Periodic Trends - Notre dame Chemistry



Periodic TrendsAtomic and Ionic RadiusNote: Coulomb’s Law is explained thourghly in question 1. In the other questions it is mentioned. Be aware that you would be able to do the same, explain fully to show understanding, then mention. Explain why atomic size decreases from Na to Cl in the periodic table.The trend, as you go across a period, is that the radius decreases. This is because all electrons are being added to the same shell (the same distance from the nucleus) and the nucleus is becoming increasingly positive due to addition of protons (effective nuclear charge increases). Since Cl has more protons in the nucleus, it pulls electrons towards it more tightly. This is the supported by Coulomb’s Law: Coulomb's law states that: The magnitude of the electrostatic force of attraction between two point charges is directly proportional to the product of the magnitudes of charges and inversely proportional to the square of the distance between them. The force is along the straight line joining themExplain why the difference between the atomic radii of Na and K is relatively large compared to the difference between the atomic radii of Rb and Cs.Sodium’s last electron is added to the 3rd energy level, K’s to the fourth. The difference between the size of Rb and Cs is between the fifth and sixth energy levels. The difference in energy of the lower shells is greater than the difference in energy of the higher shells. Also, due to electron shielding the effect of the nucleus on the outer shell electrons is much greater for sodium than it is for Rb and Cs. Coulomb's lawExplain why a Ca atom is larger than a Zn atom.Though Zn has more electrons than Ca, they were added to the third energy level, not the fourth. Both have the same number of outer electrons in the fourth energy level and Zn has a greater effective nuclear charge due to the larger number of protons in its nucleus. This makes Zn smaller than Ca. Coulomb's lawThe radius of the Ca atom is 0.197 nm; the radius of the Ca2+ ion is 0.099 nm. Account for this difference.Ca has two electrons in the fourth energy level (4s2), while the Ca2+ ion has lost these two electrons and has a full third energy level. Both have the same number of protons. Therefore, with fewer electron shells, the calcium ion will be smaller. Coulomb's lawExplain why the ionic radius of N3- is larger than that of O2-.N3- has 7 protons and 10 electrons, O2- has 8 protons and 10 electrons. Both have identical electron configurations, but oxygen ion has a more positive nucleus and therefore pulls the electrons in a little more tightly. Coulomb's lawCa2+ and Cl- are isoelectronic. Which has the larger radius? Explain why.Ca2+ has 20 p and 18 e, Cl- has 17 p and 18 e. Electron configurations and shielding are the same, Ca2+ has a greater effective nuclear charge and therefore its radius will be a little smaller.Coulomb's lawIonization EnergyIonization Energy (kJ / mol First SecondK4193050Ca5901140Explain the difference between Ca and K in regard to:Their first ionization energiesIt takes more energy to remove the first electron from Ca than it does to remove one electron from K. This is because both electrons are being added to the same energy level (same distance from nucleus, same amount of shielding), but Ca has a more positive nucleus due to the additional proton so it holds its electron to it more tightly. Coulomb's lawTheir second ionization energiesThe second ionization energy of K is much higher than the second ionization energy of Ca. This is because Ca has two electrons in the 4th energy level while K has only one. The 2nd electron removed from K is removed from the third energy level. This energy level is closer to the nucleus so the electron is held much more tightly to the nucleus than the fourth energy level electron is. There is also more electron shielding between the nucleus and Ca’s 4s electron than there is between the nucleus and the 3rd energy level electron removed from K. Coulomb's lawThe first ionization energy of Mg is 738 kJ/mol and that of Al is 578 kJ/mol. Account for the difference.It takes more energy to remove the most energetic electron from Mg than it does to remove the most energetic electron from Al. It is slightly easier to remove the first p orbital electron than it is to remove the s orbital electron. This is because the p orbital is a little higher energy than the s orbital is those they are in the same n. Coulomb's lawWhy is the first ionization energy of K lower than that of Li?It takes less energy to remove an electron from K than it does from Li because the outer electon in K is in the fourth energy level and the outer electron of Li is in the second energy level. The K electron experiences a lower effective nuclear charge to the increased electron shielding and its greater distance from the nucleus. Coulomb's law Below are ionization energies for third period elements. First Ionization Energy (kJ/mol)Second Ionization Energy (kJ/mol)Third Ionization Energy (kJ/mol)Element 1125123003820Element 249645606910Element 373814507730Element 4100022503360Which element is most metallic in character?The most metallic element is located furthest to the bottom left of the periodic table. Since all elements are in the same period, the most metallic is the one closest to the left. Element 2 has the lowest first ionization energy so it must have the fewest protons in its nucleus and is further to the left of the period. Also it has a very large increase between its first and second ionization energy which means the second electron is removed from a lower energy level and it has only one electron in its outer shell.Identify element 3. Explain your reasoning.Element 3 must be Mg. The very large increase in ionization energies is between removal of the second and third electrons. This means the element has two electrons in its outer energy level and the third electron is removed from the second energy level. Mg has 3s2 electrons in the third energy level.Write the electron configuration for element 3.1s22s22p63s2What would be the expected valence number for the most common ion of element 2?Element 2 has one electron in the third energy level (see part a) and it will lose this electron to form an Na+1 ion which is isoelectronic with Ne.What is the chemical symbol for element 2?NaA neutral atom of which of the four elements has the smallest radius?The element to the right of the period will have the most protons in its nucleus and this greater effective nuclear charge