Chapter 2. Atomic Structure and Periodicity

Transcription

1 Chapter 2 Atomic Structure and Periodicity

2 Chapter 2 Table of Contents (2.1) (2.2) (2.3) (2.4) (2.5) (2.6) (2.7) (2.8) (2.9) Electromagnetic radiation The nature of matter The atomic spectrum of hydrogen The Bohr model The quantum mechanical model of the atom Quantum numbers Orbital shapes and energies Electron spin and the Pauli principle Polyelectronic atoms

3 Chapter 2 Table of Contents (2.10) (2.11) (2.12) (2.13) The history of the periodic table The aufbau principle and the periodic table Periodic trends in atomic properties The properties of a group: The alkali metals

4 Chapter 2 Questions to Consider Why do we get colors? Why do different chemicals give different colors?

5 Section 2.1 Electromagnetic Radiation Electromagnetic Radiation - Characteristics One of the ways by which energy travels through space Three characteristics: Wavelength (λ) is the distance between two peaks or troughs in a wave Frequency (ν) points to the number of waves (cycles) per second that pass a given point in space Speed The speed of light is m/s

6 Section 2.1 Electromagnetic Radiation Relationship between Wavelength and Frequency There exists an inverse relationship between λ and ν λ 1/ n or λn = c The way I remember this n=c/l

7 Section 2.1 Electromagnetic Radiation Figure The Nature of Waves

8 Section 2.1 Electromagnetic Radiation Figure Classification of Electromagnetic Radiation

9 Section 2.1 Electromagnetic Radiation Interactive Example Frequency of Electromagnetic Radiation The brilliant red colors seen in fireworks are due to the emission of light with wavelengths around 650 nm when strontium salts such as Sr(NO 3 ) 2 and SrCO 3 are heated. Calculate the frequency of red light of wavelength nm. Solution We can convert wavelength to frequency using the equation: λn = c or n = c λ

10 Section 2.1 Electromagnetic Radiation Interactive Example Frequency of Electromagnetic Radiation Where c = m/s. In this case λ = nm Changing the wavelength to meters, we have nm 9 1m 7 = m 10 nm and ν c = = λ m m / s = s = Hz

11 Section 2.1 Electromagnetic Radiation Pickle Light On applying an alternating current of 110 volts to a dill pickle, a glowing discharge can be observed The Na + and Cl ions in the forks cause the sodium atoms to get into an excited state When the atoms reach their ground state, they emit visible light at 589 nm

12 Section 2.2 The Nature of Matter An Introduction Max Planck observed that energy can be gained or lost only in whole-number multiples of hν Where h is Planck's constant with a value of J s The change in energy for a system ΔE can be represented by the equation ΔE = nhn, where n is an integer h is Planck s constant ν represents the frequency of electromagnetic radiation

13 Section 2.2 The Nature of Matter Quantized Energy Planck also observed that energy can be quantized and can occur only in discrete units called quanta A system can transfer energy only in whole quanta This proves that energy does have particulate properties

14 Section 2.2 The Nature of Matter Interactive Example The Energy of a Photon The blue color in fireworks is often achieved by heating copper(i) chloride (CuCl) to about 1200 C. Then the compound emits blue light having a wavelength of 450 nm. What is the increment of energy (the quantum) that is emitted at nm by CuCl? Solution The quantum of energy can be calculated from the equation: ΔE = hn

15 Section 2.2 The Nature of Matter Interactive Example The Energy of a Photon The frequency ν for this case can be calculated as follows: 8 c m n = = 7 λ m / s 14 1 = s So, Δ E = hn = J. s s 1 19 = J A sample of CuCl emitting light at 450 nm can lose energy only in increments of J, the size of the quantum in this case

16 Section 2.2 The Nature of Matter The Concept of Photons Electromagnetic radiation is a stream of particles called photons The energy of each photon can be expressed by: E = = hc photon hn λ Where h is Planck s constant, ν is the radiation frequency, and λ is the radiation wavelength

17 Section 2.2 The Nature of Matter The Photoelectric Effect The phenomenon whereby electrons are emitted from the surface of a metal when light strikes it Characteristics: No electrons are emitted by any given metal below a specified threshold frequency, ν 0 When ν < ν 0, no electrons are emitted, regardless of the intensity of light When ν > ν 0, the number of electrons increases with the intensity of light When ν > ν 0, the kinetic energy of emitted electrons increases linearly with the frequency of the light

