Unit 1 Chemical changes and structure Dyslexia.pptx

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• Chemistry
• Upper Secondary (Key Stage 4)
• BTEC, GCSE

Chemical Changes and Structure

Unit 1

Periodicity

1. Elements are arranged in the periodic table in order of increasing atomic number.

1. The periodic table allows chemists to make accurate predictions of physical properties and chemical behaviour for any element, based on its position.

1. Groups: vertical columns within the table contain elements with similar chemical properties resulting from a common number of electrons in the outer shell

1. Periods: rows of elements arranged with increasing atomic number, demonstrating an increasing number of outer electrons and a move from metallic to non-metallic characteristics

Periodic Pattern - Atomic Size

There is no definite

edge to an atom.

However, bond lengths

can be worked out.

Covalent radius,

½ the distance between nuclei.

Values can be found on pg 7 of the data booklet

The atomic size of the elements in any period decrease with increasing atomic number.

The atomic size of the elements in any group increase with increasing atomic number.

Going across a period, the attraction between the outer shell and the positive nucleus increases due to the increasing nuclear charge.

Going down a group, the attraction between the outer shell and the positive nucleus decreases due to the increased distance and more occupied electron shells ‘shielding’ the nuclear charge.

Ionisation energy

This is defined as "the amount of energy required to remove one mole of electrons from one mole of gaseous atoms (only for 1st I.E) or ions.”

Values can be found on pg 12 of the data booklet

M (g)  M+ (g) + e- 1st ionisation

M+ (g)  M 2+ (g) + e- 2nd ionisation

M 2+(g)  M 3+ (g) + e- 3rd ionisation

Periodic pattern - Ionisation energy

1st Ionisation energy decreases going down a group, as the attraction between the outer negative electron and the positive nucleus decreases due to the increased distance and more occupied electron shells ‘shielding’ the nuclear charge.

The general increase in 1st ionisation energy that occurs as we move across a period is due to the increase in nuclear charge, its attraction for the outermost electron/s increases and consequently ionisation energy increases.

Going across the period, the nuclear charge increases. As a result, the Electronegativity increases.

Going down the group, the nuclear charge increases but the number of electron shells also increases. As a result of ‘shielding’ and an increased distance the outer shell is from the nucleus, Electronegativity decreases.

Periodic Patterns - Electronegativity

Electronegativity is the measure of attraction of an atom for the bonding electrons. Electronegativity is useful at predicting how electrons will be shared. The Pauling scale is used for Electronegativity values the values are found on page 12 of the data booklet.

Structure and bonding

Types of chemical bond

Metallic bonding (only in elements)

The outer shell in metals is not full and so metal electrons can move randomly between these partially filled outer shells. This creates what is sometimes called a ‘sea’ or ‘cloud’ of electrons.

The greater the number of electrons in the outer shell the stronger the metallic bond. So order of decreasing melting point Al>Mg>Na

Strong electrostatic forces exist between the positive nuclei (ions) and the delocalised outer shell electrons.

These electrostatic attractions are known as metallic bonds.

Increasing difference in electronegativity

Equal sharing

of electrons

Increasing unequal sharing of electrons

Transfer of electrons

Non-polar slightly polar covalent very polar covalent ionic

4.0 F

F4.0

1.0 Li

F 4.0

3.4 O

H 2.2

The greater the difference in electronegativity the greater the polarity between two bonding atoms and the more ionic in character.

Bonding continuum

Covalent bonding

A covalent bond is formed when a pair of electrons are shared. The atoms in a covalent bond are held together by electrostatic forces of attraction between positively charged nuclei and negatively charged electrons.

Non polar (pure) covalent bonding

• In a non polar covalent bond, the bonding electrons are shared equally between the two atoms, there are no charges on the atoms.

Polar Covalent Bond –only in compounds

In a polar covalent bond the bonding electrons are not shared equally between the two atoms, this is due to a difference in electronegativities of the atoms.

The chlorine nucleus has more protons and has a stronger pull and thus greater share of the bonding electrons than the hydrogen nucleus.

The atoms will have partial charges, the hydrogen slightly positive(δ+), the chlorine slightly negative(δ-).

The dipole produced is permanent.

Covalent bonding - molecular

A covalent molecular structure consists of discrete molecules held together by weak molecular forces

1. low melting and boiling points as only the weak intermolecular forces are being broken

1. Non conductors

Covalent Network

A covalent network structure consists of a giant lattice of covalently bonded atoms

1. High melting and boiling points as the strong covalent bonds are being broken

1. Non conductors except carbon in the form of graphite

Ionic bonds are formed between atoms with a large difference in electronegativities. There is a transfer of electrons which produces oppositely charged ions.

They are often (though not always) between metals and non-metals.

An ionic bond is the electrostatic force of attraction between positively and negatively charged ions.

