1. Topic 3 - periodicity
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2. 3.1.1 Describe the arrangement of elements in the periodic table in order of increasing atomic number
The elements in a periodic table are placed in order of increasing atomic number (Z), which we know is a fundamental property of the element - number of protons in the nucleus of its atom. To read the periodic table, simply start from the top left and read across. You will find that the atomic number increases.

3. 3.1.2 distinguish between the terms group and period
Groups are the vertical columns of elements. Periods are the horizontal rows of elements.

4. Groups state how many valence electrons there are Periods state how many electron shells there are. Valence electron is an electron in the outer shell of an atom
3.1.3 apply the relationship between the electron arrangement of elements and their position in the periodic table up to Z=20.
Example, -Carbon has an electron arrangement of 2, It is in the 2nd period and 6th group Calcium has an electron arrangement of 2, 8, 8, It is in the 4th period and 2nd group.

5. 3.1.4 apply the relationship between the number of electrons in the highest occupied energy level for an element and its position in the periodic table
An atom with one valence electron will be in which group?
An atom with a full outer shell will be in which group?
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An atom with the electron arrangement 2,8,7 is located where on the periodic table?
An atom with the electron arrangement 2,8,8,1 is located where on the periodic table?

6. 3.2 physical properties

7. 3.2.1 define the terms first ionization energy and electronegativity
Electronegativity defined as the relative attraction that an atom has for the shared pair of electrons in a covalent bond.
First Ionization Energy is the minimum energy required to remove an electron from a neutral gaseous atom in its ground state.

8. 3.2.2 Describe and explain the trends in atomic radii, ionic radii, first ionization energies, electronegativities and melting points for alkali metals and the halogens
Atomic Radii Since the outer electrons are difficult to locate, in practice, the atomic radius is measured as half the distance between two bonded atoms. For this reason noble gases are given no value as they do not bond with other atoms. On descending a group, the atomic radius increase. This is because the outer electrons are getting further from the nucleus. This applies for both alkali metals and halogens.

9. Ionic Radii
It is defined by the distance between nucleus and outer most electrons of a positive metal cation or negative non-metal anion. Both metal cation and non-metal anion increase in size on descending the group
Metal cations tend to be smaller than their atom as they have lost their outer most electrons
Non-metal anions tend to be larger than their atom as they have gained electrons into their outer energy levels

10. 1st Ionization Energy It is defined by the minimum energy required to remove an electron from a neutral gaseous atom in its ground state. Ionization energy decreases on descending a group
The outer most electrons are easier to remove as they are a greater distance from the nucleus and are shielded from the positive nucleus charge by the core electrons.

11. Electronegativity Electronegativity is defined as the relative attraction that an atom has for the shared pair of electrons in a covalent bond On descending a group, electronegativity decreases. While the atomic radius and the nuclear charge increases, the level of shielding increases and the effective nuclear charge decrease.
The most electronegative elements are in the top right of the periodic table. (F, O and N). Which are involved in hydrogen bonding.

12. Melting point trends Page 57
What is the melting point trend for group 1 and group 7?
Why is this the trend?

13. Melting point trend for Alkali metals and Halogens

14. First Ionization Energies Ionization energy generally increases across a period. This is due to the increase effective nuclear charge as the number of protons increases with no increase in shielding. Electronegativity Across a period, electronegativity increases. This is the result of an increased number of protons and thus an increased effective nuclear charge with no greater level of shielding.
3.2.3 Describe and explain trends in atomic radii, ionic radii, first ionization energies and electronegativities for elements across period 3
Atomic radii Across a period, the atomic radius decreases. This is because of the increase in effective nuclear charge and no increase shielding, therefore the electrons get pulled closer towards the nucleus. Ionic radii Across a period, the metal cation and the non-metal anion get smaller. This is because of an increase in protons (Effective Nuclear Charge) and all the ions have the same electronic configuration.