18 Section 2.2 The Nature of Matter The Photoelectric Effect Minimum energy required to remove an electron = E 0 = hν 0 When ν > ν 0, the excess energy that is required to remove the electron is given as kinetic energy (KE): Where m is the mass of the electron 1 2 KE electron = = 0 2 mv hn hn v is the velocity of the electron ν is the energy incident of the photon ν 0 is the energy required to expel the electron

19 Section 2.2 The Nature of Matter The Photoelectric Effect Greater intensity of light means that more photons are available to release electrons, which gave rise to the equation: 2 E = mc The theory of relativity signifies that energy has mass This equation can be used to calculate mass associated with a quantity of energy The mass of a photon of light with wavelength λ is given by: E hc / λ h m = = = c 2 c 2 λc

20 Section 2.2 The Nature of Matter Dual Nature of Light The phenomenon whereby electromagnetic radiation (and all matter) exhibits wave properties and particulate properties de Broglie s equation allows for the calculation of the wavelength of a particle h λ = mv

21 Section 2.2 The Nature of Matter Diffraction It is the result of light getting scattered from a regular array of points or lines This scattered radiation produces a diffraction pattern on bright and dark areas Scattered light can: Interfere constructively and produce a bright area Interfere destructively to produce a dark spot This phenomenon occurs best when the spacing between scattering points is similar to the wavelength of the diffracted wave

22 Section 2.2 The Nature of Matter Figure A Diffraction Pattern of a Beryl Crystal

23 Section 2.3 The Atomic Spectrum of Hydrogen Significance of the Hydrogen Emission Spectrum Continuous spectrum occurs when white light is passed through a prism Contains all the wavelengths of visible light Hydrogen emission spectrum is called a line spectrum Displays only a few lines, each line corresponding to discrete wavelengths Indicates that the energy of the electron on the hydrogen atom is quantized

24 Section 2.3 The Atomic Spectrum of Hydrogen Significance of the Hydrogen Emission Spectrum Change in energy from a high to lower level of a given wavelength can be calculated by: Δ E = hn = hc λ Wavelength of light emitted Change in energy Frequency of light emitted

25 Section 2.3 The Atomic Spectrum of Hydrogen Figure (a) Continuous Spectrum and (b) Line Spectrum

26 Section 2.3 The Atomic Spectrum of Hydrogen Concept Check Why is it significant that the color emitted from the hydrogen emission spectrum is not white? How does the emission spectrum support the idea of quantized energy levels?

27 Section 2.4 The Bohr Model Quantum Model for the Hydrogen Atom Electron in a hydrogen atom moves around the nucleus only in certain allowed circular orbits Bohr s model gave the hydrogen atom energy levels consistent with the hydrogen emission spectrum The expression for the energy levels available to the electron in the hydrogen atom can be expressed as: 2 18 Z E = J 2 n Where n is an integer and Z is the nuclear charge

28 Section 2.4 The Bohr Model Figure Electronic Transitions in the Bohr Model for the Hydrogen Atom: Part (a) (a) An energy-level diagram for the first three electronic transitions

29 Section 2.4 The Bohr Model Figure Electronic Transitions in the Bohr Model for the Hydrogen Atom: Part (b) and (c) (b) An orbit-transition diagram, which accounts for the experimental spectrum (c) The resulting line spectrum on a photographic plate

30 Section 2.4 The Bohr Model Quantum Model for the Hydrogen Atom When the hydrogen atom returns to its lowest possible energy state, it is called the ground state When the electron falls from n=6 to n=1, ΔE can be computed by: ΔE = energy of final state energy of initial state = J = E E = J J The negative sign indicates that the atom has lost energy and is in a more stable state

31 Section 2.4 The Bohr Model Quantum Model for the Hydrogen Atom The energy lost is taken away from the atom by the emission of a photon whose wavelength can be calculated from: c hc Δ E = h or λ = λ ΔE ΔE, the change in energy of the atom, is equal to the energy of the emitted photon hc λ = = ΔE 34 ( J.s 8 )( m / s J ) 8 = m

32 Section 2.4 The Bohr Model Interactive Example Energy Quantization in Hydrogen Calculate the energy required to excite the hydrogen electron from level n = 1 to level n = 2. Also calculate the wavelength of light that must be absorbed by a hydrogen atom in its ground state to reach this excited state. Solution Using the equation E E 2 18 Z E = J 2 n 1 = J = J = J = J with Z = 1, we have:

