get rid of any lurking anions that you don't want.
Do your tests in this order to avoid mix-ups:
1) Test for carbonates,
2) Test for sulfates,
3) Test for halides.
52
How to write an electron configuration?
So the first row (helium and hydrogen) is 1s2. This is kinda odd cause helium is in the p-block of the periodic table but it actually has s-orbitals.
Sequence:
1s2, 2s2, 2p6, 3s2, 3p6, 3d10, 4s2, 4p6.
52
What do the numbers and letter in the electron configuration mean?
1 = This is the number of shells. These numbers are called principal quantum numbers.
S/P/D/F = this is the type of sub-shell. The highest energy sub-shell is the last one in the electron configuration. E.g. 4d2 is a higher electron configuration than 1s2.
S block = there is only one s orbital (s orbitals can only hold 2 electrons on outer shell).
P block = there is 3 p-orbitals in the p block (can only hold 6 electrons on outer shell - 2 in each).
D block = there is 5 d-orbitals in the d block (10 electrons can fit in here - 2 in each).
F block = there are 7 f-orbitals in the f block (14 electrons can fit - 2 in each).
The last little number = the amount of electrons on the outer sub-shell. E.g. 1s1 = there is one electron on the outer shell of the s-orbital.
Shells further from the nucleus (higher energy level) have a greater energy then those closer to the nucleus.
Remember: we fill up p-orbitals a bit differently to a orbitals. The fill up 1 in each, left to right. Then the second electron goes in, left to right again. S orbitals just fill up 2 in each, left to right.
54
What sub-shell has a lower energy level even thought it’s higher than other sub-shells?
The 4s sub-shell has a lower energy level than the 3d sub-shell even though it’s got a higher principal quantum number.
This means we fill the 4s sub-shell up before we fill the 3d sub-shell up. Confusing, I know.
So when writing configurations, always remember this.
The sequence should go:
1s2, 2s2, 2p6, 3s2, 3p6, 3d10, 4s2.
Remember this when working out ions too. This can get complicated.
55
How to draw diagrams of electron configurations?
This is the little squares we draw with each arrow in.
We can also draw shapes of s and p orbitals. This is page 30 in revision book.
56
How to work out the electron configuration of an ion?
This is easy.
E.g. “work out the electron configuration of Mg+1”
It’s +1 so it’s lost one electron.
Count back one square behind Mg. For Mg, this is on Na. Do the electron configuration for Na.
This is the same for all types of ion.
57
The highest energy outer shell configuration of thallium is similar to that of both boron and aluminium. Use ideas about electron configuration of thallium to discuss how thallium can form a 1+ ion rather than a 3+ ion in some of its compounds?
Boron - 1s2, 2s2, 2p1.
Aluminium - 1s2, 2s2, 2p6, 3s2, 3p1
Thallium - (axe) 6s2, 6d10, 6p1.
Thallium has a d-block and so it has d-orbitals. This means it can lose an electron to form a 1+ ion because it can be stable at this structure.
Boron and aluminium are not stable if they lose 1 electron because they do not have d-orbitals to support this structure.
58
How to change the concentration of a standard solution?
Divide the concentration you want by the concentration you have.
Then multiply this by the volume you want.
E.g. you have a concentration of 6 mol dm-3 and you want a 250cm3 of 3 mol dm-3 solution.
(3/6) x 250 = 125cm3
Carefully transfer out 125cm3 of the solution you have it put it into a volumetric flask. Then top up the volume with distilled water to volume you want.
59
What is a titration?
Allow you to understand exactly how much acid is needed to neutralise a quantity of alkali.
How:
Measure out alkali using a pipette (the long glass tube with a bubble in the middle) and put it in a flask (triangle glass) along with indicator, e.g. phenolphthalein.
Fill the burette with the acid. Make sure to take an initial reading of how much acid you have in the burette. Make sure to note if you’ve measured it from top or bottom of the meniscus.
Do a rough titration to get idea of when the end point is (where the alkali is exactly neutralised).
Do it three more times, staring from 2cm3 from your rough end point reading.
Swirl the flask whilst adding acid to alkali. Record the final reading from your burette when the alkali changes colour in the flask (its neutralised).
Work out the amount of acid used. Final reading - initial reading. The volume is known as the titre.
