Structure & Bonding (Triple) quiz

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2020-02-16T12:32:52+00:00Categories: Uncategorized|Tags: , |

Structure & Bonding (Double) quiz

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Flashcards: Structure & bonding

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Q&A slides – Structure & Bonding

2019-11-12T07:39:33+00:00Categories: Uncategorized|Tags: , , , |

Dogs teach ionic versus covalent bonding – video

This should give you a ruff idea of what the difference is between ionic and covalent bonding.

1:40 draw dot-and-cross diagrams to show the formation of ionic compounds by electron transfer, limited to combinations of elements from Groups 1, 2, 3 and 5, 6, 7 only outer electrons need be shown

Sodium chloride, NaCl

Magnesium chloride, MgCl2

Potassium oxide, K2O


Calcium oxide, CaO

Aluminium oxide, Al2O3
Magnesium nitride, Mg3N2

1:41 understand ionic bonding in terms of electrostatic attractions

Ionic bonding: a strong electrostatic attraction between oppositely charged ions.

1:42 understand why compounds with giant ionic lattices have high melting and boiling points

Ionic compounds have high melting and boiling points because they have a giant structure with strong electrostatic forces between oppositely charged ions that require a lot of energy to break.


Giant 3D lattice of sodium chloride (NaCl)

1:43 Know that ionic compounds do not conduct electricity when solid, but do conduct electricity when molten and in aqueous solution

Ionic compounds do not conduct electricity when solid.

However, ionic compounds do conduct electricity if molten or in solution.



Ionic bonding – Tyler de Witt videos

These videos are a really good introduction to ionic bonding:

2019-02-10T14:36:34+00:00Categories: Uncategorized|Tags: , , , , |

1:44 know that a covalent bond is formed between atoms by the sharing of a pair of electrons

A covalent bond is formed between two non-metal atoms by sharing a pair of electrons in order to fill the outer shell.

1:45 understand covalent bonds in terms of electrostatic attractions

Covalent bonding: a strong attraction between a shared pair of electrons and two nuclei.

1:46 understand how to use dot-and-cross diagrams to represent covalent bonds in: diatomic molecules, including hydrogen, oxygen, nitrogen, halogens and hydrogen halides, inorganic molecules including water, ammonia and carbon dioxide, organic molecules containing up to two carbon atoms, including methane, ethane, ethene and those containing halogen atoms

1:47 explain why substances with a simple molecular structures are gases or liquids, or solids with low melting and boiling points. The term intermolecular forces of attraction can be used to represent all forces between molecules


Carbon dioxide (CO2) has a simple molecular structure. This just means that it is made up of molecules.

Within each molecule are atoms bonded to each other covalently. These covalent bonds inside the molecules are strong.

However, between the molecules are weak forces of attraction that require little energy to break. These forces are not covalent bonds. This is why simple molecular substances such as carbon dioxide have a low boiling point.

So when carbon dioxide changes from a solid to a gas, for example, that process can be represented as:

CO₂ (s) → CO₂ (g)

Notice that even though there has been a dramatic change of state from solid to gas, the substance before and after the change is always made up of carbon dioxide molecules. During the change of the state the covalent bonds within each molecule remain unbroken. Instead it is the weak forces of attraction between the molecules which have been overcome.


1:48 explain why the melting and boiling points of substances with simple molecular structures increase, in general, with increasing relative molecular mass

Larger molecules tend to have higher boiling points.

This is because larger molecules (molecules with more mass) have more forces of attraction between them. These forces, although weak, must be overcome if the substance is to boil, and larger molecules have more attractions which must be overcome.

1:49 explain why substances with giant covalent structures are solids with high melting and boiling points

Diamond has a high melting point because it is a giant covalent structure with many strong covalent bonds that require a lot of energy to break.

1:50 explain how the structures of diamond, graphite and C60 fullerene influence their physical properties, including electrical conductivity and hardness

Allotropes are different forms of the same element. Three different allotropes of carbon are shown here as examples: diamond, graphite and C60 fullerene.


Diamond is made up of only carbon atoms, where each of those atoms has a strong covalent bonds to 4 other carbon. Every one of carbon’s 4 outer electrons is involved in one of these strong covalent bonds.

Diamond is extremely hard because it is a giant covalent structure with many strong covalent bonds.

Because it is hard, diamond is used in high speed cutting tools, eg diamond-tipped saws.



Graphite is also made of only carbon atoms, and is also a giant structure, but it is formed of layers where each carbon atom has a strong covalent bond to 3 other carbons. This means each carbon atom has one electron not involved in a covalent bond, and these electrons form a sea of delocalised electrons between the layers.

Even though it is a non-metal, graphite can conduct electricity because the delocalised electrons are free to move.

Each layer is a giant structure, with weak forces of attraction between the layers. These layers can easily slide over each other.

Graphite is soft and slippery because it has weak forces of attraction between layers. It is used as a lubricant and in pencils because it is soft and slippery.




