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Notes

Elements

There is a little bit of history about the discovery of the elements that you have to learn, it is boring but youjust have to learn who did what. You only really need to know their second names.

Robert Boyle : said an element was something that couldn't be broken down into a simpler materialHumphry Davy : invented electrolysis and isolated a few elements.Henry Moseley : found the atomic number using x-rays.

Periodic TableJohann Dobereiner : elements can be arranged into groups of 3 called triads, with similar propertiesand related atomic weights.John Newlands : octaves, properties repeat themselves every 8 elements when arranged by atomicmassDmitri Mendelev : made a periodic table very similar to the one we use now.

Element: a substance that can't be broken down into simpler materials through chemicalmeans.

The periodic table of the elements is broken into metals (left) and non-metals (right) by the stairs line.

The AtomThere is also history about the discovery of the atom that you need to learn. You only really need to knowtheir second names.

John Dalton : All matter is made up of small indivisible particles called atoms, they cannot be createdor destroyed.William Crookes : cathode rays travel in straight lines, maltese cross, also proved they are particles byadding a wheel and a rail to the inside and watched them movingJ. J. Thomson : cathode ray particles are negatively charged, he also found the charge to mass ratioJohnstone Stoney : named the cathode ray particles the electronRobert Millikan : oil drop experiment, measured the charge on an electronJames Chadwick : found the neutrons in the atom.

Thomson's Plum pudding model

Thomson believed that the atom was a cloud of positive charge and had the little negative particles floatingaround in it.

Rutherford's Gold Foil ExperimentErnest Rutherford disproved this model with his gold foil experiment.He fired alpha particles at a thin layer of gold foil.He expected them all to pass straight through with a slight deflection according to the Plum PuddingModel.

What actually happened :

Most passed straight through with no deflectionSome were deflected at small anglesVery few were rebounded at large angles.

Rutherford's conclusions/new model :

The atom is mostly empty space with a small dense nucleus at the centre.Positive nucleusElectrons move around the atom's empty space.Negative electrons cancel out the positive and overall neutral atom.

The atom so far

So Proton mass = Neutron Mass Proton Charge = Electron Charge (except one is plus the other is minus) Electron is 1836 times lighter than a proton or a neutron.

Atomic Number : The atomic number of an element is the number of protons in the nucleus of that atom.

Mass Number: The mass number is the number of protons and neutrons in the atom of an element. MassNumber = number of protons + number of neutrons.

The mass number is always bigger than the atomic number.

Isotopes : Isotopes are atoms with the same number of protons but different numbers of neutrons.

An example of an isotope is Carbon-12 and Carbon-14. They are both the same element, howeverCarbon-12 has 6 neutrons and 6 protons. Carbon-14 has 6 protons and 8 neutrons.

An element can't have a different number of protons and still be the same element. If two substanceshave a different number of protons then they are different elements.

Relative Atomic Mass: This is the average mass of all the isotopes in an element, taking their abundanciesinto account, measured relative to the mass of Carbon-12.

Mass SpectrometerBasically what a mass spectrometer does is it finds the relative atomic mass of an element. So it findshow heavy it is, what isotopes it has, and how abundant each isotope is.You need to know the different steps.

1. Vaporisation : The sample is vaporised- turned into gas.2. Ionisation : It is ionised by an electron gun, this means it loses some of its electrons and becomes

positively charged3. Accelaration** :** The ion is accelerated to a high speed by using an electric field.4. Deflection : The different isotopes of the element are deflected using a magnetic field, the lighter ones

are deflected more, and this way the sample is separated.5. Detection : A detector detects the ion samples and then gives you a chart so you can find the relative

atomic mass.

Energy Levels

Energy Levels: An energy level is a region of definite energy within an atom that electronscan occupy.

When element are given energy, they emit light. Examples include the flame test experiments and fireworks.We will look at the reason for this now, it is closely related to energy levels.

Light SpectrumWhite light is made up of loads of different types of light that mix to form white light. When white lightis broken up we see a continuous spectrum.

However when we burn an element and analyse the light that comes from it we see a different type ofspectrum called a line spectrum.

This is the line spectrum of Hydrogen. Every single element has a different and unique line spectrum.There is another way of doing line spectrums and it's called an absorption spectrum, this is when youshine light through hydrogen and it absorbs light in the exact same places where the lines are on theline spectrum.

One other thing about the line spectrum. The lines we see there are the lines that are visible, there are

some ultraviolet and infrared lines too, they can't be seen by the human eye.

