Flashcards in Unit 2 - Chemistry: Compounds and the periodic table Deck (41):
Non-metallic elements compound is called..
Non-metallic elements form non-ionic compounds using a different type of bonding called covalent bonding involving shared pairs of electrons.
Describe lithium, sodium and potassium in Group 1
A collection of relatively soft metals showing a trend in melting point and reaction with water. Reactivity and colour increases down the group.
All elements in group 1 react with water to form an alkali.
Describe transition elements
A collection of metals having high densities, high melting points and forming coloured compounds, and which, as elements and compounds, often act as catalysts.
Describe noble gases
Unreactive. Full outer electron shell.
Describe the composition of clean air
A mixture of 78% nitrogen, 21% oxygen and small quantities of noble gases, water vapour and carbon dioxide.
Describe the formation of carbon dioxide
-A product of complete combustion of carbon-containing substances
-A product of respiration
-A product of the reaction between an acid and a carbonate.
Relate the terms exothermic and endothermic to the temperature changes observed during chemical reactions
During all chemical reactions, an energy change occurs. In the reaction heat is either released or absorbed. When a reaction releases heat to the surroundings, we call that reaction an Exothermic Reaction. The reaction that absorbs energy from the surroundings are called Endothermic Reactions.
The reactants have more energy than the products here, so a small amount of energy is required to activate the reaction.
-Release of heat
-Energy needed for the reaction to occur is less than the total energy released.
-Extra energy is released, usually in the form of heat.
-The release of heat means that an exothermic reaction increases temperature of the surroundings.
-Heat absorbs energy from the surroundings.
-Temperature of surroundings decreases during an endothermic reaction because energy from surroundings is required to drive the reaction, hence decreasing the temperature of the surroundings.
Demonstrate understanding that exothermic and endothermic changes relate to the transformation of chemical energy to heat (thermal energy), and vice versa.
Chemical -> heat
Heat -> chemical
In order to actually start a reaction, a certain amount of energy will be provided to the reactants; We often call this the Energy of Activation because this energy is essentially required to start the reaction.
The energy here is used to break the bonds between the molecules of the atoms of the reactants. The bonds then subsequently rearrange and bond again, which releases energy.
However, if the energy provided to activate the energy is less than the energy released when the bonds form together, the reaction gave out more than it took/absorbed, which makes this a exothermic reaction. If the energy given to activate is more than the energy released during the bond formation, the reaction is endothermic.
The total energy change is called enthalpy.
Describe the formation of ions
If an atom loses or gains an electron we call it an ion.
An atom is pretty stable. It has an equal number of electrons and protons.
However, atoms often have the potential to become unstable. The truth is, most atoms don’t like to have outer electrons. If you look at the noble gases in group 8, they don’t have any outer electrons. The elements in group 1-7 all want to become like a noble gases in terms of electron configuration.
Describe the formation of ionic bonds between elements from group I and VII
Ionic Bonding is basically a process where a non-metal and metal each donate the appropriate number of electrons to each other to form a compound with a stable electron configuration.
Eg Na has an electron configuration of 2, 8, 1. It needs to lose one electron
Chloe has an electron configuration of 2, 8, 7 so it needs to gain one electron.
The overall charge should be 0
Explain the formation of ionic bonds between metallic and non-metallic elements
The IONIC BOND results as a balance between the force of attraction between opposite plus and minus charges of the ions and the force of repulsion between similar negative charges in the electron clouds. In crystalline compounds this net balance of forces is called the LATTICE ENERGY. Lattice energy is the energy released in the formation of an ionic compound.
Metals always form positive ions
Non-metals always form negative ions
Draw dot-and-cross diagrams to represent the sharing of electron pairs to form single covalent bonds in simple molecules, exemplified by H2, Cl2, H2O, CH4 and HCl.
Use the symbols of the elements to write formulae of simple compounds
Some common chemical compounds
H2O - water
HCl - hydrochloric acid
H2SO4 - sulphuric acid.
