Section 2 The Electronic Structure of Atoms and the Periodic Table

Modern College F.6 Chemistry (2009 – 10)

Section 2

Name: ______________________________ Class: _______________ Class No.: ____________ Prepared by Mr. Chau Chi Keung, Richard

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Modern College F.6 Chemistry (2009 – 10)

Section 2

Prepared by Mr. Chau Chi Keung, Richard

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Modern College F.6 Chemistry (2009 – 10)

Section 2

2.1 Atomic Emission Spectrum ( 原 子 發 射 光 譜 ) and Electronic Structure of Atoms (Bohr Model) 2.1.1. Atomic emission spectrum of hydrogen and its origin 

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When a sample of hydrogen gas receives a high energy spark, the H2 molecules absorb energy, and some of the H–H bonds are broken (H—H 鍵接收高能量後分離). Originally, the hydrogen atoms are in the ground state (基 態 ), which the electron inside stays at its lowest possible energy level (能階). When receiving high energy, the hydrogen atoms are excited and contain excess energy. Their electrons are excited from the ground state to excited state (i.e. some higher energy levels) (原子內的電子由基態激發至較高能階). The excited atoms must go back to ground state in the end (∵ electrons are unstable in the excited state) Electron will fall back to its original energy level (電子最終失去能量返回基態)  Energy will be given out as electromagnetic radiation (e.g. UV, visible light, infra red) (釋出能量，以電磁波表達)

An atomic emission spectrum can be obtained by passing the light emitted from the discharge tube through a prism. Light consists of a wide range of radiation of different frequencies is spilt by the prism. Coloured lines are seen on the dark background. The hydrogen emission spectrum is seen to be a number of series of lines. These groups of lines were named after scientists who discovered them.

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Modern College F.6 Chemistry (2009 – 10)

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Note: Electromagnetic waves: c (speed of light) = fλ ♦ Ultraviolet

Visible light

higher frequency(energy), shorter λ

Infra-red

lower frequency, longer λ

2.1.2. Bohr’s model of hydrogen atom 

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In 1913, Neils Bohr, a Danish physicist, suggested the model of hydrogen atom with reference to quantum theory.  Hydrogen atom consists of a nucleus with a single positive charge  Electron is moving around the nucleus in orbits with discrete amount of energy The electrons can only exist in certain fixed energy levels. It cannot possess energy of intermediate magnitudes. When hydrogen atoms absorb or emit energy, electrons move from one energy level to another. Only transitions from one level to another are possible. Transitions of intermediate energy are not allowed. The energy of an electron can only be gained or lost in small energy packets called quanta (量子). The amount of energy involved is exactly equal to a photon (光 子 ). (Based on the particle property of light, light consists of discrete particles. Each discrete particle is an energy packet called photon.) According to the quantum theory, the energy contained in a photon of light of frequency v

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Modern College F.6 Chemistry (2009 – 10)

Section 2

can be calculated by the Planck’s equation E = hν =

hc λ

where h is the Planck constant (普 朗 克 常 數 ) (h = 6.63 × 10-34 Js) and c is the speed of light (c = 3 × 108 ms-1) For an electron to move from an orbit of energy E1 to another orbit with energy E2, the light  absorbed must have a frequency give by Planck’s Equation (i.e. v is fixed): When electrons which have been excited (raised to orbits of higher energy) drop back to  orbits of lower energy, they emit energy as light with a frequency given by Planck’s equation (again, v is fixed) n=∞ E∞ n=4 E4

Bohr’s Atomic Model

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E1 Energy levels E2 (excited state)

Electron absorbs energy, jumps to a higher energy level

Electron emits energy as photon: ΔE = hv = E2 – E1

E1 (ground state) 

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From the above concepts, we should bear in mind that: Electrons can only exist in fixed energy levels with exact amount of energy given by  E1, E2…E∞ etc. Other values are impossible – Energy is quantized. If an electron gains energy and promote to an excited state, the electron should gain an  exact amount of energy (= energy difference between two energy levels). Once it falls back to ground state, it also emits an exact amount of energy (as photon).  For this reason, only electromagnetic radiation with particular wavelengths can be  detected (∴Line spectrum instead of continuous spectrum). An analogy of the quantized theory of emission by the excited electron vs. the old continuous theory.

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(a) In the quantized theory Mr. Electron can go up or jump down some definite steps of higher energy than the ground state. The amount of energy absorbed or emitted is discrete and quantized. Note that the steps are not equal in size, meaning that the energy levels are unequal and closer at higher energies. (b) The old theory (wrong) Mr. Electron may go up or down the slope of higher energy in any infinitesimally small step.

2.1.2. Characteristic features of the hydrogen emission spectrum 

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The atomic emission spectrum of hydrogen consists of several series of discrete lines which converge in different parts of the electromagnetic spectrum: Lyman series (賴曼系) in the ultraviolet region,  Balmer series (巴爾麥系) in the visible region,  Paschen series in the infrared region, and  Brackett and Pfund series in the far infrared region.  Within each series, the lines get closer together and converge towards high frequency end of the spectrum (在每一線系，各譜線頻率之間的差距在高頻率處越來越少). Finally, the lines merge into a continuum of light (light of the highest frequency). This is the convergence limit (譜 線 最 終 結 集 成 一 連 續 區 ， 稱 為 會 聚 極 限 ). Take the Balmer series as an example:

2.1.3. Interpretation of the hydrogen emission spectrum A. Quantum theory and Bohr’s atomic model Electrons around the nucleus in an atom are moving in discrete orbits/shells.  Electrons can only exist in fixed energy levels (energy is quantized).  When hydrogen atoms absorb or emit energy, electrons move from one energy level to  another. Only transitions from one level to another are possible. Prepared by Mr. Chau Chi Keung, Richard

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Modern College F.6 Chemistry (2009 – 10)  

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The energy of an electron can only be lost or gained in small packets called quanta. The amount of energy emitted can be calculated by the Planck’s equation. ∆E = hν =

hc λ

Where ΔE = quantum of energy = energy difference between two energy levels ν = frequency of light (Hz or s-1) h = 6.63 x 10-34 Js (Planck’s constant) c = 3.00 x 108 m s-1 (velocity of light) λ = wavelength of light (usually in nm) 

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When an electron falls from a higher energy level to a lower energy level, it emits a quantum of energy (ΔE) in radiation of definite frequency. Since ΔE for the change is always the same in a given atom, ν must be a constant. Therefore, radiation always has the same energy and is always of the same frequency for that particular electron transition. Hence, the atomic emission spectrum consists of discrete lines.