Your 3 readings should be within 0.1cm3 of each other. Calculate a mean, ignoring any anomalous results. Remember to wash the conical flask between each titration to remove any acid or alkali in.
60
Indicators that are appropriate for titrations?
They change colour to show pH.
Methyl orange - yellow at acid and red at alkali.
Phenolphthalein - turns red at acid and colourless at alkali.
Universal indicator is not appropriate here because the colour change is too gradual.
61
Types of bonding?
Bonding is when two elements join together.
Ionic and covalent.
62
Ionic bonding?
Ions are formed when electrons from two different atoms are transferred from one to another.
Electrostatic attractions are the forces that hold positive and negative ions together - this is called ionic bonding.
Generally, the charge on a metal ion is equal to its group number. The charge on a non-metal ion is equal to its group number -8.
Ionic bonds have giant ionic lattice structures.
63
Dot and cross diagrams?
To show ionic bonding - draw separate atoms that don’t touch.
Remember to draw a bracket around each circle drawing to show the overall charge. All the shells are drawn.
To show covalent bonding - draw circles that cross over each other in the middle. Only the outer shells are drawn. Most of the time, the atom ends up with eight electrons in its outer shell - its stable.
Show double bonds by placing 2 dot and cross pairs into the middle bit of the circles. Show triple bonds by placing 3 pairs in the middle.
64
Giant ionic lattice structure?
Ionic bonds have giant ionic lattice structures.
Ionic crystals are giant lattices of ions.
A lattice is just a regular structure.
It’s ‘giant’ because it is made up of the same repeated unit over and over again.
The ions with different charged attract each other, and ions with the same charge repel each-other.
The ions arrange themselves to maximise the attractions and minimise the repulsions. This overall attraction in the lattice is ionic bonding.
Sodium chloride is an example.
65
Properties of ionic structures?
Ionic compounds conduct electricity when they’re molten or dissolved - but not when solid.
- the ions in the liquid or a solution are free to move (and their charge moves with them) when they’re in a dissolved or molten state.
- in a solid, they’re in a fixed positive by the strong ionic bonds.
Ionic compounds have high melting points.
- the giant ionic lattices are held together by strong electrostatic forces so it takes loads of every to overcome them.
Ionic compounds are often soluble in water.
- water molecules are polar - part of the molecule had a small negative charge, and the other bits have small positive charges.
- the water molecules pull the ions away from the lattice and cause it to dissolve.
66
Covalent bonds?
They hold atoms that are in molecules together.
Molecules are formed when 2 or ore atoms bond together, and are held together by covalent bonds.
H2O is a molecule and is held together by covalent bonds.
There is electrostatic attraction between the nuclei and the electrons.
There is also a repulsion between the positive nuclei.
To maintain a covalent bond, there must be a balance between these forces.
67
Properties of substances made up of molecules?
Substances made up of molecules have fairly low melting points and boiling points (usually liquids and gases at standard temperature and pressure).
They don’t have giant structures.
To melt or boil a simple molecular compound, you only have to overcome the intermolecular forces of the molecules. These are weak compared to ionic or covalent bonds.
You don’t need to overcome the covalent bonds that hold the atoms together to melt/boil a substance.
They don’t conduct electricity because there are no charge carriers that are free to move.
They’re usually insoluble in water (sometimes a little soluble). The polar water molecules are more attracted to each other than the molecular substances, and so tend to leave it alone.
68
Dative covalent bonding?
Dative bonds can also be called ‘coordinate’ bonds.
In covalent bonding, electrons are exchanged (shared).
In dative covalent bonding, one atom donates both electrons to the bond.
When the ammonium ion (NH4+) forms, a dative bond occurs when the nitrogen atom in the ammonia molecule donates a pair of electrons to a proton, H+.
69
Which compounds don’t need 8 electrons in their outer shell?
Boron trufluoride only wants 6 electrons in its outer shell.
In a few compounds, atoms can use d orbitals to ‘expand the octet’ - this means they have more than 8 electrons in their outer shell. E.g. sulfur hexafluoride.
70
How can covalent bonds form giant structures?
Covalent bonds form when atoms share electrons with other atoms. This often leads to the formation of small molecules, like CO2, N2.
It can also lead to big lattices.
They have electrostatic attractions holding the structures together and they are much stronger than the electrostatic attractions between simple covalent molecules.
Carbon and silicon can form these giant networks this is enter side they can form four, strong covalent bonds.
