My Chemistry as Rg . Dc

7/31/2019 My Chemistry as Rg . Dc

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By Danii Cole

AS Unit F321:Atoms,

Bonds and Groups

I'm a great believer inluck, and I find the

harder I work the more

I have of it. ~Thomas

Jefferson

Work, look for peace and

calm in work: you will find

it nowhere else.Dmitri Mendeleev


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By Danii Cole

You need to know the following before we get started

with A level chemistry .

The nucleus contains protons (positively charged) and neutrons (neutrally charged i.e. no charge).The atomic number (proton number) is equal to the number ofprotons in the atoms nucleus.

The mass number is the total number of protons and neutrons in the nucleus.

Atom: smallest particle of matter that takes up space.

All elements are made up from atoms. Atoms are made up from 3 types of particles, protons, neutrons

and electrons. Electrons have a -1 relativecharge, and Protons have a +1 charge, whilst Neutrons are

neutral. Protons and neutrons have a relative mass of 1, whilst electrons have a relative mass of

1/2000.

Proton: a positively charged particle in the nucleus of the atom, represented by a capital P.

Neutron: a particle in an atomic nucleus with no electrical charge, represented by a capital N.

Electron: a negatively charged particle orbiting around the nucleus of an atom, represented by an "e"

or an " e-."

Now onto the serious stuff ...

IsotopesIsotopes are atoms of the same element with different numbers of neutrons

They have different masses, the same number of protons and electrons . However, a different numberof neutrons in the nucleus .

A good example is Carbon.

Carbon 12 is the element form as found on the periodic table of elements.

Carbon 14 is an Isotope of Carbon (It has a different number of neutrons).

Typical Exam questions


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REMEMBER!! 1 mole of gas

occupies 24dm 3at room

temperature and pressure.

Relative mass has NO UNITS

Relative Atomic Mass has NO UNITS

Relative molecular Mass has NO

UNITS

Relative Formula Mass has NO UNITS

An atom or molecule with a net electric charge due to the loss or gain of one or more electrons.

For example

K ---> K (potassium has lost an electron here, and becomes positively charged, (called a cation).

Cl ---> Cl (chlorine gains an electron, becomes negatively charged, called an anion).

Now . The following is key. You MUST know these definitions for your exam

12C is used as the standard measurement of relative masses, the mass of one atom of carbon-12 is 12 amu.

Relative Atomic Mass, ArThe weighted mean mass of an atom of an element compared to 1/12th of the mass of one atom of12C

Relative Isotopic MassThe mass of one atom of an isotope compared to 1/12th of the mass of one atom of12C

Relative molecular mass, Mr

The weighted mean mass of a molecule compared to 1/12

th

of the mass of one atom of12

C

Relative formula mass

The weighted mean mass of a formula unit compared to 1/12th of the mass of one atom of12C

To calculate Relative Atomic Mass of this europium sample you do the following :

You multiply the relative isotopic mass and abundance (%) of your first isotope and add that to the

relative isotopic mass multiplied by the abundance (%) of the other isotope and you divide that

number by 100 . In this case your answer would be : 152.04

Now, Relative Molecular Mass is easy

For example :

Mr (Cl2) = 35.5 x 2 = 71.0

Relative Formula Mass is just as easy

CaBr2 : 40.1 + (79.9x2) = 199.9

Amount of Substance - The Mole

The amount of substance is measured using a unit called the mole and given the symbol mol.

A mole of a substance is the amount of substance that contains 6x1023 particles,

which is the same number of particles as there are atoms in exactly 12 g of12C.

We call this Avogadro's Constant .

One mole of a substance is simply the unit of measurement used to express amounts of a

chemical substance, defined as an amount of a substance that contains as manyelementary entities (e.g. atoms, molecules, ions, electrons) as there are atoms in 12


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grams of pure carbon-12 (12C), the isotope of carbon with atomic weight 12. Thiscorresponds to a value of 6.02214179(30)1023 elementary entities of that substance.

The empirical formula is the simplest whole-number ratio of atoms of each elementpresent in a compound.

