Topic: Simple Molecules & Covalent Bonding

Simple Molecules & Covalent Bonding quiz

1. How many electrons can the second shell hold?

Question 1 of 16

2. Where are the transition metals on the Periodic Table?

Question 2 of 16

3. Why do elements in the same group of the periodic table have similar chemical properties?

Question 3 of 16

4. If a substance has a simple molecular structure, what physical state might it be at room temperature?

Question 4 of 16

5. What is special about the elements in group 0?

Question 5 of 16

6. Which particle(s) in an atom has the smallest mass?

Question 6 of 16

7. How many protons in a potassium atom?

Question 7 of 16

8. What is meant by the term molecule?

Question 8 of 16

9. Explain how the atoms are held together in a hydrogen bromide molecule

Question 9 of 16

10. In the dot and cross diagram of the outer electrons showing the covalent bonding in a molecule of methane (CH₄), how many electrons should be shown in areas 1, 2, 3, and 4?

Question 10 of 16

11. Do larger molecules have higher or lower boiling points than smaller molecules?

Question 11 of 16

12. Describe the test for oxygen gas

Question 12 of 16

13. Describe the formation of a covalent bond

Question 13 of 16

14. On the Periodic Table what is the meaning of the word Period? What does that tell us about the electron configuration of the atom?

Question 14 of 16

15. On the Periodic table what is the meaning of the word Group?

Question 15 of 16

16. Describe the chemical test for water

Question 16 of 16


2020-02-22T20:57:55+00:00Categories: Uncategorized|Tags: , |

1:03 understand how the results of experiments involving the dilution of coloured solutions and diffusion of gases can be explained

Diffusion is the spreading out of particles in a gas or liquid. There is a net movement of particles from areas of high concentration to areas of low concentration until a uniform concentration is achieved.


i) dilution of coloured solutions

Dissolving potassium manganate(VII) in water demonstrates that the diffusion in liquids is very slow because there are only small gaps between the liquid particles into which other particles diffuse.

The random motion of particles cause the purple colour to eventually be evenly spread out throughout the water.

Adding more water to the solution causes the potassium manganate(VII) particles to spread out further apart therefore the solutions becomes less purple. This is called dilution.


ii) diffusion experiments

When ammonia gas and hydrogen chloride gas mix, they react together to form a white solid called ammonium chloride.

ammonia                  +              hydrogen chloride                 –>            ammonium chloride

NH3(g)                     +              HCl(g)                                     –>            NH4Cl(s)

A cotton wool pad was soaked in ammonia solution and another was soaked in hydrogen chloride solution. The two pads were then put into opposite ends of a dry glass tube at the same time.

The white ring of ammonium chloride forms closer to the hydrochloric acid end because ammonia particles are lighter than hydrogen chloride particles and therefore travel faster.

Even though these particles travel at several hundred metres per second, it takes about 5 min for the ring to form. This is because the particles move in random directions and will collide with air particles in the tube.

1:04 know what is meant by the terms: solvent, solute, solution, saturated solution

When a solid dissolves in a liquid:

  • the substance that dissolves is called the solute
  • the liquid in which it dissolves is called the solvent
  • the liquid formed is a solution
  • a saturated solution is a solution into which no more solute can be dissolved


1:05 (Triple only) know what is meant by the term solubility in the units g per 100g of solvent

Solubility is defined in terms of the maximum mass of a solute that dissolves in 100g of solvent. The mass depends on the temperature.

For example, the solubility of sodium chloride (NaCl) in water at 25⁰C is about 36g per 100g of water.

1:06 (Triple only) understand how to plot and interpret solubility curves

The solubility of solids changes as temperature changes. This can be plotted on a solubility curve.

Image result for solubility curve

The salts shown on this graph are typical: the solubility increases as temperature increases.

For example, the graph above shows that in 100g of water at 50⁰C the maximum mass of potassium nitrate (KNO₃) which will dissolve is 80g.

However, if the temperature were 80⁰C a mass of 160g of potassium nitrate (KNO₃) would dissolve in 100g of water.