71
Properties of giant structures of covalent bonding?
This info can also be used as evidence for covalent bonds.
Strong covalent bonds in the giant molecular structures mean that:
- They have high melting points,
- Extremely hard solids (also due to strong bonds),
- Are good thermal conductors due to vibrations being able to easily travel through the stiff lattices,
- Won’t dissolve - the covalent bonds mean that atoms are more attracted to their neighbours in the lattice then to solvent molecules. The fact they are insoluble in polar solvents like water shows that they don’t contain ions.
- Cant conduct electricity - since there are (in most giant lattice structures) no charged ions or free electrons (all the bonding electrons are held in localised electron bonds.
72
Metallic structures?
Metal elements exist as giant metallic lattice structures.
1. In metallic lattices, the electrons in the outermost shell of the meg atoms are delocalised - free to move. This leaves a positive metal ion.
2. The positive metal ions are attracted to the delocalised negative electrons. They form a lattice of closely packed positive ions in a sea of delocalised electrons - this is metallic bonding.
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How does metallic bonding explain metal properties?
1. The melting points are metals are generally high because of the strong metallic bonding, with the number of delocalised electrons per atom affecting the melting point. The more electrons there are, the stronger the bonding will be and the higher the melting point. Mg2+ has a higher melting point than Na+ because it has 2 delocalised electrons and Na only has 1. The size of the metal ion and the lattice structure also effect the melting point.
2. There are no bonds holding specific ions together and so the metal ions can slide over eachother when the structure is pulled, so metals can be shaped and are ductile (can be drawn into a wire).
3. The delocalised electrons can pass kinetic energy to each other, making metals good thermal conductors.
4. Metals are good electrical conductors because the delocalised ever rings are free to move and can carry a current. Any impurities can dramatically reduce conductivity by reducing the number of electrons that are free to carry charge via moving. This is because the electrons transfer to the impurities and form anions.
5. Metals are insoluble, expect in liquid metals, because of the strength of the metallic bonds.
74
Molecular shape depends on what?
Molecules and molecular ions come in loads of different shapes.
The shape depends on the number of pairs of electrons in the outer shell of the central atom.
E.g. ammonia, the outer shell of nitrogen has four pairs of electrons. 3 hydrogens and one lone pair.
Lone pairs are not shared.
75
Why and how do electron pairs repel each other?
Electrons are negatively charged and so electron pairs repel each other as much as possible.
The type of electron pair changes how much repelling happens.
So the shape of a molecule depends of the type of electron pairs surrounding the central atom as well as the number.
Lone pairs repel more than bonding pairs - take off 2.5 degrees for every lone PAIR.
This means the greatest angles are between lone pairs of electrons, and bond angles are often reduced because they are pushed together by lone pair repulsion.
Lone pair -lone pair angles repel the most.
Lone pair - bonding pair angles repel the second most.
Bonding pair - bonding pair angles are the smallest.
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What is electron pair repulsion theory?
The way of predicting molecular shape is known as ‘electron pair repulsion theory’.
Some examples:
The central atoms in these molecules have four pairs of electrons in their outer shells, but they’re all different shapes.
Methane, CH4 - no lone pairs so all bond angles are 109.5 degrees.
Tetrahedral shape.
Ammonia, NH3 - has one lone pair so the angles are 107 degrees.
Water - has 2 lone pairs so the bond angle is 104.5
77
How to predict the shapes of molecules?
To predict the shape of a molecule, you first need to draw a dot and cross diagram.
You can use this diagram to work out how many bonding and non-bonding electron pairs are on the central atom.
E.g. draw NCl3. Make sure all the atoms (not just the central atom, all of them) have 8 electrons in their outer shells. Sometimes, you might have to use a lone pair to make this true.
NCl3 has a lone pair on central atom.
Then, you can work out the shape of the molecule.
From the dot and cross diagram, you can see there are four electron pairs. 3 are bonding pairs and 1 is a lone pair.
There is more repulsion between the lone pair and the bonding pairs. This means the bonding pairs will be closer together. The shape will therefore be trigonal pyramidal.
Angle is 107 degrees because it’s a trigonal pyramidal.
78
Types of shapes of molecules?
```
Linear,
Trigonal planar,
Non-liner (‘bent’) with one lone pair,
Tetrahedral,
Trigonal pyramidal,
Non-linear (‘bent’) with 2 lone pairs,
Octahedral,
Trigonal bipyramidal.