33 Section 2.4 The Bohr Model Interactive Example Energy Quantization in Hydrogen Δ E = E2 E1 = J J = J The positive value for ΔE indicates that the system has gained energy The wavelength of light that must be absorbed to produce this change is: hc λ = = ΔE 7 = m 34 ( J.s 8 ) ( m / s J )

34 Section 2.4 The Bohr Model Importance of the Bohr Model The model correctly fits the quantized energy levels of the hydrogen atom It postulates only certain allowed circular orbits for the electron As the electron becomes more tightly bound, its energy becomes more negative relative to the free electron The free electron is at infinite distance from the nucleus As the electron is brought closer to the nucleus, energy is released from the system

35 Section 2.4 The Bohr Model Energy Change Between Levels in a Hydrogen Atom The general equation for the electron moving from n initial to n final can be derived by using the following equation: 2 18 Z E = J 2 n Δ E = energy of level n energy of level n = E final E initial 18 = J final 2 2 final initial initial = ( J) J 2 nfinal ninitial 1 1 n n

36 Section 2.4 The Bohr Model Drawbacks of the Bohr Model This model only works for hydrogen Electrons do not move around the nucleus in circular orbits

37 Section 2.5 The Quantum Mechanical Model of the Atom Erwin Schrödinger and Quantum Mechanics Standing waves are stationary waves that do not travel along any length Only certain orbits have a circumference into which whole number wavelengths of standing electron waves will fit Other waves produce destructive interference of the standing electron wave The mathematical representation for a standing wave is: Hˆ = E

38 Section 2.5 The Quantum Mechanical Model of the Atom Erwin Schrödinger and Quantum Mechanics ψ represents the wave function, which is a function of the coordinates of the electron s position in 3-dimensional space Ĥ represents an operator A specific wave function is termed as an orbital Wave function does not provide information about the pathway of the electron

39 Section 2.5 The Quantum Mechanical Model of the Atom Heisenberg Uncertainty Principle There is a fundamental limitation to just how precisely we can know both the position and momentum of a particle at a given time Mathematically, this principle can be represented by: Where Δx is the uncertainty in a particle s position Δ(mν) is the uncertainty in a particle s momentum h is Planck s constant Δ x. Δ( mv) h 4

40 Section 2.5 The Quantum Mechanical Model of the Atom Physical Meaning of the Wave Function The square of the function, represented as a probability distribution, indicates the probability of finding an electron near a particular point in space The intensity of color is used to indicate the probability value near a given point in space The more time the electron visits a particular point, the darker the negative becomes This diagram is known as an electron density map

41 Section 2.5 The Quantum Mechanical Model of the Atom Figure Probability Distribution for the 1s Wave Function - Part (a) and (b) (a) The probability distribution for the hydrogen 1s orbital in threedimensional space (b) The probability of finding the electron at points along a line drawn from the nucleus outward in any direction

42 Section 2.5 The Quantum Mechanical Model of the Atom Physical Meaning of a Wave Function A radial probability distribution graph plots the total probability of finding an electron in each spherical shell versus the distance from the nucleus Probability of finding an electron at a particular position is greatest near the nucleus Volume of the spherical shell increases with distance from the nucleus The size of the 1s orbital can be stated as the radius of the sphere that encloses 90% of the total electron probability

43 Section 2.5 The Quantum Mechanical Model of the Atom Figure Radial Probability Distribution (a) Cross section of the hydrogen 1s orbital probability distribution divided into successive thin spherical shells (b) Radial probability distribution plot

44 Section 2.6 Quantum Numbers An Introduction to Quantum Numbers They express the various properties of the orbital Principal quantum number (n) has integral values and is related to the size and energy of the orbital Angular momentum quantum number (l or l) has integral values from 0 to n 1 It is related to the shape of atomic orbitals (sometimes called a subshell) Magnetic quantum number (m l ) has integral values +l to -l It is related to the orientation of the orbital in space relative to the other orbitals in the atom Electron spin quantum number (m s ) can be + ½ or ½ It means that the electron can spin in either of the two opposite directions

45 Section 2.6 Quantum Numbers

46 Section 2.6 Quantum Numbers

47 Section 2.6 Quantum Numbers

48 Section 2.6 Quantum Numbers Interactive Example Electron Subshells For principal quantum level n = 5, determine the number of allowed subshells (different values of l), and give the designation of each Solution For n = 5, the allowed values of l run from o to 4 (n 1 = 5 1 ) Thus, the subshells and their designations are: l = 0 l = 1 l = 2 l = 3 l = 4 5s 5p 5d 5f 5g

