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Zumdahl • Zumdahl • DeCoste
World of CHEMISTRY
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Chapter 11 Modern Atomic Theory
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Chapter 11 Overview Describe Rutherford’s model of atom
Electromagnetic radiation See how atoms emit light Quantized nature of energy demonstrated by emission spectrum of hydrogen Bohr’s model of hydrogen atom Wave mechanical model of electron position Shapes of s, p, and d orbitals Electron spin Electrons filling principle energy levels Valence and core electrons Electron configurations Periodic table trends Copyright © Houghton Mifflin Company
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Figure 11.1: The Rutherford atom.
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Rutherford Atom Review
Alpha particle/Gold foil experiment Nuclear Atom Nucleus composed of protons & neutrons Nucleus small compared to atomic size Electrons account for rest of atom Unanswered questions: What are electrons doing? – How are they arranged & how do they move? Thought electrons revolved around nucleus like planets orbit the sun Couldn’t explain why electrons aren’t attracted to protons causing atom to collapse Copyright © Houghton Mifflin Company
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Electromagnetic Radiation
Energy and Light Electromagnetic Radiation Transmits energy Heat from light bulb Solar energy (energy from sun) Warmth from fireplace Many kinds: X-rays, microwaves, etc. Differ in their wave characteristics Copyright © Houghton Mifflin Company
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Figure 11.2: A seagull floating on the ocean moves up and down as waves pass.
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Wavelength (λ): distance between two consecutive wave peaks
Wave Properties Wavelength (λ): distance between two consecutive wave peaks Frequency (ν): how many waves pass a certain point per given time period Speed: how fast a given peak travels Copyright © Houghton Mifflin Company
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Figure 11.3: The wavelength of a wave.
Crest Trough Copyright © Houghton Mifflin Company
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Electromagnetic Radiation
Travel as waves Have different wavelengths See page 325 Gamma rays – shortest, radio waves – longest Important means of energy transfer Solar energy – visible & ultraviolet radiation Heat from fireplace – infrared radiation “Light” – wave that carries energy through space Copyright © Houghton Mifflin Company
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Electromagnetic radiation has particle characteristics
Photons: tiny packets of energy that travel in a stream Wave-particle nature of light: consists of both waves and particles of energy Copyright © Houghton Mifflin Company
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Figure 11.5: Electromagnetic radiation.
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Figure 11.6: Photons of red and blue light.
Different wavelengths of electromagnetic radiation carry different amounts of energy In general – the longer the wavelength, the lower the energy of the photons (red less energy than blue) Copyright © Houghton Mifflin Company
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Emission of Energy by Atoms
Recall Flame Test Laboratory (different elements gave off different colors) Color resulted from atoms releasing energy by emitting visible light of specific wavelengths (specific colors) Atoms became excited: absorbed heat energy from flame Some of excess energy released as light – carried away by photon Energy of photon = energy change of atom Short wavelength = high-energy photons Long wavelength = low-energy photons Copyright © Houghton Mifflin Company
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Figure 11.8: An excited lithium atom emitting a photon of red light to drop to a lower energy state.
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The Energy Levels of Hydrogen
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Figure 11.9: A sample of H atoms receives energy from an external source.
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Figure 11.9: The excited atoms release energy by emitting photons.
Excited atom can release some or all of its energy by emitting a photon (electromagnetic radiation “particle”) Copyright © Houghton Mifflin Company
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Figure 11.10: An excited H atom returns to a lower energy level.
Energy contained in photon = change in energy of atom Copyright © Houghton Mifflin Company
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Figure 11.11: Colors and wavelengths of photons in the visible region.
Visible light photons emitted by Hydrogen – always the same Because only certain photons are emitted, only certain energy changes are occurring Hydrogen atom has certain discrete energy levels Energy levels of Hydrogen are quantized – only certain values allowed Copyright © Houghton Mifflin Company
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Figure 11.12: The color of the photon emitted depends on the energy change that produces it.
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Figure 11.13: Each photon emitted corresponds to a particular energy change.
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Figure 11.14: Continuous (a) and discrete (b) energy levels.