```
79
What molecules can expand its octet?
Any molecules that have the electron configuration number beginning with 3 or above?
This means that there can be more than 8 electrons in the outer shell.
They have an expanded octet because they can use d-orbitals.
80
Linear molecules?
These have 2 electron pairs on the central atom.
No lone pairs.
The shape of the molecule is just a straight line across.
E.g. Cl — Be — Cl
BeCl2
And CO2 is a linear molecule too.
The angle is 180 degrees.
Doesn’t matter if the molecules are using double bonds or single bonds. The molecule still stays linear even though there is slightly more repulsion from a double bond.
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Trigonal planar?
There are 3 electron pairs around a central atom.
No lone pairs.
The shape of the molecule is a triangle with no atoms infront or behind. Literally just in a 360 degree 2D circle.
E.g. BF3 and CO3,2-
The angle is 120 degrees.
In CO3,2-, the bonds are all midway between single and double bonds. They’re not single, they’re not double, they’re in the middle.
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Non-linear molecules with 1 lone pair?
These have 3 pairs on the central atom.
1 lone pair.
The shape of the molecule is a triangle again, 2D with no atom in front or behind.
E.g. SO2
The angle is 120 degrees.
In SO2, the extra electron density from the double bonds cancels out the extra repulsion from the lone pair, so you still get 120 degrees.
83
Tetrahedral?
These have 4 electron pairs on the central atom.
No lone pairs.
The shape of the molecule is 3 molecules in a 3D triangle shape on the bottom and then a molecule on the top. One of the molecules is behind, one is in front and one is on the paper. The molecule on the top is also on the paper.
E.g. CH4 and NH4+
In the NH4+ example, the N is an ion with a + sign on it. The H is giving both the electrons in the bond. They’re not shared. This is shown with an arrow pointing towards the H.
The angle is 109.5 degrees.
84
How is a dative bond shown on a molecule?
A dative bond is shown with an arrow pointing toward the molecule giving both the electrons (dative bond).
85
Trigonal pyramidal?
These have 4 electron pairs on the central atom.
1 lone pair.
The shape of the molecule is like tetrahedral but the bond on the top is replaced with a lone pair.
E.g. NH3 and SO3,2-.
The angle is 107 degrees.
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Non-linear molecule with 2 lone pairs?
These have 4 electron pairs on the central atom.
2 lone pairs.
The shape of the molecule is not 3D. It’s just 4 pairs around the molecule in a square shape but 2 of the pairs are not bonds, they’re lone pairs.
E.g. H2O
The angle is 104.5 degrees.
87
Trigonal bipyrimidal?
5 electron pairs around central atom.
No lone pairs.
The shape of the molecule is awkward. Look in book at page 40. The molecule has two different bond angles.
E.g. PCl5
The angle is 120 degrees between the 4 atoms on the side of the central atom. The bond angle is 90 degrees around the other side. Look at book.
1 atom is behind, one is infront and 3 are on the paper. The 3 on the paper have 2 90 degree angles between them. The ones that are behind and infront have the 120 degrees between them.
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Octahedral?
These have 6 electron pairs on the central atom.
No lone pairs.
The shape of the molecule is 6 pairs around the central atom. 2 molecules are back and 2 are infront and 2 are on the paper.
E.g. SF6
The angle is 90 degrees.
If one of the atoms is taken off, it becomes squared pyramidal.
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Similarities between tetrahedral, trigonal pyramidal and non-linear with 2 lone pairs?
These molecules are all based around the same tetrahedral shape.
The reason they change is the lone pairs. As lone pairs increase, the repulsion increases and so the shape changes.
90
How to know whether an atom is behind or in front when drawn?
If it’s behind, the triangle is striped or dotted.
If it’s infront, the triangle (‘wedge’) is coloured in black.
91
How is the periodic table arranged?
Groups elements with similar properties.
It’s arranged according to atomic number.
Arranged into periods (rows) and groups (columns).
All the elements within a period have the same number of electron shells:
- the elements of period 1 have one electron shell.
- elements in period 2 have 2 electron shells.
All the elements within a group have the same number of electrons in their outer shell:
- group 1 elements have one electron in outer shell.
The one exception is group 0 - they have a full outer shell.
Elements in the same group have similar physical and chemical properties.
