WARM-UP
•Draw each of the three models of the atom that we
learned about last unit.
•Who came up with each?
•What was wrong with each?
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LIGHT
•Light is a particle AND a wave, at the same time
•Electromagnetic radiation – energy that exhibits
wavelike behavior as it travels through space.
–Visible light is only one example
PROPERTIES OF WAVES
•All light has a speed (c) of 3.0x108 m/s
•Wavelength (λ) – length of one complete wave
•Frequency (ν) – number of waves that pass a
point during a certain time period
–Hertz (Hz) = 1/s
•Amplitude (A) – distance from the origin to the
trough or crest
PRACTICE
•Looking at EM spectrum, which form of
radiation has the:
–Longest wavelength?
–Highest frequency?
–Highest energy?
PRACTICE
•Which form has the longer wavelength?
–Violet or green
–Blue or red
–Ultraviolet or infrared
–Ultraviolet or visible
–Infrared or visible
–Orange or yellow?
WAVE CALCULATIONS
c = λν
•c = speed of light (3.0 x 108m/s)
•λ = wavelength (m, nm, etc.)
•ν = frequency (Hz or /s or s-1)
WAVE CALCULATIONS
Microwaves are used to transmit information. What is the
wavelength of a microwave having a frequency of 3.44 x
109 Hz?
WAVE VARIABLE RELATIONSHIPS
•How is frequency related to wavelength?
•How is frequency related to energy?
•How is wavelength related to energy?
QUANTUM THEORY
Max Planck (1900)
•Observed the emission of light from hot objects
•Photoelectric effect – the emission of electrons
from a metal when light shines on the metal
•Light has to have a minimum frequency in order
for the photoelectric effect to occur!
QUANTUM THEORYMax Planck (1900)
•He suggested that an object emits energy in small,
specific amounts, called quanta
•Quantum – the minimum quantity of energy
that can be lost or gained by an atom
QUANTUM THEORYEinstein(1905)
•Light is a particle AND a wave!
–“wave-particle duality”
•Photon – particle of light having zero rest mass
and carrying a quantum of energy
MORE CALCULATIONS
Ephoton = hν
•E = energy (J, joules)
•h = Planck’s constant (6.6262 x 10-34 J· s)
•ν = frequency (Hz or /s or s-1)
WAVE CALCULATIONS
What is the energy of a photon from the violet portion of
the rainbow if it has a frequency of 7.23 x 1014 s-1?
BOHR’S MODEL OF THE ATOM
Bohr accounted for that problem!
•Electrons exist only in orbits with specific amounts of
energy called energy levels
ATOMIC EMISSION SPECTRA
Some definitions:
•Ground state: the lowest state or
energy of an atom
•Excited state: a state in which an
atom has a higher potential energy
than its ground state (this is when
an atom GAINS energy)
BOHR’S MODEL OF THE ATOM
•The smaller the electron’s orbit, the lower the atom’s
energy state
•Each orbit has a quantum number (n)
•When an electron moves to a HIGHER energy level
energy is put IN
•When an electron “drops” to a LOWER energy level, a
photon is EMITTED (light)
BOHR’S MODEL OF THE ATOM
•Each element has a unique bright-line emission spectrum
• Sad news? Bohr’s calculations only worked for hydrogen
MODERN QUANTUM MODEL OF THE ATOM
•DeBroglie’s Hypothesis (1923):
–If waves can behave like particles, then maybe
particles (e-) can behave like waves. Wave-particle
duality
MODERN QUANTUM MODEL OF THE ATOM
•Heisenberg’s Uncertainty Principle (1927)
–It is impossible to know both the position and
velocity of an e- at the same time.
MODERN QUANTUM MODEL OF THE ATOM
•Schrodinger’s Wave Equation (1926)
–Treat the electron as a wave
–He developed an equation used to determine the
probability of finding the e- in
any given place around the
nucleus.
–If these probabilities are
plotted in 3D, the probability
area becomes a cloud.
QUANTUM NUMBERS
•What are they?
