White Light. Chapter 7 Electron Structure of the Atom

Transcription

1 Chapter 7 Electron Structure of the Atom Electromagnetic Radiation and Energy The Bohr Model of the Hydrogen Atom The Modern Model of the Atom Periodicity of Electron Configurations Valence Electrons for Main-Group Elements Electron Configurations for Ions Periodic Properties of Atoms Copyright The McGraw-Hill Companies, Inc. Permission required for reproduction or display. 7-1 Is composed of different colors that can be separated by a prism Water acts as a prism for sunlight, giving the effect of a rainbow Sources of white light Sun Regular (incandescent) light bulbs White Light 7-2 Electromagnetic Radiation A form of energy Travels through space at the speed of light (3.0 x 10 8 m/s) as oscillating waves Is both an electric and a magnetic field Also called radiant energy Some examples of EM radiation Light X-rays 7-3 1

2 Differentiating the Kinds of Electromagnetic Radiation Two main characteristics Wavelength (λ) Frequency (ν) Figure Differentiating the Kinds of Electromagnetic Radiation Wavelength (λ) The distance between two corresponding points on a wave Units are same as length - m, or commonly nm (10-9 m) 7-5 Differentiating the Kinds of Electromagnetic Radiation Frequency (ν) A measure of the number of wave cycles that move through a point in space in 1 s Units are hertz (Hz) which are the same as inverse seconds (1/s) 7-6 2

3 Frequency, Wavelength, and the Electromagnetic Spectrum Frequency and wavelength are inversely proportional i.e. as one increases the other decreases c = λν Where c = speed of light (3.0 x 10 8 m/s), λ = wavelength (in meters), and ν = frequency (in Hz) 7-7 Frequency, Wavelength, and the Electromagnetic Spectrum Which has a greater frequency, the red light or the green light? Which has the greater wavelength? Figure Duality of Light Light exists as both waves and particles (or packets of light called photons) Characteristics of waves Frequency Wavelength c = λν Characteristic of photons Energy of a photon Is directly proportional to the frequency and inversely proportional to the wavelength E photon = hν Where E photon = energy of the photon (in Joules), h = Planck s constant (6.626 x Js), and ν = frequency (in Hz) 7-9 3

4 λ, ν, and E photon c =λν E photon = hν Using algebra, we can manipulate these two equations several ways: For c =λν, We can solve for λ: λ = c / ν or ν: ν = c / λ For E photon = hν We can substitute c / λ forν, giving us the equation: E photon = (hc) / λ This equation shows the inverse proportionality between E photon and λ (wavelength) 7-10 Practice λ, ν, and E photon If the wavelength of a microwave beam is 11.5 cm, then what are the frequency of the radiation and the energy of its photons? 7-11 Practice Solutions λ, ν, and E photon If the wavelength of a microwave beam is 11.5 cm, then what are the frequency of the radiation and the energy of its photons? λ = 11.5 cm First, convert to meters: λ = 11.5 cm x (1 m/100 cm) = 1.15 x 10-3 m ν = c / λ ν = (3.0 x 10 8 m/s)/1.15 x 10-3 m = 2.6 x s

5 Line Spectra Continuous spectrum Contains all the wavelength of light in the visible spectrum Produced by white light Line Spectrum Contains a pattern of distinct colored lines, each representing a single wavelength of light Produced by an element that has been heated or given an electric charge Each element has a distinct line spectra, which is also called atomic fingerprint 7-13 Energy is Quantized! Max Planck first hypothesized that energy produced by atoms can only have certain values and is therefore quantized. That s the reason why only distinct lines are seen in element line spectras. Energy is quantized and can only exist at certain wavelengths Bohr Model Niels Bohr hypothesized that electrons orbit the nucleus just as the planets orbit the sun (planetary model). He labeled the electron orbits with a number, starting with 1 closest to the nucleus and increasing as the orbits get further away from the nucleus. The number is known as the Principal Quantum Number (n)

