in this video we're going to go over periodic trends such as atomic and ionic radius electron negativity ionization energy electron infinity and metallic character so let's start with atomic radius the sizes of atoms increases as you travel in this Chan no direction of the periodic table that is as you go to the left the size of atoms increases and as you go down the atoms get bigger so let's take two examples hydrogen and helium on a periodic table hydrogen is to the left of helium so that means that hydrogen is bigger than helium but now why is that the case because if you look at the atomic mass for helium in the periodic table it has an atomic mass of four and for hydrogen the atomic mass is one so how is it that hydrogen is bigger when helium is more massive to understand this we need to jump into physics for a moment you know that like charges repel and opposite charges attract so if you have a negative charge and a positive charge these two will feel a force of attraction now if you increase the number of charges let's say if you have two protons and one electron the four force will be two times as large because the charge was increased the more protons you have the greater the force of attraction between the protons and electrons the second factor to consider is distance if you have a proton and an electron that are very far apart the force that they fill will be relatively weak but let's say if you cut the distance in half so let's say currently they're about 4 M apart but if you bring the proton and the electron closer let's say if they're 2 met apart if you decrease the Distance by a factor of two the force of attraction between these two particles will increase by a factor of four so the electrostatic traction is much larger if the charges are brought closer to each other so if you increase the charge of any of the protons or electrons the force of attraction increases so the more protons that a nucleus has the greater the force of attraction between the nucleus and the electrons and the second factor is distance if you increase the distance the force of attraction will decrease and if you decrease the distance the electrostatic force of attraction will increase now let's think about what this means inner core electrons that are very close to the nucleus they will feel a very strong force of attraction between itself and the nucleus whereas electrons that are away from the nucleus like valence electrons the force of attraction between those electrons and the nucleus will be relatively weak since they're further way now hydrogen has only one proton and its nucleus so the charge of the nucleus is plus one and it has one electron now helium has two protons in its nucleus and it has uh two electrons so because the nuclear charge of helium is greater than that of hydrogen the force of attraction between those electrons and the nucleus is much larger than helium as a result the size of the electron cloud shrinks as those electrons are brought closer to the nucleus and that's one reason helium is smaller than hydrogen even though it has a greater atomic mass because the charge of the nucleus increases the size decreases so as you travel from left to right across the periodic table the atomic size will decrease as the number of protons increases now let's say if we were to compare lithium versus hydrogen which atom is bigger now lithium has three protons and hydrogen has one yet lithium is bigger than hydrogen why is that I mean don't lithium have more protons now the trend as you go from left to right the size decreases because the number of protons increases but the number of shells is the same here hydrogen and lithium are not in the same row they're in different rows here's hydrogen here's helium and lithium is below hydrogen now for two elements in the same Row the one that's on the right is usually the smaller atom but now for two atoms that are in the same group the one that's below it is usually the bigger one remember atomic size increases as you go down so now let's talk about why why is it that lithium is bigger than hydrogen so hydrogen has one proton in its nucleus and it has one electron now lithium has an atomic number three so it has three protons in his nucleus now in the first shell it can only fit two electrons these are core electrons so it has to put the third electron in the second shell and because that electron is further away from the nucleus the force of attraction between this electron in a nucleus is reduced remember if you increase the distance the force of attraction decreases so the reason why lithium is bigger than hydrogen is because it has two electron shells as opposed to one and that's why as you go down the group in a periodic table atomic increases because for every row that you add you add a new shell of electrons and so the atom gets bigger now there's another reason that you need to consider and that is the concept of shielding the inner two core electrons partially Shield the outer veence electron from the nucleus now this electron is attracted to the three protons in the nucleus as we said before whenever you have opposite charges they attract each other so what if you have like charges let's say two electrons like charges repel each other so this electron is repelled by these two electrons so even though it feels a force of attraction between itself and the nucleus it's also repelled by the other electrons between it and the nucleus so you have to consider also the shielding effect of the inner core electrons now let's work on some examples let's say if we have three elements chlorine magnesium and phosphorus rank the following elements in order of increasing atomic size so that's from small to large so which of these atoms is the smallest and which one is the largest so the first