IGCSE PRINCIPLES OF CHEMISTRY Topic 1 Bonding Revision Notes 1.28 describe the formation of ions by the gain or loss of electrons The formation of ions: Ions are atoms or molecules with an electric charge due to the gain or loss of electrons. If electrons are lost, the ion has a positive charge. Metals tend to do this, so they form cations (positive ions) Normally elements from group 1-3 will form cations. If electrons are gained, the ion has a negative charge. Non-metals tend to do this, and they form anions (A-Negative-ION - ANION). So elements from group 5-7 will form anions. Group 0/8 are the noble gases and are inert + unreactive, so they do not form ions.
1.29 understand oxidation as the loss of electrons and reduction as the gain of electrons
OILRIG - Oxidation Is Lost, Reduction Is Gain Oxidation is the process of an atom losing an electron/s and becoming a positively charged ion. Reduction is the process of an atom gaining an electron/s and becoming a negatively charged ion. You can remember this using OIL RIG. Depending on their electronic configurations, atoms lose or gain electrons in order to achieve a full outer shell. Losing electrons The sodium atom has one electron in its outer shell. If it loses this one electron it will achieve a full outer shell. By losing the one electron to another atom, it becomes a sodium ion.
IGCSE PRINCIPLES OF CHEMISTRY Topic 1 Bonding Revision Notes
Gaining electrons A chlorine atom has seven electrons in its outer shell. It can reach a full outer shell by gaining one electron. It will then become the chloride ion, Cl-. A negative charge is assigned to the ion to signify that the ion contains one more electron than proton.
1.30 recall the charges of common ions in this specification
Charge 1+
Positive ions/Cations Name of ion Formula Ammonium NH4+ Copper (I) Cu+ Hydrogen H+ Lithium Li+ Potassium K+ Silver Ag+ Sodium Na+
Charge 1-
Negative ions/Anions Name of ion Bromide Chloride Hydroxide Fluoride Iodide Nitrate Hydrogencarbonate
Formula BrClOHFINO3HCO3-
2+
Barium Calcium Copper (II) Iron (II) Lead (II) Magnesium Nickel (II) Strontium Zinc
Ba2+ Ca2+ Cu2+ Fe2+ Pb2+ Mg2+ Ni2+ Sr2+ Zn2+
2-
Carbonate Sulphate Sulphite Sulphide Oxide
CO32SO42SO32S2O2-
3+
Aluminium Iron (III)
Al3+ Fe3+
3-
Nitride Phosphate
N3PO43-
IGCSE PRINCIPLES OF CHEMISTRY Topic 1 Bonding Revision Notes 1.31 deduce the charge of an ion from the electronic configuration of the atom from which the ion is formed So if the electronic configuration is 2.8.1, you can see that the atom has one outer shell electron only. And so it only needs to lose that one to have a full outer shell, making an ion with the electron configuration 2.8 with a positive 1 charge. Another example, if the electronic configuration is 2.8.7, then the atom only needs to gain 1 outer shell electron to have a full outer shell. So the ion formed would have the electronic configuration of 2.8.8, and so the charge would be negative 1.
Here, the Sodium (Na) has lost one electron. It doesn't have equal numbers of protons and electrons anymore; it has one less electron than protons (or you can think of it as one more proton than electrons), so it has a 1+ charge. The chlorine, on the other hand, has gained one electron. So it has one more electron than proton, thus it has a 1- charge. The formula for sodium chloride is NaCl.
The magnesium atom loses 2 electrons to an oxygen atom, and they both have full outer shells now. It is common that ions form noble gas structures like this to become more stable and unreactive like the group 0/8 elements. The magnesium oxide is held together by very strong attractions between the ions. The ionic bonding is stronger here than in sodium chloride as this time you have 2+ ions attracting 2ions. The greater the charge, the greater the attraction between the ions. The formula for magnesium oxide is MgO.
IGCSE PRINCIPLES OF CHEMISTRY Topic 1 Bonding Revision Notes 1.32 explain, using dot and cross diagrams, the formation of ionic compounds by electron transfer, limited to combinations of elements from Groups 1, 2, 3 and 5, 6, 7
This is a common example, it doesn't matter which element has dots/crosses for their electrons, the important thing is to make it clear that the electrons are transferred from one atom to another to make 2 ions. The one losing an electron here is sodium, so it becomes a positive ion called a cation, and the chlorine atom gains an electron and becomes a chloride ion (an anion). Dot and cross diagrams Dot and cross diagrams are used to show electron transfers, and the creation of ions.
As you can see, the magnesium atom gives 2 electrons to the oxygen atom, so an ionic bond is created and therefore, 2 ions are created; the magnesium atom becomes a positive ion, losing 2 electrons and the oxygen atom becomes a negative ion, gaining 2 electrons.
