hi everyone today we're going over chapter two of general chemistry which covers the periodic table chapter 2.1 says that there are two levels of organization in the periodic table the first is a group which goes vertically down a periodic table so this is group one and this would be group two so what's similar about all the elements in a group is that they have the same valence shell configuration for example in group one all of these elements have one electron in their valence shell and in group two all these electrons have two electrons in the valence gel the second level of organization is called a period so this would be a period and this would also be a period what the period tells you is the energy level that the valence electrons are at so hydrogen is in period one and so the valence electron of hydrogen is in the energy level n equals one and lithium and beryllium are in the second period so their valence electrons would be in the energy level n equals two there are two different types of elements one is called the representative elements and the other is called the non-representative elements so group 1 and 2 as well as 13 to 18 are known as the representative elements and so this means that the behavior of these groups can be predicted based off of where they are on the periodic table so for example um sodium and potassium can both react with chlorine to form a salt and this is very predictable because sodium and potassium have the same number of valence electrons and chlorine would behave fairly similarly in a reaction to bromine because they have similar amounts they have the same amount of valence electrons however the non-representative elements are all of the transition metals and the lanthanides and actinides and this is because the transition metals are more fluid in the amount of valence electrons they can have so you can't necessarily say that um two transition metals in the same group will react similarly chapter 2.2 talks about the characteristics of three broad groups of elements the first are the metals which are shown in yellow on this periodic table to the right and metals are lustrous which means that they're shiny they're malleable which means they can be hammered or pulled into different shapes and they're ductile which means they conduct electricity well all of these characteristics are because metals have multiple stable oxidation states which means a metal can exist as a plus one ion plus two plus three and so on um depending on the specific metal some metals can have up to seven different stable oxidation states which means they're very flexible in how many electrons they can give up or receive so this characteristic makes metals malleable because metals gen are considered to be swimming in a sea of electrons as in if you have a lump of metal all of the metals are sharing all of the electrons together so the bonds are not as rigid and therefore they can slide past each other in the sea of electrons this also makes metals ductile um because the electrons are free to move around and therefore they're free to conduct um throughout a thin wire of metal metals have low electro negative negativity which means that they don't like to receive electrons as much as they like to give up electrons so you'll more frequently see metals in the positive form um as in you'll see them as a cation and almost never as an anion non-metals are shown in blue on this periodic table and non-metals have high ionization energy which means that it takes a lot of energy to pull an electron off of a nonmetal they have high electron affinity which means they're very likely to grab an electron from a metal and they have very high electronegativity which means that in a covalent bond um non-metals are more likely to be hogging the electrons more of the time than a metal wood nonmetals are brittle so non-metals tend to form covalent bonds and therefore their bonds are very much locked into a specific angle so they're not malleable and if you try to hammer a non-metal compound it would most likely crack nonmetals are also poor conductors because they're not swimming in a sea of electrons they have very defined bonds and therefore electricity doesn't really pass through them easily metalloids or semi-metals are shown in green and so an easy way to remember where the semi-metals are and where the metals and non-metals separate is with the staircase rule so if you draw a staircase on the periodic table here all of the elements that touch the staircase are semi-metals with the exception of aluminum and it's easy to remember aluminum as an exception because we know aluminum as a very common metal so metalloids it's harder to predict the characteristics of metalloids some metalloids share more non-metal characteristics and some metalloids share more metal characteristics but in general the characteristics of semi-metals are somewhere in between that of metals and non-metals chapter 2.3 talks about some of the trends that can be found on the periodic table and goes more into depth on the characteristics that we discussed in the last slide so first we have to define a term called the effective nuclear charge which is denoted as z sub eff and the affected nuclear charge is basically the amount of attraction that the valence electrons of an atom feels so i have two elements here as examples i have lithium and beryllium and they're in group one and two respectively so lithium is the third element which means that it has three electrons and it has a nucleus with three protons giving it a charge of positive three so um so with this charge of positive three there are two electrons in the first shell which means that this first gel has a charge of negative two so if you look at the valence electron um you take