will cause it to pull its electrons in a little more closely and have the smallest size. Since element 1 has the highest first ionization energy, it has the most positive nucleus.Electron Affinity11. Which element has the most negative electron affinity: B, Al, C, Si? Explain why.The element with the most negative electron affinity releases the most energy when an electron is added to the neutral atom and therefore has the greatest electron affinity. The element with the greatest attraction for an electron, is pulling an electron in at a lower energy level, closer to the nucleus where there is less electron shielding. It will also have more protons in the nucleus. Since Al and Si are adding an electron to energy level 3, they will have lower electron affinity than B and C (energy level 2). Carbon has more protons in its nucleus than B and this greater effective nuclear charge will give it a higher electron affinity. Coulomb's law12. Why are the electron affinities of group 2A elements lower than 1A elements (why don’t they follow the expected trend)?Adding an electron to group 2A elements would be placing the first electron in the p orbital. Since p orbital electrons are slightly higher energy, the electron affinity for these electrons is a little lower than it is for the s electron. Coulomb's law13. Why are the electron affinities of group 5A elements lower than 4A elements (why don’t they follow the expected trend)?Adding an electron to group 5A elements would be adding the first paired p electron. Because there is a slight repulsion between like charged electrons, the electron affinity for this electron is a little lower than it is for the unpaired p electrons. 14. . Trend Chart Make a outline of the Periodic Table on your paper. Draw in the trends on (or near) the periodic table: Ionization energyelectronegativityatomic radiuselectron affinity shielding effectNote: Ionization energy and electron affinity have same trend.15. Explain why noble gases are inert and do not form ions.First, make sure you know what inert means: nonreactive. Noble gases have all orbitals completely filled with electrons. Atoms become ions by giving away or receive electrons in order to obtain a full set of orbitals (octet rule) because energetically this is the most stable electron configuration for an atom. Since noble gases already have this, they will not react with other atoms and will not form ions, hence they are inert.16. Will the shielding effect be more noticeable in metals or non metals? Explain your answer.The shielding effect is the result of non-valence electrons dampening the positive charge of the nucleus as felt by the valence electrons. The more non-valence electrons, or electrons on the inner energy levels, the stronger the shielding effect. That being said, electrons in d-orbitals and f-orbitals do not provide as much of a shielding effect as electrons in s and p orbitals. When comparing metals and nonmetals across the same period, the metals will have the more noticeable shielding affect. As you move across the PT, the number of p+ in the nucleus increases thus increasing the nuclear charge of the atom, without adding more s or p electrons to the inner energy levels. The shielding effect of the inner core becomes less noticeable with the nonmetals because the valence e- are pulled more tightly towards the nucleus.17. Why do elements in the same family generally have similar properties? Choose one family as an example to support your reasoningHalogens Family – The number of valence e- determine the element’s chemical properties such as reactivity. Each member of the halogen family has 7 e- in its valence shell making them very reactive since only one more valence electron is needed to obtain an octet. Therefore, the members of the halogen family all have similar chemical properties. As far as physical properties, the halogens are a difficult group to examine since they represent all three states of matter, but they are similar in odor and being a distinct color. 18. When going down a column (group, family) that has a mixture of states of matter, why are the metals always at the bottom of the group?To become a metal, an element has to be able to become a cation, meaning a loss of an electron. A nonmetal is opposite. To lose an electron requires a lot of energy unless the valence electrons are far from the positive charge of the nucleus. The atoms at the bottom of the periodic table have 5, 6 or 7 energy levels which puts a lot of distance between their nuclei and valence electrons (Coulomb’s Law). Therefore, they have low ionization energies and will more easily become cations than anions making them metals. The atoms at the top of the table have a greater effective nuclear charge with very high ionization energies and large (more negative) electron affinities because there is not much space between their nuclei and valence electrons. They will be come non-metals. Almost ever metal is a solid but no metals are gases. To be a gas, a substance must be a non-metal with the ability for covalent bonding, something only elements with high electron affinities are able to do.19. What element am I? Support your final answer with facts to show it is not merely a guess Clue #1) I have a high electron affinity, (highly negative value), and my atomic number is X. Must be a non-metal since nonmetals have high electron affinities Clue #2) The element with atomic number X-1 has a lower ionization energy and a lower electron affinity. The element to the left is less reactive based on the lower electron affinityClue #3) The element with atomic number X+1 has a higher ionization energy and basically no electron affinity (positive value). The element to the right has no electron affinity. Must be next to the noble gases.d) Within my group, I have the second highest ionization energy. In the halogens, only fluorine has a higher ionization energy. What element am I? chlorineChlorine20. In terms of electron configurations and shielding, why do atoms get smaller as you move across a period?This is similar to question 16. Explain what shielding is and why adding additional electrons does not help with shielding since the additional electrons are not inner electrons. [however, d&f electrons are considered inner electrons but do not assist with shielding]. This makes the nuclear charge increase with increasing atomic number without an increase of shielding, hence, the atoms get smaller across a period. ................
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