Molar mass is the mass per mole of a substance. Its unit is g mol-1

.

A chemical reaction starts with reactants and ends with products. An equation is a symbolicrepresentation of the reaction taking place.

MolesAmount of substance and the mole:

For SolidsAmount (mol) = mass(g)/molar mass M

Both Elements and Compounds

For gasesAmount of gas (mol) = volume(dm)/24 at RTP

or

Amount of gas (mol) = volume(cm)/24 000 at RTP

For Solutions...Amount(mol) = volume(dm3) x concentration (mol dm-3)

orAmount(mol)=[volume(cm3)/1000] x concentration(mol dm-3)

T pical Exam questions

Tip

All of these

equations can be

rearranged

EXAMPLE

What is the mass of 2 moles of calcium sulphate, CaSO4?

Molar mass of CaSO4 = 40.1+32.1+(16.0x4)= 136.2

Mass (g) = amount(mol) x M = 2 x 136.2=272.4g


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By Danii Cole

Acids and BasesAn acid is a proton donor. All acids contain hydrogen. When added towater, acids release this hydrogen as H+ ions (protons).

Whereas a base is a proton acceptor.

They neutralise acids, and will readily accept H+ ions,e.g. NH3 forming NH4

+.

Alkalis

An alkali is any chemical compound that gives a solution with a pH of greater than 7.0 when

dissolved in water. It is a type of base that dissolves in water forming hydroxide ions, OH- (aq) ions.

Common alkalis include sodium hydroxide (NaOH), Potassium hydroxide (KOH) and aqueous

ammonia (NH3).

Salts

Salts are formed from an acid when the H+ ion from the acid is replaced by a metal ion, or another

positive ion.

A salt is an ionic compound with the following features: the positive ion (cation) is usually a metal

ion or ammonium ion, NH4+

and a negative ion (anion) usually from an acid.

Salts can be produced by neutralising acids with carbonates, bases and alkalis.

When reacted with carbonates a salt, water and CO2 are formed.

On reacting with bases or alkalis a salt and water are formed.

Hydrochloric acid (HCl),

sulfuric acid (H2SO4) and

nitric acid (HNO3) are

common acids.

Bases include metal

oxides (e.g. MgO), metal

hydroxides

(e.g. NaOH) and

ammonia (NH3).


Did You Know

The word acid comes from the

Latin 'acidus' meanin sour.

acid + carbonate --> salt + H2O + CO2

acid + base --> salt + water

acid + alkali --> salt + water


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By Danii Cole

Water Of Crystallisation

Water of crystallisation refers to water molecules that form an essential part of the crystalline

structure of a compound.

hydrated = contains water molecules

anhydrous = form without water

A hydrate salt is composed of anions (negative ions) and cations (positive ions) which are

surrounded by and weakly bonded water molecules. Each hydrate salt has a fixed number of water

molecules associated with it, called waters of hydration or water of crystallization. When a salt

holds waters of hydration, we call it a hydrated salt or a hydrate (hydrate from hydor, the Greek

word for water).

Barium chloride dehydrate, BaCl22H2O, has two waters of crystallization, or two waters of

hydration. Other hydrates have waters of hydration ranging from one to twelve. Upon heating, a

hydrate decomposes and produces an anhydrous salt and water (in the form of steam).

BaCl22H2O (s ) BaCl2(s) + 2 H2O (g)

Solution: The unknown hydrate is 30.6 % water. subtracting from 100 %, the hydrate must be 69.4

% anhydrous salt. If we make the sample 100 g, the mass of water is 30.6 g and the anhydrous salt

69.4 g. to calculate the moles of water and the anhydrous salt (AS)

30.6 g H2O x 1 mole H2O = 1.70 moles H2O

18.0 g H2O

69.4 g AS x 1 mole AS = 0.283 mole AS

245 g AS

To find the water of crystallization, simply divide the mole ratio of water to anhydrous salt.

1.70 moles H2O = 6.01 6

0.283 mole AS

The water of crystallization is always a whole number, therefore the formula for the unknownhydrate is AS6H2O.

