C4 Revision Cards

1. C4 Revision Cards Hope you find these useful - V

2. Atomic Structure • An atom has a nucleus, surrounded by electrons. • The nucleus contains protons and neutrons. • The nucleus is positively charged. • This is because protons are positively charged. • The neutrons have no charge. • Protons: • Positive charge • Mass = 1 • Neutrons • No charge • Mass = 1

3. Atomic Structure • Electrons: • Negative charge • Mass = 0.0005 • The electrons circle the nucleus in shells. • They have a negative charge. • Their mass is almost nothing compared to protons or neutrons (0.0005 in relation to them). • Overall, an atom has no charge at all. • This is because the number of protons equals the number of electrons. • This balances the charge, leaving it as neutral. 2 protons + 2 electrons = neutral charge.

4. Atomic Mass 56 Fe • Each element on the periodic table comes with 2 numbers. • The MASS NUMBER = protons + neutrons. • The ATOMIC NUMBER = number of protons alone. • The atomic number also tells you the number of electrons around the nucleus. • To find out the number of neutrons, just subtract the atomic number from the mass number. 26 So for the example above: Neutron number is 56-26 = 30 neutrons neutron number = mass number – atomic number

5. Isotopes • Some elements have more than one form. • Their atomic number (protons) stays the same. • Their mass number is different. This meansisotopes have a different number of NEUTRONS • Because the number of protons is the same, isotopes of an element have the same chemical characteristics. • Some may be radioactive though (C-14).

6. Ions A normal, neutral Lithium atom • Don’t get ions and isotopes mixed up. • Atoms become ions if they gain or lose electrons. • This means that their overall charge is now unbalanced. • Losing an electron would mean there is now an overall positive charge. • Gaining an electron would mean the opposite. An ion is an atom with a charge If it loses an electron, there is one more proton overall – so the atom is now charged (an ion)

7. Electronic Structure This helium atom has a full 1st shell: 2 • Electrons occupy shells around the nucleus. • The shells fill from the inside, outwards. The 1st shell fits: 2 electrons The 2nd shell fits: 8 electrons The 3rd shell fits: 8 electrons • Atoms with full electron shells are stable. • Those with partly-filled shells will react in order to fill their shells. This neon atom has a full 2nd shell: 2,8 This argon atom has a full 3rd shell: 2,8,8

8. Ionic Bonding • An atom that has 1 or 2 electrons in it’s outer shell will be very keen to get rid of them. • An atom that 1 or 2 spaces in it’s outer shell will be very keen to fill them. • A sodium atom will give up it’s outer electron to become a Na+ ion. • A chlorine atom will pick up that electron, becoming a Cl- ion. • The now oppositely charged ions are strongly attracted to each other. • They bond together – ionically.

9. Ionic Bonding - Continued • Ionic bonding produces substances with giant ionic structures. • The ions in the structure have a lattice arrangement. • The strong attraction of ions to each other means that the melting and boiling points of giant ionic structure is very high. • Their tightly packed nature also means they don’t conduct electricity. • However if you dissolve them in water to form a solution, the ions separateand can carry current. • If you can heat them enough to melt, they’ll conduct electricity as once again, the ions will be able to move. • Remember, this requires a lot of heat. (mp for NaCl = 804 Celsius)

10. Covalent Bonding • Ionic bonding involves the transferring of electrons. • Covalent bonding involves the sharing of electrons. • Non-metals bond covalently. Non-metal atoms can share electrons in order to simulate full outer shells. Dot-cross diagrams can be used to show how the electrons are shared. Each pair of shared electrons represents one covalent bond.

11. Covalent Bonding Examples A hydrogen molecule (H2) consists of two H atoms sharing a pair of electrons. They ‘feel’ as if their 1st shell is full. A CO2 molecule consists of a carbon atom covalently bonded to two oxygen atoms. The carbon, with 4 electrons in it’s outer shell, shares 2 electrons with each oxygen atom.

12. Groups The columns are known as groups. The group number is the same as the number of electrons in the outer shell. EG: Group 1 elements all have one outer electron. Groups & Periods Periods The rows are known as periods. The period that an element is in, indicates how many electron shells it has. If an element is in period2, it will have 2 electron shells. Elements in the same groups or period have similar chemical properties. For example, group 1 elements all react in a similar way with water.

13. The Group 1 Elements • Group 1 elements are called alkali metals. • When reacted with water, they form an alkaline solution. • Their reactions with water follow the same pattern: • Eg: • The reactivity of the alkali metals increases as you go down the group. • This is because the atoms get larger and the outer electron is more easily lost, as it is further from the positively charged nucleus. alkali metal + water --> hydroxide + hydrogen lithium + water --> lithium hydroxide + hydrogen potassium + water --> potassium hydroxide + hydrogen

14. The Group 1 Elements lithium + water --> lithium hydroxide + hydrogen (word) • The symbol equations for the reactions on the last card need to be balanced. Li + H2O --> LiOH + H2 (symbol) 2Li + 2H2O --> 2LiOH + H2 (balanced) • When flame tests are carried out on alkali metal compounds, the flame changes colour:

15. Take a break! mmmm…. break

16. The Group 7 Elements • The group 7 elements are known as the halogens. • The halogens react vigorously with the alkali metals you just met. • When the halogens and alkali metals meet, a metal halide is formed. EG: lithium + chlorine  lithium chloride potassium + bromine  potassium bromide sodium + chlorine  sodium chloride 3 different metal halides Balanced: 2Li + Cl2  2LiCl 2K + Br2  2KBr 2Na + Cl2  2NaCl Use the group 1 + 7 elements to think of other combinations and write out the reactions.