49 Section 2.7 Orbital Shapes and Energies An Introduction Areas of zero probability are called nodal surfaces or nodes The number of nodes increase as n increases The number of nodes for the s orbital is given by n 1

50 Section 2.7 Orbital Shapes and Energies Figure Representations of the Hydrogen 1s, 2s, and 3s Orbitals

51 Section 2.7 Orbital Shapes and Energies p Orbitals Not spherical like s orbitals Have two lobes separated by a node at the nucleus Labelled as per the axis of the xyz coordinate system

52 Section 2.7 Orbital Shapes and Energies The d Orbitals Do not correspond to principal quantum levels n = 1 and n = 2 First level occur in level n = 3 They possess two fundamental shapes: d xz, d yz, d xy, and d x 2 y 2 d z 2 Have four labels that are centered in the plane that appears in the orbital label Possesses a unique shape with two lobes that run along the z axis and a belt centered in the xy plane d orbitals, where n>3, appear as 3d orbitals and have larger lobes

53 Section 2.7 Orbital Shapes and Energies Figure Representation of the 3d Orbitals - Part (b) (b) The boundary surfaces of all five 3d orbitals, with the signs (phases) indicated

54 Section 2.7 Orbital Shapes and Energies f Orbitals and Degenerates The f orbitals first occur in level n = 4 These orbitals are not involved in bonding in any compounds All orbitals with the same value of n have the same energy and are said to be degenerate In the ground state, the single hydrogen electron can be found in the 1s orbital This electron can be excited to higher-energy orbitals

55 Section 2.7 Orbital Shapes and Energies Figure Representation of the 4f Orbitals in Terms of Their Boundary Surfaces

56 Section 2.8 Electron Spin and the Pauli Principle Electron Spin Electron spin quantum number (m s ) can be + ½ or ½ It means that the electron can spin in either of the two opposite directions Pauli exclusion principle states that in a given atom no two electrons can have the same set of four quantum numbers An orbital can hold only two electrons, and they must have opposite spins because only two values of m s are allowed

57 Section 2.8 Electron Spin and the Pauli Principle Figure The Spinning Electron - Part (a) and (b) Spinning in one direction, the electron produces the magnetic field oriented as shown in (a) Spinning in the opposite direction, it gives a magnetic field of the opposite orientation, as shown in (b)

58 Section 2.9 Polyelectronic Atoms An Introduction Polyelectronic atoms are those atoms with more than one electron Electron correlation problem Since the electron pathways are unknown, electron repulsions cannot be accurately calculated When electrons are placed in a particular quantum level, the orbital levels vary in energy as follows: E ns < E np < E nd < E nf

59 Section 2.9 Polyelectronic Atoms

60 Section 2.9 Polyelectronic Atoms The Penetration Effect A 2s electron on average is closer to the nucleus (penetration) than one in the 2p orbital This causes an electron in a 2s orbital to be attracted to the nucleus more strongly than an electron in a 2p orbital The 2s orbital is lower in energy than the 2p orbitals in a polyelectronic atom The same occurrence can be noticed in other quantum levels

61 Section 2.9 Polyelectronic Atoms Figure Part (b) The Radial Probability Distribution for the 3s, 3p, and 3d orbitals

62 Section 2.10 The History of the Periodic Table The periodic table was originally constructed to represent the patterns observed in the chemical properties of elements Mendeleev s periodic table: Emphasized on how the table could help estimate the existence and properties of unkown elements Rectified several values of atomic masses The current periodic table lists elements by their atomic number rather than atomic mass

63 Section 2.11 The Aufbau Principle and the Periodic Table Aufbau Principle As protons are added one by one to the nucleus to build up the elements, electrons are similarly added to hydrogen-like orbitals Example - An oxygen atom has an electron arrangement of two electrons in the 1s subshell, two electrons in the 2s subshell, and four electrons in the 2p subshell

64 Section 2.11 The Aufbau Principle and the Periodic Table Orbital Diagrams They represent the number of electrons an atom has in each of its occupied orbitals Example - The orbital diagram of oxygen: O: 1s 2 2s 2 2p 4 1s 2s 2p