NH4+ - ammonium
NO3- - nitrate
OH- - hydroxide
Deduce the formula of a simple compound from the relative numbers of atoms present
It is possible to work out the chemical formulae of a compound when given the elements present.
For example, if we are told that a certain compound contains both sodium and chlorine, we can deduce its formulae to be NaCl.
Eg Mg has 2 outer electrons, Chlorine needs 1 more outer electron
Deduce the formula of a simple compound from a model or a diagrammatic representation
A simple way would be to simply count the number of electrons in each atom!
The red dots represent the electrons of the element on the left and the blue crosses represent the electrons of the element of the right.
Since the diagram features an ionic compound, we have to take in account the single electron given from the positive ion to complete its full outer shell.
In total, the element on the left has 11 red dots, or electrons; if we look at the periodic table, the element with 11 electrons is sodium.
Likewise, the element on the right has 17 crosses, or electrons. Using the periodic table, the element with 17 electrons is Chlorine
Determine the formula of an ionic compound from the charges on the ions present
Eg if element x has 3 outer electrons, and element y has 6 electrons (needs two). Then it would be X2Y3
To work out the formulae of an ionic compound using the ions present simply remember that the charge on any ionic compound must be neutral, or 0.
For example, we are told that a compound has Fe2+ ions present as well as Cl- ions.
Using that information, for every Fe2+ present in the compound there must be 2 Cl- ions to balance out the charge. Therefore, the formula of the ionic compound must be FeCl2.
Construct and use word equations
Matter cannot be created or destroyed merely changed from one form into another. ie the number of each type of atom must be the same on both sides.
Sulphuric acid + sodium hydroxide -> sodium sulphate + water.
Construct and use symbolic equation with state symbols
Deduce the balanced equation for a chemical reaction, given relevant information
To balance an equation, you have to remember that there must be same number of atoms of each element present on both sides of the equation
Describe the way the Periodic Table classifies elements in order of proton number
Across the table in rows
Describe the change from metallic to non-metallic character across a period.
Elements on the left, in Group 1, are all metallic.
Elements in Group 2 are also metallic, but their metallic properties are less apparent than the elements in Group 1.E.g. They are less reactive.
As you go across the group, elements slowly become less metallic, and elements in Group 4 become non-metals. However, they are still generally in the solid form.
As you progress group 6,7,8 elements tend to be in the gaseous form.
Describe the relationship between group number, number of outer-shell (valency) electrons and metallic/non-metallic character
Group 1 has 1 outer electron
Group 2 has 2
Group 7 has 7
Metallic elements lose electrons to form positive ions
Non-metallic elements gain electrons to form negative ions
-In Group 7, as you go down the group, the substances progress from Gas to solid. e.g. Chlorine is a gas whilst iodine is a solid.
-In Group 1 and 2 and possibly 3, even as you down down the group, they are mostly metals, however melting points and boiling points tend to decrease for these substances
Predict the properties of other elements in group 1 given data where appropriate
Reactivity increases down the group due to the lower melting and boiling points. Because the outer electrons are easier to remove than the atoms at the top of the group (more shells between inner shell and outermost shell).
Describe the trends in properties of chlorine, bromine and iodine in group VII including colour, physical state and reactions with other halide ions.
-Halogens referred to as diatomic molecules because they go around in pairs and form covalent bonds. 7 electrons in outershell -> wants 8.
-Form acidic compounds with hydrogen
-all have low melting points and boiling points when compared to other elements. Their melting points and boiling points increase down the group.
-all are reactive elements. Their reactivity increases UP the group (opposite to group 1)
-all are coloured elements. Their colour increases down the group.
Physical state at room temperature
Fluorine - gas
Chlorine - gas
Bromine - liquid
Iodine - solid
Astatine - solid
Predict the properties of other elements in Group VII, given data where appropriate
Reactivity increases up the group. Colour increases down the group. Gas gas liquid solid solid. Density increases. Heavier.