thing you want to do is you want to look at the periodic table and you want to place them in order magnesium is in the third row of the periodic table and it's in group two it's an alkaline earth metal phosphorus is also in the same role and it's in group 5 a and chlorine is a halogen in group 7A now because these three are in the same row this is going to be fairly easy we know that the atomic size increases or rather it decreases from left to right so that means that magnesium is bigger than phosphorus and phosphorus is bigger than chlorine so if we want to rank it in order of increased in atomic size we need to put the smallest first and then the largest last so magnesium is bigger than phosphorus which is bigger than chlorine let's try another example let's say if we have calcium burum and strontium rank the following elements in order of decreasing Atomic radiat so notice that these three elements are all alkaline earth metals they're found in group two so let's place them in order so first is be below be is magnesium and then below that is calcium and then strontium now we know that atomic size increases as you travel down a group now we want to rank it in order of decreased in atomic size so we know that strontium is the largest brillium is the smallest and calcium is in between so to put it in decrease in order we want to rank it from large to small so we're going to start with strum then calcium and then calcium is bigger than burum and so that's how you can rank it in order of decrease in atomic size now let's try one more example so this time let's say if we have four elements Fe CF sulfur and helium rank the following elements in order of increasing atomic size so the first thing you want to do is you want to place every element in the respective positions based on the periodic table cesium is an alkaline metal in group one but it's in the sixth row so it's like over here Fe is a transition metal and it's in a fourth row so it's somewhere in this vicinity sulfur is a calogen found in group 6A and helium is the noble gas that's in the first row so remember atomic size increases as you go to the left and as you go down so basically it increases in this general direction which tells us that cesium is the largest then it's Fe then it's s and then it's helium so we wish to rank it in order of increase in atomic size so that's from small to large so helium is smaller than sulfur and sulfur is less than Fe and Iron Metal is smaller than an atom of cium so this is the answer now let's move on to ionic radi so consider an atom of lithium and a lithium ion an atom is a particle that have an equal number of electrons and protons so atoms are electrically neutral ions have unequal number of protons and electrons and so they have a charge an ion with a positive charge means that there are more protons than electrons so lithium is an atom with three protons so the charge of the nucleus is three and as we mentioned before it has two electrons in its first shell and one valence electron in its second shell the lithium plus ion which is also known as a cation cat are positively charged ions it has three protons but it lost an electron so basically it lost an entire shell so therefore positively charged cats are significantly smaller than their parent atoms because they have less electrons and therefore less electron shells they will be uh significantly smaller now what about anion or negatively charged ions how do the sizes of these ions compare with their parent atoms so let's use nitrogen as an example and the nitride ion nitrogen has seven protons in its nucleus and an atom of nitrogen has seven electrons so in the first shell it's going to have two and in the second shell it's going to have five so nitrogen has two core electrons or inner electrons and it has five valence electrons the valence electrons are the electrons in the outermost energy level now the nitride ion has three more electrons than protons it has seven protons and 10 electrons protons have a positive charge electrons have a negative charge if you add these two numbers you're going to get a net charge of minus three so let's draw the picture for the nitride ion now the first shell contains two electrons but the second shell has a total of eight electrons so that the total number of electrons is 10 so because of the extra three electrons there's more electron repulsion which causes the second shell to expand so negatively charged ions or anions are significantly larger than apparent atoms so make sure you remember this for Ionic radi ions with a positive charge will be relatively small and ions with a negative charge will be relatively large now the trend for Ionic radi is very similar to that for Atomic rad meaning ionic rad increases as you go down and it increases as you go towards the left this is especially true for ions with similar charges if you're comparing two ions with a positive charge or two ions with a negative charge let's look at the elements of ions in a row the sodium ion is bigger than the magnesium ion both these two ions are is electronic they have the same number of electrons but because magnesium has more protons than sodium magnesium is going to be smaller and aluminum which has a plus three charge is even smaller than magnesium so as you can see as the charges increases the size of the ion decreases and as you travel to the right it decreases to the right but ionic radi increases towards the left now as you jump from the positive ions towards the negative ions the size will greatly increase for example phosphide which has a neg3 charge it is significantly bigger than almost all of these positively charged ions on the left and then the size decreases as you go towards the right so the sulfide ion is smaller than the phosphide ion and then