–Values that represent different electron
energy states and the most probable
place to find an electron
PRINCIPAL QUANTUM NUMBER
•n
•Represents the cloud size (distance from the nucleus)
•Main energy level
•n≤7
•See row on the periodic table
•As n↑, size↑, Energy↑
SECOND QUANTUM NUMBER
•𝑙
•𝑙 = 0 → 𝑛 − 1
•Number of sublevels in a level = “n”
•Represents sublevels within an electron
cloud
•s,p,d,f (order of increasing energy)
THIRD QUANTUM NUMBER
•𝑚𝑙
•𝑚𝑙 = +𝑙 → −𝑙
•Represents the orbitals within sublevels
•s 1 orbital 𝑚𝑙 = 0p 3 orbitals 𝑚𝑙 = −1, 0, +1d 5 orbitals 𝑚𝑙 = −2,−1,0, +1,+2f 7 orbitals 𝑚𝑙 = −3,−2,−1,0, +1,+2,+3
FOURTH QUANTUM NUMBER
•𝑚𝑠 = ±1
2
•Only two electrons can occupy an orbital at
the same time
•Represents the spin of the electron (clockwise
or counterclockwise)
•Electrons in the orbital must have opposite
spin
Principal Energy
Level
Number of
Sublevels
Number of
Orbitals per
sublevel
Number of
Electrons per
Sublevel
Maximum
Number of
Electrons per
Energy Level
EXAMPLES
•Name the orbitals described by the following quantum
numbers
–𝑛 = 3, 𝑙 = 1
–𝑛 = 4, 𝑙 = 2
–𝑛 = 6, 𝑙 = 0
•Give the n and 𝑙 values for the following orbitals
–2p
–4f
WARM-UP
• What are the possible 𝑚𝑙values for 𝑙 = 2? Which type of
orbital is this?
• What are the four quantum numbers for the last electron
of potassium (K)?
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PAULI’S EXCLUSION PRINCIPLE
•No two e- in the same atom can have the same set of
four quantum numbers.
• If 2 electrons occupy the same orbital, they must have
opposite spin
AUFBAU PRINCIPLE
• “building up”
•Electrons occupy lowest energy level available
•Begin with the 1s orbital
• Sublevels overlap beginning with 3d…
•How do you know the order of energy?
HUND’S RULE
• “fair share”
•Each orbital in a sublevel must have one electron before
any orbital in that sublevel receives a 2nd.
ORBITAL NOTATION
•Used to show exactly where all electrons are in an atom
•An orbital is represented by a circle, line, or box
•Up and down arrows represent electrons with opposite
spin
WARM-UP• List all the quantum numbers for n=1
•What about n=2
•Write out the electron configuration and orbital
notation for phosphorus
•What is the noble gas configuration for iodine?
ELECTRON CONFIGURATION
•The numbers in front of the sublevel letters represent
the energy level
•The “exponents” refer to the number of electrons in that
sublevel.
•The sum of the exponents should equal the atomic
number of the element
VALENCE ELECTRONS
•Electrons found in the outermost energy level
•These electrons are involved in bonding
• Since the d and f orbitals have higher energy due to
overlap, s and p orbitals contain valence electrons
•Maximum and desired amount of valence e- is 8
–Except H, He, Li, and Be (only have 2)
ELEMENT SUPERHERO PROJECT
Three Easy Steps.
1. Choose one element from the periodic table and
research it.
2. Create a super hero that represents that element and its
unique qualities.
3. Design a comic depicting your superhero and name.
4. Print out instructions and follow closely
5. Complete the required worksheet
DUE: Wednesday, September 21, 2016
LAVOISER (1790’S)
•Composed a list of known elements (23 elements)
–Including oxygen, carbon, gold and silver
MENDELEEV (1869)
•Organized the first periodic table
by increasing atomic mass
•Connection between atomic mass
and elemental properties
•Predicted properties of
undiscovered elements (like
scandium, gallium and germanium)• http://ed.ted.com/lessons/the-genius-of-mendeleev-s-periodic-
MOSELEY•Henry Mosley- arranged the Periodic Table by
atomic number.
–Atomic number = number of Protons
•TODAY! The periodic table is still organized by
atomic number.
•Periodic Law: the statement that there is a
periodic repetition of chemical and physical
properties of elements when they are arranged by
increasing atomic number
GROUPS
• Elements in the same groups have the same number of
VALENCE ELECTRONS!
• Valence electrons are the outer most electrons on an
atom. They control how the atoms attach to each other.