the number positive three and you subtract two because these two electrons are effectively shielding the valence electron from the positive three charge of the nucleus so lithium has a effective nuclear charge of positive one which is the charge that this electron here would feel so beryllium is the fourth element having four electrons and four protons so these two electrons would shield the two valence electrons giving the effective nuclear charge of positive two so these two electrons would feel a positive charge of positive two um so the effective nuclear charge increases from left to right in a period and you can see why that is here because um because from left to right all of these elements are at the same energy level or they're at the same shell so as you go from left to right on the periodic table you keep adding electrons in this valence shell here and the the amount of positive charge in the nucleus keeps increasing but there isn't an increase in shielding because this first energy level only has two electrons and more electrons cannot be added to it so the effective nuclear charge increases from left to right on the periodic table another definition is the noble gas configuration so the noble gas the noble gases are this group in blue on this periodic table and the noble gases are especially stable because the noble gases have um all of their shells filled so as in like the valence shell of the noble gases is completely full which makes them the most stable so all of the other elements um seek to gain this noble gas configuration so the first periodic property is atomic radii an atomic radii decreases from left to right atomic radii is how big um the electron is so from here to here this distance is the atomic radii so atoms get smaller as you go from left to right and this is because the effective nuclear charge increases from left to right so as you know positive and negative charges attract so this electron here feels a positive charge of plus one so this electron is only vaguely attracted to the nucleus beryllium on the other hand has an effective nuclear charge of positive two so these valence electrons are more attracted to the nucleus which means they come closer to the nucleus giving it a smaller atomic radii um atomic radii increases from top to bottom so it increases down a period and this is because as you go down a period you increase in the principal energy level or you increase in the number of shells you have so if you have electrons in outer shells this makes the radius larger and it makes the atom larger and in addition when you have more electrons the electrons will repulse each other and so when the electrons repulse each other the only way they can go is out um on the other hand ionic radii which is the radius of an atom if it became an ion is slightly more complicated because the effective nuclear charge is the same for example if i removed this electron here this electron would still feel an effective nuclear charge of plus two however there are now less valence electrons so there is less repulsion so cations get smaller and when you remove electrons from a valence shell the atom gets smaller um however if you have an anion which means if you add an electron here the effective nuclear charge is still positive too however there's more repulsion now so the atom gets larger as the electrons move out to try to get further away from each other the next periodic property is ionization energy ionization energy is how hard it is to remove an electron from its valence shell so this property increases from left to right and this is also because of the effective nuclear charge so for example lithium has an effective nuclear charge of plus one which is not very strong and fluorine has an effective nuclear charge of plus seven which is really strong which means that the valence electrons in fluorine are feeling a strong attractive force from the nucleus so you'll need to add a lot of energy in order to remove an electron from fluorine on the other hand lithium doesn't feel nearly the valence electrons of lithium doesn't feel nearly as much attraction so it's much easier to remove a valence electron from lithium and another reason why it's easy to remove valence electrons from lithium is because if you remove one electron lithium becomes the electronic configuration of helium which is a noble gas and it's very stable so lithium actually wants to have an electron removed so it can become the stable noble gas form and this is also true for sodium and potassium which is why lithium tends to exist as lithium plus one and sodium tends to exist as sodium plus one um and magnesium tends to exist as mg plus two because if you remove two electrons magnesium becomes the electronic configuration of neon which is very stable so this trend of ionization energy also decreases going down because as you go down the valence electrons get further and further away from the nucleus so if the valence electrons are further away they feel less of an attractive force and so it's easier to remove them the next periodic property is electron affinity an electron affinity is defined as the energy released when an electron is added to an atom and this increases from left to right because um of the same reason because of effective nuclear charge and because um because the noble gas configuration is most stable so using the example of lithium again lithium doesn't want to gain an electron because if lithium gained an electron it would get further away from its most stable noble gas configuration but fluorine on the other hand not only does it have a plus seven effective nuclear charge which is already very attractive to an electron fluorine if it gained an electron would become the electronic configuration