Lyman series: Ultraviolet line seriesBalmer series: Visible line seriesPaschen series: Infrared line series

Bohr's TheoryNeil's Bohr made a theory about the atom that explained the emission spectrum of Hydrogen.

Electrons have energy and occupy energy levels in the atom depending on how much energy they have,they can't be between energy levels. Energy levels are n=1,2,3,4 etc.When an electron moves from a higher energy level to a lower energy level it has to get rid of theexcess energy, this is done in the form of light emission.The energy of the light is governed by this formula

The same is true that if an electron absorbs a certain amount of energy, then it jumps up the amount ofenergy levels that it can with the energy it has absorbed.Normally the electron in a hydrogen atom is in energy level n=1, this is called the ground state. When itabsorbs energy and goes to a higher energy level it is in the excited state.It can't stay up there forever though because it is unstable, so eventually it falls back down and emitsthe light to get rid of the excess energy.Depending on what level it starts at and what level it falls down to, it will emit a different colour of light,this is effectively what the line spectrum is.The diagram is the best way of showing this.The line spectrum of each element is unique in that element because the spacings between the energylevels in each atom are also unique.

Louis de Broglie : Matter can have wave and particle properties simultaneously

The Heisenberg Uncertainty Principle : It is not possible to determine at the same time the exactposition and the velocity of an electron.

Limitations of the Bohr Theory:

Only worked well for hydrogen, didn't work well for larger atoms with more electronsDidn't allow for Heisenberg's uncertainty principleDidn't take into account that electrons have properties of waves as well as particles.

Electron Configuration

(this just means how the electrons are laid out in the atom)

Energy Levels : An energy level is a region of definite energy within an atom that electrons can occupy.

Energy Sub-levels: A group of atomic orbitals within an energy level, that all have the same energy.

Atomic Orbital : A region in space where the probability of finding an electron is relatively high.

(it is incorrect but still acceptable and easier to think of these as paths around the nucleus that theelectron follows like the moon going around the earth)There are a number of different types of orbitals.

s-orbitals and p-orbitals are the two most basic, their shapes are shown below.s-orbitals are spherical, and p-orbitals are dumbbell shaped.p-orbitals are orientated along the axes, they are all at right angles to each other. (it is hard to showthis on a 2D photo.There are also d-orbitals and f-orbitals, but their shapes aren't important.

All orbitals can hold only two electrons.

s-orbitals are sublevels and orbitals at the same time. It can hold 2 electron.p-orbitals can hold 2 electron each, so the p sublevel can hold 6 electrons altogether.Okay so to recap all of that:Firstly, we have energy levels around the nucleus.Within energy levels there are energy sub-levels.Within those energy sub-levels, there are atomic orbitals.

Energy levels: n=1,2,3,4… Energy sub-levels: 1s, 2s, 2p, 3s… Atomic Orbitals: 1s, 2s, 2px, 2py, 2pz, 3s…

In some cases above the sub-levels are the same as the atomic orbitals. But to distinguish them look atthe p-orbitals.2p refers to the entire energy sub-level, whereas 2px refers to a specific atomic orbital within thatsublevel.

They can ask you to write the electron configuration for any atom. This is easy if you just follow a fewsimple rules.

Aufbau principle: Electrons will occupy the lowest energy sublevel available to them

Hund's rule: Electrons occupy orbitals singly before doubling up, the electrons have to be of opposite spin.

(Hund's rule is known as the empty bus seat rule because on a bus everyone fills the single seats beforedoubling up. You don't need to worry about what the electron spin means.

You need to learn the order in which the sub-levels are filled and you need to know thecapacity of each sublevel.

*Notice how the 4s sublevel has a lower energy than the 3d sublevel. The reason this is important iselectrons occupy the 4s sublevel before occupying the 3d sublevel.

You also need to know the capacity of each sublevel.

Now for some examples:Example 1 : What is the electronic configuration of Mg – Magnesium?

The first thing you do is you find the element in the periodic table. Mg has atomic number 12, thismeans it has 12 electrons. Then you assign the electrons into orbitals according to the rules.

Mg = 1s22s22p63s23p64s23d6

2 + 2 + 6 + 2 = 12 (it's as easy as that)

Example 2: What is the electronic configuration of Fe – iron?

Iron has 26 electrons.Fe = 1s22s22p63s23

Example 3: What is the electronic configuration of Mg2+? - Mg2 is an ion, remember an ion is a chargedatom because it has the wrong amount of electrons. Mg has 12 electrons, and Mg2 has a positive charge of2+ which means it lost 2 electrons. This means that Mg2 has 10 electrons. - Mg2+ = 1s22s22p6

There are 2 electron configurations they ask more than any other ones and that is because they have adifferent electronic configuration than you'd expect.