Quantized nature of energy surprised scientists (b) Previously assumed atom could exist at any energy level (a) Copyright © Houghton Mifflin Company
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Figure 11.15: The difference between continuous (a) and quantized (b) energy levels.
Ramp – can be at any elevation Staircase – can move from one step to another or even skip, but must be on a step Copyright © Houghton Mifflin Company
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Figure 11.17: The Bohr model of the hydrogen atom.
Electrons moved in circular orbits like planets Electrons could jump from one orbit to another by emitting/absorbing a photon Didn’t work for other atoms Showed experimentally to be incorrect Paved way for other theories Copyright © Houghton Mifflin Company
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We do not know exactly how the electrons move in an atom!
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The Wave Mechanical Model of the Atom
Louis Victor de Broglie & Erwin Schrödinger: since light has both wave and particle characteristics, an electron might also exhibit these characteristics New hydrogen model applied to other atoms (Bohr’s did not) Electron states are described by orbitals (which are nothing like orbits) Copyright © Houghton Mifflin Company
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Figure 11.18: A representation of the photo of the firefly experiment (lightning bugs).
Shows probability (or likelihood) of where firefly will be found Usually near the center, but can be found in any of the shaded areas at any time Copyright © Houghton Mifflin Company
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Figure 11.19: The orbital that describes the hydrogen electron in its lowest possible energy state. Darker pink = greater probability Copyright © Houghton Mifflin Company
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Drawbacks of wave mechanical model:
Gives no information about when the electron occupies a certain point in space or how it moves We will probably never know the details of electron motion Confident that Bohr model is incorrect Copyright © Houghton Mifflin Company
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Figure 11. 20: The hydrogen 1s orbital
Figure 11.20: The hydrogen 1s orbital. (Lowest Energy State or Ground State) 1s orbital – spherical representation Probability map (more accurate) Copyright © Houghton Mifflin Company
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The Hydrogen Orbitals Size defined as the sphere that contains 90% of the total electron probability Spends 90% of its time somewhere within the sphere Spends 10% of its time somewhere outside of the sphere Electron can absorb energy & move to higher energy state Bohr model – orbit with larger radius Wave mechanical model – different kinds of orbitals with different shapes Copyright © Houghton Mifflin Company
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Figure 11.21: The first four principle energy levels in the hydrogen atom.
Further from nucleus Closer to nucleus Copyright © Houghton Mifflin Company
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Figure 11.22: How principal levels can be divided into sublevels.
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Figure 11.23: Principal level 2 shown divided into the 2s and 2p sublevels.
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Figure 11.24: The relative sizes of the 1s and 2s orbitals of hydrogen.
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Figure 11.25: The three 2p orbitals.
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Figure 11.26: Diagram of principal energy levels 1 and 2.
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Number tells principal energy level Letter indicates shape
Orbital Labels Number tells principal energy level Letter indicates shape s = spherical p = two-lobed (x, y, & z indicates axis) 2s orbital Principle energy level 2 Spherical shape Copyright © Houghton Mifflin Company
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Hydrogen Orbitals Why does hydrogen have more than one orbital if it only has 1 electron? Orbital is potential space for an electron Can only occupy 1 orbital at a time, but can be transferred to another by adding energy Copyright © Houghton Mifflin Company
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Figure 11.27: Relative sizes of the spherical 1s, 2s, and 3s orbitals of hydrogen.
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Orbital summary As principle energy level increases, the number of sublevels increases n = 1 – 1 sublevel n = 2 – 2 sublevels n = 3 – 3 sublevels , etc. Further from nucleus = more space = more room for orbitals Sublevels are s (1 orbital), p (3 orbitals), d (5 orbitals), and f (7 orbitals) Orbitals keep same shape, but get larger as n increases Copyright © Houghton Mifflin Company
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Figure 11.28: The shapes and labels of the five 3d orbitals.
