Chapter 5 and 6 Notes
Electromagnetic Radiation and Light
Models of the Atom

There were many different models over time
— Dalton-billiard ball model (1803)
— Thompson – plum-pudding model (1897)
— Rutherford – __________ model of the atom (1911)
— Bohr – uses quantized ________ of the atom (1913)
— Quantum Mechanical Model of the Atom (1

Each new model contributed to the model we use today.
Even our current Quantum Mechanical model, does not give us an exact model of how _____________behave.
The Bohr Model of the Atom

Bohr used the simplest element, _____________, for his model

He proposed an electron is found in specific circular paths, or orbits around the nucleus

Each electron orbit was thought to have a fixed _____________level

Lowest level-ground state

Any Higher level-_____________state
The Bohr Model of the Atom cont.

One electron is capable of many _____________excited states (whenever an electron jumps to higher level)

Quantum: specific amount of ______________ an electron can ________ or lose when moving energy levels

You can excite an electron with energy like electricity, ________________, or magnets
Problems with the Bohr Model


OOPS!-Model only works with _______________.

Model did not account for the _________________ behavior of atoms
WRONG: _________________ do not move around the nucleus in circular orbits

STILL VERY HELPFUL!!!
How do Neon Signs Work? They have __________________gases in them.
Explanation


Step 1: an electron ______________ energy and moves to a __________________ energy level
Step 2: electron drops back down to a ________ energy level
 During drop it gives off _______________ called a “photon”

Sometimes this energy is ______________ light (ROYGBIV)

When a photon is emitted, energy is released. We can calculate the energy released using the equation:
____________________
Application: Atomic Emission Spectrum

Used to determine which elements are present in a sample

Used to determine which elements are present in a star (because stars are gases)

Each element has a _________________ spectrum

Only certain _________________ are emitted because the energy released relates to a specific frequency
Spectroscope

A spectroscope is needed to see the atomic emission spectra, which acts similar to a prism, separating different
_________________ of light
1

Electromagnetic Spectrum

Electromagnetic spectrum is the range of all energies emitted from photons acting like _____________.
Electromagnetic Spectrum with Visible Light Spectrum
Light


Behaves like a ______________.
Behaves like a ______________.
Characteristics of a Wave

Wavelength

(lambda) – shortest ____________________ between equivalent points on a continuous wave
[Unit = meters]
 Frequency  (nu) – the ____________________ of waves that pass a given point per second
[Unit = 1/second = s -1 = Hertz (Hz)]

Crest – _____________________ point of a wave

Trough – ____________________point of a wave

Amplitude (a)– height from its origin to its crest (highest point) or trough (lowest point) [Unit = meters]
Wavelength and Frequency
 Wavelength (  ) and frequency (  ) are related
 As wavelength goes up, frequency goes down

As wavelength goes down, frequency goes up

This relationship is
______________proportional
Wavelength and Frequency cont.
 c = 

Speed of light (c) = 3 x 10 8 m/s
Question Time

Calculate the wavelength (

) of yellow light if its frequency (

) is 5.10 x 10 14 Hz.

What is the frequency of radiation with a wavelength (

) of 5.00 x 10 -8 m? What region of the electromagnetic spectrum is this radiation?
How Much Energy Does a Wave Have?

Energy of a wave can be calculated

Use the formula E= h



E= Energy

= frequency
 h = Planck’s constant = 6.626 x 10 -34

Joule is a unit for energy (J)
Joule .
Sec

Energy and frequency are directly proportional, as frequency increases, energy _________________
Question Time

Remember that the energy of a photon is E =h


How much energy does a wave have with a frequency (

) of 2.0 x 10 8 Hz? ( h = 6.626 x 10 -34 J .
s)
2

Visible Light, Frequency, and Energy

Red _________________ wavelength (

), smallest frequency (

)
 Red frequency smallest (  ), least amount of energy (E)
 Violet smallest wavelength (  ), largest _________________ (  )