6 Bohr Model Orbits have a fixed radius. The orbit with the lowest energy is closest to the nucleus. The energy of each orbit increases as the orbits get further away from the nucleus. When an electron jumps from one orbit to another, it absorbs or emits energy according to the equation: E = E f E i 7-16 Modern Model of the Atom The modern model of the atom is based on Schrodinger s mathematical model of waves This model describes electrons as occupying orbitals, not orbits. Orbitals Three dimensional regions in space where electrons are likely to be found, not a circular pathway Principal energy level Orbitals of similar size Figure Modern Model of the Atom Figure

7 Orbitals Figure 7.13 Come in different shapes and sizes. Lower energy orbitals are smaller. Higher energy orbitals are larger and extend further away from the nucleus. Four most common types are s, p, d, and f. Also known as sublevels Consists of just one type of orbital at a specific energy level The number of sublevels is equal to n, the Principal Quantum Number 7-19 s Orbitals Figure p Orbitals

8 d Orbitals 7-22 Hydrogen Orbital Diagram Orbital diagrams Show the sublevels and orbitals that can exist at each principal energy level Each box represents an orbital Groups of boxes represent sublevels In the hydrogen atom only, the sublevels within a principal energy level all have the same energy Multielectron Orbital Diagram In the multielectron atoms, the sublevels within a principal energy level have different energy levels

9 Orbital Diagram Rules Two principles and 1 rule determine how the electrons are filled in the principal energy levels and sublevels. Aufbau principle Electrons fill orbitals starting with the lowest-energy orbitals. Pauli exclusion principle A maximum of two electrons can occupy each orbital, and they must have opposite spins. Hund s rule Electrons are distributed into orbitals of identical energy (same sublevel) in such a way as to give the maximum number of unpaired electrons. Electrons are always filled in their ground state, or lowest energy state Filling Orbital Diagrams Pg. 251 (Carbon s orbital diagram) 7-26 Orbital Diagrams for the 1 st Ten Elements

10 Electron Configurations Shorthand notation which shows the distribution of electrons among sublevels When we write electron configurations, we write the number of the principal quantum number followed by a symbol for the sublevel, and then add a superscript to each sublevel symbol to designate the number of electrons in that sublevel. Carbon has 6 electrons. Therefore, using the orbital diagram we obtain: 1s 2 2s 2 2p Periodicity of Electron Configurations Can you tell the patterns among the following groups of elements? Alkali Metals (Group IA (1)) Li 1s 2 2s 1 Na 1s 2 2s 2 2p 6 3s 1 Alkali Earth Metals (Group IIA (2)) Mg 1s 2 2s 2 2p 6 3s 2 Ca 1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 Halogens (Group VIIA (17)) Cl 1s 2 2s 2 2p 6 3s 2 3p 5 Br 1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 3d 10 4p 5 Noble Gases (Group VIIIA (18)) Ne 1s 2 2s 2 2p 6 Ar 1s 2 2s 2 2p 6 3s 2 3p Using the Periodic Table for Electron Configurations The periodic table can be used to fill orbital diagrams or to find electron configurations. First, we need to separate the periodic table into blocks. Blocks contain elements with the same highest-energy sublevel

11 Using the Periodic Table for Electron Configurations Example: 1s 2 Principal quantum number (1) Same as the period number for s and p Ex. 3s and 3p are both in Period 3 (Period number 1) for d Ex. 3d is in Period 4 (Period number 2) for f Ex. 4f is in Period 6 Sublevel (s) also called the secondary quantum number Labeled according to the block you re in Number of electrons - superscript (2) The number of elements in the block in the period 7-31 The Principal Quantum Number and Sublevel on the Periodic Table Figure An Example of Electron Configuration Manganese (Mn) has 25 electrons. Start at the top left corner of the periodic table and move from left to right, top to bottom. The period # is 1, groups 1 and 2 are the s block, and there are two elements in the s block in Period 1 (when we move He next to H on the periodic table), so we write 1s

