About the Atom

• in 1774 Antione Lavoisier discovered the law of conservation of matter”

• in 1779 Jospeh Proust discovered the law of constant composition

• Both of these laws explain that no matter can be created or destroyed

• These laws contributed to John Dalton’s Atomic Theory

  • “All matter is composed of tiny, indivisible particles, called atoms, that cannot be destroyed or created. Each element has atoms that are identical to each other in all of their pro­perties, and these properties are different from the properties of all other atoms. Chemical reactions are simple rearrangements of atoms from one combination to another in small whole-number ratios”

    • Excerpt From AP Chemistry Premium, 2022-2023 Neil D. Jespersen

  • All scientific theories have new ideas and predictions that can and should be tested by experimentation

  • The results of these experiments will prove or disprove the theory

  • No theory can be “true” it can only be “supported”

• “Law of Multiple Proportions” was John Dalton’s idea and it stemmed from the atomic theory

• Dalton then used the “Law of Multiple Proportions” to test the atomic theory

Atomic Models

• Timeline of important models

  • 400 BC- solid particle model

  • 1909- Plum Pudding model

  • 1910- Nuclear model Rutherford

  • 1913- solar system model Bohr

  • 1927- wave-mechanical Schrödinger

• Light and Matter became an interest to scientists in mid 1800’s

  • Each elementary when sparked or heated shows colors related to their wavelength

    • Not all the rainbow

  • 1885- Johannan Balmer found mathematical relation between the spectrums and the wavelength of each element

  • Model from the book pg 148

• 1913 Niels Bohr finished the theory on the construction of the hydrogen atom

  • Used a solar system type model

  • He assumed that electrons moved orbitally around the nucleus

  • Most important: electrons could only be in certain orbits

• 1924 Louis de Broglie said that light could be considered as a particle

• 1927 Erwin Schrödinger used the equation for waves to the electrons inside of the atom

  • Wave-mechanical theory

• 1920’s Werner Heisenburg created the uncertainty principle

  • States- momentum and position or any one particle can’t be known at the exact same time as another

    • As you know one’s location you can’t know the others as well

Atomic Structure

• Light and the Atom

  • An atom usually stays in the lowest energy level/ state

    • Ground state

  • Atoms not in their lowest energy state is called an excited state

  • When the atom drops from the excited state the atom can emit light

• The Electromagnetic Spectrum

  • Visible light is 1 part of the electromagnetic spectrum

  • Microwaves have wavelengths between 10^-1 and 10^-4

  • INSERT FIGURE 1.1

• Wavelength, Frequency, and Energy of Light

  • Electromagnetic radiation is all considered as a wave

    • Defined by wavelength  λ

    • Frequency (v)

  • Wavelength- distance between 2 points that repeat on a sine wave

  • Frequency- number of waves that pass a single point in a single second

  • INSERT FIGURE 1.2

  • Wavelength and frequency and proportional

  • INSERT FIGURE 1.2 part 2

  • Speed of light- 3.0• 10^8 m/s

  • Wavelength units- meters

  • Frequency units- reciprocal seconds (s^-1) or hertz (Hz)

  • Max Planck discovered that frequency and energy of electromagnetic are proportional and inversely proportional to wavelength

  • h which is proportionally constant is equal to 6.63• 10^ –34 joule second

  • INSERT FIGURE ON PAGE 153

  • Frequency, wavelength, and energy are all related to each other

    • If you know the speed of light and Planck’s constant then only one of the 3 variables are needed to figure out the others (mathematically)

• EXAMPLE 1.1 from the book

  • “What are the frequency and the energy of blue light that has a wavelength of 400. nm? (Planck’s constant = 6.63 × 10–34 J s)”

    • 400. nm(v)= 3.0• 10^8 m/s^-1

    • 400•10^-9m(v)=3.0*10^8 m/ s^-1

    • v=(3.0•10^8 ms^-1)/(400.*10^-9m)=7.5•10^14 s^-1

    E=hv=(6.63•10^-34 Js)(7.7• 10^14 s^-1)= 49.725• 10^-20 J= 5.0 • 10^-19 J

The Bohr Model of the Atom

• Bohn rocked the the world with his solar system like model

• The model said that electrons were confined to certain orbitals

• The equation in question:

  • En=-22me4n2 h2=-2.178•10-18n2joule

  • m=mass of electron

  • e= charge on the electron

  • h=Planck’s constant

  • n=orbit

    • Later became principal quantum number

• INSERT FIGURE 1.3

• In the figure the numbers are the orbitals in ascending order from the nucleus

• Light is energy that is emitted from an atom when the electrons move down from the atom (in figure 1.4)

