Metal + acid → salt + hydrogen
Some metals react with water or steam but others do not. When reactions occur with water the products are hydrogen gas and the metal hydroxide. Bubbles of colorless gas form and the solution produced is clear.
Metal + water → metal hydroxide + hydrogen
Most metals react with oxygen to form oxides, the changes occur over a significantly long period of time. When they slowly react at room temperature they lose their shiny lustrous appearance. And some metals such as aluminium become coated with a dull layer of adhering oxide that prevents further reaction and some metals such as iron form a powdery surface layer of oxides, which impedes further reactions.
Metal + oxygen → metal oxide
- Describe and justify the criteria used to place metals into an order of activity based on their ease of reaction with oxygen, water and dilute acids.
The activity series of metals was established by comparing the reactivity of metals with oxygen, water, dilute acids and many other chemicals. Activity series lists the metals in order of decreasing ease of loosing electrons: the metals on the left loose electrons more easily then metals on the right. Ions of metals on the right accept electrons more readily than do ions on the left.
K Na Li Ba Ca Mg Al Cr Zn Fe Co Ni Sn Pb Cu Hg Ag Pt Au
- Identify the reaction of metals with acids as requiring the transfer of electrons.
When a dilute acid reacts with a metal it is the hydrogen ions with the metal atoms. Electrons are transferred from the metal atoms to the hydrogen ions producing metals ions and hydrogen molecules.
Metal + Dilute Acid → Salt Solution + Hydrogen
The anions in the dilute acid solution are the anions in the salt solution because these anions take no part in the reactions, they are called spectator ions.
When an toms loses on or more electrons we sat that it has been oxidized. When an atom gains one or more electrons we say that it has been reduced. In other words:
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Oxidation means loss of electrons
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Reduction means gain of electron
Half equations are reactions, which describe the oxidation and reduction processes separately in terms of electrons lost or gained.
Examples
For the reaction of zinc with dilute acid, the overall reaction is:
Zn(s) + 2H+(aq) → Zn2+(aq) + H2 (g)
It is made up of the half reactions
Zn → Zn2+ + 2e- (oxidation)
2H+ + 2e- → H2 (reduction)
In combining half reactions into complete reactions, it is necessary to balance the number of electrons. This is because there can be no electrons left over on either side of a complete reaction.
For the reaction of aluminium with dilute acid, the half reactions are:
Al → Al3+ + 3e- (oxidation)
2H+ + 2e- → H2 (reduction)
To combine these into an overall equation we multiply the oxidation reaction by 2 and reduction reaction by 3 (to give six electrons in each) and add to get:
2Al(S) + 6H+(aq) → 2Al3+(aq) + 3H2 (g)
- Outline examples of the selection of metals for different purposes based on their reactivity, with a particular emphasis on current developments in the use of metals
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Calcium: the ability of calcium to react moderately with water is used to dehydrate many organic solvents
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Zinc: zinc’s reactivity is used to protect iron and steel
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Aluminium: corrosion resistance due to rapid reaction of aluminum and oxygen which forms a surface passivating layer, used for house gutter
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Copper: copper’s low reactivity with oxygen and water make it useful for water pipes
Note- a passivating layer is an impervious layer that forms on a metal surface and thus prevents any corrosion of the metal beneath.
- Outline the relationship between the relative activities of metals and their positions on the Periodic Table.
- In general as you go across the period reactivity decreases and as you go down a group reactivity increases.
- The most active metals (and Na) are located in group one on the left side of the table. In general their reactivity increases down the group.
- The next most active metals belong to group two. Once again reactivity increases down the group.
- Metals of moderate reactivity such as Zn lie at the edge of the metal zone of the periodic table.
- There is no general pattern for the remaining elements, except that the least reactive such as gold are located in the lower central region called the transition metal.
These differences in reactivity can be explained in terms of the electronic structure of the atom.
- Identify the importance of first ionisation energy in determining the relative reactivity of metals
The first ionization energy of elements in the periodic table is a function of atomic number.
The first ionization energy of an element is the energy required to remove completely the electron of an atom in the gaseous state.
Ionization energy increases across a period and decreases down a group.
The reactivity of metals increases as their ionization energy decreases.
The first ionization energy increases across a period as electron shells go from near empty to full. The ionization energy decreases down a group as the outer electrons become further remove from the positive nucleus.
The second ionization energy is always greater than the first since the electron is now being removed from a positive ion.
