ELECTRONIC ENERGY LEVELS

Transcription

1 ELECTRONIC ENERGY LEVELS Electrons are grouped in energy levels around the nucleus. Electrons in the same energy level: Have about the same energy. Are about the same distance from the nucleus. SHELLS The main energy levels are referred to as shells, and are numbered outwards from the nucleus. The energy of each successive shell increases as we move away from the nucleus. Therefore, an electron in the second shell possesses more energy than an electron in the first shell, but less energy than that of an electron in the third energy level. The maximum number of electrons that a shell contains can be represented by the formula 2n 2 where n represents the shell number. QUESTION 11 Which element has uncharged atoms that contain 4 occupied electron shells? A B C D Be C K F Solution The School For Excellence 2016 Summer School Unit 1 Chemistry Book 1 Page 10

2 SUBSHELLS Each shell contains further energy levels known as subshells, of which there are 5 types: s, p, d, f and g. The order of increasing energy levels, and the maximum number of electrons associated with each subshell, is given below. Not all shells contain all 5 subshells. Subshell s p d f g Maximum electrons Increasing Energy The number of subshells in a shell is fixed and is equal to the shell number. The first shell contains 1 subshell: s The second shell contains 2 subshells: s and p The third shell contains 3 subshells: s and p and d The fourth shell contains 4 subshells: s and p and d and f ORBITALS Sub shells are made up of orbitals (regions in which electrons move). Orbitals within a particular subshell have the same energy. The number of orbitals in a sub-shell is fixed, and depends upon the sub-shell type: The s subshell contains 1 orbital. The p subshell contains 3 orbitals. The d subshell contains 5 orbitals. The f subshell contains 7 orbitals. The g subshell contains 9 orbitals. The maximum number of electrons that can be accommodated in any atomic orbital is two (the Pauli Exclusion Principle). Therefore: The s subshell can hold up to 2 electrons. The p subshell can hold up to 6 electrons. The d subshell can hold up to 10 electrons. The f subshell can hold up to 14 electrons. The g subshell can hold up to 18 electrons. Note: Orbitals that are located closer to the nucleus have lower energy levels than those that are located further away. The School For Excellence 2016 Summer School Unit 1 Chemistry Book 1 Page 11

3 QUESTION 12 Complete the following table. Shell Number of Subshells Subshell Types Number of Orbitals Maximum Electrons The School For Excellence 2016 Summer School Unit 1 Chemistry Book 1 Page 12

4 ELECTRON CONFIGURATIONS The way in which electrons are arranged around the nucleus is known as the electron configuration of the atom. Electrons fill orbitals in a particular sequence. Electrons occupy orbitals with the lowest energy levels first, i.e. those energy levels closest to the nucleus (the Aufbau Principle). The order of increasing energy levels for vacant orbitals is obtained by following the arrows in the below diagram (from tail to head). QUESTION 13 Write the electronic configurations for each of the following atoms. (a) An oxygen atom with 8 electrons. (b) A potassium atom with 19 electrons. (c) An iron atom with 26 electrons. The School For Excellence 2016 Summer School Unit 1 Chemistry Book 1 Page 13

5 EXCEPTIONS TO THE USUAL ELECTRON CONFIGURATION PATTERN There are two important exceptions in the usual electron configuration patterns, and they happen for the elements Chromium and Copper. These elements occur in the section of the periodic table known as the transition metals, whose electron configurations allow them special chemical properties such as multiple valence states. This all occurs due to the fact that the 4s and 3d orbitals are extremely close in energy. Usually, the 4s orbital will fill (contain two electrons) before the 3d orbitals begin to fill. This is not the case for Chromium and Copper. The table below shows their predicted electron configurations versus their actual electron configurations. Element Predicted electron configuration Actual electron configuration Chromium, Cr 1s 2 2s 2 6 3s 2 3p 6 3d 4 4s 2 1s 2 2s 2 6 3s 2 3p 6 3d 5 4s 1 Copper, Cu 1s 2 2s 2 6 3s 2 3p 6 3d 9 4s 2 1s 2 2s 2 6 3s 2 3p 6 3d 10 4s 1 It has been proposed, and then mathematically verified, that in the case of Chromium, having one electron in each of the energetically similar orbitals is more favourable than having the 4s orbital filled and one of the 3d orbitals left unfilled. In the same way, it has been shown that Copper is more stable with a completely full 3d subshell and partially filled 4s than it is the other way around. QUESTION 14 Write the electronic configurations for each of the following atoms. (a) A vanadium atom with 23 electrons. (b) A cobalt atom with 27 electrons. (c) A chromium atom with 24 electrons. The School For Excellence 2016 Summer School Unit 1 Chemistry Book 1 Page 14

6 CHEMICAL PROPERTIES Electrons in the outermost shell of an atom are called valence electrons. The valence electrons are those that are involved in chemical reactions and help explain that element s chemical properties. In general, atoms with similar outer electron configurations display similar chemical properties. QUESTION 15 Which two atoms are likely to display similar chemical properties? A B C D 1 p s 2s 3s 3 and 1 p s 2s 2 and 1 p s 2s 2 and s 2s s 2s 1 s s 2s s 2s 3s 3p 3d 4 and s 2s 3s 3p QUESTION 16 Which of the following configurations represents an atom in its ground state? A B C D s 2s 3s 3p 4s s 2s 3s 3p 3d 4s s 2s 3s 3p 3d s 2s 3s 4s QUESTION 17 The electron configuration of the likely ion of chlorine will be: A B C D s 2s 3s s 2s 3s 3p s 2s 3s 3p s 2s 3s 3p 4s The School For Excellence 2016 Summer School Unit 1 Chemistry Book 1 Page 15