the chloride ion which let me see if I could fit it here is even smaller than the sulfide ion so the size generally decreases as you go from left to right but if you're comparing a negatively charged Ion with a positively charged ion generally speaking the negatively charged ion will be bigger than the positively charged ion now what about as we go down a group ionic size increases as you go down for example the lithium ion is smaller compared to the sodium ion the sodium ion is larger and the potassium ion is even bigger than the sodium ion so as you go down the number of electron shells increases and therefore the ions will increase in size so if you compare the nitride Ion with the phosphide ion the nitride ion is pretty big but the phosphide ion is even bigger so ionic radi increases as you go down a group now let's try a problem let's say if we have the following ions the fluoride ion the Magnesium cation an atom of neon the sodium cation and the oxide ion now rank the following atoms and ions in order of increasing size now to do that we need to determine if any of these atoms or ions are ISO electronic meaning that if they have the same number of electrons if they're ISO electronic then all you need to do to rank them in order of increase in size is look at the atomic number for particles that have the same number of electrons the ones that have more protons will be smaller in size as you increase the amount of protons in the nucleus The increased nuclear charge will cause the electrons to move towards nucleus decreasing the size of the atom so let's count the electrons first the number of electrons is equal to the atomic number minus the charge so if we consider Florine first Florine has an atomic number of N and a mass number of 19 so it's 9 minus1 so Florine has 10 electrons now what about magnesium if you look at the periodic table magnesium has an atomic number of 12 and a mass number of 24 the smaller of these two numbers is the atomic number so for magnesium it's going to be 12 minus a charge of plus two which is 10 so magnesium has 10 electrons now for neon it's an atom which has an atomic number of 10 a mass of 20 and for atoms they're neutral so they don't have a charge so it's 10 minus a charge of zero Neon 2 has 10 electrons now what about sodium sodium has atomic number of 11 and for oxygen it has an atomic number of eight so for sodium it's 11 - +1 which is equal to 10 and for oxygen is 8 - -2 so oxygen has 10 electrons so all of these particles are isol electronic with each other they have the same electron configuration so now all we need to do is look at the atomic number the atomic number of fluide is N9 for magnesium it's 12 for neon it's 10 for sodium it's 11 and for oxygen it's 8 and as you mentioned before negatively charged ions are generally larger than positively charged ions and one of the most protons or the highest atomic number is going to be the smallest one so oxide we want to rank it in order of increase in size so let's start with the the smallest one magnesium has the most or the greatest number of protons so it's going to be the smallest it's smaller than sodium and sodium is smaller than neon which is smaller than fluoride and the oxide ion is going to be the biggest so let's draw the relative sizes so this is the smallest that's a little bit bigger then this is going to be bigger and then the size just continue to increase so as we can see oxide has the least number of protons so that's why it's bigger keep in mind each of these atoms and ions they have only two shells of electrons so for particles that have the same number of shells the one with the most protons will be the smallest the one with the least number of protons in this case oxide is going to be the biggest so let's compare the magnesium ion and the oxide ion just to get a better understanding of this concept so we said that magnesium has an atomic number of 12 which means that the charge of the nucleus is 12 since it has 12 protons and it has a total of 10 electrons so here's the first shell and here is the second shell which I think I could draw a better Circle so the first shell contains two and the second shell has eight now oxide also contains 10 electrons but the charge is eight since oxygen only has eight protons in its nucleus so the oxide ion is going to be significantly bigger than the magnesium ion and we could see why it's bigger because magnesium has more protons the electrostatic force between the nucleus and the electrons is stronger because of the increased charge and so magnesium it shrinks in size the nucleus it pulls the electrons toward the nucleus making the ion smaller now in the case of oxygen the force of attraction between the outer electrons and the nucleus is relatively weaker because the charge of the nucleus is less plus since the net charge is negative there's a lot more electron repulsion and so the atom expands when you have a lot when you have extra electron repulsion so to summarize what we just went over remember negatively charged ions are usually bigger than positively charged ions and for isoelectronic species that's particles that have the same number of electrons which is going to have the same number of electron shells the ones with more protons will be smaller than the ones with less protons so as you increase the nuclear charge and if the number of shells remain constant the particle size will decrease that extra nuclear charge will cause the electrons to contract towards the nucleus so now let's move on to electr negativity electr negativity is the ability of an atom to attract an electron to itself electro negativity increases towards Florine Florine is the most electr negative element in the periodic table on the upper right corner of the periodic table you have the nonmetals and