C60 fullerene (also known as a buckyball) is also made of only carbon atoms, but it forms molecules of 60 carbon atoms. The molecule has weak intermolecular forces of attraction between them which take little energy to overcome. Hence C60 fullerene has a low melting point, and it is soft.

C60 fullerene cannot conduct electricity. Although in each molecule every carbon is only covalently bonded to 3 others and the other electrons are delocalised, these electrons cannot jump between different molecules.



1:51 know that covalent compounds do not usually conduct electricity

Electric current is a flow of charged particles that can move.

Covalent compounds do not conduct electricity.

1:52 (Triple only) know how to represent a metallic lattice by a 2-D diagram

When metal atoms join together the outer electrons become ‘delocalised’ which means they are free to move throughout the whole structure.

Metals have a giant regular arrangement of layers of positive ions surrounded by a sea of delocalised electrons.

1:53 (Triple only) understand metallic bonding in terms of electrostatic attractions

Metallic bonding is the strong electrostatic attraction between positive metal ions and a sea of delocalised electrons.

1:54 (Triple only) explain typical physical properties of metals, including electrical conductivity and malleability

Metals are good conductors because they have delocalised electrons which are free to move.


Metals are malleable (can be hammered into shape) because they have layers of ions that can slide over each other.

(Triple only) Alloys and different forms of steel – video

This video explains the structure of alloys, including different forms of steel.

Note that a slight improvement could be made in the explanation: the structure of metals should be described as “layers of metal IONS in a sea of delocalised electrons”.

2019-11-05T20:08:54+00:00Categories: Uncategorized|Tags: , , , |

1:55 (Triple only) understand why covalent compounds do not conduct electricity

Electrical conductivity is the movement of charged particles.

In this case, charged particles means either delocalised electrons or ions.

These particles need to be free to move in a substance for that substance to be conductive.

Covalent compounds do not conduct electricity because there are no charged particles that are free to move.

1:56 (Triple only) understand why ionic compounds conduct electricity only when molten or in aqueous solution

Ionic compounds only conduct electricity only when molten or in solution.

When solid the ions are not free to move.


When molten or in solution the ions are free to move.

2:25 (Triple only) explain the uses of aluminium, copper, iron and steel in terms of their properties the types of steel will be limited to low-carbon (mild), high-carbon and stainless

AircraftLow density / resists corrosion
Power cablesConducts electricity / ductile
Pots and pansLow density / strong (when alloyed) / good conductor of electricity and heat

Aluminium resists corrosion because it has a very thin, but very strong, layer of aluminium oxide on the surface.

Electrical wiresvery good conductor of electricity and ductile
Pots and pansvery good conductor of heat / very unreactive / malleable
Water pipesunreactive / malleable
Surfaces in hospitalsantimicrobial properties / malleable
SaucepansConducts heat / high melting point / malleable
Type of steelIron mixed withSome uses
Mild steelup to 0.25% carbonnails, car bodies, ship building, girders
High-carbon steel0.6%-1.2% carboncutting tools, masonry nails
Stainless steelChromium (and nickel)cutlery, cooking utensils, kitchen sinks

Mild steel is a strong material that can easily be hammered into various shapes (malleable). It rusts easily.

High-carbon steel is harder than mild steel but more brittle (not as malleable).

Stainless steel forms a strong, protective oxide layer so is very resistant to corrosion.

2:26 (Triple only) know that an alloy is a mixture of a metal and one or more elements, usually other metals or carbon

An alloy is a mixture  of a metal with, usually, other metals or carbon.

For example, brass is a alloy of copper and zinc, and steel is an alloy of iron and carbon.

2:27 (Triple only) explain why alloys are harder than pure metals

Alloys are harder than the individual pure metals from which they are made.

In an alloy, the different elements have slightly different sized atoms. This breaks up the regular lattice arrangement and makes it more difficult for layers of ions to slide over each other.


Select a set of flashcards to study:


     Skills and equipment

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Section 1: Principles of chemistry

      a) States of matter

      b) Atoms

      c) Atomic structure

     d) Relative formula masses and molar volumes of gases

     e) Chemical formulae and chemical equations

     f) Ionic compounds

     g) Covalent substances

     h) Metallic crystals

     i) Electrolysis

 Section 2: Chemistry of the elements

     a) The Periodic Table

     b) Group 1 elements: lithium, sodium and potassium

     c) Group 7 elements: chlorine, bromine and iodine

     d) Oxygen and oxides

     e) Hydrogen and water

     f) Reactivity series

     g) Tests for ions and gases

Section 3: Organic chemistry

     a) Introduction

     b) Alkanes

     c) Alkenes

     d) Ethanol

Section 4: Physical chemistry

     a) Acids, alkalis and salts

     b) Energetics

     c) Rates of reaction

     d) Equilibria

Section 5: Chemistry in industry

     a) Extraction and uses of metals

     b) Crude oil

     c) Synthetic polymers

     d) The industrial manufacture of chemicals

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