65 Section 2.11 The Aufbau Principle and the Periodic Table

66 Section 2.11 The Aufbau Principle and the Periodic Table Hund s Rule The rule states that the lowest energy configuration for an atom is the one having the maximum number of unpaired electrons allowed by the Pauli principle in a particular set of degenerate orbitals Unpaired electrons are represented as having parallel spins Example - The orbital diagram for neon: Ne: 1s 2 2s 2 2p 6 1s 2s 2p

67 Section 2.11 The Aufbau Principle and the Periodic Table Valence Electrons They are electrons in the outermost principal quantum level of an atom Example - For the sodium atom, the valence electron is that in the 3s orbital Inner electrons are termed core electrons The elements in the same group on the periodic table have the same valence electron configuration and display similar chemical behavior

68 Section 2.11 The Aufbau Principle and the Periodic Table

69 Section 2.11 The Aufbau Principle and the Periodic Table Groups in the Periodic Table Transition metals are those whose configuration is obtained by adding electrons to the five 3d orbitals The configuration for chromium is: Cr: [Ar]4s 1 3d 5 After lanthanum, the lanthanide series occurs, which corresponds to the filling of the seven 4f orbitals The actinide series corresponds to the filling of the seven 5f orbitals

70 Section 2.11 The Aufbau Principle and the Periodic Table Main-Group or Representative Elements The labels for Groups 1A, 2A, 3A, 4A, 5A, 6A, 7A, and 8A indicate the total number of valence electrons Each member of these groups has the same valence electron configuration

71 Section 2.11 The Aufbau Principle and the Periodic Table Figure The Orbitals Being Filled for Elements in Various Parts of the Periodic Table

72 Section 2.11 The Aufbau Principle and the Periodic Table Interactive Example Electron Configurations Give the electron configurations for sulfur (S) and cadmium (Cd) Solution Sulfur is element 16 and resides in Period 3, where the 3p orbitals are being filled Since sulfur is the fourth among the 3p elements, it must have four 3p electrons Its configuration is: S: 1s 2 2s 2 2p 6 3s 2 3p 4 or [Ne]3s 2 3p 4

73 Section 2.11 The Aufbau Principle and the Periodic Table

74 Section 2.11 The Aufbau Principle and the Periodic Table Interactive Example Electron Configurations Cadmium is element 48 and is located in Period 5 at the end of the 4d transition metals It is the tenth element in the series and thus has 10 electrons in the 4d orbitals, in addition to the 2 electrons in the 5s orbital The configuration is: Cd: 1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 3d 10 4p 6 5s 2 4d 10 or [Kr]5s 2 4d 10

75 Section 2.12 Periodic Trends in Atomic Properties Commonly Observed Periodic Trends Atomic radius Ionization energy Electron affinity

76 Section 2.12 Periodic Trends in Atomic Properties Atomic Radius Atomic radii can be obtained by measuring the distance between atoms in chemical compounds They are also called covalent atomic radii due to the manner in which they are determined Since nonmetallic atoms do not form diatomic molecules, atomic radii are estimated from their covalent compounds Metallic radii are obtained by calculating half the distance between metal atoms in solid metal crystals

77 Section 2.12 Periodic Trends in Atomic Properties Atomic Radii Trends The atomic radius decreases in going across a period from left to right Effective nuclear charge increases Valence electrons are drawn closer to the nucleus, decreasing the size of the atom Atomic radius increases in going down a group Orbital sizes increase in successive principal quantum levels

78 Section 2.12 Periodic Trends in Atomic Properties Interactive Example Trends in Radii Predict the trend in radius for the following ions: Be 2+, Mg 2+, Ca 2+, and Sr 2+ Solution All these ions are formed by removing two electrons from an atom of a Group 2A element In going from beryllium to strontium, we are going down the group, so the sizes increase: Be 2+ < Mg 2+ < Ca 2+ < Sr 2+ Smallest radius Largest radius

79 Section 2.12 Periodic Trends in Atomic Properties

80 Section 2.12 Periodic Trends in Atomic Properties Ionization Energy It refers to the energy required to remove an electron from a gaseous atom or ion The atom or ion is assumed to be in its ground state + X( g) X ( g) + e The energy required to remove the highest-energy electron of an atom is called the first ionization energy (I 1 ) The value of I 1 is generally smaller than that of I 2, which is the second ionization energy