on the left side you have the metals metals tend to be electropositive they like to give away electrons and form positively charged ions non-metals like to acquire electrons and form negatively charged ions so nonmetals tend to be electronegative electr negativity increases as you go up and as you travel towards the right so Florine is more electronegative than oxygen and oxygen is more electronegative than nitrogen so let's say if you have a Florine atom and if you add an electron to it Florine will turn into the fluoride ion the non-metal will become a negatively charged anine so nonmetals tend to be electronegative they like to acquire electrons now if you have an atom of sodium sodium wants to give away an electron to form a positively charged cation so tend to be electropositive they like to give away electrons so those are some things to know now let's say if you have four elements silicon magnesium chlorine and aluminum rank the following elements in order of increasing electr negativity so look at the periodic table and place them in order so each of these elements are in the same row magnesium comes first it's in group two then aluminum which is in group 3A then silicon that's in 4 a and chlorine is in group 7 a so electro negativity increases towards the right so if we want to rank it in order of increasing electro negativity that's from low to high magnesium is going to be the least electronegative it's a metal and then it's aluminum which is also metal and then silicon which is a metaloid and then chlorine is the most electronegative that's a nonmetal so non-metals are usually more electronegative than Metals so this is the answer now what about these let's say if we have tin germanium lead and carbon rank the following elements in order of decreasing electronegativity now each of these elements are found in group 4 a of the periodic table so let's put them in order first we have carbon and then below that g geranium and then tin metal and then lead electro negativity increases as you travel up a group so we want to rank it in decrease in order so we need to put the most electronegative element first we need to rank it from high to low so the highest is carbon which is is a non-metal and then is geranium which is a metaloid and then tin metal and then lead metal so the metals they have the lowest electr negativity so this is the answer that's how you can rank it in order of decreasing electr negativity now which element is more electronegative silicon or nitrogen if you place these elements in a respective position silicon is in group 4 a in a third row nitrogen is in group 5A in the second row electr negativity increases towards Florine so it increases as you go up and to the right so therefore you can clearly see that nitrogen is more electronegative than silicon now what about nitrogen versus sulfur if you place them in their respective positions nitrogen is in group 5 a in the second row so sulfur is in group 6A but in the third row but electro negativity increases this way so which one is more Electro negative so as you travel up electro negativity increases but as you travel to the right it decreases so sulfur is to the right of nitrogen but it's also below nitrogen so how can you tell which one is more electronegative in this case it turns turns out that the electro negativity increases more when you go up than when you go to the left so the increase for going up one row is greater than a decrease from traveling one unit to the right now if you ever unsure at this point you can look at the electro negativity table and see which one is higher the electro negativity for nitrogen according to most textbooks is roughly about 3.0 for sulfur it's 2.5 so as you can see going up has more priority than going to the right so let's go over some common electr negativity values for elements like boron carbon nitrogen oxygen Florine phosphorus sulfur chlorine bromine and iodine or iodine the electro negativity value for Bon is about 2.0 for carbon it's 2.5 for n it's 3.0 for o is 3.5 and for Florine it's about 4.0 now for phosphorus it's roughly around 2.1 for sulfur 2.5 and for chlorine it's about 3.0 for Bromine 2.8 iodine 2.5 so as you can see there's a large increase as you go from row two to row three here the increase is .9 and here it's about one so that's why nitrogen was significantly higher than that for sulfur so if you need to go this way generally speaking this is the one that's going to win but not always though for example bromine is higher than sulfur so there are some exceptions so you may just need to know some of these values if you ever get an unusual question like that by the way the electro negativity for hydrogen is about 2.1 so even though hydrogen is to the left of boron Boron is in row two hydrogen is in row one so hydrogen is a little bit more electronegative than Boron now the next train that we need to talk about is metallic character metallic character increases in this general direction towards the metals the non-metals are located in the upper right corner of the periodic table and the metals is just to the left but metallic character increases as you travel to to the left and down across the periodic table so let's go over some examples so let's say if we have elements such as silicon sodium sulfur aluminum and chlorine rank the following elements in order of increase in metallic character so as we've been doing before let's place it based on their respective positions in a periodic table now each of these elements are located in the same Row in the periodic table sodium is found in group one and then aluminum in group 3A silicon in group 4 a sulfur in group 6A and then CL in group 7 a so metallic character increases as you travel towards the left on a periodic table so if we wish to rank