Describe the uses of the noble gases in providing an inert atmosphere
Argon in lamps
Helium for filling balloons.
Use the Periodic table to predict properties of elements by means of groups and periods
Elements in a group have the same number of outer electrons. Same sort of reactivity and properties.
Elements in a period have the same number of electron shells.
Distinguish metals and non-metals by their general physical and chemical properties
-form positive ions
-all conduct electricity
-good conductors of heat
-ductile (able to be drawn into a wire)
-malleable (can be hammered into shapes)
-sonorous (rings like a bell or a gong)
-form basic oxides when react with oxygen
(Basic means it will dissolve in water to give an alkali
-mostly form negative ions
-can form compounds with each other
-only carbon conducts electricity
-good insulators of heat (mostly)
-form covalent compounds
-tends to be brittle (snaps easily)
-form acidic oxides when reacted with oxygen
Identify and interpret diagrams that represent the structure of an alloy
Alloys is the mixture of two or more metallic elements.
As we can see in the diagram, we see two different types of atoms incoherently mixed together, without any apparent order. In GCSE, if you see this, you can almost assume that the diagram is suggesting an alloy.
What makes alloys special is that since the atoms are all jumbled together of different sizes, it is much more difficult for alloy layers to slide over each other, so alloys are harder than pure metals.
Different coloured dots in the cube solid thing
Explain why metals are often used in the form of alloys
An alloy is a mixture of two elements, one of which is a metal. Alloys often have properties that are different to the metals they contain. This makes them more useful than the pure metals alone. For example, alloys are often harder than the metal they contain.
Alloys contain atoms of different sizes, which distorts the regular arrangements of atoms. This makes it more difficult for the layers to slide over each other, so alloys are harder than the pure metal.
Metals such as Copper or iron are too soft for many uses. Therefore, these metals are often mixed with other methods to acquire make it harder. Additionally, an alloy has the properties of both metals, therefore it is beneficial when two metals can mix to negate the weaknesses of each other.
Examples of Alloys used:
An example of a bronze structure
Brass is used in electrical fittings, 70% copper and 30% zinc.
Bronze is used for bearings and bells, and it often composed composed of 80% copper and 20% Tin.
Duralumin is used for airplane manufacture, 96 % aluminium and 4% copper and other metals.
Place in order of reactivity: potassium, sodium, calcium, magnesium, zinc, iron, hydrogen and copper, by reference to the reactions, if any, with water or steam, dilute hydrochloric acid (not alkali metals)
Potassium Reacts with water
Sodium Reacts with water
Calcium Reacts with water
Magnesium Reacts with acids
Can be displaced by carbon:
Zinc Reacts with acids
Iron Reacts with acids
Hydrogen Included for comparison
Copper May react with strongly
Compare the reactivity series to the tendency of a metal to form its positive ion, illustrated by its reaction, if an, with
-the aqueous ions of other listed metals
-the oxides of the other listed metals
Elements at the top form positive ions the easiest, and this tendency decreases as you go down the group. Valence electrons are more easily lost up in the reactive series to form ionic bonds.
Reaction of Potassium with Water
2K (s) + 2H2O (l) —-> 2KOH (aq) + H2 (g)
Potassium + Water —-> Potassium Hydroxide + Hydrogen
Reaction of Magnesium with Water
2Mg (s) + 2H2O —> 2Mg(OH)2 (aq) + H2
Magnesium + Water —> Magnesium Hydroxide + Hydrogen
Reaction of Magnesium with Steam
Magnesium + Steam —> Magnesium Oxide + Hydrogen
Deduce an order of reactivity from a given set of experimental results
Let’s take three unknown metals X,Y and Z
We will put these three metals in three separate beakers immersed with Hydrochloric Acid
In the exam, you are likely to see these type of questions.
The beaker which produces the most bubbles, effervescence is likely to contain the most reactive metal.
However, for this to be a successful experiment, we have to ensure some variables are controlled. We call this controlled variables.