81 Section 2.12 Periodic Trends in Atomic Properties Ionization Energy Trends in the Periodic Table While going across a period from left to right, the first ionization energy increases Electrons added to the same principal quantum level cannot completely shield the increasing nuclear charge and are generally more strongly bound from left to right on the periodic table While going down a group from top to bottom, the first ionization energy decreases The electrons being removed are farther from the nucleus

82 Section 2.12 Periodic Trends in Atomic Properties Figure The Values of First Ionization Energy for the Elements in the First Six Periods

83 Section 2.12 Periodic Trends in Atomic Properties Interactive Example Ionization Energies Consider atoms with the following electron configurations: 1s 2 2s 2 2p 6 1s 2 2s 2 2p 6 3s 1 1s 2 2s 2 2p 6 3s 2 Which atom has the largest first ionization energy, and which one has the smallest second ionization energy? Explain your choices.

84 Section 2.12 Periodic Trends in Atomic Properties Interactive Example Ionization Energies Solution The atom with the largest value of l 1 is the one with the configuration 1s 2 2s 2 2p 6 (this is the neon atom), because this element is found at the right end of Period 2 Since the 2p electrons do not shield each other very effectively, l 1 will be relatively large The other configurations given include 3s electrons, which are effectively shielded by the core electrons and are farther from the nucleus than the 2p electrons in neon Thus l 1 for these atoms will be smaller than for neon

85 Section 2.12 Periodic Trends in Atomic Properties Interactive Example Ionization Energies The atom with the smallest value of l 2 is the one with the configuration 1s 2 2s 2 2p 6 3s 2 (the magnesium atom) For magnesium, both l 1 and l 2 involve valence electrons For the atom with the configuration 1s 2 2s 2 2p 6 3s 1 (sodium), the second electron lost (corresponding to l 2 ) is a core electron (from a 2p orbital)

86 Section 2.12 Periodic Trends in Atomic Properties Electron Affinity It refers to the energy change associated with the addition of an electron to a gaseous atom X (g + e g ) X ( ) While going across a period from left to right, electron affinities become more negative Electron affinity becomes more positive in going down a group Electrons are added at increasing distances from the nucleus

87 Section 2.12 Periodic Trends in Atomic Properties Concept Check Explain why the graph of ionization energy versus atomic number (across a row) is not linear. Where are the exceptions?

88 Section 2.12 Periodic Trends in Atomic Properties Concept Check Which of the following would require more energy to remove an electron? Why? Na Cl

89 Section 2.12 Periodic Trends in Atomic Properties

90 Section 2.12 Periodic Trends in Atomic Properties Concept Check Which of the following would require more energy to remove an electron? Why? Li Cs

91 Section 2.12 Periodic Trends in Atomic Properties Concept Check Which element has the larger second ionization energy? Why? Lithium Beryllium

92 Section 2.12 Periodic Trends in Atomic Properties

93 Section 2.12 Periodic Trends in Atomic Properties Concept Check Which of the following should be the larger atom? Why? Na Cl

94 Section 2.12 Periodic Trends in Atomic Properties Concept Check Which of the following should be the larger atom? Why? Li Cs

95 Section 2.12 Periodic Trends in Atomic Properties

96 Section 2.12 Periodic Trends in Atomic Properties Concept Check Which is larger? The hydrogen 1s orbital The lithium 1s orbital Which is lower in energy? The hydrogen 1s orbital The lithium 1s orbital

97 Section 2.13 The Properties of a Group: The Alkali Metals Information in the Periodic Table The number and type of valence electrons primarily determine the chemistry of an atom Electron configurations can be determined from the organization of the periodic table Certain groups in the periodic table have special names Elements in the periodic table are basically divided into metals and nonmetals Elements that exhibit both metallic and nonmetallic properties are termed metalloids of semimetals

98 Section 2.13 The Properties of a Group: The Alkali Metals Figure Special Names for Groups in the Periodic Table

99 Section 2.13 The Properties of a Group: The Alkali Metals

100 Section 2.13 The Properties of a Group: The Alkali Metals Figure Special Names for Groups in the Periodic Table (Contd.)

101 Section 2.13 The Properties of a Group: The Alkali Metals The Alkali Metals Li, Na, K, Rb, Cs, and Fr are the most chemically reactive of the metals Hydrogen exhibits a nonmetallic character due to its small size Going down a group: Ionization energy decreases Atomic radius increases Density increases Melting and boiling points decreases smoothly