IGCSE PRINCIPLES OF CHEMISTRY Topic 1 Bonding Revision Notes
Magnesium oxide: MgO Even though only one magnesium ion and one oxide ion is shown, the actual equation is: 2Mg + O2 Ă 2MgO (remember that oxygen is diatomic)
Calcium chloride: CaCl2 However, here you have to show 2 chloride ions because calcium loses 2 electrons, and 2 chlorine atoms gain an electron each to form 2 chloride ions.
1.33 understand ionic bonding as a strong electrostatic attraction between oppositely charged ions 1.34 understand that ionic compounds have high melting and boiling points because of strong electrostatic forces between oppositely charged ions Electrostatic attraction in ionic bonding The electrostatic attraction keeps the ions bonded, and as it is very strong, it takes a lot of energy to break these bonds. For this reason, ionic compounds have very high melting and boiling points, because of the amount of energy needed to break these strong bonds between oppositely charged ions.
IGCSE PRINCIPLES OF CHEMISTRY Topic 1 Bonding Revision Notes TRIPLE: 1.35 understand the relationship between ionic charge and the melting point and boiling point of an ionic compound The relationship between the ionic charge and boiling/melting points For ions of similar size, the strength of the forces of attraction between the ions will depend on the size of their charge. For this reason, magnesium oxide (Mg²+O²-) has a higher melting and boiling point than sodium chloride (Na+Cl-). The more electrons that have been transferred, the stronger the bond because of the stronger electrostatic forces. TRIPLE: 1.36 &1.37 describe an ionic crystal as a giant three-dimensional lattice structure held together by the attraction between oppositely charged ions Giant ionic lattice structure
As you can see from the diagram, this is a giant ionic lattice. When an ionic compound forms, the positively charged ions attract the negatively charged ions and arrange themselves into a three-dimensional structure called an ionic lattice. The arrangement of the ions in the ionic lattice of sodium chloride is shown above.
1.38 describe the formation of a covalent bond by the sharing of a pair of electrons between two atoms What is a covalent bond? This is a shared pair of electrons between two non-metal atoms. If each atom shares 2 electrons, it is called a double covalent bond. E.g. in oxygen, as each oxygen atom has 6 outer shell electrons it shares 2 electrons to have 8 as a full outer shell, hence oxygen is diatomic. O2 Each of the positively charged nuclei is attracted to the same negatively charged pair of electrons, which is why covalent bonds are so strong. The atoms come close enough together for their outer electron (valence) shells to overlap.
IGCSE PRINCIPLES OF CHEMISTRY Topic 1 Bonding Revision Notes
Why does hydrogen form molecules? Whenever a bond is formed (of whatever kind), energy is released, and that makes the things involved more stable than they were before. The more bonds an atom can form, the more energy is released and the more stable the system becomes. In the case of hydrogen, each hydrogen atom has only one electron to share, so it can only form one covalent bond. The H2 molecule is still much more stable than two separate hydrogen atoms.
The significance of noble gas structures in covalent bonding The formation of covalent bonds producing noble gas structures is quite common. When atoms bond covalently, they often produce outer electronic structures the same as noble gases-full outer shell. This is so they become stable and unreactive. The more electrons shared, the more covalent bonds there are, the more stable the molecule is.
1.39 understand covalent bonding as a strong attraction between the bonding pair of electrons and the nuclei of the atoms involved in the bond Covalent bonding During covalent bonding, the sharing of electrons between 2 non-metals, the electrons become attracted to the nucleus of each ion in a bond, this is electrostatic attraction.
IGCSE PRINCIPLES OF CHEMISTRY Topic 1 Bonding Revision Notes 1.40 explain, using dot and cross diagrams, the formation of covalent compounds by electron sharing for the following substances: i hydrogen ii chlorine iii hydrogen chloride iv water. v methane vi ammonia vii oxygen viii nitrogen ix carbon dioxide x ethane xi ethene Hydrogen: Each hydrogen atom has only one electron and needs one more to complete its first shell. When two hydrogen atoms get close together their shells can overlap and then they can share their electrons.
Since, electrons are being shared, there is a strong force of attraction between them. This force is a covalent bond. The bonded atoms form molecules. Hydrogen's molecular formula is H2. Chlorine: A chlorine atom needs a share of one other electron to obtain a full outer shell. If two chlorine atoms are placed together the result is as shown below:
IGCSE PRINCIPLES OF CHEMISTRY Topic 1 Bonding Revision Notes
Oxygen: Each oxygen atom requires a share of two electrons.
Since each oxygen atom has a share of two pairs of electrons, we call this a double covalent bond.
