Chemistry I Electrons in Atoms Chapter 5

Rutherford’s nuclear model did not provide enough detail about how electrons occupy the space around the nucleus. In this chapter we will learn how electrons are arranged around the nucleus and how that arrangement effects chemical behavior.

In what ways did many scientists find Rutherford’s model to be incomplete?

1.)In what ways did many scientists find Rutherford’s model to be incomplete?  It did not explain: 1.How the electrons were arranged around the nucleus 2.Why they were not pulled into the nucleus 3. The differences in chemical behavior of the different elements.

2.)In the early 1900’s what did scientists observe about certain elements when heated in a flame?  Every element gives off specific wavelengths of light that can be used like a fingerprint to identify the element.

spectroscope

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An element’s chemical behavior is related to the arrangement of electrons in its atoms.

Wave Nature of Light

What is electromagnetic radiation?

3.)What is electromagnetic radiation?  A type of wave that is part electrical and magnetic energy acting at right angles.  Visible light is a type of electromagnetic radiation.

What are the 4 characteristics of waves?

4.) What are the 4 characteristics of waves?  1.) wavelength – distance from crest to crest  2.) frequency- # of waves that pass a given point per second  3.) Amplitude – height of wave from the normal resting to wave crest  4.) speed – the speed of light is a constant 3.0 x 10 8 m/s.

5.)What unit do we use to express frequency?  Hertz (Hz)  The hertz unit is 1/second ( the inverse of a second)

6.) What is the speed of light?  The speed of light is a constant 3.0 x 10 8 m/s  The formula for the speed of light is C = λ v λ = wavelength (measured in meters) v = frequency ( measured in 1/sec) The unit for speed is m/s (distance/time)

#6 continued  When wavelength (λ) increases then frequency (v) decreases because the speed of light ( c) is a constant 3.0X10 8 m/s.  Also  When λ decreases, V increases

7.) List types of electromagnetic radiation from the lowest to the highest in energy ( page 139)

List types of electromagnetic radiation from the lowest to the highest in energy ( page 140)  Radio waves ↓decreasing wavelength  Microwaves ↓increasing frequency  Infrared  Visible  Ultraviolet (U.V.)  X rays  gamma

Particle Nature of Light  When an object is heated only certain wavelengths of light are emitted. The instrument used to separate light into its different wavelengths is called a spectroscope. The pattern of wavelengths is called the spectrum.

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 When a substance is heated the spectrum of light that is given off can be used to identify the substance.

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 The wave model of light could not explain why different substances emit particular wavelengths of light when heated

8.)What are some problems with the wave model of the atom?  It doesn’t explain why only certain wavelengths of light are emitted when an object is heated.  It also does not explain the photoelectric effect.( we will discuss the photoelectric effect in just a little bit)

Who is Max Planck?

9.)Who is Max Planck?  A German physicist who, in 1900, began searching for the reason that only certain wavelengths of light are given off for every element.

10.)What did Planck conclude?  Matter can gain or lose energy in only small specific amounts called quanta.

As metal is heated, it glows. In fact, it glows different colors as it gets hotter (white hot being the hottest). By studying glowing metals, Max Planck discovered that only certain wavelengths of light are emitted at each specific temperature.

11.) What is a quantum?  The minimum amount of energy that can be gained or lost by an atom.  The amount of energy it takes for a electron to get from one energy level to the next.

12.) If energy can only absorbed in quanta, specific amounts, why does energy appear to continuous to us?  b/c quanta are extremely small.

13.) What is Planck’s equation that demonstrates that energy is related to the frequency of radiation?  Equation E = hv  E – energy of a quantum  h- plancks constant  v- frequency

14.) What is Planck’s constant?  6.62 x Js

15.) How are energy and frequency related?  They are directly related  As frequency increase, energy increases.

What is the photoelectric effect?

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16.)What is the photoelectric effect?  When light of a certain frequency shines on the surface of a metal, electrons are ejected. Blue light always results in electrons being ejected. Red light of any brightness does not eject electrons.