Titrations

Titration is a laboratory method of quantitative analysis used to determine unknown concentration of

a known substance.

Analysis is performed using a burette - kind of laboratory glass made for exact measurement of

volume of solution used and other glassware.

The most popular titrimetric experiment is a determination of amount of acid.

Imagine you have a solution of a sulfuric acid of unknown concentration.


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Pour exactly measured volume of sulfuric acid into a beaker and add few drops of alcoholic

phenolphthalein solution. Solution will be colorless, as phenolphthalein becomes pink only in basic

solutions (color becomes visible at pH above 8.2).

Now use burette to slowly add NaOH solution (called titrant) of known molar concentration. pH

slowly goes up. Once all sulfuric acid becomes neutralized one excess drop of strong base is enoughto rapidly change pH of the solution and change its color to pink.

Color change of phenolphthalein during titration - on the left, colorless solution before end point, on

the right - pink solution after end point.

When the colour of the solution changes you know that you have neutralized all acid present - you

have reached a titration end point. Using burette scale you may read volume of the titrant used.

We know that one mole of H2SO4 reacts with exactly two moles of NaOH:

2NaOH + H2SO4 Na2SO4 + 2H2O

Titrations

IndicatorColour in Acid Colour in base End point colour

Methyl orange

Red Yellow Orange

Bromothymol blueYellow Blue Green

PhenolpthaleinColourless Pink Pale Pink


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By Danii Cole

Electron Configurations

The 1st shell can hold 1 maximum of 2 electrons, the 2nd shell 8, the 3rd shell 18 and the4th shell can hold 32 electrons.

Shell Shell 1 Shell 2 Shell 3

A sub-shell is a group of the same type of atomic orbital's (s,p,d or f) within a shell

For example

Br has 35 Electrons: electron configuration is 1s 2s 2p6 3s 3p6 3d104s 4p5

An s orbital looks like this :

A p Orbital looks like this :

Sub-shell s s p s p d

Max no. of electrons

in sub-shell

2 2 6 2 6 10

5f

4f

6d

5d

4d

3d

6p

5p

4p

3p

2p

7s

6s

5s

4s

2s

1s

7p

3s

This diagram helps you to work out the order in which orbitals fill:1s, 2s, 2p, 3s, 3p, 4s, 3d, 4p, ..

However, it can be easier to read across the periodic table, but rememberthat the first transition metal row is 3d:

1s 1s

2s 2p

3s 3p

4s 3d 4p

Its centre is in the

nucleus . It is 3-dimentional. It has a

shape like a dumb-bell.


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By Danii Cole

Chemical Bonding :

There are three types of bonding you will have to know .It is essential that you know the definitions too .

Ionic Bonding

Bonding Usually found in compounds of a metal (Groups 1 and 2) with a non-metal (Groups5-7).

Ionic Bonding is the electrostatic attraction between oppositely charged ions.It appears when an electron is transferred from one atom to another, forming ions.

The metal ion is positive .

The non-metal ion is negative.

Example 1: A Group 1 metal + a Group 7 non-metal - e.g. sodium + chlorine ==> sodium chloride

NaCl or ionic formula Na+Cl

-In terms of electron arrangement, the sodium donates its outer electron

to a chlorine atom forming a single positive sodium ion and a single negative chloride ion. The atoms

have become stable ions, because electronically, sodium becomes like neon and chlorine like argon.

Na (2.8.1) + Cl (2.8.7) ==> Na+

(2.8) Cl-(2.8.8)

ONE combines with ONE to form

NaCl (Table Salt) is an ionic compound


Complete the electronic configuration of an atom of24Mg.

1s2.................................................................................... .........................................................

Complete the electronic configuration of a bromine atom.1s22s22p63s23p6........................................................................................................................

Why is bromine classified as a p-block element?....................................................................................................................................................


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By Danii Cole

The covalent bond

and ionic bond are

both very strong

chemical bonds.

Covalent Bonding

This is the bonding usually found in compounds of a non-metal with another non-metal.

Covalent Bonding is a bond formed by a shared pair of electrons.

We can show covalent bonding by "dot-and-cross" diagrams .