12 An Example of Electron Configuration Since there are no more elements in period 1, we move to the next period #, which is 2, groups 1 and 2 are the s block, and there are two elements in the s block in period 2, so we write 2s 2. 1s 2 2s 2 Moving across the periodic table, the period # is still 2, groups are the p block, and there are 6 elements in the p block in period 2, so write 2p 6. 1s 2 2s 2 2p Copyright The McGraw-Hill Companies, Inc. Permission required for reroduction or display. 34 An Example of Electron Configuration There are no more elements on that period, so let s go to period 3. In period 3, starting at the left, groups 1 and 2 are the s block, and there are 2 elements in the s block on period 3, so we write 3s 2. 1s 2 2s 2 2p 6 3s 2 In the p block in period 3, there are 6 elements so we write 3p 6. 1s 2 2s 2 2p 6 3s 2 3p An Example of Electron Configuration 1s 2 2s 2 2p 6 3s 2 3p 6 Adding together the superscripts to make sure what we still need to go on: = 18 electrons Remember, Manganese (Mn) has 25 electrons, so we need to go on

13 An Example of Electron Configuration 1s 2 2s 2 2p 6 3s 2 3p 6 Since there are no more elements in Period 3, we go to Period 4, where on the left, we are again in the s block, and there are 2 elements in the s block in Period 4. Moving right, we come across the d block, which is 3d (period # - 1), and there are 5 elements in 3d, including Mn. 1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 3d An Example of Electron Configuration 7-38 Practice Electron Configurations Write electron configurations for the following: 1. Al 2. Sc 3. K 4. Br 5. Zn 6. Hg

14 Practice Solutions Electron Configurations Write electron configurations for the following: 1. Al 1s 2 2s 2 2p 6 3s 2 3p 1 2. Sc 1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 3d 1 3. K 1s 2 2s 2 2p 6 3s 2 3p 6 4s 1 4. Br 1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 3d 10 4p 5 5. Zn 1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 3d Hg 1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 3d 10 4p 6 5s 2 4d 10 5p 6 6s 2 4f 14 5d Valence Electrons for Main-Group Elements Valence level (shell) Last-filled principal energy level Highest energy level Contains orbitals that are larger than orbitals in lower energy levels Valence electron An electron that occupies the valence level Elements in the same group have the same number of valence electrons Example: Br 1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 3d 10 4p Valence Electrons for Main-Group Elements Valence electrons The Roman numeral group number (which is paired with an A or B). The number of valence electrons is also equal to the number of s and p electrons in the valence level for any main-group element. Core electron An electron in a principal energy level below the valence level Inner electron

15 Valence Electrons for Main-Group Elements Figure Practice Valence Electrons For each of the following, determine the number of valence electrons. 1. Magnesium 2. Carbon 3. Boron 4. Chlorine 5. Selenium 7-44 Practice Solutions Valence Electrons For each of the following, determine the number of valence electrons. 1. Magnesium In group IIA (or 2): 2 valence electrons 2. Carbon In group IVA (or 14): 4 valence electrons 3. Boron In group IIIA (or 13): 3 valence electrons 4. Chlorine In group VIIA (or 17): 7 valence electrons 5. Selenium In group VIA (or 16): 6 valence electrons

16 Abbreviated Electron Configuration Starts the electron configuration at the last noble gas before the element in question. Write down the noble gas in brackets, then fill in the rest of the electron configuration until you reach the element in question. Example: Phosphorus (P) has 15 electrons. In the long notation, the electron configuration for P would be: 1s 2 2s 2 2p 6 3s 2 3p 3 The abbreviated electron configuration for P would be: [Ne] 3s 2 3p Practice Abbreviated Electron Configuration Write the abbreviated electron configuration for the following: 1. Magnesium 2. Carbon 3. Boron 4. Chlorine 5. Selenium 7-47 Practice Solutions Abbreviated Electron Configuration Write the abbreviated electron configuration for the following: 1. Magnesium [Ne] 3s 2 2. Carbon [He] 2s 2 2p 2 3. Boron [He] 2s 2 2p 1 4. Chlorine [Ne] 3s 2 3p 5 5. Selenium [Ar] 4s 2 3d 10 4p