B. The Balmer series The following figure shows the energy levels in a hydrogen atom and the electron transitions  which produce the lines in the visible region of the atomic emission spectrum of hydrogen (i.e. Balmer series):

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It should be noted that: The electronic energy levels are numbered (n = 1, n = 2, n = 3, etc.). These numbers are  sometimes called the principal quantum numbers (主 量 子 數 ), which can represent the electron shells of an atom. The level of lowest energy is given the principal

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Modern College F.6 Chemistry (2009 – 10)

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quantum number 1, the next lowest 2, and so on. The coloured lines in the Balmer series are caused by electron transitions from the higher level to the level n = 2.

For example, transitions from n = 3 to n = 2 result in the first line of the series which is a red line (at frequency 4.571 × 1014 Hz) in the hydrogen spectrum, while transitions from n = 4 to n = 2 produce the second line (a blue line at frequency 6.172 × 1014 Hz). (Question: Calculate the energy difference between energy levels of n = 3 and n = 2.)

At higher energies, the energy levels get closer and eventually come together, it follows that the spectral lines also get closer and eventually come together. This particular frequency is called the convergence limit.

C. The Lyman series Besides Balmer series, the complete hydrogen emission spectrum consists of many series.  The origin of the other series is similar to that of Balmer series, the only difference is the destination of the electrons. When transition occur from the higher levels to the lowest energy level (n = 1), more energy  is released than with transitions to the n = 2 level. Consequently, these lines appear at higher frequencies in the spectrum. In the case of hydrogen atom, transitions of electron from any given energy level (n ≥ 2) to  the n = 1 level result in lines in the ultraviolet region of the emission spectrum. The series of spectral lines is known as the Lyman series. 

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D. Convergence limit and ionization energy (電離焓 ) If sufficient energy is given to an atom, it is possible to excite an electron just beyond the  highest energy level. In this case the electron will escape and the atom becomes a cation. Ionization has taken place. In other words, we can say that the electron is promoted to an energy level n = ∞.  To be more specific, ionization refers to a process that the electron is excited from ground  state (lowest possible energy level) to an energy level n = ∞. The energy involved in this process is called ionization energy or ionization enthalpy. By determining the frequency at which the converging spectral lines come together (i.e.  convergence limit), the ionization energy of an atom can be found. For hydrogen atom, the ionization energy can be determined by measuring the frequency  the convergence limit (υ∞→1) of Lyman series (如 果 我 們 知 道 賴 曼 系 的 會 聚 極 限 ， 就 可以計算氫原子的電離焓). (Question: An accurate value for the frequency at the convergence limit of Lyman series for hydrogen is 3.27 x 1015 Hz. Calculate the ionization energy of hydrogen.)

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Hint: When working on calculations about ionization energy, you should be very careful about whether you need to calculate the I.E of “one hydrogen atom” or “one mole of hydrogen atom”. (For the latter case, you have to multiply the value by the Avogadro constant)

E. Uniqueness of atomic emission spectra of elements Prepared by Mr. Chau Chi Keung, Richard

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Modern College F.6 Chemistry (2009 – 10)  

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The line spectrum of hydrogen is characteristic of the element. Higher elements also produce similar, though more complicated spectra of their own. When being excited or strongly heated, each element can produce a unique emission line spectrum. (e.g. the brightest emission light of Na is yellow in colour.) This is because atoms of each element have a unique arrangement of electrons with definite energy levels. For this reason, each element produces a completely individual pattern of emission spectral lines which can be used to identify it (a kind of atomic fingerprint). As a result, these line spectra can be used to identify elements (e.g. flame test).

2.1.4. Using ionization enthalpies to predict electronic structures A. Ionization enthalpy of other elements One important property that determines an element’s chemical behaviour is the ease with  which its outer electrons can be removed.  Ionization enthalpies of various elements provide us useful information to investigate that. Ionization enthalpy is defined as the energy absorbed when 1 mole of electrons is  removed from 1 mole of atoms or ion to infinity at gaseous phase (一 摩 爾 氣 態 的 中 性 原子的最外層移走一摩爾電子所需的最低能量). 

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Energy is always absorbed in this process because work has to be done in overcoming the attractive force of the positively charged nucleus for the negatively charged electrons. (電離 焓必定是正值，是吸熱反應。因為要克服原子核和電子之間的吸引力) The first ionization energy is the minimum energy required to remove the most loosely held electron. That is, the electron in the outermost shell (∵smallest attraction from the nucleus).  More specifically, the 1st I.E. is the energy required to remove an electron from a  free atom in gaseous state. ∆ H = 1st I.E. X(g) → X+(g) + e– The energy needed to remove the next electron is the second ionization enthalpy, and so on.

∆ H = 2nd I.E. X+(g) →X2+(g) + e– 2+ 3+ – ∆ H = 3rd I.E. X (g) →X (g) + e ∆ H = 4th I.E. X3+(g) →X4+(g) + e– ………………………………………... and so on. 

