CHEMISTRY - CLUTCH CH.8 - PERIODIC PROPERTIES OF THE ELEMENTS

Transcription

1 !!

2 CONCEPT: ELECTRON CONFIGURATIONS In this chapter we will focus on how an element s - the distribution of electrons within the orbitals of its atoms relates to its chemical and physical properties. History Lesson: In 1870, Dmitri Mendeleev arranged 65 elements into a. He summarized their behavior in the. When arranged by atomic mass, the elements exhibit a periodic recurrence of similar properties. The Electron Configuration According to the Principle you first have to totally fill in the lowest energy level before moving to the next. 1s 2s 2p 1s 2s$$$$$$2p 3s$$$$$$3p$$$$$$3d 4s$$$$$$4p$$$$$$4d$$$$$4f 5s$$$$$$5p$$$$$$5d$$$$$5f$$$$$5g 6s$$$$$$6p$$$$$$6d$$$$$6f$$$$$6g$$$$6h$ 7s$$$$$$7p$$$$$$7d$$$$$7f$$$$$7g$$$$7h F (9 electrons) 1s 2s 2p Hund s Rule states that electron orbitals that are are first half-filled before they are totally filled. Page 2

3 CONCEPT: CONDENSED ELECTRON CONFIGURATION EXAMPLE: Write the condensed configuration for each of the following elements: a. Co (27 electrons) b. Se (34 electrons) PRACTICE: Write the condensed configuration for each of the following elements: a. Ag (47 electrons) Page 3

4 CONCEPT: INNER CORE & VALENCE ELECTRONS EXAMPLE: How many core (inner) and valence electrons are present in each of the following elements? a. P b. Al c. Mn Page 4

5 CONCEPT: PARAMAGNETISM Vs. DIAMAGNETISM EXAMPLE: Write the condensed electron configuration of each ion and state if the ion is paramagnetic or diamagnetic. a. Ni 3+ b. S 2- PRACTICE: Write the condensed electron configuration of each ion and state if the ion is paramagnetic or diamagnetic. a. Cu + Page 5

6 CONCEPT: EFFECTIVE NUCLEAR CHARGE & SLATER S RULES When looking at any particular electron within an atom it experiences two major forces. A(n) force from the nucleus and a(n) force from the surrounding electrons. Now the electron can become shielded from the full force of the nucleus because of the other surrounding electrons. Effective Nuclear Charge (Zeff) measures the force exerted onto an electron by the nucleus, and can be calculated using Rules. e - e - e - e - e - Z = Nuclear Charge e - e - Z eff = Z S S = Shielding Constant e - e - e - e - Guidelines for Determining S for an electron: 1. The atom s electronic configuration is grouped as follows, in terms of increasing n and l quantum numbers: (1s) (2s,2p) (3s,3p) (3d) (4s,4p) (4d) (4f) (5s,5p) (5d) etc. 2. Electrons in groups to the right of a given electron do not shield electrons to the left. 3. The shielding constant S for electrons in certain groups. For ns and np valence electrons: a) Each electron in the same group will contribute to the S value, with exception to 1s electrons, which contribute to the S value. b) Each electron in n 1 group contributes to the S value. c) Each electron in n 2 group or greater contributes to the S value. For nd and nf valence electrons: a) Each electron in the same group will contribute to the S value. b) Each electron in groups to the left will contribute to the S value. EXAMPLE: Using Slater s Rules calculate the effective nuclear charge of a 3p electron in argon. Page 6

7 PRACTICE: EFFECTIVE NUCLEAR CHARGE & SLATER S RULES 1 EXAMPLE 1: Using Slater s Rules calculate the effective nuclear charge of the 4s electron in potassium. EXAMPLE 2: Using Slater s Rules calculate the effective nuclear charge of a 3d electron in bromine. Page 7

8 CONCEPT: THE FOURTH QUANTUM NUMBER An electron in an atom is described completely by a set of four quantum numbers. The first three describe its and the fourth describes its. The quantum number (ms) helps to discuss the rotational spin of the electron and has values of either and.!! According to the : no two electrons in the same atom can have the same four quantum numbers. EXAMPLE: State the electron configuration of boron and list the four quantum numbers of the 1 st and the 5 th electron. Page 8