1:07 (Triple only) practical: investigate the solubility of a solid in water at a specific temperature

At a chosen temperature (e.g. 40⁰C) a saturated solution is created of potassium nitrate (KNO₃) for example.

Some of this solution (not any residual solid) is poured off and weighed. The water is then evaporated from this solution to leave a residue of potassium nitrate which is then weighed.

The difference between the two measured masses is the mass of evaporated water.

The solubility, in grams per 100g of water, is equal to 100 times the mass of potassium nitrate residue divided by the mass of evaporated water.

 solubility (g/100g) = \frac{mass Of Solute}{mass Of Solvent} \times 100

1:11 understand how a chromatogram provides information about the composition of a mixture

Paper chromatography can be used to investigate the composition of a mixture.

A baseline is drawn on the paper. The mixture is spotted onto the baseline alongside known or standard reference materials. The end of the paper is then put into a solvent which runs up the paper and through the spots, taking some or all of the dyes with it.

Different dyes will travel different heights up the paper.

The resulting pattern of dyes is called a chromatogram.

In the example shown, the mixture is shown to contain the red, blue and yellow dyes. This can be seen because these dots which resulted from the mixture have travelled the same distance up the paper as have the red, blue and yellow standard reference materials.

1:12 understand how to use the calculation of Rf values to identify the components of a mixture

When analysing a chromatogram, the mixture being analysed is compared to standard reference materials by measuring how far the various dyes have travelled up the paper from the baseline where they started.

For each dye, the Rf value is calculated. To do this, 2 distances are measured:

  • The distance between the baseline and the dye
  • The distance between the baseline and the solvent front, which is how far the solvent has travelled from the baseline

The Rf value is calculated as follows:

 R_f=\frac{distance\:of\:dye\:from\:baseline }{distance\:of\:solvent\:front\:from\:baseline}

If the Rf value of one of the components of the mixture equals the Rf value of one of the standard reference materials then that component is know to be that reference material. 

Note that because the solvent always travels at least as far as the highest dye, the Rf value is always between 0 and 1.

Dyes which are more soluble will have higher Rf values than less soluble dyes. In other words, more soluble dyes move further up the paper. The extreme case of this is for insoluble dyes which don’t move at all (Rf value = 0). The other aspect affecting how far a dye travels is the affinity that dye has for the paper (how well it ‘sticks’ to the paper).

1:13 practical: investigate paper chromatography using inks/food colourings

  1. A pencil line (baseline) is drawn 1cm from the bottom of the paper. Pencil will not dissolve in the solvent, but if ink were used instead it might dissolve and interfere with the results of the chromatography.
  2. A spot of each sample of dye is dropped at different points along the baseline.
  3. The paper is suspended in a beaker which contains a small amount of solvent. The bottom of the paper should be touching the solvent, but the baseline with the dyes should be above the level of the solvent. This is important so the dyes don’t simply dissolve into the solvent in the beaker.
  4. A lid should cover the beaker so the atomosphere becomes saturated with the solvent. This is so the solvent does not evaporate from the surface of the paper.
  5. When the solvent has travelled to near the top of the paper, the paper is removed from the solvent and a pencil line drawn (and labelled) to show the level the solvent reached up the paper. This is called the solvent front.
  6. The chromatogram is then left to dry so that all the solvent evaporates.

Common solvents are water or ethanol. The choice of solvent depends on whether most of the dyes are soluble in that solvent.

1:15 know the structure of an atom in terms of the positions, relative masses and relative charges of sub-atomic particles

An atom consists of a central nucleus, composed of protons and neutrons.

This is surrounded by electrons, orbiting in shells (energy levels).

Atoms are neutral because the numbers of electrons and protons are equal.

Electronnegligible (1/1836)-1

1:16a know what is meant by the terms atomic number, mass number and relative atomic mass (Aᵣ)

Atomic number: The number of protons in an atom.

Mass number: The number of protons and neutrons in an atom.

Relative atomic mass (Ar): The average mass of an atom compared to 1/12th the mass of carbon-12.

1:18 understand how elements are arranged in the Periodic Table: in order of atomic number, in groups and periods

The elements in the Periodic Table are arranged in order of increasing atomic number.