When an electron is removed, the atom has a net positive charge so that the next electron is more difficult to remove because of large attraction between positive ion and negatively charged electron. i.e. first ionization energy < second ionization energy and in general: 1st I.E. < 2nd I.E. < 3rd I.E.

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Modern College F.6 Chemistry (2009 – 10)  

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When outer electrons are removed, the electrons to be removed afterwards come from lower energy levels (closer to the nucleus). I.E. is normally measured in electron volts (電子伏特) (eV) or kJmol-1. (1eV = charge of electron × 1V = 1.602 × 10-19 C × 1V = 1.602 × 10-19 J)

B. Successive ionization enthalpies of an element (Evidence of electron shells) The variation in the successive ionization enthalpies is a good evidence for the existence of  electron shells (or energy levels). 

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Take potassium as an example, the ionization of potassium can take place in steps 1st ionization K(g) → K+(g) + e– ΔHI1 = 419 kJmol–1 2nd ionization K+(g) → K2+(g) + e– ΔHI2 = 3051 kJmol–1 3rd ionization K2+(g) → K3+(g) + e– ΔHI3 = 4412 kJmol–1 Consider the plot of log ionization enthalpy against the number successive ionization enthalpy of potassium atom:

There is a general increase in successive ionization with the no. of electrons removed. It is due to the nucleus becoming more positively-charged as each electron is being lost. The electrostatic attraction between the ion and electron increases as the charge of  cation increases. So more energy is needed to remove the electron. The 1st I.E. of sodium is much less than the 2nd I.E. This is due to the 1st electron is removed

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Modern College F.6 Chemistry (2009 – 10)

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from the outermost shell which experience the least attraction form the nucleus. The second electron is removed from the second electron shell which is closer to the  nucleus with stronger electrostatic attraction between them. So the energy needed is larger. Similar reason explains the difference between the 10th I.E. and 11th I.E. or between 17th  I.E. and 18th I.E. It can be observed from the graph that the I.E. of K atom can be divided into 4 groups. 1 in the first group, 8 in the second group, 8 in the third group and remaining 2 in the fourth. This indicates that a K atom has four electron shells and electronic configuration of 2,8,8,1.

C. Evidence of sub-shells (電子亞層 ) The figure below shows the 1st I.E. of the first 20 elements plotted against atomic number. 

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The 1st I.E. increases roughly across a period due to the increase in nuclear charge. Besides the main trends, some irregularities are found within a period. For instance, in period 2 and 3, the increasing trends can further be broken up into 3 parts, in a 2-3-3 manner. This suggests the presence of sub-shells (s, p, d and f sub-shells) within an electron shell (n = 1, 2, 3 ...) Those sub-shells may have small difference in energy which makes the 1st I.E. of some elements in the same shell different. Conclusion: The pattern of 1st I.E. of various elements indicates that there are further subdivisions of energy within some shells called sub-shells.

When we consider the 1st I.E. down a group, the trend will be as follows:

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Going down a group, the 1st I.E. decreases as the number of electron shells or the size of the atom increases. The attraction between the outermost electron and the nucleus decreases with  increasing atomic size. A prevailing explanation for these variations in ionization energy is that there are subdivisions of energy within some sub-shells, which are called orbitals (軌 態 ). Different electrons occupy different orbitals in an atom. Take the third energy level (n = 3) as an example: 

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It is suggested that different orbitals have different energies, the electron in different orbital shows difference in energy. A more energetic electron would be removed more readily and has lower ionization energy. The details about the atomic orbitals will be discussed in the following part.

2.2 Atomic Orbitals 2.2.1. Limitations of Bohr’s atomic model 

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In Bohr’s atomic model electron in the H atom (a one-electron system) moves around the nucleus in circular orbits. Basing on classical mechanics, Bohr calculated values of frequencies of light emitted for electron transitions between such ‘orbits’. The calculated values for the frequencies of light matched with the data in the atomic emission spectrum of hydrogen Although Bohr’s atomic model successfully explained the atomic emission spectrum of hydrogen and its main concepts such as “ground state”, “excited state” and “discrete energy levels in an atom” are still used more than 110 years after the establishment of the model, it

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is found that Bohr’s model has some major limitations: It only explain the behaviour of electron in hydrogen atom or other one-electron species  which can only be made in laboratory such as He+, Li2+, Be3+, B4+ and C5+. It fails to predict the emission spectrum of any other elements, even helium which is the  second simplest element (∵The calculated values for the frequencies of light based on Bohr’s model did not match with the data in the emission spectra of the elements.) The Bohr model speculated that the electron moves around the nucleus in certain fixed  circular orbits. However, the reality is that the movement of an electron is far less clearly defined.

2.2.2. The wave nature of electron – Building up a new model 

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In 1924, a French Physicist Louis de Broglie postulated that electrons could have both particle and wave properties. This postulate was later confirmed by the experiments performed by C. Davisson with his partner L. Germer, and by G. P. Thomson. When they passed a beam of electrons through a crystal and through a gold foil, they obtained a diffraction pattern on a screen similar to that observed when X-ray was used. This piece of evidence demonstrated that electrons have wave properties.