it in order of let's say increasing metallic character we need to rank it from low to high so chlorine which is a non-metal is going to have the lowest metallic character selfer is also a non-metal silicon is a metaloid aluminum is a metal and sodium is a metal now let's think about what this means sodium really wants to get rid of its electrons more than aluminum both of these elements are metals but sodium has a greater metallic character than aluminum so sodium is more electropositive it wants to get rid of its electrons with a a stronger Force than aluminum now between sulfur and chlorine chlorine is more electronegative so it's less metallic than sulfur silicon is a metaloid it's in between a metal and a non-metal so this is the answer that's how you can rank it in order of increas in metallic character now metallic character increases as you go down in group let's consider the group 4 a elements like carbon silicon tranium tin and lead as you go down the periodic table the elements change from non-metals to metals carbon is a nonmetal silicon is a metaloid and geranium is a metaloid but tin is a metal and Lead is a metal so as you can see metallic character increases as you travel down a group let's try one more example let's say if we have gallium manganese nitrogen helium and francium rank the following elements in order of decreasing metallic character francium is located in group one in the seventh row and then you have maganese which is a a transition element and then gallium is to the right of that and then there's nitrogen and helium so metallic character increases in this general direction so if we want to rank it in decreas in order we need to rank it from high to low so francium has the greatest metallic character then it's manganese and then GA and then n and then h e the next topic that we need to talk about is ionization energy I ization energy is the energy required to remove an electron from a gaseous atom now generally speaking it's easier to remove an electron from a metal than it is to remove from a nonmetal metals like to give away electrons so it doesn't require that much energy to remove an electron from it non-metals like to acquire electrons so it's harder to remove an electron from a nonmetal the trend for ION is ization energy is in this direction ionization energy increases towards helium it increases as you go up across the periodic table and towards the right so the main reason why it increases as you travel from left to right is due to the increase in nuclear charge from left to right the number of electron shells or the principal quantum number Remains the Same however the number of protons in the nucle increases and as you add more protons to the nucleus those protons will have a tighter grip on the electrons so it requires more electrons I mean it requires more energy to remove those electrons so let's compare lithium and burum lithium has three protons in its nucleus and buril has a nuclear charge of four now keep in mind burum is smaller than lithium because burum has a higher nuclear charge so here's the question for you is it easier to remove the electron from lithium or a valence electron from burum it's going to be easier to remove the electron from lithium the ionization energy for lithium is about 520 but for brillium it's about 899 so it requires much more energy to remove an electron from brillium than to remove it from lithium for one thing the force of attraction between the electron and the nucleus is weaker compared to the force of attraction between buril an electron in buril then and the nucleus the nuclear charge of burum is much greater than that of lithium so the nucleus has a tighter grip on this electron so that's one reason why it's harder to remove it but also burum is smaller than lithium and so the distance between the electron and the nucleus is much less than the veence electron and the nucleus within lithium so remember if you decrease the distance the force of attraction increases so because burum is smaller and plus the fact that it has a higher nuclear charge means that burum holds on to that veence electron with a tighter grip than lithium and so that's why it requires so much more energy to remove that veence electron in Brum it has to do with the increased nuclear charge and the smaller size both of those factors increases the ionization energy of be now what about between Lithium and sodium which element has a higher ionization energy sodium is below lithium and ionization energy increases as you go up the ionization energy for sodium is about 495 but for lithium it's 520 so it doesn't vary much now lithium as we mentioned before has a nuclear charge of three and sodium has 11 protons so sodium is bigger than lithium it has three electron shells so why is it easier to remove a veence electron from sodium than it is to remove it from lithium considering that sodium has a higher nuclear charge now granted as the charge increases we know that the electrostatic force of attraction increases so an atom with a higher nuclear charge which means that it's it has electrons that are harder to remove however you have to take into account distance the distance between this electron and a nucleus is significantly larger than the distance between this electron and the nucleus so if you increase the distance the force greatly decreases and so because this electron is very far away from the nucleus is relatively easy to remove that electron and that's why sodium has a slightly lower ionization energy than uh lithium is because of the increased distance between that electron and a nucleus even though it has a higher nuclear charge now we could see that distance plays a greater role in affecting the force of attraction than the charge because the electron is further away the