Special Topics for SOL 2 3 rd Power Point Periodic Trends (Chap 14)

Shorthand Electron Configurations Shorthand configurations are a useful tool. Shorthand configurations are a useful tool. Let’s look at an example for Y, Z=39 Let’s look at an example for Y, Z=39 The electron configuration for yttrium is 1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 3d 10 4p 6 5s 2 4d 1 The electron configuration for yttrium is 1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 3d 10 4p 6 5s 2 4d 1 To do a shorthand configuration, we use the noble gas preceding the element and we put that in brackets (the bold and italics part) To do a shorthand configuration, we use the noble gas preceding the element and we put that in brackets (the bold and italics part) That’s Kr and then we also just write whatever is left over. That’s Kr and then we also just write whatever is left over. [Kr] 5s 2 4d 1 [Kr] 5s 2 4d 1

You Try… Do a shorthand configuration for Do a shorthand configuration for Fe Fe Br Br Rb Rb

The Answers… Do a shorthand configuration for Do a shorthand configuration for Fe = [Ar]4s 2 3d 6 Fe = [Ar]4s 2 3d 6 Br = [Ar]4s 2 3d 10 4p 5 Br = [Ar]4s 2 3d 10 4p 5 Rb = [Kr]6s 1 Rb = [Kr]6s 1

Objective B Notice that the halogens all have an ending configuration of ns 2 np 5. That means they have 7 valence electrons. Notice that the halogens all have an ending configuration of ns 2 np 5. That means they have 7 valence electrons. Similarly, alkali metal have 1 valence electron. Noble gases have 8, etc. All transition metals have 2. Similarly, alkali metal have 1 valence electron. Noble gases have 8, etc. All transition metals have 2. F[He]2s 2 2p 5 Cl[Ne]3s 2 3p 5 Br[Ar]3d 10 4s 2 4p 5 I[Kr]4d 10 5s 2 5p 5 At[Xe]4f 14 5d 10 6s 2 6p 5

Objective B All of the transition metals have 2 valence electrons, with 2 exceptions. “d” electrons are not valence electrons. Why not? All of the transition metals have 2 valence electrons, with 2 exceptions. “d” electrons are not valence electrons. Why not? Transition metals are where the d orbitals are being filled up. Here are the electron configurations for all of them. Transition metals are where the d orbitals are being filled up. Here are the electron configurations for all of them. Sc[Ar]3d 1 4s 2 Ti[Ar]3d 2 4s 2 V[Ar]3d 3 4s 2 Cr[Ar]3d 5 4s 1 Mn[Ar]3d 5 4s 2 Fe[Ar]3d 6 4s 2 Co[Ar]3d 7 4s 2 Ni[Ar]3d 8 4s 2 Cu[Ar]3d 10 4s 1 Zn[Ar]3d 10 4s 2

Objective B Notice that Cr and Cu are “exceptions.” Notice that Cr and Cu are “exceptions.” They both have 1 valence electron. They do this because in the case of Cr, moving an electron from the 4s level to the 3d level gives us a half full set of d orbitals. They both have 1 valence electron. They do this because in the case of Cr, moving an electron from the 4s level to the 3d level gives us a half full set of d orbitals. That’s more stable than if Cr would have followed the pattern, and ended with “4s 2 3d 4 ” That’s more stable than if Cr would have followed the pattern, and ended with “4s 2 3d 4 ” Cr[Ar]3d 5 4s 1

Objective B Similarly, Cu has 1 electron in the 4s energy level and 10 in the 3d level, because having a full set of d electrons is also more stable. Similarly, Cu has 1 electron in the 4s energy level and 10 in the 3d level, because having a full set of d electrons is also more stable. Cu[Ar]3d 10 4s 1

Objective B The “inner transition metals” are the lanthanide and actinide series. The “inner transition metals” are the lanthanide and actinide series. That’s where the f electrons are filled up. That’s where the f electrons are filled up. That’s about all I’m going to say about that. That’s about all I’m going to say about that.

Objective C The periodic table allows you to predict trends in certain properties. The periodic table allows you to predict trends in certain properties. Get out a periodic table and put these trends as notes on your periodic table. Get out a periodic table and put these trends as notes on your periodic table. The first trend is Atomic radius. The first trend is Atomic radius. Atomic radius is the size of the atom. It’s defined as ½ the distance between two nuclei which are bonded together. Atomic radius is the size of the atom. It’s defined as ½ the distance between two nuclei which are bonded together.

Objective C Ionic radius is another property Ionic radius is another property It is the size of an ion. Ionic radius is fairly similar to atomic radius. It is the size of an ion. Ionic radius is fairly similar to atomic radius. A positive ion is also called a CATION. A positive ion is also called a CATION. A negative ion is also called an ANION. A negative ion is also called an ANION. A cation is always smaller than the atom it is formed from. A cation is always smaller than the atom it is formed from. An anion is always larger than the atom it is formed from. An anion is always larger than the atom it is formed from.

Objective C Since cations lose electrons to form positive ions and anions gain electrons to form negative ions, it should make sense that they are SMALLER than the atom. Since cations lose electrons to form positive ions and anions gain electrons to form negative ions, it should make sense that they are SMALLER than the atom.

Objective C Ionization energy is the amount of energy required to remove an electron from a gaseous atom. Ionization energy is the amount of energy required to remove an electron from a gaseous atom. The energy required to remove the first electron is called the FIRST IONIZATION ENERGY. The energy required to remove the first electron is called the FIRST IONIZATION ENERGY. The energy required to remove the second electron is the second ionization energy. And so on… The energy required to remove the second electron is the second ionization energy. And so on… Metals always have LOWER ionization energies than nonmetals. Metals always have LOWER ionization energies than nonmetals. That is because metals tend to lose electrons and nonmetals tend to gain them. That is because metals tend to lose electrons and nonmetals tend to gain them.

Objective C It is VERY MUCH easier to remove a valence electron (an electron in the highest energy level) than an “inner core” electron. It is VERY MUCH easier to remove a valence electron (an electron in the highest energy level) than an “inner core” electron. The inner core electrons are ANY electrons which are not VALENCE electrons. The inner core electrons are ANY electrons which are not VALENCE electrons. Na = 1s 2 2s 2 2p 6 3s 1 White = inner core electrons and Blue = Valence electrons

Objective C Electronegativity is measured on a scale from 0.0 to 4.0. Electronegativity is measured on a scale from 0.0 to 4.0. By definition, F is the most electronegative element at 4.0. By definition, F is the most electronegative element at 4.0. Nonmetals have a high electronegativity. Nonmetals have a high electronegativity. Metals have a low electronegativity. Metals have a low electronegativity.

Electronegativity Think of this as the “greediness” of an atom not only holding on to it’s own electrons, but ALSO wanting to “steal” electrons from other atoms. Think of this as the “greediness” of an atom not only holding on to it’s own electrons, but ALSO wanting to “steal” electrons from other atoms.

The Trends Atomic Radius AND Ionic Radius increase as you go down a group. Atomic Radius AND Ionic Radius increase as you go down a group. Atomic Radius AND Ionic Radius decrease as you go from left to right across a period. Atomic Radius AND Ionic Radius decrease as you go from left to right across a period. Electronegativity AND Ionization Energy decrease as you go down a group. Electronegativity AND Ionization Energy decrease as you go down a group. Electronegativity AND Ionization Energy increase as you go from left to right across a period. Electronegativity AND Ionization Energy increase as you go from left to right across a period. Note the trends are opposites. Draw some arrows on your periodic table to help you remember the trends.

The End