Make sure to know these exceptions off by heart.

The reason that these are different is that half filled sub-levels are more stable than partly filled, so theelectrons arrange themselves to be the most stable (and this means breaking the rules a little)

Ionisation Energies

Atomic radius: The atomic radius of an element is half the distance between the nuclei of two atoms of theelement that are joined together by a single covalent bond.

Trend for atomic radii:

Increasing as you go down in a group because:

A larger number of energy levels occupiedScreening effect

Explanation:

What the first bullet point means is that every time you go down an element within a certain group, thenthere will be a new energy level in the lower element. For example Hydrogen has 1 energy level, Lithium has2 energy levels, and sodium had 3. So as you get more energy levels they will obviously be spaced furtheraway from the nucleus and thus making the atom bigger.

Screening effect: so the nucleus has a positive charge and the electrons have a negative charge. So thenucleus tries to pull the electrons in. But the lower energy level electrons actually take a lot of that pull andprotect (or screen) the higher energy levels from the pull of the positive charge.

Decreasing as you go from left to right across a period:

Increased nuclear chargeScreening effect doesn't work going across.

Explanation:

The first bullet point means that as you go across a period the atomic number of the atom increases whichmeans there are more protons in the nucleus. Remember the protons pull on the electrons so the moreprotons in the nucleus the closer the electron will be pulled and the smaller the radius will get.

The second one means that although the screening effect does happen, it doesn't make a difference whenyou are just going across a period because there are no new energy levels starting, so all of the newelectrons that come are at the same distance as the previous ones. And the first bullet point shows how theyare pulled closer in.

First Ionisation Energy : the minimum energy in kilojoules required to remove the most loosely boundelectron from each isolated atom in a mole of the element in its ground state.

Basically how hard it is to steal an electron from an atom.

Trend in first ionisation energies:

Increasing as you go from left to right across a period because:

Increase in Nuclear chargeDecrease in Atomic Radius

Explanation : It is the opposite of the atomic radius trend, which makes sense. The atomic radius decreasesbecause there is a greater pull on the electrons, that means that the energy needed to steal an electron willdefinitely increase.

Decreasing as you go down a group because:

Increase in atomic radiusScreening effect

Explanation : The same as above, if the radius increases there will be a lower pull on the outer electrons,therefore they will be easier to steal.

The outer electrons also won't be affected by the pull of the nucleus as much because of the screeningeffect.

Exceptions to the trend:

So generally speaking the ionisation energies get bigger as you go across the period, however there are twoexceptions.

In the second period:

Boron has a lower energy than Beryllium.Oxygen has a lower energy than Nitrogen.

The reason for the first one is because Be has an electronic configuration of 1s22s2, and Boron has anelectronic configuration of 1s22s22px1</sup. Beryllium is actually more stable because it has each sublevel fully filled,and Boron's last electron is on its own in the 2p sublevel and therefore is easy to take.The reason for the second one is pretty much the same. The electronic configuation of Nitrogen is1s22s22px12py12pz1. And the electronic configuration of Oxygen is 1s22s22px22py12pz1. Again thereason is more or less the same, in Nitrogen all of the p-orbitals are half full (which makes them quitestable). In Oxygen the 2px orbital has two electrons in it, this second electron is quite easy to take andtherefore Oxygen has a lower ionisation energy than Nitrogen.

The second and third ionisation energies refer to the removal of the second and thirdelectrons from the element.

This chart above shows the different ionisation energies of Beryllium. You can see that the firstionisation energy and the second are quite small. Then the third one is much much bigger.This is used as evidence to support the idea of energy levels.

Look at the electron configuration of Be 1s2 2s2, the first two electrons are in the same energy level sothe first two ionisation energies should be roughly the same. Then the next electron is in a lower energylevel which means it is much closer to the nucleus and there is no screening effect. This and the factthat the third ionisation energy is massive all support the idea of energy levels in the atom.

This second part of the chapter comes up a lot less often than the above part butyou have to know it as well.

Ionic and Covalent bonding

There are two main types of chemical bonding, covalent and ionic, we will look at each in turn.

Octet Rule: Atoms want to have 8 electrons in their outermost energy level.

Ion : A charged particle.

So for example if sodium reacts with another element it will normally lose its electron to obey the octetrule, but in doing this it becomes an ion of sodium. Because it loses something with a negative charge,the ion will have a positive charge. Na+.

Ionic Bond: The electrostatic force of attraction betweenoppositely charged ions.