Dative Covalent Bonding

This is a particular type of covalent Bonding

In a dative covalent Bond, both the electrons in the bond are supplied by one atom only.Dative covalent bonds are sometimes called co-ordinate bonds.


CH4


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By Danii Cole

Oxidation and Reduction

Oxidation is the loss of electrons or an increase in oxidation number.Reduction is the gain of electrons or a decrease in oxidation number.

Simple definition :Oxidation is the gain of oxygenReduction is the loss of oxygen

OXIDATION REDUCTION

Gain/addition of oxygen by an atom, molecule orion e.g.

1. S ==> SO2o Burning sulphur - oxidised.

2. CH4==> CO2 + H2Oo Burning methane to water and

carbon dioxide, C and H bothgain O.

3. NO ==> NO2o Nitrogen monoxide oxidised to

nitrogen dioxide, rapid in air withNO from car exhaust fumes.

4. SO32-

==> SO42-

o Oxidising the sulphite ion to the

sulphate ion e.g. bychlorine/bromine oxidisingagents.

Loss/removal of oxygen from a compound or ione.g.

1. CuO ==> Cuo Loss of oxygen from copper(II)

oxide to form copper atoms inmetal extraction with or

displacement by more reactivemetal.

2. Fe2O3==> Feo Iron(III) oxide reduced to iron in

blast furnace using carbon orcarbon monoxide.

3. NO ==> N2o Nitrogen(II) oxide reduced to

nitrogen.4. SO3==> SO2

o Sulphur trioxide reduced tosulphur dioxide.

Loss/removal of electrons from atom, ion or

molecule e.g.

1. Fe ==> Fe2+ + 2e-o An iron atom loses 2 electrons to

form the iron(II) ion e.g. in theinitial chemistry of iron rusting orin an iron-acid reaction.

2. Fe2+==> Fe3+ + e-o The iron(II) ion loses 1 electron

to form the iron(III) ion, e.g. via

chlorine or manganate(VII)oxidising agents.

Gain/addition of electrons by an atom, ion, or

molecule e.g.

1. Cu2+ + 2e-==> Cuo The copper(II) ion gains 2

electrons to form neutral copperatoms

2. Fe3+ + e- ==> Fe2+o The iron(III) ion gains an electron

and is reduced to the iron(II) ione.g. by adding zinc to acidified

iron(III) salt solution.

An oxidising agent is the species that givesthe oxygen or removes the electrons

A reducing agent is the species that removesthe oxygen or acts as the electron donor

REDOX REACTIONS - in an overall reaction, oxidation and reduction must go together e.g. interms of oxygen and these are complete and balanced equations


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By Danii Cole

1. copper(II) oxide + hydrogen ==> copper + watero CuO(s) + H2(g)==> Cu(s) + H2O(g)o Heated copper oxide is reduced to copper (O loss) when hydrogen is passed over it,o hydrogen is oxidised to water (O gain),o hydrogen is the reducing agent (removes O from CuO),o and copper oxide is the oxidising agent (donates O to hydrogen).

2. iron(III) oxide + carbon monoxide ==> iron + carbon dioxideo Fe2O3(s) + 3CO(g)==> 2Fe(l) + 3CO2(g)o In the blast furnace the iron(III) oxide is reduced to liquid iron (O loss),o the carbon monoxide is oxidised to carbon dioxide (O gain),o CO is the reducing agent (O remover/acceptor from Fe2O3),o and Fe2O3 is the oxidising agent (O donator to CO)]

Oxidation Number

An oxidation number is a measure of the number of electrons that an atom uses to

bond with atoms of another element . Oxidation numbers are worked out from a set ofrules.

An oxidation number indicates the formalcharge of a chemically combined particle in acompound.

The oxidation number of metals usually equals the group number (as a positive value) andminus (8 group number) for non-metals.

An element has been oxidised if the oxidation number increases, and reduced if theoxidation number decreases.

When they react, metals are normally oxidised (they lose electrons), whereas non-metals

gain electrons and are reduced.