17. The Group 7 Elements • The reactivity of the halogens increases as you go up the group. • A more reactive halogen, will displace a less reactive halogen. • Eg: If you bubble chlorine through potassium bromide solution, the following reaction will occur: • The chlorine has taken the place of the bromine (displacement). • A less reactive halogen cannot displace a more reactive one. • No reaction would occur in this case. chlorine + potassium bromide  potassium chloride + bromine Cl2 + 2KBr  2KCl + Br2

18. Transition Elements • Transition elements are found in the middle of the periodic table. • Transition metals have the following properties: • Conduct heat. • Are shiny (lustrous). • Conduct electricity. • Are sonorous (ring when struck). • Are malleable (bendy). • Ductile (stretchable). • Transition metal compounds are often coloured: • Copper compounds are often blue. • Iron(II) compounds are pale green. • Iron(III) compounds are orange/brown. • Transition metals are often used as catalysts in reactions: • Iron in the Haber Process. • Nickel in Margarine Production.

19. Thermal Decomposition • When a transition metal carbonate is heated, it decomposes to form a metal oxide and carbon dioxide gas: • Notice that the pattern is the same for all of the reaction. • Sometimes chemistry is just about learning patters… look at displacements. • That’s why chemistry’s easy. FeCO3 decomposes into Iron Oxide and CO2 CuCO3 decomposes into CopperOxide and CO2 MnCO3 decomposes into ManganeseOxide and CO2 ZnCO3 decomposes into ZincOxide and CO2 On heating…

20. Precipitation Reactions Precipitation is a reaction between solutions that make an INSOLUBLE SOLID. • In other words, when you mix two solutions, a cloud of solid particles appears. • If transition metal compounds are mixed with sodium hydroxide, precipitation reactions occur. • Colourful clouds of solids form:

21. Metal Structure • Metals have some very characteristic properties: lustrous, hard, high density, high boiling point, good heat/electricity conductors, high tensile strength. • The uses of a metal depend on its properties. • Iron is stronger and harder than aluminium, but it rusts and is heavy. Aluminium is therefore used to make planes. • Metal atoms are held together with metallic bonds. • Metallic bonds are strong forces of attraction between close-packed positive metal ions and a ‘sea’ of delocalised electrons.

22. Metallic Bonding & Free Electrons Their delocalised nature means: They can conduct electricity. They allow the metal to bend (malleable) They allow the formation of strong metallic bonds, giving metals very high melting and boiling points. It’s the sea of delocalised (free) electrons that give metals most of their properties. It’s important to understand that the electrons can move.

23. Superconductors • Most metals are great electrical conductors. • However, all metals show electrical resistance. • If a metal is cooled to a very low, critical temperature, it becomes a superconductor. • Superconductors are materials that conduct electricity with little or no electrical resistance. Superconductors do not have a magnetic field. They repel permanent magnets and cause them to levitate. Super-fast electronic circuits Loss-free power transmission Powerful electro-magnets The potential benefits of superconductors….

24. Superconductor Problems • Remember - metals conduct electricity due to delocalised electrons. • At the moment, superconductors work only at unreasonably low temperatures – this limits their use. • Superconductors that work at room temperature need to be developed. Maglev trains are kept levitating above the track with help of very powerful superconducting electromagnets. This eliminates contact friction, allowing very high speeds. At the moment, superconductors have to be kept very cold with the help of liquid nitrogen.

25. Water Purification • Water is not only essential to life, but has many industrial uses: • Cheap raw material • Used as a coolant (reduces the temp of other substances) • A valuable solvent (can dissolve substances). • Water in a river is rarely pure and needs to be purified before we can drink it. • It may contain the following: • Salts and minerals • Pollutants • Insoluble materials • Microorganisms

26. Water Purification • There are 3 main stages to water purification: • These steps are expensive. • It is therefore important to conserve water. • Pollutants can sometimes get into water after it has left the waterworks. • Seawater is completely undrinkable, and requires processes such as distillation to remove dissolved substances. 2. Filtration Very fine particles are filtered using sand and gravel 3. Chlorination The water in treated with chlorine to kill microbes 1. Sedimentation Solid particles and bacteria settle out from the water.

27. Testing Water • Precipitation reactions are excellent for finding out if the following are present in water: • SULPHATE IONS (add barium chloride) • CHLORIDE, BROMIDE or IODIDE IONS (add silver nitrate). 2. To test for chloride, bromide or iodide ions, add silver nitrate: If chloride ions are present, a white precipitate will form. If bromide ions are present, a creamprecipitate will form. If iodide ions are present, a yellowprecipitate will form. 1. Adding barium chloride to water will produce a white precipitate if sulphate ions are present.

28. These cards are for revision purposes. They were made in a single night! Some of the content has been condensed! Make sure you check out anything you don’t understand by looking at your revision guide or using the internet. Good luck.