Following this, studies in quantum mechanics have shown that electrons are not localized in fixed orbits, and as a matter of fact, we can find only describe the location of an electron in terms of probability of finding it in a certain position at any time. The probability, considered over a period of time, gives an average picture of how the electron behaves. The picture does not show the actual location of the electron at any given time. Instead, it shows where the electron is most likely to be at any time. To sum up, the electrons are moving in all directions within the atom, forming an electron cloud (電子雲). 電 子在原子核外很小的空間內作高速運動，其運動規律跟一般物體不同，它 沒有明確的軌道。根據量子力學中的不確定性原理 (Uncertainty principle)，我 們 不 可 能同 時準 確地 測定 出電 子在 某一 時刻 所處 的位 置和 動量 (運動 速度 )， 也 不能描畫出它的運動軌跡。因此，人們常用一種能夠表示電子在一定時間 內 在核外空間各處出現機會的模型來描述電子在核外的運動。在這個模型裡 ，

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Modern College F.6 Chemistry (2009 – 10)

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某 個點附近的密度表示電子在該處出現的機會的大小。密度大的地方，表明 電 子在核外空間單位體積內出現的機會多；反之，則表明電子出現的機會少 。 由於這個模型很像在原子核外有一層疏密不等的「雲」，所以，人們形象地稱 之為「電子雲」。 

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The following figure shows the probability distribution of the electron in a hydrogen atom:

The denser the dots, the more likely the electron will be in that region. However, even at points far away from the nucleus there is some probability, although it is small. More specifically the volume of the space in which there is a high probability (≧ 90%) of finding the electron is called the atomic orbital. An atomic orbital can be viewed as a representation of a region within which there is a very high probability of finding an electron. It can also be regarded as the space in an atom occupied by an electron, a subdivision of the available space within an atom for an electron to orbit the nucleus. This new atomic model is called wave-mechanical model (波動模型) and it was developed by Erwin Schrödinger, an Austrian physicist in 1926. In this model, electrons are considered as both particles and waves:  Not localized in fixed orbits  Their arrangements and motions are described by mathematical equations We describe an electron in terms of the probability of finding it 

2.2.3. Quantum numbers – The I.D number of electron 

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In quantum mechanics, the electron is treated as wave. A wave function is used to describe its motion. Solutions of the wave equation for electrons show electronic energy levels are quantized and described by quantum numbers. The four quantum numbers are: Quantum numbers Function Allowed values Principle quantum number Specifies the principal Any positive integer energy level and the (n) (主量子數) distance of electrons from the nucleus. Subsidiary quantum number Specifies various subshell Depends on n, in the range of and the shape of the orbital 0, 1,2 … (n-1) (l) (副量子數)*1

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Modern College F.6 Chemistry (2009 – 10) Magnetic quantum number (m) (磁量子數) Spin quantum number (s) (自旋量子數)

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Specifies the *2direction Depends on l, in the range of (spatial orientations) of the -l, -l + 1...l – 1 and l orbitals (e.g. px, py and pz for p orbitals) 1 Specifies the spin of an + electron (either clockwise *3Either 2 (which is or anti-clockwise) 1 indicated by ↑) or (which is indicated by ↑)

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*Notes: 1. If a magnetic field is supplied, the subshell is shown to consist of “sub-subshells”, each having a quantized spatial orientation in a magnetic field. 2. Electrons with different subsidiary quantum numbers can be denoted by the letters s, p, d and f, or s orbital, p orbital, d orbital and f orbital. Value of l 0 1 2 3 Orbital s p d f 3. Each orbital only maximum holds 2 electrons. Both of them have opposite spin (will be discussed later).

2.2.4. More about orbitals 

The term electron shell is used for a group of orbitals with the same principal quantum number. A subshell is a group of orbitals with the same energy level (called degenerate orbitals (同階軌態)).

Electron shell

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subshell

orbitals

The number of orbitals in each subshell: Subshell Number of orbitals s 1 p 3 d 5 f 7

individual electron

Maximum number of electrons in the subshell 2 2×3 = 6 2 × 5 = 10 2 × 7 = 14

The relative energies of different orbitals increase as the principal quantum number increases. For a one-electron system like hydrogen atom, the division will be as follows:

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Modern College F.6 Chemistry (2009 – 10)

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The situation for many-electron systems will be more complex. The energies of subshells increase in the order s < p < d < f. The relative energy level diagram of orbitals is shown in the following figure:

The details and division of electrons in an atom can be summarized as follows: Shell Subshells Orbitals Maximum number of electrons per shell (2n2) n = 1 1s 1s 2 n = 2 2s, 2p 2s, 2px, 2py and 2pz 8 n = 3 3s,3p and 3d 3s, 3px, 3py, 3pz and five d orbitals 18 n = 4 4s, 4p, 4d and 4f 4s, 4px, 4py, 4pz, five 4d orbitals 32 and seven 4f orbitals Different orbitals have different shapes. (Reference website: http://winter.group.shef.ac.uk/orbitron/ ). 

For example, s orbitals are spherical in shape (not circular)(形 狀 是 球 形 ， 不 是 圓

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形) Therefore, the charge cloud, being non-directional is not concentrated in any direction. Atomic nucleus is located at the centre of the ‘sphere’.

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Modern College F.6 Chemistry (2009 – 10)

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The p orbitals are all dumb-bell in shape (啞 鈴形 ). They are directional and lie along any of the three coordinate axes (The figure below shows the three 2p orbitals).

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The shape and orientation of d orbitals will be even more complex.

Having the knowledge on atomic orbitals, we can go further to build up a new concept on electronic configuration of atoms.

2.3 Electronic Configurations of Elements 

Building up of electronic configurations based on three principles:

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Modern College F.6 Chemistry (2009 – 10) 

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The Aufbau (Building-up) principle (構 築 原 理 ): Electrons will enter the possible orbitals in the order of ascending energy. (電子首先填入能量較低的軌態，然後才填入能量較高的軌態。) The order of energy levels needs not to be exactly parallel to the principal ♦ quantum numbers. When the atom becomes heavy, the high energy levels tend to overlap with each other. For example, the energy of 4s orbital is lower than that of 3d orbitals. Hence ♦ electrons would enter 4s first instead of 3d. However, after filling electrons in 3d orbitals, the 4s orbital becomes higher ♦ in energy than 3d orbitals. This occurs for d-block elements (Sc to Zn in the Periodic Table)

Pauli’s exclusion principle ( 泡 利 不 相 容 原 理 ): Electrons occupying the same orbital must have opposite spins and no orbital can accommodate more than two electrons. (每個原子軌態最多可容納兩個自旋方向相反的電子。) Hund’s rule ( 洪 特 規 則 ): Electrons must occupy each energy level singly before pairing takes place because of their mutual repulsion. Moreover, all of the unpaired electrons would have parallel spins in order to minimize mutual repulsion. (當 電 子 分 佈 於 相 同 能 量 的 軌 態 時 ， 由 於 電 子 相 互 間 的 排 斥 作 用 ， 所 以 它 們 首先會單獨地分佈於各個軌態中，然後才開始配對。)

The electronic configuration can be expressed in two ways: Notations using 1s, 2s, 2p…etc. (refer to the table on the next page) Electron-in-box diagram ( 格 中 電 子 圖 ), which tells how electrons occupy different  orbitals with spins shown. 