• The same amount of energy is needed to go up an orbital as is released when it comes back down

• Amount of energy emitted can be calculated by subtracting the lower energy level from the higher energy level

• INSERT FIGURE 1.4

  • Insert the text as well

• INSERT FIGURE 1.5

• EXAMPLE 1.3- from figure 1.5

  • “Determine the energy and wavelength of light associated with an electron moving from the second to the fourth energy level in a hydrogen atom”

The Size Of The Atom

• Bohr suggested that the momentum (velocity•mass) of the electron had to be related to the size of the orbit of the electron

  • The equation used

    • mv=nh2r

      • m= electron mass

      • h=Planck’s constant

      • v= velocity

• 1st energy level is radius=53 picometers

• 2nd energy level is radius=106 picometers

• The radii of all the other orbital are multiples of Bohr radius

• Bohr radius presented chemists with a theoretical value for the size of a hydrogen atom

  • It also confirmed that the atomic size that had been found by experimentation were generally correct

The Wave-Mechanical Model of the Atom

• Louis de Broglie proposed that an electron could be like a wave and also be like a particle

• His argument involved the equation for the energy of the wave and Einstein's equation for energy

  • They both had to be equal to each in the case of the energy of light other since electrons can only have one energy at a time

    • This does not mean that electrons always are equal to e=mc2

• These two equations proved Broglie's theory

• We use this equation for particle that do not move at the speed of light

  • h=mv

• Wave equations:

1. All require 3 numbers and they’re called quantum numbers

   1. n → principal quantum number

   2. l → azimuthal quantum number

   3. ml → magnetic quantum

   4. If you want to specify an electron you  need a fourth number

      1. ms→ spin quantum number

2. Change the way we view the atom

   1. Electrons aren’t fix in position or orbit neatly

   2. It's more of a space they occupy

3. Circular orbits became spherical electron clouds

   1. The clouds are more complicated but still geometrical shapes

4. The arrangement of electrons is deduced from the wave equations and it works with the periodic table

   1. The physical elements are more understood because of our prior knowledge

5. The wave equations agree with the Bohr model except for the cloud vs orbit

   1. The energy change matches almost perfectly

6. Heisenberg uncertainty principle is a key part of the wave model

   1. Heisenberg uncertainty principle: both the position and momentum of an electron can’t be known at the same time

   2. The more you know position then the less you know momentum and vise versa

   3. INSERT FIGURE 1.6

Heisenberg Uncertainty Principle

• “The uncertainty in the position times the uncertainty of the momentum is greater than or equal to Planck’s constant divided by four pi.”Excerpt From AP Chemistry Premium, 2022-2023 Neil D. Jespersen

Structure of The Atom

• Principal Energy Levels (shells)

  • Current model of the atom is that the positive charged nucleus is encompassed by energy levels

  • The energy level closest to the nucleus is called the number 1 energy level

    • Each energy level after ascends in numerical order

  • The biggest electron only has 7 energy levels

  • Number of energy levels has the symbol ‘n’

  • As the energy levels increase the size also increases

    • Each holds a max number of electrons

      • This is equal to 2n2

      • n is energy level

    • 1st level: 2

    • 2nd level: 8

    • 3rd level: 18

    • 4th level: 32

    • The last three could have 50,72, and 98 respectively but they aren’t fully filled

• Sublevels (subshells)

  • Each energy level has one or more sublevels

  • Number of sublevels possible is equal to the value of n for that particular energy level

  • Ex. third principal energy level (n=3) may have up to 3 sublevels

  • Only 4 sublevels are actually used

    • 5,6,7 are possible theoretically but not needed with our current Periodic Table

  • Numbered consecutively starting with 0

    • Numbers are azimuthal quantum numbers →l

    • Value of l can't be more than (n-1)

  • Sublevels are given the letters for each level:

    • 1-s

    • 2-p

    • 3-d

    • 4-f

Orbitals

• Each sublevel an have 1+ orbitals

• Orbital is a space that has a bunch of electrons

• Each orbital can have up to 2 electrons

• They must have opposite spins

• After the electrons start sharing an orbital they are ‘paired’

• Orbitals are categorized as s,p,d, or f depending on which sublevel that are in

• This shows why the energy levels have 2, 8, 18, and 32 electrons

• 1st sublevel only has one s orbital so it only can have 2 electrons

• Each orbital has a number and its called magnetic quantum number,ml

• Possible values of ml can range from -l to +l

  • Can be zero

• The sublevels tell the chemist the orbital shape

• In the s sublevels the shape is spherical

• INSERT FIGURE 1.7

• P orbitals are shaped like a dumbbell

  • This is because the electron density is greeted us the two lobes on both sides of the nucleus