Ionization energy trends to provide evidence that chemical bonding attempts to reach noble gas configuration.
- Identify an appropriate model that has been developed to describe atomic structure.
Bohr Model
A modified version of the Bohr 1913 model is used to explain the properties of atoms and elements.
Aspects of the Bohr model are
- A single electron moves around the nucleus in a circular orbit
- The energy of the electron is restricted to certain values/levels
- Atom is composed of a small, central nucleus made up of protons and neutrons
- Electrons occupy stationary levels and do not radiate energy
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First energy level → lowest energy level
- The inertia of the electron would tend to move it away from the nucleus but the inertia is counterbalanced by the electrostatic attraction which exists between the positive nucleus and the negative electron
- Electron can absorb a quantum of energy and move to an excited state. It can radiate energy by “falling” back to a more stable state.
- Outline the history of the development of the Periodic Table including its origins, the original data used to construct it and the predictions made after its construction.
The modern periodic table classifies elements in groups and periods in order of increasing atomic number. There are seven horizontal periods and eighteen vertical groups (or eight groups + transition metals). Groups II and III are separated by a large sub-group called transition metals. Groups I to VIII form the main block- their group numbers is equal to the number of electrons in their valence shell).
- Explain the relationship between the position of elements in the Periodic Table, and:
-Electrical conductivity
Electrical conductivity increases as we move across a period. This is a result of more freely available delocalised electrons, to conduct the electricity. The trend in electrical conductivity also increases for as we go down a group (as the number of shells increases). This is because the delocalised electrons are less attracted to the positive nucleus they get further away, hence the larger the atomic number; generally the more electrical conductivity
-Ionisation energy
Ionisation energy refers to the energy needed to move the least strongly held electron (first electron) from the atom in a gaseous state. As we progress down a group, less energy is required to for ionization of energy to occur (decreases). This is because of the attraction between the delocalised electrons and the positive nucleus. However, as we move across a period, ionization energy increases. This is because nuclear charge of the atom increases (the number of protons) and therefore atomic radius increases.
-Atomic radius
The radius of the atom decreases across a period due to the increasing nuclear charge causing the valence electron to be attracted closer to the nucleus. Down a group the atomic radius increases as more electrons shells are added to the atom.
-Melting point
Melting points refers to the value in temperatures required for a substance to melt. As you go across a period, the melting point increases from group I-IV and then decreases.
A periodic trend in melting point is shown with the noble gasses (almost no tendency to form bonds) at thoughts, and carbon and silicon in-group 4, forming peaks (due to infinite covalent bonds). In-group 8 (monatomic gasses) the only forces between atoms are week dispersion forces, so they have very low melting points. In-group 7 (diatomic covalent), the forces between molecules are weak, so melting points are low.
In-group 4 are strongly bonded covalent network solids, so melting points are very high. Groups 1 and 2 metallic bonding have moderate to high melting points.
Down a group, I-IV decreases and increases for groups V-VIII and transition metals.
-Boiling point
Boiling points refers to the value in temperatures required for a substance to boil. Boiling points show a similar variation to melting points. As the atomic number increases across a period, the boiling points increase (reaching a peak in group 4) and then decrease (reaching a thought in group 8).
-Electronegativity
Electronegativity of an element is a measure of the ability of the atom of that element to attract bonding electron towards itself when it forms compounds (electron attracting ability of atoms). As you go across a period, electronegativity increases. This is because more electrons are added to the shell so they are more willing to gain electron. As you go down a group the electronegativity decreases.
Down a group, I-IV decreases and increases for groups V-VIII and transition metals.
-Combining power (valency)
The valance of a substance is the amount of electrons the substance is willing to wither gain or receive. The valance of a substance stays constant as we go down the periodic table of elements. As we progress across a period the valance of the substance increases until it reaches +-4 in-group 4 and then starts to decrease. This is due to the fact the all atoms desire full electron shell configuration and want to give up or gain electrons to become and become ions.
- Define the mole as the number of atoms in exactly 12g of carbon-12 (Avogadro ’s number)
A mole of any substance is the amount that contains the same number of chemical unit (atoms molecules or ions), as there are atoms in exactly 12g of carbon 12 or 6.02 x 1023 chemical units.
- Compare mass changes in samples of metals when they combine with oxygen.
- Describe the contribution of Gay-Lussac to the understanding of gaseous reactions and apply this to an understanding of the mole concept.