ionization energy is less even though the nuclear charge is greater and this is consistent with colum's law his law describes the relationship between electrostatic force between two charged particles so let's say if you have a proton and if you have an electron the force of attraction between these two particles is proportional to the charge and inversely related to the distance between them which is R so if you double the value of Q the force will double in value however if you double the distance because it's squared 2 squar is four the force will reduce by a factor of four so the distance between the charges has a greater impact than the magnitude of the charge based on the equation so to summarize what we've learned as we travel down a group the ionization energy decreases because the distance between the protons and the valence electrons increases and as we travel to the right ionization energy increases due to the increase in the effect of nuclear charge as you go to the right the principal quantum number Remains the Same so the number of electron shells is roughly about the same now there are some exceptions that you need to be aware of for example between burum and Boron the ionization energy for Boron is 800 and for burum the first ionization energy is 8.99 NOW Boron is to the right of burum in the periodic table so typically we should expect that Boron should have an higher ionization energy but it doesn't the question is why now the last electron in buril is the 2s2 electron however the last electron in boron is the 2p1 electron 2p on average is farther away from the nucleus than 2s if you write the electron configuration it's 1 S2 2 S2 and then 2 P6 but if you draw the orbital diagram energy levels it looks like this so 2p is higher in energy which means that it's further away from the nucleus and as you mentioned before if you increase the distance between the electrons and the nucleus on average the ionization energy tends to decrease and so that's why we see a slight drop as we go from be to Boron the same is true from magnesium to aluminum as you go from the S2 level to the the first P level there's a temporary decrease even from calcium to gallium so from magnesium to aluminum it changes from 735 to 580 which is pretty significant and from calcium to gallium from group 2 a to group 3A it drops from 590 to 570 so I mean 579 so the change between these two is not that great but these are exceptions that you want to be aware of if you're tested on it now there are some other exceptions that you need to be aware of as well it's not over yet and that's going from nitrogen to oxygen or phosphorus to sulfur or arsenic to selenium the first ionization energy for nitrogen is 1402 and for oxygen is 1314 at according to the textbook I'm using now why does it temporarily decrease from n to the electron configuration for nitrogen ends in 2p3 and for oxygen it ends in 2p4 now for nitrogen it has three unpaired electrons but for oxygen it has a paired electron this paired electron is the one that's being removed and due to the electron repulsion between the two electrons in this orbital that allows this electron the one being removed it allows it to be removed with ease because that electron is being repelled by the other electron in that orbital it doesn't take much energy to remove that veence electron and that's why we see uh this drop in ionization energy is due to the electron repulsion between these two paired electrons now let's try some examples rank the following elements in order of increasing ionization energy that is first ionization energy so let's say if we have the elements gallium bromine potassium chromium and arsenic so once again you want to look at the periodic table and you want to place them in their respective positions so each of these elements are in the first not the first row but the the fourth row the first one is potassium and then after potassium it's uh chromium which is a transition metal and then you have gallium and then Arsenic and then bromine which is a nonmetal the first ionization energy increases as you travel left to right on a periodic table so let's say if we want to rank it in order of increasing ionization energy we could simply write it like this so This Is The Answer now what about these three phosphorus arsenic nitrogen and antimony go ahead and rank these elements in order of increasing or rather decrease in first ionization energy now these elements are all found in group 5A of the periodic table so if we place them in order first we have nitrogen below that is phosphorus then Arsenic and then after that antimony now the first ionization energy increases as you go up a group within the periodic table so if we want to rank it in decrease in order from high to low we need to start with the highest which is nitrogen and that's greater than phosphorus which is larger than Arsenic and which has a higher first ionization energy than animony so that's how you can rank it in order of decrease in first ionization energy now let's try another example let's say if we have Florine phosphorus helium francium and vadium so let's put this in order so francium is all the way at the bottom towards the left it's in the seventh row First Column vadium is a transition metal which is in this region in the periodic table and then we have phosphorus Florine and helium is to the upper right of Florine so first ionization energy increases towards helium so therefore let's say if we want to rank it in order of increase in ionization energy from low to high let's start with francium Which is less than vadium that's less than P Which is less than F and that's less than h e so it helps if you can place the elements based on where they're located in a periodic table and then use the trend to rank them in