Violet frequency largest (

), greatest amount of energy (E)
Flame Test

The flame test is a way to determine the _________________ present in a sample

When placed in a flame, each element gives off a ________________ color

Operates the same as neon signs; electrons are excited by _________ and fall back down and give off different colors
Current Model of the Atom
Quantum Mechanical Model of the Atom
• Quantum Mechanical Model is the current description of electrons in atoms.
– It does not describe the electron’s ______________________ around the nucleus
• Quantum Mechanical Model is based on several ideas including:
– Schrodinger wave equation (1926) treats electrons as _______________.
– Heisenberg uncertainty principle (1927) states that it is impossible to know both the ____________________ and
______________________________of a particle at the same time.
Where do electrons “live”?
Principal Energy Levels
1.
Principal energy levels n =1 to _______. (Row # on the periodic table)
• The electron’s principal energy level is based on its location around the nucleus.
• Electrons closer to the nucleus are at a __________________energy level and have lower energy than those farther away from the nucleus
Atomic Orbital
• An __________ ____________is a region of space in which there is a _________ ___________ of finding an electron
– Orbitals ____ _____ necessarily spherical
Energy Sublevels (also called orbitals) and Orbitals
1. Energy sublevels
– assigned letters ______, _______, ________, or f (smart people do fine)
– Energy sublevels correspond to a _______________ where the electron is likely to be found.
2. Orbitals – describes the electron’s ________________________
(maximum of ______ electrons per orbital)
– s sublevel has 1 orbital (2 electrons total) - spherical
– p sublevel has 3 orbitals (6 electrons total) – dumb-bell shaped
– d sublevel has 5 orbitals (10 electrons total )-double dumb-bells
– f sublevel has 7 orbitals (14 electrons total)
Electron Configurations
Energy Levels, Sublevels, and Orbitals
1.
Principal energy levels – n, assigned values _______________ (Like floors in a hotel)
2.
Energy sublevels- s, p, d, f (Type of suite in a hotel) (Orbitals are like the number of rooms in a suite)
1.
s sublevel – 1 orbital
2.
p sublevel – 3 orbitals
3.
d sublevel – 5 orbitals
4.
f sublevel – 7 orbitals
3.
Orbitals – ___________ electrons per orbital (Two people per room)
Electron Configurations
• Electron configuration – the ______________________ of electrons in an atom.
• Example Sodium (Na) – 1s 2 2s 2 2p 6 3s 1
• Three rules determine electron configurations
– the Aufbau Principle,
– the Pauli Exclusion Principle
– Hund’s rule
The Aufbau Principle
• Each electron occupies the ______________________ energy orbital available
3

• Like filling the hotel from the bottom up
Pauli Exclusion Principle
• A maximum of ______________electrons may occupy a single orbital
• Like only two people sharing one hotel room
Hund’s Rule
• If two or more orbitals of _________ energy are available, electrons will occupy them ______________ with the same spin, before filling them in pairs with opposite spins
• A spin is denoted with an up  or down  arrow to fill orbitals
• This is like trying to find your own room in the same suite before having to share a room with someone else
Writing Electron Configurations
• Aufbau diagram for sodium (Na) which has 11 electrons
• Na electron configuration1s 2 2s 2 2p 6 3s 1
Exceptions to Electron Configurations
• Copper and chromium are exceptions to the ___________________ principle.
Element
Copper
Should be
1s 2 2s 2 2p 6 3s 2 3p 6 3d 4 4s 2
1s 2 2s 2 2p 6 3s 2 3p 6 3d 9 4s 2
Actually is
1s 2 2s 2 2p 6 3s 2 3p 6 3d 5 4s 1
1s 2 2s 2 2p 6 3s 2 3p 6 3d 10 4s 1 Chromium
• Some configurations violate the Aufbau Principle because half-filled sublevels are not as stable as filled sublevels, but they are more stable than other configurations
Valence electrons
• Valence electrons are electrons in the ______________________ orbitals.
• For A group elements the ________________________ number corresponds to number of valence electrons.
• Electron-dot structures – Element’s symbol surrounded by ___________representing the valence electrons
Noble Gas Configuration
What are Noble Gases?
• Noble gases are found in group _____________
• The elements are called noble because they are non-reactive and very __________________.
• The do not tend to form compounds
Complete Electron Configuration
• What is the electron configuration for Ne?
• Ne: ________________________
• What is the electron configuration for Mg?
• Mg: _________________________
• What do both electron configurations have in common?
• 1s 2 2s 2 2p 6 = [Ne]
Noble Gas Configuration (Abbreviated Configuration)
• Using neon’s configuration and then adding magnesium’s extra electrons we can get the noble gas configuration.
• Ne: 1s 2 2s 2 2p 6 = [Ne]
• Mg: 1s 2 2s 2 2p 6 3s 2
• Noble gas configuration Mg: __________________________
• Only use noble gases in the brackets.
Which Noble Gas is Used?
• To figure out which noble gas to use find the noble gas that is closest to the element without going over in atomic number
• Which noble gas is closest without going over?
• Rb : ____
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• Cl : ____
• Ra : ____
What About the Other Electrons?
• To know what to write for the other electrons that are not included in the noble gas, understanding the periodic table is important.
• The periodic table is organized by blocks according to the energy ____________________
• Blocks of the Periodic Table
• There are s, p, f, and d blocks of the periodic table which correspond to the energy sublevels. s Block Elements
• Write the closest noble gas without going over in brackets.
• Use the row number to get the energy level.
• Count the number of electrons until you get to the element in the s block.
• Mg _______________________________
Question Time
• Try other s-block elements. Write the noble gas configuration of the following elements
• Cs _______________________________
• Ca _______________________________
• Ba _______________________________ p block elements (Between 5-18)
• Write the closest noble gas without going over in brackets.
• Use the row number to get the energy level.
• Write s 2 after the row number because you have to go through the s-block to get to the p-block.
• Write the row number again
• The write “p” and then count the number of p electrons you must get through to get to your element as a superscript
• Si: ________________________________
Question Time
• Try other p-block elements. Write the noble gas configuration of the following elements
• N : ________________________________
• S : ________________________________
• Cl : ________________________________ d block elements (Between 21-48)
• Write the closest noble gas without going over in brackets.
• Use the row number to get the energy level.
• Write s 2 after the row number because you have to go through the s-block to get to the d-block.
• Write one less than the row number (d-block elements are always one less than the row number)**d for down one row number
• Then write “d” and count the number of d electrons you must get through to get to your element as a superscript
• Co: ____________________________
Question Time
• Try other d-block elements. Write the noble gas configuration of the following elements
• Ti : ________________________________
• Zn : ________________________________
• Mn : ________________________________ p block elements (Between 31-53)
• Write the closest noble gas without going over in brackets.
• Use the row number to get the energy level.
• Write s 2 after the row number because you have to go through the s-block to get to the p-block.
5