Basically what this means is that in an ionic bond, one atom will give an electron to another atom sothat they both obey the octet rule. They will both be ions now because they will be electrically charged.One of them will be negatively charged (the one with the extra electron), and one of them will bepositively charged (the one who lost the electron).Then as we know negative and positive attract, so the two ions are attracted together. This is an ionicbond.The best example is sodium chloride, NaCl, this is the table salt we use.

Cation – positive ion.

Anion – Negative ion.

They can also ask in the exam to show a bond on a dot and cross diagram, this diagram above shows adot and cross diagram. Basically just let one atom's electrons be crosses and the other one's be dots,this was you can tell which electron moves to which atom.The square brackets mean that it is an ion.

Valency : The number of bonds an atom forms when it reacts.An example to show this is that Carbon has a valency of 4, so it forms 4 bonds when it reacts. So a commoncompound of carbon is CH4 (methane). You can see the 4 bonds it has with Hydrogen atoms in the diagrambelow.

Finding the valency of an element is simple enough. The valency is almost always going to be theamount of electrons the element has to lose or gain so that it can obey the octet rule.

So if we look at the periodic table of the elements:

Everything in Group 1, will lose one electron when it reacts, therefore it will have a valency of 1.Everything in Group 2, will lose two electrons when it reacts, therefore it will have a valency of 2.

Everything in Group 3, will lose three electrons when it reacts, therefore it will have a valency of 3.

This trend continues until it becomes easier for elements to start gaining electrons rather than losing

them to fulfil the octet rule. The amount of electrons they gain still equals their valency.

So the valencies of elements in different groups will follow this table here:

Covalent bonding

Covalent Bond : The bond that is formed when two atoms share a pair of electrons.

You can have single, double or triple covalent bonds.

The next bit is a bit weird, but it comes up a lot in question 4 as a short question.

Single bonds: are always sigma (�σ) bonds. The orbitals overlap end on. They are between two sorbitals, an s and a p orbital, or two p orbitals.Double bond: the first bond is a sigma and the second is a pi bond. This is a side on overlap of orbitals.Triple Bond: the first bond is a sigma, and the second and third are pi bonds. Again pi-bonds are side onoverlaps of orbitals.

Polar and Non-Polar bonds

When Hydrogen and Hydrogen have a covalent bond, they share the electrons equally because theyhave the same pull on the electrons.However for example in H2O, since H and O have different pulls on the electrons, the electrons aren'tshared equally.

Polar bond : a bond where the electrons are shared unequally. Non-polar bond : a bond where the electrons are shared equally.

In a polar bond the molecule will be negative on one side and positive on the other side, because of theelectrons not being dispersed equally.

Like dissolves like. This means that polar substances will dissolve in polar solvents. And non-polarsubstances will dissolve in non-polar solvents.This means that for cleaning you need to use the right solvent.You can't use water to clean permanent marker, you have to use a non-polar solvent like cyclohexane,because the ink in permanent marker is also non-polar.However, if you want to dissolve sugar into a liquid to make a sugary drink, water is perfect.

Electronegativity : The relative attraction of an atom for shared pairs of electrons in acovalent bond.

Measured on the Pauling scale from 0-4, the bigger the number the more pull it has on the electrons

The trend of increasing electronegativity is:

It is in the same direction as Ionisation energies, and for the same reasons:

Increase in Nuclear chargeDecrease in Atomic RadiusScreening effect

To predict the type of bonds between two elements we look at the difference between their electronegativityvalues.

The electronegativity values for each element can be found in the tables book.

Shapes of Molecules and Intermolecular BondingThe shape of molecules is literally the way they are arranged, and the shapes they make, exactly like in theatom toys you might have seen.

Electron Pair Repulsion Theory (EPRT)

The shapes aren't totally random though, they are dictated by the theory named above.

The electron pairs that are being used in covalent bonding repel each other, and end up as far apart asis geometrically possible.Lone pairs (pairs not in bonding), have a greater repelling effect than bonding pairs.

The four different shapes we will look at are:

LinearTrigonal PlanarTetrahedralPyramidalV-Shaped

Linear – 180** °**

Trigonal Planar 120 degrees: The example we will look at is Boron trifluoride – BF3.

The shape will not be like the one shown in the dot diagram above, this is only to show how theelectrons are shared. To find the shape we have to think of the EPRT. There are three bonds above, sothree pairs of electrons, one with each F atom. The furthest away these three pairs can be is if they areat the three corners of a triangle.

They have a bonding angle of 120°.

