Periodic Table
Electron Spin
the clockwise or counterclockwise motion of an electron. Electrons are magnets, they have magnetic fields. Those fields have only two possible orientations, and a single orbital can only be occupied by two electrons if those orientations are mutually opposed. The magnetic moment of electrons is due to a property called "spin"; the spin value of an electron is 1/2, so it can adopt one of two spin quantum states, +1/2 and -1/2, which correspond to the "up" and "down" orientations of the magnetic moment. Since it has a half-integer spin, an electron is a member of a class of sub-atomic particles called "fermions" which obey rules called the Pauli Exclusion Principle and Fermi-Dirac statistics - one key result of these rules is that no two identical fermions can simultaneously occupy the same quantum state: you cannot have two electrons in the same orbital if they have the same spin orientation, so if one is +1/2, the other must be -1/2, and no more can be added because a third would have to adopt the same quantum state as the one of the first two. In other words, the two clouds have opposite magnetic properties (spins) so they want to pair-up and balance the magnetic effects, which also explains why there can only be 2 electrons in a single orbital.
Electron Affinity
Adding an extra electron either requires or releases energy. Electron affinity describes the energy required/released if you add an electron to an element. For example, lithium would release energy if we add an electron because it has space in its 2s orbitals. The configuration for lithium is 1s² 2s¹. It still has space for one more electron to reside in the s-orbital, so adding an electron *would be favorable since an extra electron to complete an orbital is making it more stable*. Any reaction that increases the stability of an atom results in the release of energy. Although this electron would be repelled by the electron already there in the s-orbital, the effective nuclear charge *overpowers* this repulsion and so the added electron would *still be able to attract to the nucleus*. This is the opposite of nitrogen, in which the effective nuclear charge *doesn't overpower* the electron-electron repulsion and thus is the reason why nitrogen has a *positive* electron affinity. You can think of it as the repulsion of the two electrons in the orbital being greater than the proton-electron attraction in the nucleus, and so the energy needed to force the electron into the atom is greater than the energy released (this reaction still releases some energy since the atom is nonetheless becoming more stable despite the repulsion). Back to lithium, lithium would release energy upon gaining an electron since it can accommodate that extra electron effectively--as opposed to nitrogen--so the electron affinity would be negative. *The more negative the value is, the greater the electron affinity*. The opposite occurs if we add an electron to an atom that's already happy with its configuration since it's an *unfavorable reaction* (noble gases). Adding an electron to a noble gas would require an extra shell that's in *higher energy* just for that electron and is therefore unfavorable. An example of this is helium, which is already happy since it has a complete outer shell. Adding an electron would mean it needs to place the added electron to a *higher energy state, which is unfavorable since you always need to put in energy to place an electron at a higher energy level*. For example, enough photons (energy) are required for an electron to jump to a higher energy level. The added electron would also not feel any attractive force from the nucleus since the other electrons *completely shield the protons from the electron*. So, noble gases would have a positive electron affinity and it would require energy to force that electron in to the noble gas.
Metallic Nature
Any chemical element that is an effective conductor of electricity and heat can be defined as a metal. A metal is also good at forming bonds and cations with non-metals. Atoms inside of a metal quickly lose electrons in order to make positive ions or cations. Metals generally have a low ionization energy, including the transition metals which fills its valence electrons (4s orbital) before backfilling its 3d orbitals. The electrons that are backfilled in the 3d orbitals provide extra shielding for the 4s electrons, so this explains why transition metals are willing to lose their electrons in an attempt to be in a more stable state by forming cations. This is why elements that reside in the lower left display the most metallic features--they have the lowest ionization energy (they are big atoms) who are very willing to give away their valence electrons (they have very few valence electrons). However, these particular metals are very reactive and so we never see them in their pure state as opposed to common metals such as iron, gold, etc. When metals do form cations, they form ionic and metallic bonds. For ionic bonds, metals have low electronegativity and so they would lose their electrons to a nonmetal, which have high electronegativity (they need more electrons to achieve an octet) to form an ionic bond. A metallic bond only occurs in metals. To form metallic bonds, metals remove their valence electrons. This doesn't require a lot of energy, since they have low ionization energies. Once this occurs, all of the metals involved in the metallic bonding share their electrons since the positive metal cations are then attracted to the sea of delocalized electrons surrounding them. Essentially, it can be described as strong forces of attraction between positively charged metal ions and the sea of electrons. This conducts electricity very well since the free electrons can move through the metals very well. A lot of energy is also needed to break these bonds which explains the high melting and boiling points. It also explains why it's generally able to maintain its structure, but when applied enough pressure the metal can change shape due to the ability of the atoms to roll over each other into new positions without breaking the metallic bond.
Atomic Radius
As we move from to left to right in the periodic table, the atomic radius gets smaller. This is because as we move right, we are adding more protons and electrons to the atom. This increase in the number of protons *increases* the force of attraction between the electrons and protons (Zₑff increases, so shielding becomes less effective), so the electrons would be closer to the nucleus and therefore have a smaller radius. As we move down in the periodic table, the atomic radius increases. This is because we are adding more shells as we move down, so the electrons would be *further away* from the nucleus and therefore have a larger atomic radius. The same logic applies when we think of anions and actions of a neutral atom. If we remove an electron from a neutral atom, the cation would be smaller since there would be a *decrease in electron-electron repulsion* (shielding) which would also enable the protons to better pull the remaining electrons towards the nucleus. An anion would be bigger than a neutral atom since an extra electron would *increase shielding* (repulsion) and would therefore have a larger atomic radius.