Electrochemical reactions involve the transfer of electrons. Mass and charge are conservedwhen balancing these reactions, but you need to know which atoms are oxidized and whichatoms are reduced during the reaction. Oxidation numbers are used to keep track of howmany electrons are lost or gained by each atom. These oxidation numbers are assignedusing the following rules:

The way it works is that the cation is written first in a formula, followed by the anion.

For example, in NaH, the H is H-; in HCl, the H is H+.

The oxidation number of a free element is always 0.The atom in He, for example, has an oxidation numbers of 0.

The oxidation number of a monatomic ion equals the charge of the ion.For example, the oxidation number of Na+ is +1; the oxidation number of N3- is -3.

The usual oxidation number of hydrogen is +1.

The oxidation number of hydrogen is -1 in compounds containing elements that are lesselectronegative than hydrogen, as in CaH2.

The oxidation number of oxygen in compounds is usually -2.

Exceptions include OF2, since F is more electronegative than O, and BaO2, due to thestructure of the peroxide ion, which is [O-O]2-.


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The oxidation number of a Group IA element in a compound is +1.The oxidation number of a Group IIA element in a compound is +2.

The oxidation number of a Group VIIA element in a compound is -1, except when thatelement is combined with one having a higher Electronegativity.

The oxidation number of Cl is -1 in HCl, but the oxidation number of Cl is +1 in HOCl.

The sum of the oxidation numbers of all of the atoms in a neutral compound is 0.

The sum of the oxidation numbers in a polyatomic ion is equal to the charge of the ion.

For example, the sum of the oxidation numbers for SO42- is -2.

Ionisation Energy

The first ionisation energy of an element is the energy required to remove one electron

from each atom in one mole of gaseous atoms to form one mole of gaseous 1+ ions.The second ionisation energy of an element is the energy required to remove one electronfrom each ion in one mole of gaseous 1+ ions to form one mole of gaseous 2+ ions.Successive ionisation energies can tell us which group an element is from . It simply meansthe first, second, third etc. ionisation energies for one particular element .

General Trend

1st ionisation energy (across period)Increases because the nuclear charge increasesAtoms get smaller

The electron shielding and atomic radius remain similar across a period.

1st ionisation energy (down group)Decreases because the atomic radius and electron shielding both increase.Atoms Get BiggerMore ShieldingHowever, as you go down the Group, the distance between the nucleus and the outerelectrons increases and so they become easier to remove - the ionisation energy falls.

Note that 2nd ionisation energy is the energy required to remove the second electron (not

both electrons).

e.g. 1st IE of Na: Na (g) Na+(g) + e-2nd IE of Na: Na+(g) Na2+(g) + e-

3rd IE of Na: Na2+ (g) Na3+(g) + e-



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By Danii Cole

Shapes of Molecules

The shape of a molecule is determined by the repulsion between bonded electron pairs and non-

bonded electron pairs (lone pairs).

Lone electron pairs repel more than bonded pairs of electrons and give rise to distortedshapes.

By reducing the number of bonded electron pairs and lone pairs of electrons, the shape of a

molecule may be predicted.

BF3 is trigonal planar; CH4 and NH4+ are tetrahedral; SF6 is octahedral; H2O is non-linear (V-

shaped/bent); CO2 is linear and ammonia, NH3, as pyramidal

The examples on this page are all simple in the sense that they only contain two sorts of atoms

joined by single bonds - for example, ammonia only contains a nitrogen atom joined to three

hydrogen atoms by single bonds. If you are given a more complicated example, look carefully at the

arrangement of the atoms before you start to make sure that there are only single bonds present.

For example, if you had a molecule such as COCl2, you would need to work out its structure, based

on the fact that you know that carbon forms 4 covalent bonds, oxygen 2, and chlorine (normally) 1.

If you did that, you would find that the carbon is joined to the oxygen by a double bond, and to the

two chlorines by single bonds.

The Valence Shell electron pair repulsion theory (VSEPRT)

Assumes that each atom in a molecule will be positioned so that there is minimal repulsion between

the valence electrons of that atom.