• Write one less than the row number (d-block elements are always one less than the row number)**d for down one row number
• Then write “d” and count the number of d electrons you must get through to get to your element as a superscript
• Write the row number again and “p” and count over the number of p electrons until you get to your element
• Br: ______________________________
Question Time
• Try other p-block elements. Write the noble gas configuration of the following elements
• Sn : ________________________________
• Se: ________________________________
The Modern Periodic Table
Early Periodic Table – Atomic Number
• In 1913 Henry Mosley discovered that each element contained a unique number of protons in the nuclei
• Arranged elements in order of atomic ___________________________.
• Resulted in a clear periodic pattern of properties.
Periodic Law
• There is a periodic repetition of chemical and physical _______________________ of elements when arranged in increasing atomic number (increasing number of protons) is called the periodic ___________
Modern Periodic Table
• Organized in columns called _________________ or families
• Rows are called ________________________
• Group A – representative elements (1A-____________)
• Group B - ___________________ elements (1B-8B)
Classification of Elements
• Three classifications for elements metals, nonmetals, and metalloids (semimetals)
Metals
• Properties of Metals
– shiny, smooth, clean solids (except mercury)
– __________________conductors of heat and electricity
– High ______________________
– High melting and boiling points
– ______________________ – bended or pounded into sheets
– Ductile – drawn into _________________
Groups of Metals
• ______________________ metals – group 1A except H
• Alkaline earth metals – group ____________
– Alkali metals and alkaline earth metals are chemically reactive
• Transition metals – group __________ elements
• Inner transition metals
– Lanthanide
– Actinide
Organizing by Electron Configuration
• Group number for group A elements represents the number of ___________________ electrons
• Atoms in the same group have similar chemical properties because they have the same number of valence electrons
Alkali Metals
• Electron configurations for alkali metals
• Lithium ________________
• Sodium 1s 2 2s 2 2p 6 3s 1
[He]2s
[Ne]3s 1
• Potassium 1s 2 2s 2 2p 6 3s 2 3p 6 4s 1
• Rubidium 1s 2 2s 2 2p 6 3s 2 3p 6 4s
[Ar]4s
2 3d 10 4p 6 5s 1 [Kr]5s 1
1
1
• What do the four configurations have in common?
• They have a _____________________ electron in their outermost energy level
• They all have one valence electron, thus similar chemical properties
Alkaline Earth Metals
• Electron configuration for alkaline earth metals
• Beryllium [He]2s 2
• Magnesium [Ne]3s 2
• Calcium
• Strontium
[Ar]4s
[Kr]5s
2
2
6