Properties often change gradually as you go down each group. E.g. metals in group 1 become more reactive as you g down the group whereas group 7 elements become gradually less reactive.
92
What are periodic trends?
Periodic trends are patterns in the periodic table.
93
Period trend of metals?
Metals (Li, Be, Na, Mg and Al) melting points increase across the period, left to right.
This is because the metal-metal bonds get stronger because of metal ions increase in the number of delocalised electrons and their ionic radium decreases.
Ionic radium decreases because the electrons are pulled closer as the number of protons in the nucleus increases.
This leads to a higher charge density, which attracts the ions together more strongly.
94
Period trends of giant covalent structures?
C and Si have strong covalent bonds making up a giant covalent structure.
A lot of energy is needed to break these bonds.
Carbon (as graphite or diamond) and silicon have the highest melting points in their period.
Structures with giant covalent structures have the highest melting points.
(Carbon sublimes from solid to gas at atmospheric pressure, it doesn’t actually melt but the rule still applies).
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Periodic trends of simple molecular substances?
Simple molecular substance (N2, O2, F2, P4, S8, Cl2).
Their melting points depend upon the strength of their intermolecular forces between their molecules. Intermolecular forces are weak and easily broken so these elements have low melting points.
More atoms in a molecule mean stronger intermolecular forces. E.g. period 3, sulfur is the biggest molecule (S8) so it’s got a higher melting point than chlorine or phosphorus.
The noble gases have the lowest melting points be said they exist as individual atoms (monatomic) resulting in very weak intermolecular forces.
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What is ionisation?
When electrons have been removed from an atom or molecules, the atom has been ionised.
The energy you need to remove an electron is called first ionisation enthalpy.
97
First ionisation enthalpy?
The first ionisation enthalpy is the energy needed to remove 1 electron from each atom in 1 mole of gaseous atoms to form 1 mole of gaseous 1+ ions.
Equation for the first ionisation enthalpy:
X(g) —> X+(g) + e-
E.g. equation for oxygen. O(g) —> O+ + e-
1st ionisation enthalpy = +1314kj mol-1
Important points about ionisation enthalpies:
- you must use the gas state symbol in equation because ionisation enthalpies are measured for gaseous atoms only.
- always refer to 1 mole of atoms, rather than to a single atom.
- the lower the ionisation enthalpy, the easier it is to remove an outer electron and form an ion.
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What effects the size of ionisation enthalpies?
Atomic radius - the further the outer shell electrons are from the positive nucleus, the less they’ll be attracted towards the nucleus. So, the ionisation enthalpy will be lower.
Nuclear charge - this is the positive charge on the nucleus caused by the presence of protons. The lore protons there are in the nucleus, the more it’ll attract the outer electrons - it’ll be harder to remove the electrons and so the enthalpy will be higher.
Electron shielding - the inner electron shells shield the outer shell electrons from the attractive force of the nucleus. Because more inner shells mean more shielding, the ionisation enthalpy will be lower.
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Periodic trends (group) of first ionisation enthalpies?
The first ionisation enthalpy decreases down a group because:
1. As you go down the group, the outer electrons are in shells further from the nucleus so they’re attracted to the nucleus less.
2. The amount of shielding increases because there are more filled inner shells. This means less nuclear attraction for the outer shell electrons.
3. Although the number of protons increases down the group, this doesn’t lead to an increase in ionisation enthalpy because its a less important factor than either shielding or the distance of the outer electrons from the shell.
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Period trends (period) of first ionisation enthalpies?
1. As you move across a period, the ionisation enthalpies increase so it’s harder to remove outer electron.
2. This is because the number of protons is increasing so the outer electrons are attracted to the nucleus more strongly.
3. All the electron outer-shells are at roughly the same energy level across a period and so there’s generally little extra shielding effect or extra distance to lessen the attraction from the nucleus.
4. There are small drops between groups 2 and 3, and 5 and 6 but we don’t need to know why.
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Ionisation enthalpies of s-block metals?
S-block metals have relatively low nuclear charges so they have low first ionisation enthalpies compared to other elements in the same period.
This means they lose their electrons easily as there is less attraction between the nucleus and the electrons.
This makes the s-block metals reactive.
The p-block metals have higher nuclear charges than s-block metals in the same period. This is down to the increase in the number of protons across each period.