7 THE PERIODIC TABLE Elements are made up of only one type of atom (with equal numbers of protons), each being identified by a symbol. Elements are arranged in the periodic table according to: (i) (ii) Increasing atomic number. Electron configuration. A new row is started wherever the outer shell electron configurations repeat so that elements with similar chemical properties can be grouped together. This pattern of similar chemical properties recurring as atomic number increases is known as periodicity. The School For Excellence 2016 Summer School Unit 1 Chemistry Book 1 Page 16

8 GROUPS The vertical columns are referred to as groups, and are numbered 1 to 18. Each element in a group has the same number of electrons present in the outermost shell (valence electrons). Many atoms are able to lose or gain electrons to form charged particles called ions. Whether an atom loses or gains electrons depends upon its position in the Periodic Table. In general, an atom will gain or lose the number of electrons required to reach the electron configuration of the nearest Group 18 element. In the process, the species becomes more stable. Due to similar outer electron configurations, elements (not ions) in a given group generally exhibit similar chemical properties. The horizontal rows are called periods. PERIODS The period number indicates the outermost shell containing electrons, i.e.: the total number of shells in each atom. For example: Na ( 1 2s 3s 2 He ( Fe ( 1 2s 3s 3p 3d 4s s ) is located in period 3. 1s ) is located in period 1. s 2 ) is located in period 4. The elements on the left hand side of the periodic table tend to be metals (species that lose electrons) and those on the right hand side non-metals (species that gain electrons). Elements that have the properties of both metals and non-metals are called metalloids (or semi-metals). The School For Excellence 2016 Summer School Unit 1 Chemistry Book 1 Page 17

9 The elements in Group 18 are called the noble gases. They contain the maximum possible number of electrons in their outer shells, are highly stable and are generally unreactive. Elements in the middle of the periodic table (Groups 3-12 or d block) are called transition elements. These elements involve the filling of the d subshell. The f block consists of the lanthanides and actinides, and involves the filling of the 4f and 5f subshells respectively. QUESTION 18 A ground state element has the electron configuration 1s 2 2s 2 6 3s 2 3p 6 3d 10 4s 2 4p 2. (a) Which period is this element in? (b) Which group is this element in? (c) What is the name of this element? QUESTION 19 Which statement about the modern Periodic Table is CORRECT? A Group 3 contains more elements than any other group. B All elements with 2 valence electrons in their uncharged atoms are in Group 2. C The elements are listed in order of increasing atomic weight. D Elements in the same Group of the Periodic Table have similar sized atoms. Solution The School For Excellence 2016 Summer School Unit 1 Chemistry Book 1 Page 18

10 QUESTION 20 Complete the following table: Element Atomic Number Mass Number Number Protons Number Neutrons Number Electrons Electron Configuration Group Number Period Number Be Ca Ar N Si The School For Excellence 2016 Summer School Unit 1 Chemistry Book 1 Page 19

11 QUESTION 21 State the charge on the ions produced by: (a) Potassium atoms (b) Sulfur atoms (c) Aluminium atoms QUESTION 22 Name the elements with the following electron configurations: (a) s 2s 3s 3p (b) s 2s (c) s 2s 3s 3p 4s QUESTION 23 Why are there six groups of elements in the p-block of the Periodic Table? A B C D The atoms of the elements in the p block have 6 p orbitals filled. The outermost shells of their atoms can only hold 6 electrons. The p sub shell is filling and it holds a maximum of 6 electrons. It was found experimentally that elements in the six groups are very similar. Solution The School For Excellence 2016 Summer School Unit 1 Chemistry Book 1 Page 20

12 QUESTION 24 The element for which the uncharged atoms have an electronic configuration s 2s 3s 3p 3d 4s 4p would be found in the Periodic Table in: A Period 4, Group 16 B Period 6, Group 4 C Period 6, Group 14 D Period 4, Group 6 QUESTION 25 Write the electron configurations of the following ions: (a) S 2 (b) Ca 2 + (c) + Na (d) 2 O The School For Excellence 2016 Summer School Unit 1 Chemistry Book 1 Page 21

13 QUESTION 26 Write the electron configuration for the calcium ion, and explain why it would be smaller than the calcium atom. Solution QUESTION 27 The reason that elements within a group have similar chemical properties is that: A B C D They have a similar configuration of outer-shell electrons. Each element in the group has one more proton than the previous element. The atoms have the same number of electron shells. The atoms have a similar radius. Solution QUESTION 28 Name an element with similar chemical properties to: (a) Magnesium (b) Silicon (c) Oxygen The School For Excellence 2016 Summer School Unit 1 Chemistry Book 1 Page 22

14 TRENDS IN THE PERIODIC TABLE Given the elegant structure of the periodic table and its underlying electronic associations, it is not surprising that there are a number of trends as we move down groups or across periods. CORE CHARGE This is the share of the nuclear charge that valence electrons experience. An understanding of core charge helps explain some of the trends in the periodic table. Consider the following 2 atoms: Magnesium Sulfur Both the Magnesium and sulfur are in the same period. The 2 valence electrons in Magnesium are less attracted to the nucleus than the 6 valence electrons in sulfur are to their nucleus. The electrons of each atom experience a different core charge. Shielding by inner shell electrons (in this case electrons in shell number 2) decreases the attraction between the nucleus and valence electrons. To calculate core charge we subtract the number of inner shell electrons from the atomic number of the atom in question. For example: Magnesium Sulfur = +2 Core Charge = +6 Core Charge So the valence electrons in sulfur experience a greater share of the nuclear charge than those of Magnesium. Core charge increases across a period and remains constant down a group. The School For Excellence 2016 Summer School Unit 1 Chemistry Book 1 Page 23