Examplewater, H2O

Water is composed of 3 atoms, 1 atom of oxygen and 2 atoms of hydrogen covalently

bonded.

Oxygen, O, is the central atom.

Oxygen has 6 valence electrons.2 of oxygen's electrons will be used to form bonding pairs of electrons with hydrogen (2

covalent bonds).4 of oxygen's valence electrons will not be used for bonding, these will remain as 2 lonepairs of electrons.lone pair-lone pair repulsion is greater than lone pair-bonding pair or bonding pair-

bonding pair repulsion, so the lone pairs of electrons push the bonding pairs of electronscloser together than in a tetrahedral arrangement of the 'electron clouds'.This distorted tetrahedral arrangement is called bent.

H2O is bent in shape.


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By Danii Cole

Electronegativityis the ability of an atom to attract the bonding electrons in a

covalent bond.

General TrendsAcross a period: the electronegativities generally increase from left to right across a period with theGroup VII element having the highest value for the period.

Down a group: the electronegativities generally decrease from top to bottom down a group.

Francium is the element with the lowest electronegativity.

Fluorine, the most reactive non-metal, is assigned the highest value since it has the greatest

attraction for the electron being shared by the other element.

Oxygen is also highly electronegative and has a strong attraction for electrons.

Metals have low electronegativities since they have weak attraction for any shared electrons.

When two unlike atoms are covalently bonded, the shared electrons will be more strongly attracted

to the atom of greater electronegativity. This is why a bond is said to be polar.

A polar bond results in the unequal sharing of the electrons in the bond.

The presence or absence of polar bonds within molecule plays a very important part in determining

chemical and physical properties of those molecules.

Some of these properties are melting points, boiling points, viscosity and solubility's in solvents.

The most commonly used electronegativity scale is Pauling's.

Typical Exam question

Define the termelectronegativity.

......................................................................................................



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By Danii Cole

Metallic bonding is the electrostatic attraction

between positive metal ions and delocalised

electrons.

Intermolecular ForcesAn intermolecular force is an attractive force between neighbouring molecules.

Permanent Dipole-Dipole forceWeak attractive force between permanent dipoles in neighbouring molecules .The facts :

Polar molecules have permanent dipoles

Van der waal's forcesAttractive forces between induced dipoles in neighbouring molecules .The facts :

exist between all molecules (polar or non-polar) van der waals' forces increase with increasing number of electrons

Hydrogen Bonding

A hydrogen bond is a strong dipole-dipole attraction between an electron-deficient Hydrogen atom

(O-H+ or N-H+) on one molecule and a lone pair of electrons on a highly electronegative atom

(H-O:-OR H-N:

-) on a different molecule.

In simpler terms:It is a weak bond between two molecules resulting from an electrostatic attraction between aproton in one molecule and an electronegative atom in the other.

Hydrogen bonding is strong enough to give water some anomalous properties: Firstly, ice is lessdense than water because ice has an open lattice with hydrogen bonds holding the water molecules

apart. It has relatively high melting and boiling point due to these extra forces and the extra

intermolecular bonding also explains the relatively high surface tension and viscosity of water.

Metallic Bonding

Water is very strange .

The hydrogen bond are extra forces, over

and above the van der waals' forces.

Typical Exam question

State and explain two anomalous

properties of water resulting from

hydrogen Bonding.

This question is worth 2 marks . So youwill need to give 2 properties with an

explanation for each property .


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By Danii Cole

The delocalised electrons are shared between more than two atoms in the metallic structure,

whilst the positive ions occupy fixed positions.

Metallic bonding is strong meaning that metals have a high melting/boiling point. Only some,some have not . Such as : The metal with the lowest melting point is mercury - the only metallic

element with a melting point (-40oC) below room temperature (20

oC). Gallium melts at 86

oF (30

oC)

which would be liquid on a warm day or in your hand.

Metals are also good conductors of electricity, due to the fact that the delocalised electrons can

move freely and can carry a charge.

Metals are ductile and malleable. Meaning they can be hammered into shape and can be drawn

out or stretched. The delocalised electrons are the reason for these properties . They can move

and atoms can slide past each other .