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Further Development of the Wave Mechanical Model
Applies to all atoms Helps explain the periodic table Electrons spin like a top – can only spin in one of two directions Use arrows to represent spin (↑or↓) Electrons must have opposite spins to occupy the same orbital Pauli exclusion principle: an atomic orbital can hold a maximum of two electrons and those two electrons must have opposite spins Copyright © Houghton Mifflin Company
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Principle components of wave mechanical model
Atoms have principal energy levels (n) Energy of level increases as n increases Each principal energy level contains one or more types of orbitals, called sublevels Number of sublevels present = n (p. 338) Label with n and shape (ex.: 3p) Orbital can have 0 to 2 electrons, 2 electrons in same orbital must have opposite spins Shape of orbital indicates probabilities, not electron movement Copyright © Houghton Mifflin Company
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Electron Arrangements of First 18 Elements
Electrons will occupy orbitals closest to nucleus first As n increases, orbital becomes larger – electron is further from nucleus Electron configuration = electron arrangement (Example: 1s22s1) Abbreviate Na: 1s22s22p63s1 = [Ne]3s1 Orbital Diagram = Box Diagram: orbitals are represented by boxes grouped by sublevel with arrows indicating electrons Copyright © Houghton Mifflin Company
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Electron Configurations and Orbital Diagrams
Atom Configuration Diagram Hydrogen 1s1 1s Helium 1s2 Beryllium 1s22s2 1s 2s Carbon 1s22s22p2 1s 2s p Oxygen 1s22s22p4 ↑ ↑↓ ↑↓ ↑↓ ↑↓ ↑↓ ↑ ↑ ↑↓ ↑↓ ↑↓ ↑ ↑ Copyright © Houghton Mifflin Company
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Valence Electrons Electrons in the outermost (highest) principal energy level of an atom Most important electrons to chemists – electrons involved when atoms form bonds (attach to each other) Atoms in same group on periodic table have same number of valence electrons in outer orbital (orbitals are at different principal energy levels) Core electrons Inner electrons Not involved in bonding Copyright © Houghton Mifflin Company
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Figure 11.30: Partial electron configurations for the elements potassium through krypton.
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Figure 11.31: Orbitals being filled for elements in various parts of the periodic table.
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After lanthanum – lanthanide series – fill 4f orbitals
Orbital Filling If energy level has d orbitals, s orbitals from next level will fill first After lanthanum – lanthanide series – fill 4f orbitals After actinium – actinide series – fill 5f orbitals Except Helium – group number indicates sum of electrons in outer s & p orbitals (number of valence electrons) Copyright © Houghton Mifflin Company
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Figure 11.34: Periodic table with atomic symbols, atomic numbers, and partial electron configurations. Copyright © Houghton Mifflin Company
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Atomic Properties & the Periodic Table
Chemistry is fundamentally based on observed properties of substances Atomic theory is attempt to help us understand why these things occur Theories may change Copyright © Houghton Mifflin Company
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Figure 11.35: Classification of elements as metals, nonmetals, and metalloids.
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Metals, Nonmetals, & Metalloids
Lustrous appearance, change shape without breaking (pulled into wire), excellent conductors Tend to lose electrons to form positive ions Nonmetals Lack properties of metals, some exceptions Tend to gain electrons to form negative ions Metalloids Have properties of metals and nonmetals Along stair step Copyright © Houghton Mifflin Company
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Not all metals/nonmetals behave exactly the same way
As you go down group – more likely to lose electrons (further from nucleus) Most chemically active – lower left corner of periodic table Nonmetals Most chemically active in upper right corner (not noble gases) Strongest attraction (closer to nucleus) Copyright © Houghton Mifflin Company
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Figure 11.36: Relative atomic sizes for selected atoms.
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Increases as you go down group: more electrons = larger atom
Atomic Size Increases as you go down group: more electrons = larger atom Decreases as you go across period: more protons in nucleus – greater pull on electron Copyright © Houghton Mifflin Company
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Ionization Energy The energy required to remove an electron from an individual atom in the gas phase Metals relatively low – easily lose electrons (small amount of energy needed) Nonmetals relatively large – prefer to gain electron, not lose Decreases going down a group, increases going across a period Bottom left – lowest (most chemically active) Upper right – highest (most chemically active) Copyright © Houghton Mifflin Company
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