Back on topic: if you need to review the basics of atomic structure, go to my post on Atomic Structure and the Periodic Table (oh my, it has the word "the" in it as opposed to the subheading written here on my course outline! Okay, I'll stop being stupid now).
Explain the structure of the atom in terms of protons, neutrons and electrons.
See my first post on Atomic Structure and the Periodic Table.
Write the electron configuration using the shell model for the first 20 elements.
Also see my first post on Atomic Structure and the Periodic Table.
Explain trends in first ionisation energy, atomic radius and electronegativity across periods and down groups (for main group elements) in the Periodic Table.
Yay! Something new!
I'm going to start with atomic radius, because I think it's the easiest to explain, and then I can explain the other two from there.
When you go down groups, atomic radius generally increases. This is because you're adding more shells and therefore the atom is getting bigger. Also, when you go up groups, atomic radius generally decreases.
When you go across periods (from left to right), atomic radius generally decreases. This is because the number of shells remains consistent across a period, but the number of protons increases, making the nucleus "more positive" and therefore pulls the electrons in the shells closer to it (if that made any sense). Therefore, the radius decreases. Again, the inverse is true: as you move left across a period, atomic radius generally increases.
I'm saying "generally" because occasionally you get some retarded element that doesn't want to follow the rules exactly. All these things are trends only. Sort of like how there might be a fashion trend where a lot of people are dressing according to that fashion trend but then there'd be someone like me who doesn't give a damn about it.
Now for electronegativity! Electronegativity is an atom's ability to attract electrons and, in doing so, become more electronegative. Now, since it's the positively-charged protons in the nucleus that are doing most of the attracting (opposites attract, remember?) electronegativity is largely based on two things: how many protons are in the nucleus, and how many shells of electrons are surrounding that nucleus (remember, likes repel, so the more shielding shells there are, the harder it is for an atom to attract an electron). Therefore, electronegativity decreases as you go down groups as there are more "shielding" shells of electrons, and increases as you go across periods since there are more protons for the same number of "shielding" shells. (The inverse is also true.)
Now, what is this "first ionisation energy" thing you might ask? (I've started like 3 sentences with the word "now" in the past few minutes... argh.) Well, "first ionisation energy" is the energy required to remove one electron from an atom. As electronegativity of an atom increases, the electrons are held more tightly to the atom and thus require more energy to remove, and vice versa. Thus trends in first ionisation energy are the same as trends in electronegativity: first ionisation energy increases as you go right across a period, or up a group.
Explain the trend in successive ionisation energies
Apart from first ionisation energy, there's also second ionisation energy, third ionisation energy, and so on, as you take more and more electrons away from an atom. Second ionisation energy is higher than the first, third is higher than the second, and so on. This is because it is harder to take an electron away from an atom which is turning into a more and more positive ion.
Ionisation energy jumps significantly when you've finished taking away all the electrons from one shell and have to start on the next. This is because this next shell is closer to the nucleus and therefore is bound more strongly by the protons in the nucleus of the atom. Analysing ionisation energy numbers, therefore, allows you to get a pretty good idea of how many electrons are in the outer shell. For example, if there's a big difference between the 3rd and 4th ionisation energies, you can assume that there's 3 electrons in the outer shell, and that the big jump is due to the 4th electron having to be taken from the shell second furthest from the nucleus.
Describe and explain the relationship between the number of valence electrons and an element's bonding capacity, position on Periodic Table and physical and chemical properties.
I've explained everything apart from position on Periodic Table back at my first post on Atomic Structure and the Periodic Table.
Position on Periodic Table is pretty simple though. Elements are arranged according to the number of electrons that they have. They are then arranged into periods according to how many shells they have, and groups according to how many valence electrons they have (1 valence electron- group 1, 2 valence electrons- group 2, 3 valence electrons- group 13, 4- group 14, 5- group 15 until 8- groups 18). Then there are the transition metals in the middle which I can't be bothered explaining, but they're also listed in order of number of protons.
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