SECTIONS IN THE PERIODIC TABLE• Alkali metals
• Alkaline earth metals
• Transition metals
• Halogens
• Noble gases
• Lanthinides
• Actinides
• Metals
• Non-metals
• Metalloids
CREATE A PAPER SLIDE VIDEO• http://www.youtube.com/watch?v=Qf6L1PTG3p4
• Each group (of approximately 4 people) will receive one of these
families (listed above) and create a presentation based on each
family. This presentation will then be presented to the class. Each
group will TEACH the class about their particular Periodic Table
Family.
• Each project will contain:
– 3 characteristics of the periodic family
– where they can be found on earth (if they are found on earth)
– how many valence electrons and/or what charge
– at least 2 ways the family is used by humans.
Number of Points Category
40 Completion
Did you turn in a completed project?
Did it include 3 characteristics, where they can be found, how
many valence electrons, and 2 ways it is used by humans?
20 Research
Did you provide hardcopies of research?
10 Complexity
How much effort (Thinking or Artistic) did you put into your
project?
10 Creativity
How creative is your idea?
How creative did you get with your project itself?
10 Presentation
Did you present your project?
Did you speak clearly, animatedly, and loudly?
10 Final Product
What does it look like?
How neat is it?
ALKALI METAL – GROUP 1A
• Soft
• Most reactive metals
• Group 1
• One valence electron
• Forms ions with a +1 charge
• Reactive violently with water.
• Francium is even radioactive!
ALKALINE EARTH METALS – GROUP 2A
•Are very reactive.
•Density greater than group 1
•Hardness greater than group 1
• Less reactive than group 1
•Two valence electrons
• Forms ions with a +2 charge
•Used in fireworks, batteries, and
your body!
TRANSITION METALS – D BLOCK
• Form colored ions in solution.
–Copper is blue or green
•More than one charge.
–Copper can be +1 or +2
•Most widely used by Humans
• Iron, Nickel, Copper, Gold, etc.
METALLOIDS
• Exhibit characteristics from both metals and nonmetals.
• Semiconductors.
• Boron, Silicon, Germanium, Arsenic, Antimony, and
Tellurium
• On staircase (except Aluminum)!
INNER TRANSITION METALS
• Lanthanides and Actinides
• Located outside of the periodic table.
• Often used with light and film.
• Uranium is in this section
– Used for nuclear power.
HALOGENS – GROUP 7A
• Most reactive nonmetals
• Called salt formers
• Forms ions with a -1 charge
• 7 valence electrons
• Bromine and iodine are used in halogen headlights
• Halogens form salts when forming an ionic bond with a
metal.
NOBLE GASES – GROUP 18
•All of these elements are gases.
•Do NOT interact with other elements - Inert
•Why?
–Because they have a FULL valence/outer shell.
OTHER GROUPS
• Other groups in the table are named by the
element at the top of the table.
• Some group are mixed with metals and non-
metals.
• Which two groups are most mixed?
ATOMIC RADIUS
• Defined by how close one atom is to another
• Rule: Atomic radius increases as you go DOWN a group
and decreases as you go to the RIGHT along a period
• Why??
–Because as you go down a group you increase the number
of energy levels
–As you go across radius decreases because you are
increasing protons causing a greater pull on the electrons
ATOMIC RADIUS
•What has the largest atomic radius: carbon,
fluorine, beryllium or lithium?
•Which has the largest radius: Mg, Si, S, Na?
IONIC RADIUS
• When atoms LOSE electrons, they become smaller
• When atoms GAIN electrons, they become larger
• The rule: Increases as you go DOWN a group and
decreases as you go to the RIGHT along a period
• Cations < neutral atoms < anions
ELECTRONEGATIVITY
•The ability of an atom to attract an electron.
• Fluorine has the highest electronegativity.
•The Rule: Decreases DOWN a group and increases
across the period.
•The noble gases have a electronegativity of zero.
Why?
ELECTRONEGATIVITY
•Rank the following elements by increasing
electronegativity: sulfur, oxygen, neon, and aluminum
IONIZATION ENERGY
• The energy it takes an atom to lose an electron.
• The alkali metals have a very low ionization energy,
because they want to lose an electron to have a full outer
shell.
• The halogens have a very high ionization energy.
• The noble gases have the highest (Helium the highest of all
of them) because they want to keep all of their electrons
so they keep their outer shell.
• **More valence electrons = higher ionization energy
REACTIVITY - METALS–The most reactive metal on the periodic table is Francium
–Metals want to give their electrons away
–The lower the ionization energy and the electronegativity,
the more reactive a metal is