Time allowed for reaction to occur.
Temperature of acid
Initial surface of metal
Volume of acid
Describe the use of carbon in the extraction of copper from copper oxide
Copper is less reactive than carbon, so it can be extracted from its ores by heating it with carbon. For example, copper is formed if copper oxide is heated strongly with charcoal, which is mostly carbon:
copper(II) oxide + carbon → copper + carbon dioxide
2CuO + C → 2Cu + CO2
Removing oxygen from a substance is called reduction. The copper oxide is reduced to copper in the reaction above.
If the carbon is more reactive than the metal it will remove the oxygen from the metal oxide and leave traces of the metal in the reaction vessel.
Describe the essential reactions in the extraction of iron in the blast furnace
1. The hot air starts burning the coke and allows it to react the oxygen in the air to produce CARBON DIOXIDE
C(s) + O2 (g) -> CO2(g)
2. CO2 reacts with more coke, forming carbon dioxide.
Carbon + carbon dioxide -> carbon monoxide
C(s) + CO2(g) -> 2CO(g)
3. The carbon monoxide meets and reacts with the iron ore. It reduces the iron ore to iron
Fe2O3(s0 + 3CO(g) -> 2Fe(l) + 3CO2
(4). Carbon is more reactive than iron, so it can push out or displace the iron from iron oxide. Here are the equations for the reaction:
iron oxide + carbon → iron + carbon dioxide
2Fe2O3 + 3C → 4Fe + 3CO2
Define oxidation and reductio in terms of oxygen loss/gain, and identify such reactions from given information
-gaining of oxygen by an atom, molecule or ion
-the loss or removal of electrons from an atom, ion or molecule.
-an oxidizing agent is the species that gives the oxygen or removes the electrons.
-the removal of oxygen in a compound
-the gain or addition of electrons to an atom, ion or molecule
-A reducing agent is the species that removes the oxygen and “donates” the electrons.
Relate the method of extraction of a metal from its ore to its position in the reactivity series limited to group I and II metals, aluminium, iron and copper.
Group 1 Extracted using electrolysis
Group 2 Electrolysis
Describe a chemical test for water
-Test the liquids boiling point. Water boils at precisely 100 degrees and freezes at exactly 0 degrees.
-If you add water to Anhydrous Copper Sulphate Powder, it forms a blue solution and may give out heat.
-Add Anhydrous Cobalt Chloride which is blue in colour. If water is present, should change to colour pink.
Describe and explain, in outline, the purification of the water supply by filtration and chlorination
Water extracted from the earth is always infested with impurities. This water might be contaminated with disease and bacteria. That is why it is crucial to “purify” the water before it is drank. This is done by two processes, Filtration and Chlorination.
Here is how it works:
-Water is extracted from reservoirs and sent to be “treated”
-The water is first passed through a filter to filter out large objects such as rocks or mud.
-Smaller particles in the water is removed by adding Aluminium Sulphate which causes the smaller particles to stick together in large pieces and settle down the filter.
-Water is now passed through sand and gravel filters which continue to filter off the smaller particles and kills bacteria.
-Now its time for chlorination
-Chlorine gas is first bubbled through the water to kill the bacteria that exists in the water.
-Sodium Hydroxide may be added in the water to prevent the water from being acidic from the chlorine.
-Water is delivered to the people that need them.
Explain why the proportion of carbon dioxide in air is increasing, and why this is important
-The Sun sends energy to the earth in two discrete forms, heat and light.
-Some of the heat is reflected back to the sun/space, but some is trapped in the earth.
-This is caused by the existence of some gases and we call this the Greenhouse effect.
-The primary Greenhouse gases are Carbon Dioxide and Methane.
Increased combustion of carbon in industries which mass produce Carbon Dioxide as a side product and the cutting down of trees which release CO2 via respiration are two major reasons why the greenhouse effect is becoming more serious.
The increase of heat trapped in the earth causes an average rise in sea level and global average temperatures, and we call this effect Global warming