Gay-Lussac’s law of combining of combining volumes states that when gasses combine at the same temperature and the same pressure, they do so in volumes that bare a simple ratio to each other (the ratio of the volumes of gasses involved in a chemical reaction can be expressed as simple whole numbers). This means that the volumes are in the same ratio as the co-efficient in the balanced equation.
Since coefficients represent the number of particles of reactants and products, and because these particles (atoms ions or molecules or ions are extremely small) a convenient way of counting particles was developed as the mole concept.
- Recount Avogadro ’s law and describe its importance in developing the mole concept.
Avogadro’s law states that under the same conditions of temperature and pressure, the same volume of different gasses contains equal number of molecules (hence equal number of moles of molecules). Because of Avogadro’s law, in gaseous reactions, gas volumes can replace the number of moles.
- Distinguish between empirical formulae and molecular formulae.
The empirical formula of a compound is the simplest whole number ratio of the atom or ions in the compound.
Molecular formula of a covalent molecular compound represents the actual number of atoms of each element present in the molecule.
- Define terms mineral and ore with reference to economic and non-economic deposits of natural resources
A mineral is a pure crystalline compound that occurs in the earth’s crust.
An ore is a compound or mixture of compounds from which it is economic (or commercially profitable) to extract a desired substance such as a metal.
- Describe the relationship between the commercial prices of common metals, their actual abundances and relative costs of production.
Commercial price of common metals is directly proportional to the abundance of the metal and the relative cost of production. Therefore if it costs more to extract the metal, (cost of production) it will therefore be passed on to the consumer in forms of higher prices. If the mineral is abundant, then the cost of the mineral is reflected in its value, as the demand will be more than the supply. The location and transportation cost also have an effect on the prices of metals.
- Explain why ores are non-renewable resources
A substance is defined to be renewable if it regenerates or reforms in one lifetime for examples, a tree. Mineral ores dot not reform or regenerate in a lifetime, but rather over millions of years, this is why they are defined to be non renewable. Our future use of metals is limited by the finite amounts of accessible natural minerals.
- Describe the separation processes, chemical reactions and energy considerations involved in the extraction of copper from one of its ores.
The extraction of copper includes, mining and crushing, froth flotation, roasting and smelting and electrolytic refining.
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Mining and crushing- the mined copper ore (-0.5% copper) is crushed and ground into ball mills to help free the copper mineral from the gangue (unwanted rocks).
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Froth flotation- uses different surface properties of sulfide and waste minerals in order to separate and extract copper from its ore. The finely milled ore is mixed with water, detergent and oils (“collectors”) in a froth flotation tank. The copper minerals adhere to the bubbles, which float to the surface and are recovered as a slurry. The percentage of copper has risen to around 30% in this concentration phase. The gangue remains at the bottom of the tank and is discarded to a tailings pond.
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Roasting and smelting-
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the dried concentrated mineral is mixed with coal, lime and sand (SiO2) and roasted (1000ºC) in an oxidizing atmosphere to form a mixture of sulfides and oxides of copper and iron. Energy is supplied from the heat released in the combustion of coal. The sand and lime act as a flux to remove the iron impurities as a molten slag of iron (II) silicate.
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At the end of the roasting stage the copper is present as molten mixture of copper (I) oxide and copper (II) sulfide [called a copper matte (50-70% copper)].
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In the final stage the copper matter is mixed with more sand and smelted in the furnace to form molten ‘blister’ copper metal (98% pure).
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Electrolytic refining- The blister copper from the smelter is made into copper anodes and a pure copper sheet from the cathode in the electrolytic cell. Copper oxidises at the anode and copper ions are reduced at the cathode to form layers of pure copper (99.9% copper).
Anode (+) -- Cu(s) → Cu2+ + 2e-
Cathode (-) – Cu2+ + 2e- → Cu(s)
- Recount the steps taken to recycle aluminium.
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Collection of Aluminium Products: Scrap Aluminium is collected be individuals, community groups and council collections.
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The Aluminium is sorted and separated according to different alloy types: As aluminium is not magnetic, steel can be removed from the scrap by the use of magnets. Used cans are separated from other aluminium scrap; this allows their alloy composition to stay constant.
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Preparation: The cans are weighed, cleaned and squashed together and baled into blocks. These blocks are taken to the furnace.
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Remelting and refining: The blocks are tipped into a rotary furnace where they are melted.
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Casting: The molten aluminium is poured into a mould to make ingots.
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Recycled aluminium ready to be reused: The ingots are taken to where they can be formed into new aluminium items.