9 CONCEPT: ATOMIC ORBITAL SHAPE The quantum number deals with the shape of the atomic orbital. Each atomic orbital has a specific shape. It uses the variable and formula. Each atomic sub-level has an L value associated with it. Sublevel s p d f g L value EXAMPLE: Based on the following atomic orbital shape, which of the following set of quantum numbers is correct: a) n = 8, l = 1, ml = 1 2 b) n = 8, l = 2, ml = -2 c) n = 8, l = 0, ml = 1 d) n = 8, l = 0, ml = 0 PRACTICE: Based on the following atomic orbital shape, which of the following set of quantum numbers is correct: a) n = 2, l = 1, ml = +1, ms = - 1 b) n = 4, l = 1, ml = - 2, ms = c) n = 3, l = 1, ml = - 1, ms = 0 d) n = 2, l = 1, ml = + 1, ms = 1 2 Page 9

10 CONCEPT: TRENDS IN ATOMIC RADIUS Atomic radius is defined as half the distance between the nuclei in a molecule of two identical elements. Generally, it going from left to right across a period and going down a group. ATOMIC RADIUS EXAMPLE: If the sum of the atomic radii of diatomic carbon is 154 pm and of diatomic chlorine is 198 pm, what is the sum of the atomic radii between a carbon and a chlorine atom. PRACTICE: Which one of the following atoms has the largest atomic radius? A) K B) Rb C) Y D) Ca E) Sr Page 10

11 CONCEPT: TRENDS IN IONIC RADIUS Ionic Size estimates the size of an ion in an ionic compound. (POSITIVE IONS) tend to be smaller than their parent atoms. Lithium ( 3 Electrons) 1s 2s 1s 2s (NEGATIVE IONS) tend to be larger than their parent atoms. Fluorine ( 9 Electrons) 1s 2s 2p 1s 2s 2p The pattern for ionic size correlates with the following trend when comparing ions with the same number of electrons: -3 > -2 > -1 > 0 > +1 > +2 > +3 EXAMPLE: Rank each set of ions in order of increasing ionic size. a) K +, Ca 2+, Ar b) Sr 2+, Na +, I c) V 5+, S 2-, Cl Page 11

12 CONCEPT: TRENDS IN IONIZATION ENERGY Metals tend to lose electrons to become positive ions called. IONIZATION ENERGY Therefore they have ionization energies. Nonmetals tend to gain electrons to become negative ions called. Therefore they have ionization energies. Ionization energy (IE) is the energy (in kj) required to remove an electron from a gaseous atom or ion. Generally, it going from left to right of a period and going down a group. Exceptions: Atom (g) ion + (g) + e E = IE1 > 0 When in the same period, Group elements have lower ionization energy than elements in Group. O 1s 2s 2p 1s 2s 2p N 1s 2s 2p 1s 2s 2p When in the same period, Group elements have lower ionization energy than elements in Group. B 1s 2s 2p 1s 2s 2p Be 1s 2s 1s 2s Page 12

13 PRACTICE: TRENDS IN IONIZATION ENERGY EXAMPLE: Of the following atoms, which has the smallest second ionization energy? a. Al b. Li c. Rb d. Mg e. Be PRACTICE 1: Of the following atoms, which has the smallest third ionization energy? a. Al b. Ca c. K d. Ga e. Cs PRACTICE 2: Which of the following statements is/are true? a. Sulfur has a larger IE1 than phosphorus b. Boron has a lower IE1 than Magnesium c. Magnesium has a higher IE1 than Aluminum PRACTICE 3: Shown below are the numerical values for ionization energies (IE s). Match the numerical values with each of the following elements provided in the boxes. Na Mg Al Si P S Cl Ar Numbers: 496, 578, 738, 786, 1000, 1012, 1251 & Page 13

14 CONCEPT: TRENDS IN ELECTRON AFFINITY Electron Affinity (EA) is the energy change (in kj) from the addition of 1 mole of e to 1 mol of gaseous atoms or ions. Generally, it going from left to right across a period and going down a group. Atom (g) + e ion (g) E = - EA1 ELECTRON AFFINITY EXAMPLE: Rank the following elements in order of increasing electron affinity. a. Cs, Hg, F, S b. Se, S, Si PRACTICE: Shown below are the numerical values for electron affinities (EA s). Match the numerical values with each of the following elements provided in the boxes. Li Be B C N O F Ne Numbers: - 328, -141, -122, -60, -27, > 0, > 0, > 0. Page 14