17 Electron Configurations for Ions Ions form because atoms gain or lose electrons Cations Positively charged ions Subtract the number of the charge from the total number of electrons Move to the left the number of spaces equal to the charge on the periodic table Anions Negatively charged ion Add the number of the charge to the total number of electrons Move to the right the number of spaces equal to the charge on the periodic table 7-49 Electron Configurations for Ions Isoelectronic Have the same number of electrons Ions and elements with the same electron configuration Example: Are the bromine ion and strontium ion isoelectronic? Bromine forms an ion with a -1 charge, while strontium form an ion with a +2 charge. Moving the correct spaces on the periodic table, we find that bromine s ion and strontium s ion share an electron configuration with krypton, and are therefore isoelectronic. Br - 1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 3d 10 4p 6 Sr s 2 2s 2 2p 6 3s 2 3p 6 4s 2 3d 10 4p 6 Kr - 1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 3d 10 4p Electron Configurations for Ions An example problem: Write the electron configuration for Na + : Na + has a positive charge of 1; therefore, we need to subtract 1 electron from the total number of electrons, 11. Na + has 10 electrons and is isoelectronic with Ne. 1s 2 2s 2 2p

18 Practice Electron Configurations for Ions Write the electron configuration in long and abbreviated notation for the following ions. 1. Br - 2. N 3-3. K + 4. Sr S 2-6. Ni Practice Solutions Electron Configurations for Ions Write the electron configuration in long and abbreviated notation for the following ions. 1. Br - 1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 3d 10 4p 6 [Kr] isoelectronic with Kr 2. N s 2 2s 2 2p 6 [Ne] isoelectronic with Ne 3. K + - 1s 2 2s 2 2p 6 3s 2 3p 6 [Ar] isoelectronic with Ar 4. Sr s 2 2s 2 2p 6 3s 2 3p 6 4s 2 3d 10 4p 6 [Kr] isoelectronic with Kr and Br S s 2 2s 2 2p 6 3s 2 3p 6 [Ar] isoelectronic with Ar and K Ni s 2 2s 2 2p 6 3s 2 3p 6 [Ar]4s 2 3d 6 isoelectronic with Fe 7-53 Periodic Trends Valence electrons are the electrons that participate in chemical reactions because they are the farthest electrons from the nucleus. Because elements in the same group have the same number of valence electrons, elements in the same group have very similar reactivities. Figure

19 Ionization Energy Ionization Energy A measure of the energy required to remove a valence electron from a gaseous atom to form a gaseous ion. In general, atoms with low ionization energies do not bind their electrons very tightly, and are therefore, very reactive. Figure Trends in Ionization Energy The general trend for ionization energy is for ionization energy to increase from bottom to top and from left to right across the periodic table. Figure Successive Ionization Energies The trend describes the first ionization energy (IE 1 ), or the amount of energy it takes to remove 1 electron from an atom. The amount of energy it takes to remove a 2 nd electron is known as the second ionization energy (IE 2 ) and is larger than IE 1. In general, IE 3 > IE 2 > IE

20 Successive Ionization Energies 7-58 Atomic Size Atomic size is often described in terms of atomic radius. Atomic radius is the distance from the center of the nucleus to the outer edge of the atom. Figure Trends in Atomic Size

21 Trends in Atomic Size The general trend for atomic size (or radius) is for atomic size to increase from top to bottom and from right to left across the periodic table. Figure Ionic size Radius of an ion Atoms change radius when they become ions Ionic Size 7-62 For any isoelectronic series, as the number of protons increases, the ion size decreases. Ionic Size

22 Trends in Ionic Size The general trend for ionic size (or radius) is for ionic size to increase from top to bottom. For cations, as the charge increases, the ionic size decreases. For anions, as the charge increases in negativity, the ionic size increases