A giant metallic lattice is a three-dimensional structure of positive ions and delocalised electrons,bonded together by strong metallic bonds.

Ionic CompoundsStructures with ionic bonding are known as giant ionic lattices, with each ion being surrounded by

oppositely charged ions. All ionic compounds exist as a giant ionic lattice in the solid state. Ionic

compounds have a high melting and boiling point due to the strong electrostatic forces that hold the

oppositely charged ions together. In a solid ionic lattice the ions are in fixed positions and cannot

move, so do not conduct electricity, but when an ionic compound is melted or dissolved the solid

lattice breaks down and the ions can move so can thus conduct electricity. Some Ionic lattices

dissolve inpolarsolvents, such as water because the solvent surrounds each ion to form a solution.

Covalent CompoundsElements and compounds with covalent bonds are either simple molecular lattices or giant covalent

lattices. Simple molecular structures are made up from small, simple molecules such as Ne, H2 and

H2O. In a simple molecular lattice molecules are held together by weak van der waals' with the

atoms within molecules being strongly bonded together by covalent bonds. Simple molecular

structures have lowmelting and boiling points because the van der Waals forces are weak; so little

energy is needed to break them. They do not conduct electricity because there are no charged

particles free to move. They are soluble in non-polarsolvents (e.g. Hexane) because van der Waals

forces form between the structure and the solvent weakening the lattice structure.Giant covalent structures are a three-dimensional structure of atoms, bonded together by strongcovalent bonds. They have high melting and boiling points because high temperatures are needed to

break the strong covalent bonds. Aside from graphite there are no free charged particles so they do

not conduct electricity. They are insoluble in both polar and non-polar solvents because the

covalent bonds are too strong to be broken by polar or non-polar solvent. Diamond has a

tetrahedral structure held together by covalent bonds. It does not conduct but is very hard. Graphite

has a strong hexagonal layer structure but only weak van der Waals forces in-between layers so the

layers slide easily. It is however, a good conductor of electricity due to the presence of delocalised

electrons between layers.


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By Danii Cole

Module 3 : The Periodic Table

PeriodicityThe periodic table is arranged by increasing atomic number.

Elements often show trends (gradual changes) in properties across a period in the table these

trends are repeated across each period.

Each vertical column is called a group and contains elements with similar properties and the same

number of outer-shell electrons.

Elements in the same group react in a similar way because they have similar electron configuration.

Across each period there is a general increase in the number of protons so there is more attraction

acting on the electrons. Meaning that the first ionisation energy increase across a period.

Electrons are added to the same shell, so the outer shell is drawn inwards slightly, reducing the

atomic radius, and the electron shielding will not change.

At the start of a next period, a new shell is formed, increasing the distance of the outermost shell

from the nucleus and increasing the electron shielding of the outermost shell by inner shell.

Going down a group the first ionisation energies decreases due to the increase in the number of

shells, shielding increases and atomic radius increase and the distance of the outer electrons from

the nucleus.

In periods 2 and 3 there is a clear change between groups 4 and 5, in both the physical structure of

the elements and the forces holding the structures together.

Johns Newlands

First Person toarrange elements in

order of their relative

atomic masses.

Aristotle

Ancient Greekphilosopher who

thought that

everything was made

up of just four

elements .

Lavoisier

Put together the firstextensive list of

elements.

Dmitri Mendeleev

Credited as being thecreator of the first

version of the

periodic table of

elements. Using the

table, he predicted

the properties of

elements yet to be

discovered.


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By Danii Cole

This change is: from giant structures to simple molecular and from strong to weak forces. The

bonding changes from metallic to covalent for C and Si to van der Waals for groups 5-0.

Group 2 Elements - The alkaline Earth MetalsMagnesium, Calcium, strontium and barium.

These elements all react with water to form a solution of the hydroxide and hydrogen gas. These

elements react with oxygen to form the oxide. The alkaline earth metals all have hydroxides which

are alkaline.

They are good conductors of electricity because they are metallically bonded .

The atomic radii increase going down the group because the size of the atom increase as a new shell

of electrons is present with each element .