of neon so a lot of energy is released when fluorine gains an electron and lithium doesn't want to gain an electron the next periodic trend is electronegativity which is basically how much an atom likes to hang on to shared electrons when it's in a bond so for example if sodium was bound with chlorine as in sodium chloride sodium is not very electronegative sodium is actually electropositive which means it likes to give up its electron because it wants to become neon and chlorine is very electronegative because it wants to become argon so because chlorine is far more electronegative than sodium is these electrons actually all go to chlorine so this becomes n a plus and cl minus and this is an ionic bond because chlorine is just that much more electronegative than sodium is um and for example if chlorine was bound to carbon this would this would form a covalent bond because the electronegativities of chlorine and carbon are not that different from each other um but however in this bond the electrons are shared unequally as in chlorine has the electrons closer to itself and further away from carbon because chlorine is more electronegative and so this um this notation here is how you denote electronegativity um this arrow means that carbon would have more of a positive charge because the electrons are being pulled away from carbon and chlorine would have more of a negative charge because the electrons are pulled toward chlorine and so these are this is basically the same trend as electron affinity these are just different definitions because electron affinity is when you add an electron and electronegativity is when two elements are bound to each other chapter 2.4 is about the chemistry of different groups on the periodic table and as you remember from before a group is defined as going up and down vertically on the periodic table and the first group are the alkali metals which are group 1 and it's this group here so alkaline metals are known for having low density and this is because they have a low effective nuclear charge which gives them a large atomic radius so this a large atomic radius means that their mass is spread out over a large area which gives them a low density alkali metals easily lose one electron because if they lose one electron they could become their respective noble gas form so sodium likes to give up one electron to become the electron configuration of neon and potassium likes to give up an electron to become potassium plus one sodium plus one um because the noble gas configuration is the most stable configuration possible and so you'll rarely ever see these um these alkali metals in their elemental form you'll see them as plus one ions and this is why they're known as active metals because in their elemental form they're very reactive and like to lose an electron group two are the alkaline earth metals and alkaline earth metals easily form what are known as divalent cations which just means they have a plus two charge so calcium likes to form calcium plus two and beryllium likes to form beryllium plus two because when they form these ions they can become their respective noble gas forms and so these are also known as active metals um so we're going to skip a bunch of groups here for later and talk about the chocogens so chalcogens have six valence electrons the chalcogens are this group here which includes oxygen and they have six valence electrons and they have a small atomic radius because they're toward the right side of the periodic table giving them a high effective nuclear charge um which means the valence electrons are pulled in very closely so the chalcogens include some non-metals some metalloids and some post-transition metals which just means that they're metals but they're past the transition metals the transition metals are the orange elements here so the chalcogens include oxygen and sulfur which are very very in biologically important elements like you see oxygen everywhere in biology and sulfur is very common as well the next group are the halogens and the halogens are this group here they're group number seven um and halogens are known for being very electronegative because first of all they have a very high effective nuclear charge and also if they gained an electron they would be at their noble gas configuration which is very stable so halogens tend to exist in either their diatomic form which means that two halogens are bound together or in their ionic form which means that an electron has been added to it so the next group are the noble gases and the noble gases are very inert which means they're unreactive and this is because the noble gases are already in the most stable possible conformation and they don't want to do anything to change that so noble gases have very low boiling points um because noble gases are like not really likely to interact with anything else so they disperse and become gases at room temperature which is why they're known as the noble gases last but not least the transition metals the transition metals are all of the elements in orange here and because they're metals they have very low electron affinity which means that they don't like to gain electrons and the transition metals have multiple oxidation states possible so for example manganese here can exist in either its plus two form or plus three or plus four or plus six or plus seven um and all of these oxidation states are stable from ankenys so you can find manganese in the wild in any of these oxidation states but these transition metals almost never become their negative form so you will never see a negative form of manganese because they have low electron can they have low electron affinity and as always thank you so much for watching and i hope you enjoyed this video