  • 3 p orbitals in each p sublevel

    • Designated as px,py,and pz

  • 5 d orbitals in the d sublevel have these shapes

  • INSERT FIGURE 1.9

  • 7 f orbitals are a more complex shape than the d orbital

• Electronic Structure of the Atom

  • The arrangement of the electrons are based on the energy of each orbital

  • Electrons fill the orbitals with the lowest energy first

    • Then they fill from lowest to highest

  • The sequence goes like this

    • “1s 2s 2p 3s 3p 4s 3d 4p 5s 4d 5p 6s 4f 5d 6p 7s 5f 6d”

    • It's called aufbau principle

    • You can figure it out from the periodic table

    • INSERT FIGURE 1.10

  • Periods fill like this:

    • First periods- 1s electrons

    • Second period- 2s on the left and 2p on the right

    • Third period- 3s then 3p

    • Fourth period- 4s, 3d then 4p

    • Fifth period- 5s, 4d, and 5p

    • Sixth period- 6s, 4f, 5d, 6p

    • Seventh period- 7s, 5f, 6d

Electronic Configuations

• When you know the energy order of the orbitals then you can describe the electron configuration

• 2 forms of configuration lists the info of the electrons

1. Complete electron configuration

   1. Every single electron

2. Abbreviated version

   1. Only lists highest energy level

• Complete Electronic Configurations, nlx

  • Complete electronic configurations electrons are listed by: designating n by number, l by level and x in each sublevel

  • Atoms are made by adding electrons to the lowest sublevel

  • Electrons must fill the lower sublevels before filling the higher ones

• Important Exceptions to Aufbau Ordering

  • The chart shows aufbau order

  • Some elements do not follow aufbau

    • Only important ones:

      • Copper

      • Silver

      • Gols

      • Chromium

      • Molybdenum

    • They do not fill the d or s sublevel (regularly)

    • It has to do with stability

• Abbreviated Electronic Configurations

  • The first reaction in the atom is at the highest energy level

  • The inner levels have very little role in chemical reactions

  • Inner electrons are mostly represented by the noble gas at the end of the period prior to the actual element

    • Fe= 1s22s22p63s23p64s23d6

    • Fe= [Ar]4s23d6

• Valence Electrons

  • Chemists care more about the highest electrons in the atoms

  • Only the valence electrons truly matter

    • s&p really only matter

• How To Count Valence Electrons

  • To find the amount of valence electrons count the groups from the left of the periodic table to the element you want to find

  • Valence electrons are shown in figure 1.11

• Hund’s Rule

  • Second valence electrons are represented as a pair

    • Indicate the s orbital

  • 3rd through 8th group fill s orbital

  • Electrons only pair up if every orbital is already filled

• Orbital Diagrams

  • Details in electronic configurations and valence electrons so we have Orbital Diagrams

  • Show each of the orbitals in the valence of the atom

    • Either a box or circle

    • Arrows represent electrons

      • Opposing directions= opposite spins

  • INSERT FIGURE 1.12

  • Describe valence since all the inner electrons should be paired

  • Sometimes these diagrams also show the d sublevel

  • Electrons must fill all the orbitals in the sublevel before pairing up

  • INSERT 1.13 and 1.14

• Quantum Numbers

  • 4 types:

1. Principal quantum number (n)

   1. can have any value starting at 1

   2. Represents principal energy level of the atom that the electron belongs to

   3. Related to avg distance of the electron from the nucleus

2. Azimuthal quantum number (l)

   1. Any number from 0 to 1 less than the value of n

   2. Designates sublevel of the electron

   3. Represents shape of the orbital in sublevel

3. Magnetic quantum number (ml)

   1. Any integer including zero from -l to +l

   2. Designates orientation of an orbital in space

4. Spin quantum number (ms)

   1. Either +12 or -12

   2. Represents spin of the electron

   3. For electrons to pair one must have +12 and the other -12

   4. Don't need this for wave equations

   5. Do need for pauli exclusion principle

• Significance of the Quantum Numbers

  • Simplified forms of the quantum numbers

1. Principal quantum numbers-n

   1. Average distance of the electron from nucleus

   2. Or size of principal energy level

2. Azimuthal quantum number-l

   1. Shape of the orbital within the sublevel

3. Magnetic number- ml

   1. Orientation of each orbital

      1. In space

4. Spin quantum number- ms

   1. Spin of electron