Orbitals
Each shell can be broken down into subshells. We have s, p, d, and f subshells. The s subshell has just one s-orbital (spherical orbital), so it can contain 2 electrons max. The p subshell (dumbbell orbital) has 3 p-orbitals, so it can contain 6 electrons max. The d subshell has 5 d-orbitals, so it can contain 10 electrons max. And the f subshell has 7 f-orbitals, so it can contain 14 electrons max. The first shell has one subshell - an s subshell. The second shell has 2 subshells - s and p subshells. The third shell has 3 subshells - s, p, and d subshells. The fourth shell has 4 subhells - s, p, d, and f subshells. This is why the first shell can only have 2 e-, second shell 8 e-, third shell 18e-, and etc. Argon is a noble gas because it fills its outer shell--even though the third shell can hold 18 electrons and argon only has 8 in that shell, any additional electrons would go to the 4s orbital (fourth shell, s-subshell) instead of the 3d orbital. This is because the 4s orbital would be filled before the 3d orbital due to lower energy states (read the electron config. card). The third shell is filled (18 e-) only after the 4s orbital is filled (2 e-). After the 3d orbital is completely filled with 18 e-, it would then go back to the 4th shell and continue to fill in the p-subshell (the 4s orbital will no longer be lower in energy once the 3d orbital is filled). This is why transtion metals lose their electrons from the 4s first even though they were once lower in energy—they become higher in energy relative to the 3d orbital electrons after the 3d orbital is filled.
Effective Nuclear Charge
Effective nuclear charge, or Zₑff, determines how effective shielding is in terms of the attraction between the outermost electrons and the protons in the nucleus. The equation is Zₑff=Z-S where Z=number of protons and S=number of shielding electrons (all of the electrons that's closer to the nucleus relative to the outermost electron). Zₑff is convenient for determining the size and ionization energy of a particular element. Zₑff can also explain the stability of noble gases. If we tried to add an extra electron to neon (which would have to go into the 3s orbital) the electron would feel a Zₑff of 0; we'd be putting it in an orbital where the attraction between the nucleus and the electron would be shielded by the intermediate electrons. Conversely, trying to remove an electron from neon is difficult because each electron feels the pull of eight protons on it (10 protons-2 shield electrons=8), which is a very powerful electrostatic interaction.
Electronegativity
Electronegativity is how bad an element wants to hog an electron from another element in a covalent bond. As you move from the left to the right in the periodic table, electronegativity increases. This is because elements on the left really want to lose their electrons to achieve a noble gas configuration as opposed to elements on the right. Lithium, for example, wants to lose its one electron in its second shell badly to be stable like helium. Fluorine, on the other hand, is only one electron from becoming stable like neon so it would really want an electron. So, lithium has a low electronegativity while fluorine has a high electonegativity. As we move from the top to bottom, electronegativity decreases even more. We know that elements get bigger as we move down (more shells), so there would be more shielding and distance between the valence electrons and the protons in the nucleus for those elements. So, those particular elements are considered to be the most electropositive elements in the periodic table. Generally, an electronegativity difference of about 1.7 would be considered an ionic bond.
Ionization Energy
Ionization energy is the energy required to take an electron away from an atom. Ionization energy is *always endothermic*. For example, lithium would have a low ionization energy since it actually wants to lose its electron in the second shell so that it could achieve a noble gas configuration. On the other hand, noble gases would have high ionization energy since they're already happy about its configuration. Removing an electron would be hard in this case since potential energy would increase if you remove an electron, making the element less stable. So, ionization energy would generally increase from left to right in the periodic table. As you go down the table however, ionization energy would be even lower since more shells are added as you move down. It would be even easier to remove an electron if an atom has more shells since that electron is further from the nucleus. Ionization energy can also be explained from the effective nuclear charge: Zₑff=Z-S where Z=number of protons and S=inner electrons that would shield the valence electrons from the nucleus. For lithium, Zₑff=1 while beryllium's Zₑff=2. This means that beryllium's shielding is *less effective* than lithium's since beryllium has one more proton than lithium. So, beryllium's valence electrons in its second shell would be *more attracted to the nucleus relative to lithium* (this also explains why beryllium is smaller), which also explains why beryllium has higher ionization energy since it would be *harder to take away its valence electrons* (it's more attracted to the nucleus as opposed to lithium due to having a higher Zₑff). As we move from beryllium to boron however, ionization energy decreases instead of increasing (against what we'd expect). This is because if we draw out the boron atom, there's now an electron in the 2p orbital in the second shell. It's not the only electron in the second shell--there are two electrons as well in the 2s orbital. Although these two orbitals--p and s--both reside in the second shell, that one electron in the 2p orbital is *on average further away from the nucleus compared to the 2 electrons in the 2s orbital. This means that the two electrons in the 2s orbital actually provides extra shielding (however small) against the one electron in the 2p orbital*. So, boron would have less ionization energy than beryllium since that one 2p orbital electron is easier to remove due to the extra shielding. As you continue to move right in the periodic table, protons continue to increase as well as the number of electrons in the 2p orbital which then brings back up the ionization energy. When we get to oxygen however, the ionization energy decreases again. To explain this, recall that a p-subshell has 3 p-orbitals, which is able to hold a total of 6 electrons since only a maximum of 2 electrons can fit into one single orbital. If we draw out oxygen, we can see that *there are now 4 electrons in the 2p orbital as opposed to nitrogen which only has 3 electrons in the 2p orbital (one electron residing in each p-orbital). This one added electron in oxygen means that it must reside in one of the p-orbitals that already has an electron in it, since there are only three orbitals each able to fit in two electrons. This means that the added electron is now paired up with another electron in the same orbital, so they both repel each other which makes it easier to take away one of them due to the repulsion*. This exception can also be explained from the fact that half-filled and fully-filled orbitals are always the most stable configurations. From there, we see the general trend increase in ionization energy as we move right.