This means they have higher ionisation enthalpies so it’s more difficult to lose their outer electron and they are less likely to lose the electron.
102
Group 2 element properties?
When group 2 elements react, they form 2+ ions. They loose the 2 electrons when they react.
Summary:
They react with water to produce hydroxides,
They burn in oxygen to form oxides,
They form alkaline solutions in water,
They neutralise acids,
Group 2 carbonates decompose to form CO2 and metal oxides.
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Group 2 reacting with water?
Group 2 metals react with water to give a metal hydroxide and hydrogen.
E.g. Ca(s) + 2H2O —> Ca(OH)2(aq) + H2
They get increasingly reactive down the group because the outermost electrons are further from the nucleus, and more easily lost.
```
Be - doesn’t react.
Mg - slow reaction.
Ca - steady reaction.
Sr - fairly quick,
Ba - rapid.
```
(If you react water with some barium in a test tube, lists of bubble are produced which shows how barium reacts really easily. If you add water to some magnesium in a test tube, you won’t see much bubbling).
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Group 2 burning in oxygen?
Group 2 metals burn in oxygen to form solid, white oxides.
E.g. 2Ca(s) + O2(g) —> 2CaO(s)
105
Group 2 reacting to form alkaline solutions?
We already know that in water, they form metal hydroxides and burn in oxygen to form oxides.
The oxides of the group 2 metals (that can be created by burning in oxygen) react readily with water to form metal hydroxides, which dissolve. This creates a strong alkaline solution because there are lots of OH- (hydroxide) ions.
(Magnesium oxide is the only exception - it only reacts slowly and the hydroxide isn’t very soluble.)
The oxides form more strongly alkaline solutions as you go down the group. This is because the hydroxides get more soluble.
Calcium hydroxide solution has a pH of 12.
E.g. CaO(s) + H2O(l) —> Ca^2+(aq) + 2OH-(aq)
106
Group 2 reacting to neutralise acids?
Both oxides and hydroxides neutralise dilute acids because they’re bases.
This forms solutions of the corresponding salts.
E.g. MgO(s) + 2HCl(aq) —> H2O(l) + MgCl(aq)
Magnesium oxide - MgO (white solid).
E.g. Mg(OH)2(s) + 2HCl(aq) —> 2H2O(l) + MgCl2(aq)
Magnesium hydroxide - Mg(OH)
107
Solubility trends of group 2 elements?
Generally, compounds of group 2 elements that contain singly charged negative ions (e.g. OH-) increase in solubility down the group.
Compounds that contain double charged negative ions (e.g. CO3^2+ and SO4^2-) decrease in solubility down the group.
Magnesium hydroxide has OH- ions and is least soluble hydroxide.
Magnesium carbonate has CO3^2- ions and is most soluble carbonate.
108
What is it called when something has a very low solubility?
‘Sparingly soluble’.
109
Group 2 carbonates decompose to form?
Group 2 carbonates decompose to form CO2 and metal oxides.
Thermal decomposition is when a substance breaks down when heated.
E.g. CaCO3(s) —(heat)-> CaO(s) + CO2
The volume of CO2 produced decreases as you go down the group because the Mr of the metal gets bigger so a smaller number of (element)CO3 moles is contained in the same mass.
110
Thermal stability of carbonates changes how as you go down the group?
The more thermally stable a substance is, the more heat it will take to break down.
The thermal stability of group 2 carbonates increases as you go down group 2.
The charge on all group 2 cations is the same, 2+ but the charge density decreases down the group.
Anions - negatively charged ions.
Cations - positively charged ions.
Mg ions polarise carbonate ions more than barium ions do, meaning magnesium carbonate is less stable.
111
Why does thermal stability increase down group 2?
Carbonate ions are large anions and can be made unstable by the presence of a cation (such as a group 2 metal ion) - formation of a group 2 carbonate.
The cation draws the electrons on the carbonate ion towards itself (it polarises it).
This distorts the carbonate ion. The greater the distortion, the less stable the carbonate ion.
Large cations cause less distortion than small cations because they have a lower charge density.
The further down the group, the larger the cations and so the less distortion caused and so they are more stable carbonate anions.
112
What does charge density mean?
This is just the charge on the ion compared to (relative to) it’s volume.
113
What is the electromagnetic spectrum?
Electromagnetic spectrums show the range of electromagnetic radiation (energy that is transmitted as waves with different frequencies).