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: represents one electron which occupies an orbital singly. : represents two electrons having opposite spins. Atomic

Element

Symbol Arrangement of

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Electronic configuration Page 19

Modern College F.6 Chemistry (2009 – 10) number 1 Hydrogen 2 Helium 3 Lithium 4 Beryllium 5 Boron 6 Carbon 7 Nitrogen 8 Oxygen 9 Fluorine 10 Neon 11 Sodium 12 Magnesium 13 Aluminium 14 Silicon 15 Phosphorus 16 Sulphur 17 Chlorine 18 Argon 19 Potassium 20 Calcium 21 Scandium 22 Titanium 23 Vanadium 24 Chromium 25 Manganese 26 Iron 27 Cobalt 28 Nickel 29 Copper 30 Zinc 31 Gallium 32 Germanium 33 Arsenic 34 Selenium 35 Bromine 36 Krypton

H He Li Be B C N O F Ne Na Mg Al Si P S Cl Ar K Ca Sc Ti V Cr Mn Fe Co Ni Cu Zn Ga Ge As Se Br Kr

electrons in shells 1 2 2, 1 2, 2 2, 3 2, 4 2, 5 2, 6 2, 7 2, 8 2, 8, 1 2, 8, 2 2, 8, 3 2, 8, 4 2, 8, 5 2, 8, 6 2, 8, 7 2, 8, 8 2, 8, 8, 1 2, 8, 8, 2 2, 8, 9, 2 2, 8, 10, 2 2, 8, 11, 2 2, 8, 13, 1 2, 8, 13, 2 2, 8, 14, 2 2, 8, 15, 2 2, 8, 16, 2 2, 8, 18, 1 2, 8, 18, 2 2, 8, 18, 3 2, 8, 18, 4 2, 8, 18, 5 2, 8, 18, 6 2, 8, 18, 7 2, 8, 18, 8

Section 2 Standard form Abbreviated form 1s1 1s1 1s2 1s2 1s22s1 [He]2s1 1s22s2 [He]2s2 1s22s22p1 [He]2s22p1 1s22s22p2 [He]2s22p2 1s22s22p3 [He]2s22p3 1s22s22p4 [He]2s22p4 1s22s22p5 [He]2s22p5 1s22s22p6 [He]2s22p6 1s22s22p63s1 [Ne]3s1 1s22s22p63s2 [Ne]3s2 1s22s22p63s23p1 [Ne]3s23p1 1s22s22p63s23p2 [Ne]3s23p2 1s22s22p63s23p3 [Ne]3s23p3 1s22s22p63s23p4 [Ne]3s23p4 1s22s22p63s23p5 [Ne]3s23p5 1s22s22p63s23p6 [Ne]3s23p6 1s22s22p63s23p64s1 [Ar]4s1 1s22s22p63s23p64s2 [Ar]4s2 1s22s22p63s23p63d14s2 [Ar] 3d14s2 1s22s22p63s23p63d24s2 [Ar] 3d24s2 1s22s22p63s23p63d34s2 [Ar] 3d34s2 1s22s22p63s23p63d54s1 [Ar] 3d54s1 1s22s22p63s23p63d54s2 [Ar] 3d54s2 1s22s22p63s23p63d64s2 [Ar] 3d64s2 1s22s22p63s23p63d74s2 [Ar] 3d74s2 1s22s22p63s23p63d84s2 [Ar] 3d84s2 1s22s22p63s23p63d104s1 [Ar] 3d104s1 1s22s22p63s23p63d104s2 [Ar] 3d104s2 1s22s22p63s23p63d104s24p1 [Ar] 3d104s24p1 1s22s22p63s23p63d104s24p2 [Ar] 3d104s24p2 1s22s22p63s23p63d104s24p3 [Ar] 3d104s24p3 1s22s22p63s23p63d104s24p4 [Ar] 3d104s24p4 1s22s22p63s23p63d104s24p5 [Ar] 3d104s24p5 1s22s22p63s23p63d104s24p6 [Ar] 3d104s24p6

Electronic structures of elements with atomic number 1-36 

Remember: When writing electronic configurations, we have to follow the principal

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Section 2

quantum number, not the order of placing electrons. It can be observed that Cr (Z = 24) and Cu (Z = 29) do not follow the general trend. Instead of [Ar]3d44s2, Cr has the electronic configuration [Ar]3d54s1  Instead of [Ar]3d94s2, Cu has the electronic configuration [Ar]3d104s1  This is because extra stability can be gained by have some special configurations – halffilled (e.g. np3, nd5) or completely/full-filled subshells (e.g. ns2, np6, nd10) (半滿或者全滿 的電子亞層能夠賦予原子額外的穩定性。). Half-filled subshells attain extra stability because electrons in those subshells are  evenly distributed in different orbitals and have parallel spin. As a result, the electronic repulsion can be reduced. Completely filled subshells attain even higher stability because extra energy is  required to break the arrangement of paired electrons with different spins (spin paired arrangement). One common example is ns2np6, which is known as the stable noble gas electronic configuration. The concept of “extra stability” can also be used to explain the irregularities in the variation of 1st I.E. across a period.’ For example, the 1st I.E. of oxygen is lower than that of nitrogen because nitrogen  atom has a half-filled 2p subshell which gives it extra stability (∴more difficult to remove electron). As mentioned before, the electronic configuration can be expressed by using electron-in-box diagram. Some examples are as follows: Carbon (Z = 6): 1s22s22p2 (2, 4) 1.