15. Atomic radii Across a period, the atomic radius decreases
Atomic radii Across a period, the atomic radius decreases. This is because of the increase in effective nuclear charge and no increase shielding, therefore the electrons get pulled closer towards the nucleus.

16. Ionic radii Across a period, the cations and the anions get smaller
Ionic radii Across a period, the cations and the anions get smaller. This is because of an increase in protons (Effective Nuclear Charge) and all the ions have the same electronic configuration.

17. First Ionization Energies Ionization energy generally increases across a period. This is due to the increase effective nuclear charge as the number of protons increases with no increase in shielding.

18. Electronegativity Across a period, electronegativity increases
Electronegativity Across a period, electronegativity increases. This is the result of an increased number of protons and thus an increased effective nuclear charge with no greater level of shielding.

19. 3.2.4 Compare the relative electronegativity values of two or more elements based on their positions in the periodic table.
By looking at the difference in electro-negativity values
We can tell whether a molecule has covalent or ionic bonds

20. 3.3 Chemical properties

21. 3.3.1 Discuss the similarities and differences in the chemical properties of elements in the same group
Group 0: Noble Gases -Colorless
-Monatomic
-Unreactive
-Their lack of reactivity can be explained by the inability of their atoms to lose or gain electrons.

22. Group 1: Alkali Metals
Good conductors of electricity
Low Density
Shiny surface 
Reactive metals
Forms ionic compounds with non-metals
They form single charged ions M+. Their low ionization energies give an indication of the ease with which the outer electron is lost.
Reaction with Water
The alkali metals react with water to produce hydrogen and the metal hydroxide.
The reaction becomes more vigorous as we descend the group.

23. Group 7: Halogens -Diatomic molecules
Group 7: Halogens -Diatomic molecules
-Colored
-Gradual change from gases (Fluorine and Chlorine) to Liquid (Bromine) to Solids (Iodine and Astatine)
-Reactive non-metals
-Reactivity decrease down the group
-Form ionic compounds with metals or covalent compounds with other non- metals

24. Halogen reaction with group 1 metals
Halogens react with Group 1 metals to form ionic halides. The halogen atoms gains one electron from the Group 1 elements to form a halide ion. The electrostatic force of attraction between the alkali metal ions and halides bonds the ions together.
Once the transfer is complete, the ions are pulled together by the mutual attraction of their opposite charges.
This would be the equation

25. Displacement reactions with Halogens
The relative reactivity of the elements can also be seen by placing them in direct competition for an extra electron.
A chlorine nucleus has a stronger attraction for an electron than a bromine nucleus because of its smaller atomic radius and so takes the electron from the bromide ion. The chlorine has gained an electron and so forms the chloride ion. The bromide ion loses an electron to form bromine.

26. Halogen Reaction with Silver
The halogens form insoluble salts with silver. Adding a solution containing the halide to a solution containing silver ions produces a precipitate which is useful in identifying the halide.
Silver Chloride = White
Silver Bromide = Cream
Silver Iodide = Yellow Silver Nitrate = Colorless

27. 3.3.2 Discuss the changes in nature, from ionic to covalent and from basic to acidic, of oxides across period 3
The acid-base properties of the oxides are closely linked to their bonding. Metallic elements, which form ionic oxides, are basic; non-metal oxides, which are covalent are acidic. Aluminum oxide, which can be considered as an ionic oxide with some covalent character, shows amphoteric properties reacting with both acids and bases

28. Sodium and Magnesium Oxide are bases
Na2O + HCl →2NaCl + H2O MgO + H2SO4 → MgSO4 + H2O
Aluminium oxide is amphoteric (base and acid) so it reacts with both bases and acids
Al2O3 + 3H2SO4 → Al2(SO4)3 + 3H2O Al2O3 + 2NaOH +3H2O → 2NaAl(OH)4
Phosphorous Oxide, Sulfur Oxide, Dichlorine Oxide reacts with water to form Acid Solutions P4O10 + 6H2O → 4H3PO4 SO3 + H2O → H2SO4 Cl2O7 + H2O → 2HClO4