Topic 2 Flashcards
(39 cards)
Understand the terms group and period
Groups:
A group is a column (goes down) in the periodic table. All elements in a group have similar chemical properties (because they have the same number of electrons on their outer shell), however the properties of each elements (such as reactivity) often gradually change as you go down a group (as the atomic number increases).
Periods:
A period is a row (goes right) in the periodic table. Properties of elements in the same period are not similar and quite often change.
Recall the positions of metals and non-metals in the Periodic Table
Red line under B, Si, As, Te, At that separates metals and non-metals
Left + Metals and Right = Non metals
The group of elements between group 1 and group 3 are called the transition metals. They don’t tend to follow normal patterns and don’t have a group number.
Explain the classification of elements as metals or non-metals on the basis of their electrical conductivity and the acid-base character of their oxides
- Metals are conductors that form metal oxides that are alkaline
- (Most) non-metals are ones that don’t conduct and form non-metal oxides that are acidic
Understand why elements in the same group of the Periodic Table have similar chemical properties
Elements in the same column have the same number of electrons in their outer shells. This means they react and bond similarly. This is because they need to lose or gain the same number of electrons to become stable and have full outer shells
Understand that the noble gases (Group 0) are a family of inert gases and explain their lack of reactivity in terms of their electronic configurations.
Nobel gases (group 0, the last column in the Periodic Table) are inert, meaning they don’t reach with many things at all. This is because the noble gases have a full outer shell, so they have no need to gain or lose electrons
Describe the reactions of these elements with water and understand that the reactions provide a basis for their recognition as a family of elements
Group 1 Metals
Lithium (Li), Sodium (Na) and Potassium (K) all react vigorously with water. The rest further down the group is not safe to experiment with.
Lithium:
- Floats
- Fizzes
- Solid eventually disappears
- Produces some heat but doesn’t melt
Sodium:
- Floats
- Heat given off, causing it to melt into a ball
- Gradually disappears
- Sometimes a white trail of sodium hydroxide can be seen, but it usually ends up dissolved in the water
- Moves around (due to hydrogen given off)
- If it gets stuck on the side of the container, it will burns with an orange flame
Potassium:
- Burns with a lilac flame
- Reacts similarly to sodium but quicker
- Gives off heat, causing it to melt into a ball
Describe the relative reactivities of the elements in Group 1
As you go down the group, the metals because more reactive
Explain the relative reactivities of the elements in Group 1 in terms of distance between the outer electrons and the nucleus.
Group 1 elements only have 1 electron in their outer shell. Electrons are held to the atom by the forces of attraction between the nucleus (which has protons, so a positive charge) and the electrons themselves (which have a negative charge). If there are more shells, the last electron (which will be lost if the metal reacts with anything) is further away from the nucleus, so the forces of attraction between the nucleus and the electron are weaker. The more shells, the further away, and it’s more reactive.
As you go down the group, they become more reactive because there is an increase in the number of shells.
Recall the colours and physical states of the elements at room temperature
Fluorine (F) - Gas, Yellow Chlorine (Cl) - Gas, Yellow-green Bromine (Br) - Liquid, Brownish red Iodine (I) - Solid, Purple Astatine (At) - Solid, Black
Make predictions about the properties of other halogens in this group (Group 7)
As you go down the group:
- Reactivity decreases
- Colour darkens
- Melting/boiling point increases
- Particles become closer together (i.e. fluorine is a gas but astatine is a solid)
Understand the difference between hydrogen chloride gas and hydrochloric acid
Hydrogen chloride gas is HCl(g)
When dissolved in water, it becomes HCl(aq) because the ions become detached: they dissociate, leaving the separate H+ and Cl- ions. H+ is acidic, which is what makes it an acid
Dissociate - Disconnect or Separate
Explain, in terms of dissociation, why hydrogen chloride is acidic in water but not in methylbenzene
If HCl is dissolved in water, it dissociates (Splits up) into H+ and Cl- ions. This solution (hydrochloric acid) is acidic because it contains H+ ions.
However, if HCl is dissolved in methylbenzene, it doesn’t dissociate, therefore no H+ ions are present so the solution is not acidic.
Describe the relative reactivities of the elements in Group 7
Group 7 elements become less reactive as you go down the group
At the top, the positive charge of the proton in the nucleus is close to the surface (as there are few shells) this makes it easy for them to pull in the one electron they need to become stable, meaning they are very reactive.
Lower down where there are more shells, the pull of the proton is further from the surface making it less easy to pull in another electron.
Describe experiments to demonstrate that a more reactive halogen will displace a less reactive halogen from a solution of one of its salts
A displacement reaction is basically a reaction where the more reactive element displaces (pushes out) a less reactive element from a compound.
Chlorine is more reactive than iodine (as it is higher up in group 7). Therefore, if you add chlorine water to potassium iodide solution, the chloride will react with the potassium to form potassium chloride (basically, it displaces the iodine). The iodine is displaced from the salt (potassium iodide) and just kind of gets left in the water solution (this turns the solution brown),
Understand these displacement reactions as redox reactions
When a more reactive halogen displaces a less reactive halogen, it’s called a redox reaction. This is because an element has gained something and the other has lost something.
OIL RIG:
Oxidisation Is Loss (of electrons)
Reduction Is Gain (of electrons
Recall the gases present in air and their approximate percentage by volume
The air is made up of approximately:
- 21% oxygen
- 78% nitrogen
- 1% other gases (including 0.9% argon and 0.037% carbon dioxide
Explain how experiments involving the reactions of elements such as copper, iron and phosphorus with air can be used to investigate the percentage by volume of oxygen in air
Copper, iron and phosphorus all react with air.
If you know the volume of air that you have, then react it with one of these, then re-measure the volume of air, What has been lost is all oxygen that reacted.