1s

2s

2p

2.

Potassium(Z = 19): 1s22s22p63s23p64s1 (2, 8, 8, 1)

3.

Chromium (Z = 24): 1s22s22p63s23p63d54s1 (2, 8, 13, 1)

4.

Copper (Z = 29): 1s22s22p63s23p63d104s1 (2, 8, 18, 1)

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Section 2

From the electronic configuration of an element, the electronic configuration of its ions can be found and the method of expression of their electronic configuration is the same. Remember: For transition metals, after filling electrons in 3d orbitals, the 4s orbital becomes higher in energy than 3d orbitals. Originally, energy of 4s < energy of 3d (∴ electrons enter 4s first)  Once 3d orbitals are filled with electrons, those 3d electrons will repel 4s electrons  further away from the nucleus. Hence, the 4s orbital will be pushed to a higher energy level (> 3d)  As a result, in case of ionization, we start from 4s instead of 3d.  Examples: Give the “electrons-in-boxes” diagram of Cu+ and Cu2+ 1.

2.

Give the “electrons-in-boxes” diagram of Fe, Fe2+ and Fe3+. Hence, briefly explain whether Fe2+ and Fe3+ would be more stable.

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Modern College F.6 Chemistry (2009 – 10)

Section 2

2.4 The Periodic Table

s-block

p-block

d-block

f-block

The Modern Periodic Table of Elements In the Periodic Table, the elements are arranged in order of increasing atomic number.  A new row (period) is started when electrons start to enter a new principal energy level. Elements whose atoms have a similar outer electronic configuration are placed in a vertical  column (group).  Metals: elements on the left of the Periodic Table  Non-metals: elements on the right of the Periodic Table  Metalloids (semi-metals): elements close to the zig-zag line (e.g. B, Si, As, etc.) 

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Section 2

s-block elements (s 棟元素 ): Elements with outermost shell electrons occupying the s orbital. Alkali metals: Group IA elements Alkaline metals: Group II A elements All the elements are active metals. Group similarities and trends within the groups are generally clear.

p-block elements (p 棟元素 ): Elements with valence electrons occupying the p orbitals. Group III to Group 0 elements Group VII (Halogens) Group O (Noble gases) Chemical behaviour in this block varies widely with the reactivity of the metals, metalloids and non-metals, and the comparative lack of reactivity of noble gases. Similarities within a group are shown by halogens and noble gases. Group IV elements carbon to lead show a dramatic change in chemical properties within a group (a transition from non-metal to metal). 



  

 

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d-block elements (d 棟 元 素 ): Also called transition elements. Elements with valence electrons occupying d orbitals. Transition elements are defined as elements which have incomplete d subshell when  combined in compounds. They frequently have coloured compounds and complex ions in various oxidation  states.  These elements and their compounds are often used as catalysts in many important industrial processes. They have high melting points, and are denser and harder than non-transition metals.  f-block elements: Lanthanide and actinide series (inner transition elements), having f orbitals occupied by electrons. Lanthanides are widely spread throughout the earth’s crust in trace amount. They are  also known as ‘rare earth elements’ (稀土元素).  Actinides are all radioactive metals

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Section 2

2.5 Atomic Properties of the Elements and their Variations 2.5.1. Variation of the first ionization enthalpies 

Ionization enthalpy is defined as the energy absorbed / enthalpy change when 1 mole of electrons is removed from 1 mole of atoms or ion of an element at gaseous phase (從 一 摩爾氣態的原子/離子的最外層移走一摩爾電子所需的能量). For example, the first ionization enthalpy of an element is the enthalpy change when 1  mole of electrons is removed from 1 mole of gaseous atoms of the element under standard conditions. X(g) → X+(g) + e– ΔH = 1st I.E. (in kJmol-1)

There are many factors affecting the magnitude of the ionization enthalpy of an atom: 1. The electronic configuration of the atom The atoms of some elements have half-filled or completely filled sub-shells which give them extra stability. Larger amount of energy is required to ionize these atoms. 



2.

3.

4.

Screening (shielding) effect (屏蔽效應) Electrons in an atom will have electronic repulsion. ♦ Electrons in the inner shells repel (screen) the electrons in the outer shells. ♦ Therefore, the attraction between the outer shell electrons and the nucleus will ♦ be weaker. (In other words, the electron in the outer shell is effectively shielded from the nucleus by the inner electrons.) Such effect will become more significant as the number of electron shell ♦ increases (∴ 1st I.E. decrease down the group). The effective nuclear charge (有效核電荷) ♦ As the atomic number increases across a period, there are more protons in the nucleus of atoms. ♦ This results in greater attraction between the nucleus and electrons. Hence, larger energy is required to ionize the atom. ♦ ♦ The effective nuclear charge is given by: E.N.C = Atomic number – Number of electrons in the inner shells The atomic radius (原子半徑)

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Modern College F.6 Chemistry (2009 – 10) ♦

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Section 2

When the atomic radius of an atom increases, the outermost shell electrons are further away from the nucleus. Then the attraction that the electrons experience from the nucleus would be less. Hence, the ionization enthalpy becomes lower.