Along the electromagnetic spectrum, the radiation increases in frequency and decreases in wavelength.
```
Radio waves,
Micro-waves,
Infrared,
Visible light,
Ultraviolet,
X-rays,
Gamma rays.
```
Radio waves have the lowest frequency/energy and the smallest wavelength.
114
Electrons absorb or release energy?
Electron shells are sometimes called energy levels.
Atoms in their ‘ground state’ have all their electrons at their lowest possible energy levels.
If an atoms electrons take in energy from their surroundings, they can move to a higher energy level, further from the nucleus. This means they are excited.
Electrons can also release energy by dropping from a higher energy level to a lower one.
The energy levels all have certain fixed values (they’re discrete).
Electrons can jump from one energy level to another by absorbing or releasing a fixed amount of energy.
An electron absorbs energy if it goes up to a higher level and emits energy if it goes down an energy level.
115
Differences between line spectra?
When electrons move between energy level, they produce line spectra.
Two types of line spectra:
- absorption spectra,
- emission spectra.
116
Atomic absorption spectra?
Energy is related to frequency. So when electromagnetic radiation is passed through a gaseous element, the electrons only absorb certain frequencies, corresponding to the differences between energy levels.
This means the radiation passing through the gaseous element has certain frequencies missing. These frequencies correspond to the differences between the energy levels.
The missing frequencies show up as dark lines on a coloured back ground.
117
Emission spectra?
When electrons drop to lower energy levels, they give out certain amounts of energy.
This produces lines in the spectrum too.
For any particular element, the frequencies in an emission spectrum are the same as those missing in the absorption spectrum (the patterns are opposites).
Each elements had a different electron arrangement, so the frequencies of radiation absorbed and emitted are different. This means the spectrum for each element is unique.
Shown as coloured lines on black background.
118
Reading the lines on a line spectra?
Each set (series) of lines in a soectra represents electrons moving to or from an energy level.
N = 1 is the ground state.
The lines on both absorption and emission spectra get closer together the higher the frequency of the light.
(Do exam questions on these because it’s confusing).
119
Energy related to frequency equation?
🔺E = hv
🔺E = change in energy (between two shells maybe). Units - jules/J.
h = Plancks constant. Units - J Hz-1. (This is always the same number and it’s given on the data sheet).
v = frequency. Units - Hz or s-1.
120
Equation for wavelength?
C = v(upside down y)
C = speed of electromagnetic radios (which is always 300m/s because it’s the same as the speed of light). Units - m/s.
v - frequency. Units - Hz or s-1.
(Upside down y) - wavelength. Units - m.
121
Equation relating the two equations for frequency and energy change?
You can combine 🔺E=hv and c=vy into one calculation.
This is:
E = hc / y
122
How to convert nm to metres?
Convert 400nm to m (for an equation, you usually have to use metres).
400 x10^-9 = (this number will be in metres).
400 times 10 to the power of -9 = (this number will be in metres).
123
Nuclear Fusion?
Nuclear fusion is when two small nuclei combine under high temperature and pressure to make one larger nucleus.
It happens naturally in stars. If we could create the same temperature and pressure on earth, we could do it (but we can’t create this temp and pressure).
You can represent nuclear fusion using a nuclear equation (the one where the two top numbers must add and the two bottom numbers must add).
A neutron has a 1 on the top and 0 on the bottom.
In stars, hydrogen nuclei combine to make helium nuclei, real easing huge amounts of energy. This happens in the suns core.
When hydrogen in a stars core runs out, the temperature and pressure of the core rises. In a big enough star, it’ll get hot enough to fuse heavier elements, starting with helium.
Large nuclei can only be made by stars (either inside them, or as a dead star explodes as a supernova).
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Salts properties summary?
They are ionic compounds,
They can be soluble or insoluble in water,
Precipitation reactions make insoluble salts,
You can use a metal or an insoluble base to make soluble salts,
You can also make soluble salts using an alkali.
125
pH scale?
7 is neutral.
Below 7 is acidic.
Above 7 is alkali.
126
What are salts?
Acids react with bases in neutralisation reactions to form a salt and water.
E.g. HCl (aq) + NaOH (aq) —> NaCl (aq) + H2O
Salts are formed from positively charged cations and negatively charged anions. The product is neutral.