17.) In 1905 how did Albert Einstein explain the photoelectric effect?  He proposed that all electromagnetic radiation (including visible light) is both wavelike and particle like in nature.  Continued next slide

17.) In 1905 how did Albert Einstein explain the photoelectric effect?  He suggested that while a beam of light has many wavelike characteristics, it is also like a stream of particles called photons.

18.) What is a photon?  A photon is a particle of electromagnetic radiation with no mass that carries a quantum of energy.

19.)How did Einstein use Planck’s idea of quantum energy to explain the photoelectric effect?  Planck proposed that E=hv. Energy is equal to his constant times the frequency of radiation.  Einstein proposed that since blue light has a larger frequency, it has a larger quantum of energy.  Continued next slide

19.) How did Einstein use Planck’s idea of quantum energy to explain the photoelectric effect?  E = hv for blue light is more than E=hv for red light because blue light has a larger frequency than red light.

Practice Problems page 143

20.)What is the atomic emission spectrum?  The atomic emission spectrum of an element is the set of frequencies of (light) electromagnetic waves emitted ( given off) by atoms to that element when it is heated.

21.) How does the atomic emission spectrum for one element compare to another element?  Each is unique and can be used to identify the element like a fingerprint.

22.) Are atomic emission spectrums of elements continuous or distinct individual wavelengths of light?  They are distinct, individual wavelengths of light.

23.) In a line spectrum which line has the highest energy?  The line farthest to the blue end of the spectrum. The line corresponding to the shortest wavelength/ highest frequency.

24.) How does the quantitization of energy ( energy only exists in specific sizes) help explain line spectrum?  The energy of the photon of light that is emitted is tied to a specific frequency/color of light  E = hv Each frequency (v) corresponds to a specific color

Questions page 145

Section 2 Quantum Theory and the Atom

 Scientists concluded that light was both wave and particle. Scientists were able to understand atomic structure, electrons and atomic emission spectra better because of a better understanding of light.

 Scientists wanted to know why atomic emission spectrums are not continuous but are discontinuous.

23.)Who was Niels Bohr?  A Danish physicist who worked in Rutherford’s laboratory. He proposed the quantum model of the atom that helped explain why atomic emission spectrums only contain certain frequencies of light.

25.)What did Niels Bohr propose?  He proposed that a hydrogen atom has only certain allowable energy levels.  The lowest level is called the ground state.  continued

25.)What did Niels Bohr propose?  All other levels are called excited states.  Electrons move around the nucleus in only certain allowed orbits.  Bohr labeled these orbits.  continued

25.)What did Niels Bohr propose?  The orbit closest to the nucleus is labeled n=1 and was lowest in energy.

26.) How did Niels Bohr explain line spectrum with the quantization of energy for electrons. ( How did Bohr explain that certain wavelengths of light are released when an element is heated by proposing that electrons can only be at certain energy levels.)

#26  He said that when an element is at ground state ( the lowest energy level) no energy (light) is released. But when an electron becomes excited, it jumps up to another energy level.

#26  When the electron returns to the lower energy level the energy is released. The energy (light) that is released is equal to the difference between energy levels.

#26  The amount of energy that is released is equal to a particular frequency of light.

27.)Do we still accept Bohr’s model?  Bohr was correct about his idea of quantization of energy and his basic explanation of atomic emission spectrum. However he could only predict the atomic emission spectrum of hydrogen.

27.)Do we still accept Bohr’s model?  His basic model of the model is considered to be incorrect. Today we do not believe that the electrons are in orbits so that we can predict their positions.

The Quantum Mechanical Model of the Atom

28.)Who was Louis de Broglie?  A French physics student that proposed an idea that eventually accounted for the fixed energy levels in Bohr’s model.

29.)What did De Broglie propose?  He proposed that all matter moves in waves. The larger the mass the smaller the wavelength. Only very small objects have noticeable wavelengths.

30.)What was De Broglie’s equation?  λ = h/ m·v  DeBroglie’s equation predicts a large mass will have a small wavelength (λ) only a very small mass will have a noticeable wavelength.

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If electrons move in waves then only electrons with matching wavelengths can be accepted into an energy level and when an electron is released to a lower energy level particular wavelengths of energy are released.