the appropriate order now sometimes you may have questions about the second ionization energy and the third and so forth the second ionization energy is the energy required to remove the second electron that second electron may be a valence electron or may be a core electron so let's consider aluminum aluminum is in group 3A and it has a total of 13 electrons its atomic number is 13 so as an atom a neutral atom has 13 protons and 13 electrons now let's draw the aluminum atom so the charge of the nucleus is 13 in the first shell it's going to have two electrons in the second energy level it's going to have eight electrons so right now we have a total of 10 and in the third energy level it has three valence electrons now the first ionization energy of aluminum is the energy required to remove the first valence electron let's say this one so it's going to take about 580 K of energy to remove one mole of electrons from aluminum and that's really one electron per one aluminum atom but it's 580 KJ per mole now for the second electron the first ionization energy is higher it's 1815 and it makes sense why it's significantly higher once you remove the first veence electron aluminum is going to have a net charge of plus one and so there's going to be less shielding between this electron and a nucleus so the nucleus is going to have a tighter grip on that electron plus there a net positive charge and as we know whenever you have a positively charged ion the size decreases so if you decrease the distance between the veence electron and the nucleus the ionization energy will increase the nucleus will have a stronger hold on that electron so the second ionization energy is always higher than the first one and the third ionization energy will be higher than the second once you remove the first two electrons the aluminum ion now has a charge of plus two the energy required to remove the third electron is 2,740 significantly higher than the last one now what about the fourth ionization energy the fourth ionization energy is associated with the removal of an inner core electron and the energy required to take off one of those electrons is much higher it's 11,600 so as you can see because the in core electrons are very very close to the nucleus and because they're less shielded from the nucleus the energy that's required to remove such an electron is significantly higher so because aluminum has three valence electrons the jump in ionization energy occurs or is associated with the fourth electron because that's that's a core electron to make sure you understand the concept of that so the core electrons are very difficult to remove they're too close to the nucleus and so the ionization energy for those electrons are very very high an element has the following ionization energies 735 1445 and 7730 identify the element and the options to choose from are potassium magnesium gallium silicon Arsenic and sulfur now let's look at the ionization energies going from the first to the second the increase is about 700 but from the second to the third the increase is over 6,000 that means the jump occurs after the second ionization energy so the third ionization energy is associated with the removal of a core electron which means the first two are veence electrons so this element has two valence electrons so we have to find out which of these elements contain two valence electrons sulfur is found in group 6A so sulfur has six valence electrons arsenic is found in group five it has five valance electrons silicon has four gallium has three magnesium has two and potassium has one so the answer is magnesium because it contains two valence electrons so the third electron um that's removed from magnesium is a core electron which is associated with this number 7730 now the next topic of interest is electron affinity electron affinity is associated with the energy change that occurs when at an electron to a gaseous atom now when you add electrons to non-metals that really wants electrons particularly electronegative nonmetals they tend to release a lot of energy and whenever energy is released in a reaction you have an exothermic reaction so chlorine has a strong desire for electrons when it acquires an electron it turns into chloride and so it releases a lot of energy so the energy that's released in a reaction is associated with the element's electron affinity ionization energy is the opposite ionization energy is the energy required to remove an electron from a gaseous atom whereas electron affinity is the energy required or the energy that the energy change that occurs when an electron is absorbed by a gaseous atom it's important to know that the halogens they are the most exothermic in terms of electron affinity because they're so electron Negative they really want electrons so if you add an electron to let's say a gaseous Florine atom as it turns into floride it's going to release 327.2 K per Mo so it's highly exothermic now the fact that it releases so much energy means that as it acquires that electron it becomes very stable so if the addition of an electron to a gaseous atom produces a stable ion it's going to be highly exothermic if it produces something unstable chances are it's not going to be very exothermic it might be even endothermic so what are the trends for electron generally speaking keyword generally electron affinity increases or becomes more exothermic as you go to the right now there's a lot of exceptions it's helpful to know these numbers 76 45 13 28 what does that mean these are group numbers group seven is the most exothermic group when you add electron to an element in group seven the halogens they will release the most amount of energy group six is next and then it's group four but not five four is more exothermic than five and then it's one