When naming salts, you put the cation bit first and then the anion. E.g. a salt formed from zinc ions and sulfate ions is zinc sulfate.
To work out the formula for a salt, you need to balance the charges.
127
Cations to memorise?
NH4+ - ammonium
Cu^2+ - copper (II)
Zn^2+ - zinc (II)
Pb^2+ - lead (II)
Fe^2+ - iron (II)
Fe^3+ - iron (III)
The Roman numeral tell yous the charge on the ion. If the Roman numberal is 2, then the charge on the ion is 2+. This only works for positive ions.
128
Anions to memorise?
NO3- - nitrate
OH- - hydroxide
HCO3- - hydrogen-carbonate
SO4^2- - sulfate
CO3^ 2- - carbonate
129
What salts are soluble or insoluble.
You need to know if a salt is soluble or insoluble so you know how to make it.
- Lithium, sodium, potassium, and ammonium salts are soluble.
- Nitrates are soluble.
- Most chlorides, bromides and iodides are soluble (some exemptions on another flashcard).
- Most sulfates are soluble (exceptions on another flashcard).
- Most hydroxides are insoluble (expeditions on another flashcard),
- Most carbonates are insoluble (exceptions on another card).
130
What chlorides, bromides and iodides are not soluble?
```
Silver halides,
copper iodide (white precipitate),
lead chloride (white precipitate),
lead bromide (white precipitate)
and lead iodide (yellow precipitate).
```
131
What sulfates are not soluble?
Barium sulfate,
calcium sulfate,
lead sulfate.
All form white precipitates.
132
What hydroxides are soluble?
```
Lithium,
sodium,
potassium,
strontium,
calcium,
barium,
ammonium hydroxides.
```
133
What carbonates are soluble?
Lithium,
Sodium,
Potassium,
Ammonium carbonates.
134
Some insoluble carbonates form coloured precipitates?
Most carbonates are insoluble.
They form coloured precipitates.
Copper carbonate is blue-green.
Silver carbonate is yellow.
Most other carbonate precipitates are white (barium, calcium, lead (II), and zinc-carbonates.
Iron (II) is off-white.
135
Making insoluble salts?
Precipitation reaction.
Occur in aqueous solution, when cations and anions combine to form an insoluble ionic salt. This insoluble salt is called a precipitate.
1. Pick two solutions that contain the ions you need.
E.g. to make lead chloride, you need a solution which contains lead ions and a solution which has chloride ions. So mix lead nitrate (nitrates are soluble) which sodium chloride solution (most chlorides are soluble).
E.g. Pb(NO3)2(aq) + 2NaCl(aq) —> PbCl2(s)+ 2NaNO3 (aq)
Once the salt has precipitated out (and is lying in bottom of flask), you use filter paper, wash it, then dry it on filter paper.
Salts can be made in other ways (not just reaction between acids and bases).
136
Making insoluble salts using a metal or an insoluble base?
You use an acid and suitable metal (or insoluble base like a metal oxide or metal hydroxide).
E.g. I’d you want to make copper chloride, mix Hydrochloric acid and copper oxide.
CuO(s) + 2HCl (aq) —> CuCl2 (aq) + H2O (l)
(Acids and reactive metals produce a salt and hydrogen instead of water. You shouldn’t do the reaction with a really reactive metal cause it will explode).
Picking the right acid:
- to make chlorides, use hydrochloric acid.
- to make sulfates, use sulfuric acid.
- to make nitrates, use nitric acid.
Then add the solid metal, metal oxide or hydroxide to the acid. It will dissolved in the acid as it reacts. You will know when all the acid has been neutralised because no more solid will dissolved, and so it just sinks to bottom.
Then filter the excess metal (or whatever you used) to get the salt solution.
Then to get solid salt, evaporate some water to make the solution for concentrated and leave to rest to evaporate slow. This is crystallisation.
137
Making soluble salts using an alkali?
Alkalis are just soluble bases in this.
You can’t use the other method for making soluble salts with alkalis (like potassium, sodium or ammonium hydroxides) because you can’t tell whether the reaction has finished because it won’t be in excess.
You have to add exactly the right amount of alkali to just neutralise the acid.
You do this by doing a titration and using an indicator.
Once you’ve found the right amount of alkali, repeat the titration by combing the volumes again but don’t add the indicator because it contaminates the salt.
Then just evaporate the water off to crystallise it as normal.