1.41 understand that substances with simple molecular structures are gases or liquids, or solids with low melting points 1.42 Explain why substances with simple molecular structures have low melting and boiling points in terms of the relatively weak forces between the molecules Simple molecular substances Both elements and compounds can exist as simple molecular substances. Simple molecular substances usually have low melting and boiling points. This is because the forces of attraction between oppositely charged ions, (i.e. ionic bonds) and therefore very little energy is required to overcome them. Although, when simple molecular substances change state, the covalent bonds between the atoms are not usually broken. Covalent bonds are strong compared to the forces of attraction between the molecules.
1.43 explain the high melting and boiling points of substances with giant covalent structures in terms of the breaking of many strong covalent bonds Giant covalent structures Some substances are made up of millions of atoms covalently bonded together to form a giant structure. These substances have high melting and boiling points, because when melting or boiling it, you are not separating inter-molecular bonds between the atoms, but the inter-molecular bonds that keep the molecule together. There are a lot of these, therefore they have high melting and boiling points.
IGCSE PRINCIPLES OF CHEMISTRY Topic 1 Bonding Revision Notes TRIPLE: 1.44 draw diagrams representing the positions of the atoms in diamond and graphite Diamond and graphite Part I
As you can see from the diagram, the diamond structure, which is made up from carbon atoms are joined to 4 other carbon atoms, this is an example of a giant covalent structure. Graphite is another example of a giant covalent structure made out of carbon atoms. Each carbon atom is joined to 3 other carbon atom. It is an example of a lattice, where weak forces exist between each layer.
TRIPLE:1.45 explain how the uses of diamond and graphite depend on their structures, limited to graphite as a lubricant and diamond in cutting Diamond and Graphite Part II Diamond is very hard and abrasive (strong covalent bonds that are difficult to break and it does not conduct electricity as there is no delocalised electrons. This makes it suitable to be used for cutting tools on drills because it is so abrasive. Diamond is also used for jewellery because shiny rocks seem attractive to humans. Graphite is soft and slippery (forces of attraction between the layers are weak so the layers easily slide over one and another so can be easily separated). Graphite also conducts electricity, as there is a fourth electron that is delocalised between the layers and are free to move parallel to the layers. It is used as a lubricant because it is slippery and soft and because it is an conductor it is perfect for electrodes for electrolysis.
IGCSE PRINCIPLES OF CHEMISTRY Topic 1 Bonding Revision Notes 1.46 understand that a metal can be described as a giant structure of positive ions surrounded by a sea of delocalised electrons Metallic crystals Metals have a giant, three-dimensional lattice structure in which positive ions are arranged in a regular pattern in a 'sea of delocalised electrons'.
The outer shell electrons are detached from the atoms and are delocalised throughout the structure. The attraction between the positive ions and the delocalised electrons is known as a metallic bond and this attraction keeps the ions together.
1.47 explain the electrical conductivity and malleability of a metal in terms of its structure and bonding. Most metals have high melting and boiling points because metallic bonds are strong and there are many of them to overcome in a giant structure, hence a lot of heat energy is required. Good conductors of electricity because the delocalised electrons are free to move when a potential difference is applied across the metal.
Malleable and ductile because the layers of positive ions can easily slide over one another and take up different positions. The delocalised electrons move with them so the metallic bonds are not broken.
IGCSE PRINCIPLES OF CHEMISTRY Topic 1 Bonding Revision Notes
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Ductile – can be pulled into a wire The delocalised electrons allow metal atoms to slide past one another without being subjected to strong repulsive forces that would cause other materials to shatter. Metals can be made harder by alloying them with other metals. An alloy is a mixture of metals-for example, brass is a mixture of copper and zinc. In an alloy, the different metals have slightly different sized atoms. This breaks up the regular arrangement and makes it more difficult for the layers to slide. A common example of an alloy is stainless steel, which is iron mixed with chromium and nickel. The chromium and nickel form strong oxide layers which protect the iron. This is why stainless steel is so resistant to corrosion. Obvious uses include kitchen sinks, saucepans, knives and forks and gardening tools. But there are also major uses for it in the brewing, dairy and chemical industries where corrosion-resistant vessels are essential. Properties Metals:
Non-metals:
Strong
Brittle
Malleable and ductile
Brittle
React with oxygen to form basic oxides React with oxygen to form acidic oxides Sonorous
Dull sound when hit with hammer
High melting and boiling points
Low melting and boiling points
Good conductors of electricity
Poor conductors of electricity
Good conductors of heat
Poor conductors of heat
Mainly solids at room temp. Exception mercury - liquid at room temp.
Solids, liquids and gases at room.temp.
Shiny when polished
Dull looking
When they form ions, the ions are positive
When they form ions, the ions are negative - except hydrogen that forms a positive ion, H+.
High density
Low density
END OF TOPIC 1 PRINCIPLES OF CHEMISTRY