Tetrahedral – 109.5** °**

The example we will use for this is Methane CH4.

There are 4 bonds here, and that means 4 pairs of electrons.

They arrange themselves in a shape that makes them as far away from each other as possible,according to our EPRT.

This shape is a tetrahedral shape. This can only really be shown by a picture rather than a diagram.

The bond angle is 109.5°. The molecule is perfectly symmetrical, between each atom of Hydrogen (thered balls) the bonding angle is 109.5°.

Pyramidal - 107 °

(In all of the above examples, all of the electrons in the outer shell are being used in bonding, in the next twoexamples there will be some lone pairs of electrons, this will be shown in the diagrams)

The example we will use is ammonia NH3.

You can see the electron pair on top belongs to Nitrogen, but they are alone, they aren't pairing withanything.

This is kind of similar to the methane example, where it is tetrahedral in shape, however lone pairs havea stronger repulsion than bonding pairs, so the three Hydrogen molecules are actually pushed a little bitcloser together than normally.

(the lone pair doesn't get shown on the kind of diagrams like the one above)

The result being a bonding angle of 107° between the Hydrogen atoms (the white balls) rather than109.5 like in the methane.

V-Shaped 104.5 °

The example we will look at is water - H20.

As we can see there are two lone pairs on top here. This means that the hydrogen atoms will be pushedeven closer together than in the ammonia.

The bond angle in this case is 104.5° between the hydrogen atoms.

If they ask you to figure out a random molecules shape that you haven't seen before, then look at thevalencies of the atoms. Whatever the big molecule in the middle is, if you know its valency then you canfind out what shape it is from the table above.Also a trick to remember the bond angles of the last three: If you remember 109.5°, the other two arejust -2.5° each time.

Polarity in MoleculesPolarity in a molecule just means if it is unevenly charged i.e. one side is negative and the other side ispositive. For example:

The electrons are more attracted to the chlorine, therefore the chlorine ends up being more negativebecause there are more electrons hanging around it.

This is because Chlorine is more electronegative than Hydrogen.

However if the molecule is symmetrical then molecules can be non-polar, even though their bonds arepolar. We will look at the examples from earlier to show this.

BeH2 is non-polar even though both the bonds are polar. This is because you need to look at the centreof negative charge – this is the middle point between all of the negative charges in the molecule.

The centre of negative charge is halfway between the two hydrogen atoms, which is at the Berylliumatom. This is however also the centre of positive charge, because it is where the positive charge is allfocused. Therefore they cancel out and overall the molecule is non-polar.

The exact same thing happens in this case. This is because the molecule is perfectly symmetrical.

Again this is hard to show in a 2D diagram, but we can see that again because the molecule issymmetrical the negative and positive charges are in the same place and they cancel out, and themolecule is overall non-polar.

In this case however… The centre of negative charge is on top where the nitrogen atom is because ofthe lone pair of electrons.

The centre of positive charge is at the bottom between all of the Hydrogen atoms.

Because the molecule isn't symmetrical and the centres of negative and positive charge aren't in thesame place…

The molecule is polar overall.

Again here the same things is true. Water is a polar molecule.

Intermolecular Forces

Intra-molecular bonding – that is the bonding between atoms inside a molecule Inter-molecular forces – that is the forces between different molecules.

Don't get those mixed up

There are three different types of intermolecular forces. (listed below from weakest to strongest)

Van der Waals' forcesDipole-Dipole forcesHydrogen Bonding

Van der Waal's

These are weak forces that act between gases.

Imagine an oxygen atom:

In this picture all of the electrons are perfectly symmetrical, so the atom is non-polar because thecentre of positive and negative charge are both in the centre of the nucleus.

But in reality they are always moving around, and they might not always be symmetrical, this meansthe atom might be temporarily polar because the centre of positive and negative charge won't be in thesame place.

This is called a temporary dipole.

If this happens in two molecules of oxygen at the same time, it can create weak intermolecular forcesbetween them. i.e. Van der Waal's forces

Another thing that can happen is that a temporary dipole in one atom can create or induce a temporarydipole in another atom, creating a knock on effect and creating these weak Van der Waal's forces.

Dipole-Dipole

This is effectively the same thing as above, except instead of temporary dipoles being the reason forthe forces between the molecules, it is permanently polar molecules that are attracted to one another.

Hydrogen Bonding

This is a special type of dipole-dipole interaction, which occurs when hydrogen is bonded to small,highly electronegative atoms such as O, N or F.

This occurs in water. The bond between O and H is highly polar. This creates hydrogen bonds between thedifferent water molecules.