Variation in 1st I.E. of the first 20 elements:

Discussion: 1. There is a general increase in 1st I.E. across both Periods 2 and 3: Moving across a period, there is an increase in the nuclear attraction due to  the addition of proton in the nucleus. (原子內的核電荷隨著原子序的上升而增 加) The added electron is placed in the same quantum shell. Thus it is only poorly shielded by other electrons in that shell. (所加入的電子是進入同一電子 

層，電子與電子之間的屏蔽效應較小) As a result, there will be an increase in effective nuclear charge across the  period (有效核電荷隨之增加，故此第一電離焓逐漸增加) Besides, going across a period, the atomic radii of elements decrease and  hence the size of electron cloud decreases.  As the electron cloud becomes closer to the nucleus, it becomes more difficult to lose an electron. 2.

In going down any group, the 1st I.E. generally decreases. For each element down the group, the outermost electron which is to be  removed is screened by an extra shell of electrons. Atomic size increases down the group. Therefore, the distance between the  nucleus and the outermost electrons increases.

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More difficult for the nucleus to hold this outermost electron (attraction ↓). Thus electron is easier to be removed. 

3. 



4. 



1st I.E. of B < 1st I.E. of Be, Be has a completely filled 2s-subshell. It has extra stability and more energy is needed to remove its outermost shell electrons. B has an additional 2p electron beyond the full 2s subshell. It is easier to remove an outermost 2p electron because this electron is further from the nucleus and is shielded from the attraction of the nucleus by a full 2s subshell. 1st I.E. of Al < 1st I.E. of Mg Mg has a completely filled 3s-subshell. It has extra stability and more energy is needed to remove its outermost shell electrons. Al has an additional 3p electron beyond the full 3s subshell. It is easier to remove an outermost 3p electron because this electron is further away from the nucleus and is shielded from the attraction of the nucleus by a full 3s subshell. (Al 有 一 粒 額外增加的 3p 電子，它離原子核更遠、亦會被已填滿的 3s 亞層遮蔽， 故此較易被遷移。)

5.  



 



6.  

  

1st I.E. of O < 1st I.E. of N N has a half-filled 2p subshell. All 2p orbitals are singly occupied and electronic charge is evenly distributed. Therefore electronic repulsion is minimized. (電 荷 平 均 分 佈 ， 故 此 電 子 之 間 的排斥力減低。) N has extra stability. Thus the 3 electrons in 2p of N need more energy to remove. (N 有額外的穩定性，故此較難發生電離。) On the other hand, O has 4 electrons in the outermost 2p subshell. The fourth electron goes into a 2p orbital which is already occupied by an electron. (其中一個 p 軌態有一對電子，因此電子之間的排斥力增加。) The increased electronic repulsion causes this electron to be less tightly held. Therefore, less energy is required to remove an outermost electron from O. 1st I.E. of S < 1st I.E. of P P has a half-filled 3p subshell. All 3p orbitals are singly occupied and electronic charge is evenly distributed. Therefore electronic repulsion is minimized. P has extra stability. Thus the 3 electrons in 3p of P need more energy to remove. On the other hand, S has 4 electrons in the outermost 3p subshell. The fourth electron goes into a 3p orbital which is already occupied by an electron The increased electronic repulsion causes this electron to be less tightly held. Therefore, less energy is required to remove an outermost electron from S.

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Modern College F.6 Chemistry (2009 – 10)



There is a large drop in 1st I.E. when moving from one period to another. The last element of each period has the stable noble gas electronic configuration (ns2np6). The ionization enthalpy will be very high. (一個週期內的最後一個元



素有穩定的貴氣體電子排佈。電離焓極高。) At the beginning of a new period, the additional s electron is added to a new electron shell which is further away from the nucleus. (一個週期內的第一個元



素有 一 粒 額外 增加 的 s 電 子， 它離 原子 核 更 遠、 亦 會 被內 層的 電 子 遮 蔽。) Also, this s electron is effectively shielded from the nucleus by the inner shells. As a result, this outermost s electron is less firmly attracted by the nucleus and easily removed, leading to a particularly low 1st I.E. in Group I elements.

7.



8. 

9. 



Section 2

1st I.E. of Ne and Ar is maximum in the 2nd and 3rd period respectively They have completely filled n = 2 and n = 3 electron shell respectively, which they are most stable. Helium has the highest 1st I.E. of all elements. The atomic size of helium is the smallest of all elements and the only 2 electrons present are not screening one another and do not experience from any screening effect. Electrons experience full positive nuclear attraction.

The following figure shows the comparison between 1st I.E. and 2nd I.E.

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Section 2

The trend of variation of 2nd I.E. is similar to that of 1st I.E. 2nd I.E. is always higher than 1st I.E. One more positive charge (∴ have to overcome the attraction force) ♦ Successive electron may be removed from a lower quantum shell (e.g. 2nd ♦ I.E. of Group I elements) which is closer to the nucleus. (連 續 遷 移 的 電 ♦

子可能是從較低的電子層電離，而該電子離原子核較近) Successive electron may belong to a half-filled or completely filled subshell which has extra stability. (連續遷移的電子可能屬於一個全滿 /半滿的電子亞層，而該電子亞層擁有額外穩定性)

2.5.2. Variation of the atomic radius 





Atomic radius is not a precisely defined physical quantity, nor is it constant in all circumstances. The value assigned to the radius of a particular atom always depends on the definition chosen for “atomic radius”, and the appropriate definition depends on the context. Actually the term “atomic radius” itself is problematic: it may refer to the size of free atoms, or it may be used as a general term for the different measures of the size of atoms. At least three ways of definition: Covalent radius (共價半徑) – Half of the distance between two nuclei in an element  where the atoms are chemically combined together (joined by covalent bonding). van der Waals’ radius (范德華半徑) – Half of the distance between two nuclei in an  element where the atoms are not chemically combined together. Metallic radius (金屬半徑 ) – Half of the distance between two nuclei in an element  where the atoms are joined by metallic bonding.