Electron Configuration
The electron configuration describes the most likely arrangement of electrons in the electron cloud. It indicates that each shell is subdivided into orbitals, which describe the shape of the most likely location around a fixed point where electrons could be found. These are indicated by the letters, s,p,d, and f which each represent a particular shape. The first two columns on the left side of the periodic table are where the s subshells are being occupied. Because of this, the first two rows of the periodic table are labeled the s block. Similarly, the p block are the right-most six columns of the periodic table, the d block is the middle 10 columns of the periodic table, while the f block is the 14-column section that is normally depicted as detached from the main body of the periodic table. It could be part of the main body, but then the periodic table would be rather long and cumbersome. Let us start with H and He. Their electron configurations are 1s1 and 1s2, respectively; with He, the n = 1 shell is filled. These two elements make up the first row of the periodic table. The next two electrons, for Li and Be, would go into the 2s subshell. For the next six elements, the 2p subshell is being occupied with electrons. On the right side of the periodic table, these six elements (B through Ne) are grouped together. The next subshell to be filled is the 3s subshell. The elements when this subshell is being filled, Na and Mg, are back on the left side of the periodic table. Next, the 3p subshell is filled with the next six elements. Instead of filling the 3d subshell next, electrons go into the 4s subshell. This is because electrons in the 4s subshell have lower potential energy than those in the 3d subshell, and the orbital filling sequence is based on the lowest potential energy of electrons in each orbital. However, after filling out the 4s subshell and we first enter the 3d subshell (scandium), the energy levels shift and the 4s subshell then has higher energy than the 3d subshell. The 4s subshell is only lower in energy if there are no electrons in the 3d subshell. According to the Aufbau principle, the 4s subshell is filled before the 3d subshell because the 4s is lower in energy. As the 3d subshell becomes populated with electrons, the relative energies of the 4s and 3d fluctuate relative to one another and the 4s ends up higher in energy as the 3d subshell fills. This is why when electrons are lost from the orbitals of the transition metals, they are lost from the 4s first because it is higher in energy. The third shell is filled (18 e-) only after the 4s orbital is filled (2 e-). After the 3d orbital is completely filled with 18 e-, it would then go back to the 4th shell and continue to fill in the p orbital.
Transition Metals
There are two definitions for transition metals: 1. All elements in groups 3-12. 2. An element whose atom has a partially filled d-shell (holds maximum of 10 electrons).
Coulomb's Law
There's a basic formula which allows you to calculate the force between two point charges. F=(kxq1*q2)/r^2 We call this relationship Coulomb's Law. For our purposes we're not really interested in quantifying the magnitude of the forces. We're more interested in understanding how different variables are related to each other. So let's boil it down a little bit and then go into detail. We'll call the first charge q1 the electron charge. This is always going to be -1. The second component q2 is the effective nuclear charge. This will vary with the number of protons in the nucleus, as we shall see. Finally, we have r, the distance between the electron and the nucleus. Finally, we'll get rid of the equals sign and the constant to highlight the proportional nature of these relationships--force is directly proportional to q1 and q1 and inversely proportional to the distance. One thing to note - the signs are opposite, so we'll get a negative number, which implies an attractive force. If the charges were the same, the force would be the same but act in the opposite direction (i.e. repulsion). Secondly, while it's true that the electrons are also going to repel each other, it's relatively small for our purposes and we're going to ignore it. Since the electronic charge (q1) is constant (-1), the magnitude of the interaction is going to be extremely dependent on the effective nuclear charge. Recall that Zeff is # of protons-all shielding electrons. So in the case of sodium, q2 would be +1 since its one valence electron in the 3s subshell experiences shielding from the 10 shielding electrons that are in between. To reiterate, Coulomb's force is directly proportional to q2 (Zeff), so if it's very high then the force between the electron and the nucleus is strong. Both concepts are great at explaining periodic trends since both concepts deal with the strengths between electron and nucleus and the distance.