Image result for periodic table groups and periods

Columns are called Groups. They indicate the number of electrons in the outer shell of an atom.

Rows are called Periods. They indicate the number of shells (energy levels) in an atom.

1:19 understand how to deduce the electronic configurations of the first 20 elements from their positions in the Periodic Table

Electrons are found in a series of shells (or energy levels) around the nucleus of an atom.

Each energy level can only hold a certain number of electrons. Low energy levels are always filled up first.

Rules for working out the arrangement (configuration) of electrons:

Example – chlorine (Cl)

1) Use the periodic table to look up the atomic number. Chlorine’s atomic number (number of protons) is 17.

2) Remember the number of protons = number of electrons. Therefore chlorine has 17 electrons.

3) Arrange the electrons in levels (shells):

  • 1st shell can hold a maximum of 2
  • 2nd can hold a maximum of 8
  • 3rd can also hold 8

Therefore the electron arrangement for chlorine (17 electrons in total) will be written as 2,8,7

4) Check to make sure that the electrons add up to the right number

The electron arrangement can also be draw in a diagram.

Electron arrangement for the first 20 elements:

1:21 identify an element as a metal or a non-metal according to its position in the Periodic Table

Metals on the left of the Periodic Table.

Non-Metals on the top-right, plus Hydrogen.

1:22 understand how the electronic configuration of a main group element is related to its position in the Periodic Table

Elements in the same group have the same number of electrons in their outer shell.

This is why elements from the same group have similar properties.

1:23 Understand why elements in the same group of the Periodic Table have similar chemical properties

Elements in the same group of the periodic table have the same number of electrons in their outer shells, which means they have similar chemical properties.

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.

1:46 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.

2:09 know the approximate percentages by volume of the four most abundant gases in dry air

Air is a mixture of different gases.

The abundance of gases in the air is as follows:

Gas% by volume
Nitrogen, N278.1
Oxygen, O221.0
Argon, Ar0.9
Carbon dioxide, CO20.04

2:10 understand how to determine the percentage by volume of oxygen in air using experiments involving the reactions of metals (e.g. iron) and non-metals (e.g. phosphorus) with air

The following 3 experiments can be used to determine that oxygen (O2) makes up approximately 20% by volume of the composition of air.


The copper is in excess and uses up the oxygen to form copper oxide (CuO).

All the oxygen in the air is therefore used up, and so the volume of the air decreases by about 20% (the percentage of oxygen in air).



The iron reacts with the oxygen in the air (rusting).

As long as the iron and water are in excess, the total volume of air enclosed by the apparatus decreases by about a fifth (20%) over several days.



The phosphorus is lit with a hot wire.

It reacts with the oxygen in the air and causes the water level in the bell jar to rise by about 20%.


2:14 Practical: determine the approximate percentage by volume of oxygen in air using a metal or a non-metal

The following 3 experiments can be used to determine that oxygen (O2) makes up approximately 20% by volume of air.


The copper is in excess and uses up the oxygen to form copper oxide (CuO).

All the oxygen in the air is therefore used up, and so the volume of the air decreases by about 20% (the percentage of oxygen in air).



The iron reacts with the oxygen in the air (rusting).

As long as the iron, oxygen and water are all in excess, the total volume of air enclosed by the apparatus decreases by about a fifth (20%) over several days.



The phosphorus is lit with a hot wire.

It reacts with the oxygen in the air and causes the water level in the bell jar to rise by about 20%.


2:44 describe tests for these gases: hydrogen, oxygen, carbon dioxide, ammonia, chlorine

Tests for gases

GasTestResult if gas present
hydrogen (H2)Use a lit splintGas pops
oxygen (O2)Use a glowing splintGlowing splint relights
carbon dioxide (CO2)Bubble the gas through limewaterLimewater turns cloudy
ammonia (NH3)Use red litmus paperTurns damp red litmus paper blue
chlorine (Cl2)Use damp litmus paperTurns damp litmus paper white (bleaches)

2:49 describe a test for the presence of water using anhydrous copper(II) sulfate

Add anhydrous copper (II) sulfate (CuSO4) to a sample.

If water is present the anhydrous copper (II) sulfate will change from white to blue.

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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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