1. An ionic structure consists of a giant lattice of oppositely charged ions.

2. High melting and boiling points as the strong ionic bonds are being broken

3. Conduct in solution and as a liquid

Ionic Bonds

The First twenty elements bonding and structure

Metallic Elements

Exist as a giant lattice of positively charged ions and delocalised outer electrons.

Covalent Network Elements

Boron, Carbon ( in the form of diamond & graphite ) and Silicon exist as giant lattices of covalently bonded atoms.

Covalent molecular - Gases

They consist of discrete (separate) molecules

London Dispersion forces exist between the molecules of the gases

Covalent molecular - solids

They consist of discrete (separate) molecules

Stronger London Dispersion forces exist between the molecules making them solids at room temperature

Monatomic elements

The Noble gases

• They consist of discrete (separate) atoms

• London Dispersion forces exist between the atoms of nobles gases

Carbon – covalent networks

• In diamond, each carbon atom forms four covalent bonds

• In graphite, each carbon atom forms three covalent bonds to form layers which can slide over each other. Delocalised electrons in graphite allow it to conduct electricity.

Fullerenes, molecules of carbon

Fullerenes are a molecular form of carbon, which are made of balls, 'cages' or tubes of carbon atoms.

Buckminster fullerene is one type of fullerene. Its molecules have 60 carbon atoms arranged in a hollow sphere.

Due to the large molecules , fullerenes have stronger London dispersion forces between their molecules, compared to elements made from smaller molecules.

Intermolecular forces of attraction

VAN DER WAALS’ FORCES

1. All molecular elements and compounds and monatomic elements will condense and freeze at sufficiently low temperatures.

1. For this to occur, some attractive forces must exist between the molecules or discrete atoms.

1. Any "intermolecular" force acting between molecules are known as Van der Waals' forces.

1. There are several different types of Van der Waals' forces such as

   1. London dispersion forces

   2. Permanent dipole-permanent dipole interactions which includes hydrogen bonding.

London Dispersion forces

These are very weak attractions between atoms/molecules and are caused when electrons are not equally distributed around the atom/molecule.

This causes temporary polarisation of the atom/molecule so that one end becomes slightly negative (δ-) and the other end becomes slightly positive (δ+).

The atom/molecule is called a temporary dipole.

The size of the London Dispersion force and hence the melting point depends on the number of electrons in the molecule, so the bigger the molecule, the bigger the London Dispersion force .

Permanent dipole to permanent dipole interactions

The differing electronegativities of different atoms in a molecule and the spatial arrangement of polar covalent bonds can cause a molecule to form a permanent dipole.

No permanent dipole

Symmetrical molecule

Permanent dipole

Asymmetrical molecule

NON-POLAR molecule

Only has London dispersion

Forces between the molecules

POLAR molecule

Permanent dipole-permanent dipole interactions are additional forces of attraction between polar molecules, these are stronger than London dispersion forces.

Hydrogen Bonding

Hydrogen bonding is a special type of permanent dipole to permanent dipole interaction, it is the strongest of the van der waals forces.

Hydrogen bonding occurs when hydrogen is covalently bonded to a highly electronegative element, either Fluorine, Oxygen or Nitrogen

Properties

Physical properties of Polar and non polar molecules – boiling point

Both methanal and ethane have the same formula mass of 30 however, methanal boils at – 21 oC while ethane boils at – 89oC

Ethane has no permanent dipoles, so only London Dispersion forces, temporary dipoles attractions between molecules. So has a lower boiling point.

Methanal is a polar molecule as it has a permanent dipole, so has permanent dipole-permanent dipole interactions as well as London Dispersion forces between molecules.

N.B. you can only compare compounds with similar molecular masses, as their London dispersion forces will also be similar

Physical properties of hydrides – boiling point

Water has a much higher b.p. than similar compounds containing hydrogen e.g. NH3, HF, CH4.

Hydrogen bonding explains

why water has a b.p. higher

than expected. It also explains

waters viscosity.

Similarly with HF b.p. 19 oC

Whereas:

HCl –85 oC

HBr –68 oC

HI –35 oC

Hydrogen bonding & density of water

Hydrogen bonding between molecules in ice results in an expanded structure that causes the density of ice to be less than that of water at low temperatures.

Solubility - Ionic

The polar water molecules are attracted to the ions in the crystal lattice.

The attraction drags the ions out of the crystal ‘into solution’, where the ions are surrounded by water molecules.

Non-polar compounds, like tetrachloromethane do not dissolve in water.

Non polar molecules cannot hydrogen bond with water molecules and if they were to go into water, they would separate the water molecules and reduce the number of hydrogen bonds.

This is energetically unfavourable so they don’t dissolve.

Solubility – Non polar molecules

Non-polar compounds are miscible with each other because the only attractions present are london dispersion forces and they can exist between different non polar molecules.

eg tetrachloromethane and benzene (two non-polar molecules) are miscible.

Small polar molecules, like ethanol or ethanoic acid, dissolve in water because the polar functional groups are able to hydrogen bond with water molecules.