If not specified, atomic radius normally refers to covalent radius. However, for those free atoms without formation of chemical bonds, atomic radius  refers to the van der Waals’ radius.

There are few factors affecting the atomic radius of an atom:

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Section 2

Atomic radii decrease across a period because The numbers of protons increase. Therefore, the nuclear charge increases  with atomic number. The additional electrons enter the same electron shell, while the electrons in  the same shell do not shield each other from the nuclear charge effectively. The effective nuclear charge of successive elements increases and the  electrostatic attraction between outer electrons and positively charged nucleus increases. The electrons are pulled closer towards the nucleus, resulting in a reduction  in atomic radius. Within a period, alkali metal atom has the largest size because the outer  electron is in a new electron shell, which is further away from the nuclear attraction. The atom of noble gas has the smallest size in a period because the electrons  in the same shell do not shield each other from the nuclear charge effectively while the effective nuclear charge is large. Atomic radii increase down a group Each successive element has one more filled inner electron shell than the  previous element in the same group. As the electrons occupy electron shells of greater quantum number in  successive period, they become further away from the nucleus. Also, as the inner electrons can effectively screen the outermost electron and  number of inner electrons increase, the outermost electrons are less attracted by the nucleus. As a result, atomic size increases down the group. 

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Section 2

2.5.3. Electron affinity (電子親合勢 /電子親和力 ) 

The electron affinity for an element is the enthalpy change when 1 mole of electrons is added to 1 mole of neutral atoms in the gaseous state to form 1 mole of gaseous anions under standard conditions. X(g) + e– → X–(g)



ΔH = E.A.1 (in kJmol-1)

The 2nd E.A. and 3rd E.A. can be defined as: X–(g) + e– → X2–(g) X2–(g) + e– → X3–(g)



















ΔH = E.A.2 ΔH = E.A.3

Electron affinities are generally obtained indirectly from thermochemical cycles because direct measurement is difficult. The value of the electron affinity depends on the attraction between the incoming electron and the nucleus and the shielding effect offered by the existing electrons. In general, the 1st E.A. is increasingly exothermic (more negative) from left to right across the periodic table. Going across a period, the extra electrons enter the same electron shell and so they are  not shielded from the nuclear attraction effectively. Moreover, as the nuclear charge increases with the atomic number, atoms have  progressively greater attraction for electrons. Hence, addition of electron becomes more favorable ( ⇒ E.A. becomes more  exothermic). Some electron affinities are given below, the values being in kJ mol–1.

When an electron is added to an atom with a fully filled or half-filled outer subshell (e.g. nitrogen), the 1st E.A. is less negative or endothermic since such electronic configuration shows enhanced stability. The halogens are expected to have the most negative 1 st E.A. because the addition of one electron to the atoms leads to a stable noble gas electronic configuration. 1st E.A. of alkali metals usually close to zero or endothermic. This is due to strong shielding effect of s-subshell. The 1st E.A. is usually being less negative when going down a group. Take Group VII as an example:

(The values being in kJ mol-1) Going down a group, the atomic radii of the elements increase.

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Section 2

The outermost electrons become further away from the nucleus and so they would experience less nuclear attraction. Besides, as the number of occupied inner electron shells increase, the screening  effect becomes more significant. This also leads to a reduction of nuclear attraction. The value for fluorine is anomalous (反常的) because it includes an electron cloud with  high repulsion. This is a consequence of a low atomic radius and a compact outer electron shell. The second electron affinity of an element relates to the uptake of a second electron, and is defined as the enthalpy required adding one mole of electrons to one mole of isolated uninegative ions. 2nd E.A. and so on are usually endothermic (+ve) because there is electronic repulsion between the anion and the incoming electron and the repulsion component outweighs the attraction component. Take oxygen and sulphur as an example: Oxygen Sulphur 1st E.A. (kJ mol-1) –142 –200 2nd E.A. (kJ mol-1) +791 +649 







The 1st E.A. of O and S are exothermic because the attraction between the incoming electron and the nucleus is stronger than the repulsion between the incoming electron and the existing electrons. The 2nd E.A. of O and S are endothermic because in the presence of the extra electron, the extra electron charge offers a stronger repulsion with the incoming electron. Also, the electron cloud of O– and S– will expand and the incoming electron will occupy a position further from the nucleus which reduces the attraction with the nucleus. 



2.5.4. Electronegativity (電負性 /電負度 ) 







 

Electronegativity is a measure of tendency of an atom in a stable molecule to attract electrons within a bond (原子在分子中吸引鍵合電子的相對能力). For example, in H–Cl molecule, the bonding electrons are much closer to the chlorine atom than to the hydrogen atom. The chlorine atom is the more electronegative(電負的/陰電的) atom in this case. There are several methods to measure the electronegativity of an element. One commonly used method is the Pauling scale of electronegativity (鮑 林 電 負 性 標 度 ). It is developed by Linus Pauling, an American physical chemist who spent most of his life in the research of the nature of chemical bonding. The scale is from 0 to 4. Pauling defined the electronegativity of an atom as the power of that atom in a molecule to attract electrons. In contrast with electron affinity, it is a measure of the power of a single gaseous atom to attract electrons. In Pauling scale, 4 is assigned to the most electronegative atom, fluorine. The electronegativity of monoatomic molecules (i.e. noble gases) is set as 0.

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Section 2

In general, electronegativity values increase from left to right across each period This can be explained in term of the increase in effective nuclear charge across a  period. Although an extra electron is also added for each element, this does not fully shield the effect of the increase nuclear charge. An increase in effective nuclear charge makes the attraction for the bonded electrons increase. Electronegativity decreases when going down each group This occurs because, in moving down a group, the screening of the atoms increases and  the attractive force of the shielded nucleus is therefore reduced.

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