29.)What did Heisenberg propose?  It is impossible to make a measurement of an object without disturbing it. For example, a thermometer changes the temperature of an object.

30.)What is the Heisenberg uncertainty principle?  It is impossible to know precisely both the velocity and position of a particle at the same time.

The Schrodinger Wave Equation

32.) The Austrian physicist Edwin Schrodinger derived an equation that treated the electron as a wave. What was significant about his equation?  It worked equally well for atoms of other elements unlike Bohr’s model which worked only for hydrogen.

33.) What is the name of the model proposed in which electrons are treated as waves?  The quantum mechanical model of the atom.

34.)What is the quantum mechanical model of the atom?  It limits the electrons energy to certain values but makes no attempt to describe the path of the electron around the nucleus

35.) What is an atomic orbital?  In the quantum mechanical model of the atom, the orbital is the region in which we expect to find an electron 90% of the time.  Orbitals are described by electron clouds.

Hydrogen’s Atomic Orbital

36.)What are principal quantum numbers?  Principal quantum numbers represent the relative sizes and energies of the atomic orbitals.  As n increases, the size and energy level increases.

37.) What are sublevels?  Each principle energy level has sublevels. The number of sublevels is equal to the n level.  n = 1 has one sublevel  n = 2 has two sublevels  etc.

38.) What type of sublevels are there?  The sublevels are named s, p, d, and f

Sublevel Shapeenergy level sspherical lowest pdumbell d cloverleaf fcomplex highest

39.)How many electrons does an orbital contain? two

40.) How many orbitals does each sublevel contains? Sublevel orbitals# e - s 1 2 p 3 6 d 5 10 f 7 14

n = 4 n= 3 n= 2 n = 1 spd sp spdf s

Principal Quantum # Sublevel types # orbitals Total # of orbitals Total # electrons 1s112 2spsp sPdsPd spdfspdf

Section 3 Electron Configurations The arrangement of electrons in atoms follows a few very specific rules.

Ground State Electron Configurations

41.)What is an electron configuration?  The arrangement of electrons in an atom

42.) What is ground state electron configuration? The most stable, lowest energy arrangement of electrons.

43.) What are the three rules that govern ground state electron configurations. 1.) Aufbau principle 2.) Hund’s Rule 3.) Pauli Exclusion Principle

44.)What is the Aufbau principle?  Electrons occupy the lowest energy orbital available.  German for building up

45.) In general what is the order of increasing energy among energy levels and sublevels?  In general n=1 is lower than n=2. The order of increasing order of sublevels is s, p, d, f.  Use the diagram

46.) What is the Pauli Exclusion Principle?  A maximum of two electrons can occupy an orbital and they must have opposite spins.  The way to indicate two electrons with opposite spin is ↓↑

47.) What is the Hund’s Rule? Single electrons with the same spin must occupy equal energy orbitals before additional electrons with opposite spin can occupy the same orbital.

48.)What are the two ways that electron configurations can be described?  Orbital diagrams – using boxes  Mg 1s 2s 2p 3s ↓↑

49.) What is noble gas configuration?  Electron configurations that are used to show just the valence electrons. The full inner core orbitals are represented by the noble gas symbol with the lower atomic number and the electrons in the valence shell.  K 1s 2 2s 2 2p 6 3s 2 3p 6 4s 1  K [Ar] 4s 1

50.) Exceptions to predicted configurations.  Cr [Ar]4s 1 3d 5  Cu [Ar]4s 1 3d 10  It is more stable to “borrow” an electron from the s orbital and put it in the d orbital

Pg 160 practice problems

51.) Which electrons determine the chemical properties of an element?  The valence electrons/ the outermost electrons.

52. How does the number of valence electrons compare to the family the element is in?  Group 1 – 1 valence electron  Group 2 – 2 valence electrons  Group 13 – 3 valence electrons  Group 14 – 4 valence electrons  Group 15 – 5 valence electrons  Group valence electrons  Group 17 – 7 valence electrons  Group 18 – 8 valence electrons except for hydrogen which only has two.