However, as we go up their homologous series, the increasing length of the non-polar hydrocarbon chain means that higher members are insoluble in water.

Some polar covalent molecules ionise when they dissolve in water: δ+H—Clδ– (g)  H+(aq) + Cl−(aq)

Solubility – Polar molecules

Solubility summary

Ionic compounds and polar molecular compounds tend to be soluble in polar solvents such as water and insoluble in nonpolar solvents.

Non-polar molecular substances tend to be soluble in non-polar solvents and insoluble in polar solvents.

Remember like dissolves like

Oxidising and reducing agents

Redox Equations - Revision

O.I.L. R.I.G.

Oxidation is loss of electrons

Reduction is gain of electrons

Redox reactions

Redox reactions include reactions which involve the loss or gain of electrons.

Both oxidation and reduction happen simultaneously, however each is considered separately using ion-electron equations.

Cu(s)  Cu2+(s) + 2e- Oxidation

O2(g) + 4e-  2O2-(g) Reduction

Cu(s) + O2(g)  CuO (s)

Writing Redox Reaction

Cu(s)  Cu2+(s) + 2e- Oxidation

O2(g) + 4e-  2O2-(g) Reduction

_____________________

Balance the number of electrons

2Cu(s)  2Cu2+(s) + 4e-

O2(g) + 4e-  2O2-(g)

_____________________

Cancel the electrons and combine

2Cu(s) + O2(g)  2Cu2+(s) + 2O2-(g)

Reducing & oxidising agents

The reactant giving away (donating) electrons is called the reducing agent (which is oxidised)

Cu(s)  Cu2+(s) + 2e-

The reactant taking (accepting) electrons is called the oxidising agent (which is reduced)

O2(g) + 4e-  2O2-(g)

1. The electrochemical series indicates the effectiveness of oxidising and reducing agents.

1. Oxidising agents can only react with a reducing agent which is above it in the ECS.

1. For example, an acid (H+ions) will react with lead atoms but not with silver atoms

1. The elements with low electronegativities (metals) tend to form ions by losing electrons (oxidation) and so can act as reducing agents

1. The elements with high electronegativities (non-metals) tend to form ions by gaining electrons (reduction) and so can act as oxidising agents.

1. The strongest reducing agents are found in Group 1 whilst the strongest oxidising agents come from Group 7.

2. Compounds, group ions and molecules can act as oxidising or reducing agents:

   1. hydrogen peroxide is a molecule that is an oxidising agent

   2. dichromate and permanganate ions are group ions that are strong oxidising agents in acidic solutions

   3. carbon monoxide is a gas that can be used as a reducing agent

Writing Red or ox equations

1. Make sure there are the same number of atoms of each element being oxidised or reduced on each side of the half equation. (ignore oxygen)

2. If there are any oxygen atoms present, balance them by adding water molecules to the other side of the half-equation.

3. If there are any hydrogen atoms present, balance them by adding hydrogen ions on the other side of the half-equation.

4. Make sure the half-reactions have the same overall charge on each side by adding electrons.

Uses of Oxidising agents

1. Oxidising agents are widely employed because of the effectiveness with which they can kill fungi and bacteria and can inactivate viruses.

1. The oxidation process is also an effective means of breaking down coloured compounds making oxidising agents ideal for use as ‘bleach’ for clothes and hair.

1. A redox titration is the same as an acid-base titration except it involves a redox reaction

1. In titrations using acidified permanganate, an indicator is not required, as purple permanganate solution turns colourless when reduced.

1. H+ ions in an equation show the reactant has to be acidified

Redox Titrations

What volume of 0.40mol L-1 iron(II) sulphate solution is needed to react completely with 25 cm3 of 0.10mol L-1 potassium manganate(VII) solution?

5Fe2+(aq) + MnO4-(aq) + 8H+(aq) Mn2+(aq) + 4H2O(l) + 5Fe3+(aq)

5 moles 1mole

n= cxv

= 0.1x0.025

= 0.0025 moles

n= 0.025x5

= 0.125

V=n/c

= 0.125/0.4

= 0.3125L

1. A redox reaction will occur if a solution containing iron(II) ions is added to a solution containing manganate(VII) ions (sometimes called permanganate ions).
   

2. The two half reaction equations are:-

Fe2+(aq)  Fe3+(aq) + e-

and
MnO4-(aq) + 8H+(aq) + 5e-  Mn2+(aq) + 4H2O(l)

1. To combine these equations multiply the oxidation half reaction by 5 to make the electrons balance and then add the equations.

5Fe2+(aq) + MnO4-(aq) + 8H+(aq) Mn2+(aq) + 4H2O(l) + 5Fe3+(aq)

1. This equation allows us to see the ratio in which Fe2+ ions react with MnO4- as shown below.

5 moles Fe2+ ions reacts with 1 mole MnO4- ions.

Redox Calculations exercise

Q1, 2 + 3

1. No indicator required as there is a colour change in the reactants

1. H+ ions in the equation show the reactant has to be acidified