Copper (google for image):
- When copper is heated, it reacts with oxygen in the air to make copper (II) oxide
- This reaction uses up oxygen
- If you heat an excess of copper in a tube and pass it over two syringes, you can use the markers on the syringes to work out how much oxygen has been used up.
- New X (after reaction is cooled and cooled)/Starting X = X% must be oxygen
Iron (google for image):
Iron will react with oxygen to produce rust, this means iron will remove oxygen from the air.
- Soak some iron wool in acetic acid (this acid will catalyse the reaction)
- Push the iron wool at the bottom of a test tube and invert the tube into a beaker of water
- Mark the level of the water in the tube at the beginning
- Leave the experiment for a set amount of time (eg
1 hour)
- Mark the level of water in the tube at the end of the experiment
In conclusion, over time, the level of the water will rise in the test tube. This is because the iron reacts with the oxygen in the air, making iron oxide (The water rises as it takes the place the oxygen took up).
To work out the percentage of air that is oxygen, mark the level of the water in the tube at the beginning and end of the experiment, then fill up the tube to each mark and pour the contents into a measuring cylinder to find out the volume of air at the start and end. Using the difference between the start and end volumes, work out the % that has been used to (should be approx 20%).
Phosphorus:
You can do a similar experiment with white phosphorus. White phosphorus smoulders in air to produce phosphorus oxide. Use the same calculation method as with iron.
Describe the laboratory preparation of oxygen from hydrogen peroxide, using manganese(IV) oxide as a catalyst
Hydrogen peroxide will decompose to form oxygen and water. However, this process is very slow so manganese (IV) oxide is added to speed up the decomposition (acts as a catalyst). The oxygen produced can be collected in two ways…
- Over water
- Connect a delivery tube to bubble the gas into a n upside-down measuring cylinder in a beaker/bowl full of water. - In a gas syringe
- You can use a gas syringe to collect pretty much any gas. Just connect it to the flash that the hydrogen peroxide is decomposing in.
NOTE: The equation for this reaction is 2(H2O2) (aq) -> 2(H2O) (l) + O2 (g)
Describe the reactions of magnesium, carbon and sulfur with oxygen in air, and the acid-base character of the oxides produced
When anything is burnt, it reacts with oxygen in the air to form oxides (which can have either acidic or basic character).
Magnesium:
When magnesium burns in air, it produces a bright white flame and a white powder is formed (this is magnesium oxide). Magnesium oxide is slightly alkali when dissolved in water.
The equation for the reaction (burning Mg in air) is 2Mg(s) + O2(g) -> 2MgO(s)
Carbon:
Carbon will only burn in air if it is very strongly heated. It burns with a yellowy-orangey flame and produces carbon dioxide (as a gas). Carbon dioxide is slightly acidic when dissolved in water.
The equation for the reaction (burning C in air) is C(s) + O2(g) -> CO2(g)
Sulfur: Sulphur burns (in air) with a pale blue flame and produces sulfur oxide which is acidic when dissolved in water.
The equation for the reaction (burning S in air) is S(s) + O2(g) -> SO2(g)
Describe the laboratory preparation of carbon dioxide from calcium carbonate and dilute hydrochloric acid
Dilute HCl will react with calcium carbonate (marble chips) to produce calcium chloride, water and carbon dioxide.
Method:
- Put marble chips at the bottom of a flash
- Fill the flash with hydrochloric acid, doesn’t have to be full, just covering the marble chips
- Immediately attach a bung with a delivery tube into an upturned test tube in water
- The carbon dioxide will be collected in this tube
Equations:
Hydrochloric acid + Calcium carbonate -> Calcium chloride + Water + Carbon dioxide
2HCl(aq) + CaCO3(s) -> CaCl2(Aq) + H2O(l) + CO2(g)
Describe the formation of carbon dioxide from the thermal decomposition of metal carbonates such as copper(II) carbonate
The thermal decomposition (heating) of metal carbonates will produce CO2 as in thermal decomposition, the substance being heated will break down into simpler substances.
Method:
- Put some copper (II) carbonate (its a green powder) into a test tube and insert a bung with a deliver tube at the top
- Camp the test tube at a 90 degree angle and insert the delivery tube into another test tube (that is positioned vertically)
- Heat the copper (II) carbonate with a bunsen burner
NOTE: Because CO2 is denser than air, the downward delivery method can be used.
Equations:
Copper (II) carbonate -> Copper oxide + Carbon dioxide
CuCO3(g) -> CuO(s) + CO2(g)
Describe the properties of carbon dioxide, limited to its solubility and density
Carbon dioxide is more dense than air (which is why it can be collected using the downward delivery method). It is also soluble in water at high pressure.
Explain the use of carbon dioxide in carbonating drinks and in fire extinguishers, in terms of its solubility and density
Carbon dioxide makes soft drinks fizzy. As CO2 is slightly soluble in water and dissolves in drinks when under high pressure, a slightly acidic solutions forms due to the formation of carbonic acid. It eventually goes flat because one the bottle is opened, the CO2 is escaping.
CO2 is more dense than air, this makes it perfect for fire extinguishers. The CO2 will sink onto flames and ‘suffocate’ them (stop the oxygen getting to the flames). As fire needs oxygen to burn, the fire will go out as no oxygen can get to it.
NOTE: CO2 fire extinguishers are only used when water extinguishers aren’t safe. For example, when putting out an electrical fire.
Understand that carbon dioxide is a greenhouse gas and may contribute to climate change.
Carbon dioxide is a green house gas. This means it traps heat in the planet and stops the sun’s rays being reflected away from the earth. Excess levels of carbon dioxide contribute to climate change, causing the earth to warm up at unnatural rates.