• All alkaline earth metals have _____________ valence electrons, thus similar chemical properties.
Nonmetals
• Gases or brittle, dull looking solids
• ______________________ conductors of heat and electricity
• Usually have lower densities, melting point, and boiling point than metals.
• Groups of nonmetals
– Halogens ____________
– Noble gases ____________
Noble Gases
• Noble gases – Group _______________
• Called inert gases because they rarely take part in a reaction
– He – 1s 2
– Ne – 1s 2 2s 2 2p 6
– Ar – 1s 2 2s 2 2p 6 3s 2 3p 6
– Kr – 1s 2 2s 2 2p 6 3s 2 3p 6 3d 10 4s 2 4p 6
• Because noble gases have completely filled s and p sublevels, they do not react with other elements
Metalloids (Semimetals)
• Physical and chemical properties similar to both metals and nonmetals
• They are metallic-looking _________________ solids
• Relatively good electrical conductivity.
• Used in glasses, alloys, and semiconductors
• The six elements commonly recognized as metalloids are boron, silicon, germanium, arsenic, antimony, and tellurium.
Polonium and astatine are sometimes classified as metalloids
Do the Trends w/s first!
Periodic Trends
Atomic Radius

Defined as ___________ of the distance between two bonding atom’s nuclei
Atomic Radius Across a Period
• Atomic radius generally ___________________________ in size as you move left to right across the period
– ___________________ positive charge in the nucleus pulls the electrons of the same energy level in.
Atomic Radius Down a Group
• Atomic radius _______________________________ as you move down a group
– Orbital size increases as you move down a group with increasing energy level
– Larger orbitals means that outer electrons are _______________________ from the nucleus. This increased distance offsets the greater pull of the increased nuclear charge.
– As additional orbitals between the nucleus and the outer electrons are occupied, the inner electrons shield the outer electrons from the pull of the nucleus this is called
__________________________.
7

Cation and Anion
• An ion is a positively or negatively charged atom that gains or loses an ___________________________.
• A cation loses electrons and produces a _________________________ charge
• An anion gains electrons and produces a _________________________ charge
Ionic Radius - Cations
• Groups 1A, 2A, 3A, and other metals _____________________ electrons and form cations.
• When atoms lose electrons they become __________________________
– The electron lost will be a valence electron leaving a completely empty outer orbital
– Protons in nucleus can pull fewer electrons tighter
Ionic Radius - Anions
• Group 5A, 6A, and 7A tend to ________________________ electrons and form anions
• When atoms gain electrons and form negatively charged ions, they become ________________________.
• Protons in nucleus have more electrons to pull and cannot pull in as tight
Do Ionization and Electronegativity w/s First!
Ionization Energy
• The energy required to _________________________ an electron from a gaseous atom
• Indication of how strongly an atom’s nucleus holds onto its __________________________electron
• Groups 1A, 2A, and 3A tend to have low ionization energies because they want to lose electrons.
Ionization Energy Trends – Across a Period
• Ionization energy generally ________________________as you move left to right
– Across a period electrons are added to the same energy level (same distance away from the nucleus), yet the nuclear charge is increasing across a period increasing the attraction to the electrons.
Ionization Energy Trends – Down a Group
• Ionization energy __________________________ as you move down a group
– Down a group electrons are added to a higher energy level (farther distance away from the nucleus), making it easier to remove an electron
Octet Rule
• Sodium atom 1s 2 2s 2 2p 6 3s
• Sodium ion 1s 2 2s 2 2p 6
1
(Sodium atom lost 1 electron)
• Neon 1s 2 2s 2 2p 6
• Sodium ion has the same electron configuration as neon
• Octet rule states that atoms gain, lose, or share electrons to acquire a full set of ___________________ valence
electrons (to be like a noble gas)
Electronegativity
• Indicates an element’s ability to _________________________ electrons in a shared chemical bond
• fluorine (F) is the most electronegative element
• Cesium (Cs) and francium (Fr)are the least electronegative
• Noble gases do not tend to have an electronegativity number since they tend not to form __________________
Trends with Electronegativity
• Electronegativity___________________________ as you move left-to-right across a period
• Electronegativity _____________________________ as you move down a group
8