and three and 2 and8 are the least exothermic in fact most elements in two and eight are endothermic you have to put in energy to add an electron now it's easy to see why group seven is the most exothermic because they're the most electronegative they only need one electron to complete their oite the why is group four more exothermic than group five for example carbon which is in group four is more exothermic than let's say nitrogen which is found in group 5A now carbon has the configuration that ends in 2p2 for nitrogen it's 2p3 so if we add an electron to carbon and one to nitrogen which one is going to be more stable currently carbon has two electrons in its 2p orbital nitrogen has three so if we wish to create the C minus gaseous ion we just need to add one electron on to carbon and notice that the orbitals 2p orbitals are still unpaired so because it has this empty space for an extra electron the C minus ion is fairly stable now what about adding an electron to an element in group 5A like nitrogen once we add it notice that there's going to be electron repulsion between the two electrons in this orbital so this is not a good stable Arrangement and it's because of that electron repulsion that's why elements in group 4 a are more exothermic than elements in group 5A when you add an electron to carbon you can create or you will create a fairly stable ion but when you add it to nitrogen the ion won't be as stable due to the electron repulsion and so that's why group 4A elements are more exothermic than group 5A elements in terms of electron Infinity now what about let's say group one versus group two notice the number 76 6 45 13 28 if you add an electron to a group one element it's going to be exothermic but if you add it to a Group Two element most of them are endothermic so let's understand why let's use sodium and magnesium as examples so the configuration for sodium ends in 3s1 for magnesium is 3s two so magnesium has two PA electrons and sodium only has one now the next orbital after 3s is 3p so if we add an electron to sodium that electron will go in this empty 3s well this half filled 3s orbital but if we add an electron to magnesium we have to put that electron at a higher 3p orbital it's more stable to put an electron in a lower energy level than a higher energy level even if there's going to be some sort of electron repulsion so by adding the electron in this half fill orbital the electron affinity for sodium is still relatively exothermic it's much less than four and five but it's still more exothermic than group two now for group two elements it's endothermic because we're putting an electron in a higher energy level and so we got to add energy to put that electron there and so that's why it's endothermic for many of the group two elements and the same is true for group a elements for example let's consider neon neon is in group 8 and it ends in 2p6 so the 2p orbital is completely filled to add an electron we need to add it to the 3s orbital and so because we're adding an electron to a higher energy level it's going to be endothermic we need to add energy to the system to add that electron to that atom so that's why group two and group 8 are endothermic with respect to electron Infinity it's because they're electrons the the electrons in their configuration are completely filled they're completely paired so to add a new electron you got to put it in a higher energy level all of the other groups 7 6 4 5 1 and three you don't have to add an electron to a higher energy level you simply have to add it to an unfilled or half fied orbital and so adding in electrons to half filled or empty orbitals is usually an exothermic process but to put an electron in a higher energy level that's usually an endothermic process so when you're dealing with electron affinity you have to ask yourself if I add this electron will it create a stable ion or an unstable ion is there room for me to put this electron somewhere is there a half filled or an empty filled orbital to put it somewhere in if there is then it's probably going to be exothermic if you got to put it in a higher energy level then chances are it's going to be endothermic so hopefully uh these principles will help you to understand the concept of electron affinity and when it will be endothermic versus when it will be exothermic now let's try one more problem this is going to be the last problem for this video so let's say if we have the following elements chlorine phosphorus argon magnesium sodium and silicon rank the following elements in terms of their Electro I mean their electron affinity values so you want to rank it from endothermic to most exothermic now it helps to know the order 76 45 13 28 so as you travel this way towards group seven it's going to be mostly exothermic towards group 2 and8 it's going to be endothermic so chlorine is in group seven phosphorus is in group five argon is in group eight magnesium is in group two sodium is in group one silicon is in group four so from Endo to EXO let's start with group eight so argon is going to be endothermic and magnesium which is in group two that's endothermic now we don't have any group three elements so we can get rid of that the next one is group one which is uh sodium and then after that we have group five which is phosphorus and then group four which is silicon we don't have any group six elements the last one is group seven which is chlorine so out of the elements listed if we add an electron chlorine is going to be the one that's most exothermic it's going to release a lot of energy whereas argon is going to be the least exothermic or the most endothermic to add an electron to a gaseous atom of argon we need to add energy to get that going so that is it for this video thanks for watching and have a great day