All group 2 elements lost their two outer electrons when they react . They form cations with a +2

charge. They are ALL oxidised when they react.

Calcium, strontium and barium oxides react with water to form hydroxides:

CaO(s) + H2O(l) --------> Ca(OH)2(s)

Calcium hydroxide is known as slaked lime. It is soluble in water and the resulting mildly alkaline

solution is known as lime water which is used to test for the acidic gas carbon dioxide.

In all their compounds these metals have an oxidation number of +2 and, with few exceptions, their

compounds are ionic. They are all ionic except beryllium chloride.

Reactivity increases as you go down the group because the outer two electrons are further from the

nucleus and are less shielded.

Halides

The Group 2 halides are normally found in the hydrated form. Anhydrous calcium chloride has such a

strong affinity for water it is used as a drying agent.

Metal hydroxides are weak alkalis and typically have a pH between 8 and 11.

Group 2 carbonates decompose with greater difficulty as the group is descended, to form the metal

oxide and carbon dioxide gas.

They have reasonably high melting/boiling points, low densities and form colourless compounds.

These properties are due largely to the presence of two valence electrons on each atom, which leads

to stronger metallic bonding than Group 1.

Three of these elements give characteristic colours when heated in a flame:

Mg brilliant white Ca brick-red Sr crimson Ba apple green

They each have two electrons in their outer shell and are strong reducing agents and reactive

metals. They are oxidised in reactions to form a 2+ ion.

Below I have included data that shows these trends in the period table.


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By Danii Cole

They all react vigorously with oxygen in a redox reaction. They all react with water to form

hydroxides (typically pH 10-12). Mg reacts very slowly, as you go down the group each metal reacts

more vigorously with water.

The solubility of the hydroxides in water increases down the group. The resulting solutions are also

more alkaline. E.g. Mg (OH)2 is only slightly soluble and forms a dilute solution. Ba (OH)2 is muchmore soluble in water, and forms a more alkaline solution.

The group 2 carbonates are decomposed by heat (thermal decomposition). The carbonates become

more difficult to decompose with heat as you move down the group.

All the metals except beryllium form oxides in air at room temperature which dulls the surface of the

metal. Barium is so reactive it is stored under oil.

Data

Atomic Number

Relative Atomic

Mass Melting Point/K

Be 4 9.012 1551

Mg 12 24.31 922

Ca 20 40.08 1112

Sr 38 87.62 1042

Ba 56 137.33 1002

Ionisation Energies/kJ mol-1

1st 2nd 3rd

Be 899.4 1757.1 14848

Mg 737.7 1450.7 7732.6

Ca 589.7 1145 4910

Sr 549.5 1064.2 4210

Ba 502.8 965.1 3600

Group 7 Elements - The Halogens

The only elements you need to know about are chlorine, bromine and iodine.

Chlorine 3s

2

3p

5

bromine 4s2 4p5iodine 5s25p5

All halogens exist as diatomic molecules in which van der Waalsintermolecular forces act between

the molecules.

Halogens dissolve in organic solvents, like hexane, to form characteristic colours, for example, iodine

forms a purple solution.

Melting Point (C)

1287

648

838

769

728


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Halogen atoms gain one electron to form halide ions, X-, and this ability becomes easier on moving

up the group. Halogen atoms become larger on descending the group, so a gained electron is only

weakly attracted due to greater shielding.

The Halogens have low melting and boiling points and exist as diatomic molecules. Because there is

an increase in electrons as you go down the group, the van der Waals forces are consequentlystronger and the boiling points increase going down the group.

The oxidising power of a halogen is a measure of the strength with which a halogen atom is able to

attract and capture an electron to form a halide ion.

The halogens become less reactive down the group as their oxidising power decreases as atomic

radii and electron shielding increases.

Redox reactions can show this by using the reactions of aqueous solutions of halide ions and

halogens in a displacement reaction.

Chlorine oxidises both bromide and iodide ions. Bromine oxidises iodide only. Iodine does neither.

And that in a nutshell is it . I hope you have enjoyed reading this.Good Luck !

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