atomic theory chemistry first yrear students

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Atomic Theory Unit Test Review
Chemistry (High School - Canada)
Studocu is not sponsored or endorsed by any college or university
Chemistry (High School - Canada)
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January 17th, 2019 Chloe Popov
History of Atomic Theory
• Dalton: Matter’s composed of small parts called atoms.
• Thomson: Matter has positive and negative charged parts.
• Rutherford: Positive parts (nucleus) is the center, negative parts (electrons) are
in empty space, and the atom is mainly empty space.
• Planck: Matter emits EM radiation in quanta called photons.
• Einstein: The photons must have a threshold energy to remove an electron from
an atom. By colliding and removing electrons, photons act as particles.
• Bohr: Explained and calculated the size of hydrogens orbital radii (energy level),
and the energy of electrons at different orbitals (energy levels).
• De Brolie: Explained that electrons, like all matter, are particles that act like
waves.
• Schrodinger: Explained and calculated a region in space where an electron at a
specific energy level (orbital) can be found.
• Born: Calculated the probability of finding an electron in a region of space
(orbital).
• Heisenberg: Explained and calculated that for electrons, it’s difficult to know their
exact position and velocity at the same time.
Quantum Mechanics
• The current theory of atomic structure based on the wave properties of electrons.
• It’s currently unable to explain what electrons are and what they’re doing, and it
also has limits as to the precision with which we can measure and know things.
The 4 Quantum Numbers
• The Principal Quantum Number (n):
o Represents the energy level and relative size of the orbital. It can be any
integer and is the main energy level for an electron.
o To determine it, a line spectra of hydrogen was observed and it was
determined that electrons have specific energy levels, quanta, and that the
energy level describes the size of the orbital in which the electron can be
found.
• The Secondary Quantum Number (l):
o Represents the shape of the orbital and the energy sublevels for each
electron at a main level. They are the range of integers from 0 to n-1 and
are most commonly represented by the letters s (0), p (1), d (2), and f (3).
o To determine it, a higher quality spectra was observed where the initial
lines produced in the initial energy levels were thought to be made up from
electrons released at slightly different energies. Later, it was determined
that the different energy levels were due to the presence of different
shaped orbitals.
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• The Magnetic Quantum Number (m):
o Represents the orientation or the orbital in space. Electrons with different
magnetic quantum numbers will have the same energy level but different
orientations. The number of possible magnetic quantum numbers is equal
to 2l+1 and the range are the integers from -l to +l.
o To determine it, spectra were observed under a magnetic field. The
energy released by the electrons was found to split again. The values
were later described as the orientation of the orbitals the electrons were
found.
• The Spin Quantum Number (ms):
o Represents the orientation which the electron spins on its axis and it’s
either +
1
2
or -
1
2
.
o To determine it, atoms/molecules were observed to have a weak magnetic
force. The difference in magnetic attraction was deemed due to the
movement of electrons in orbitals. Each orbital can hold a maximum of 2
electrons which must have opposite spins.
Equations
Planks Equation:
E= hf
E- minimum energy needed to remove an electron.
h- Planck’s constant which equals 6.6 x 10-34 J•s
f- frequency of light in Hz (s-1)
Rydberg Equation:
1
𝜆
= 𝑅𝐻(
1
𝑛𝑓
2 −
1
𝑛𝑖
2)
λ- wavelength of photon
RH- Rydberg constant which is 1.10 x 10-7 m-1
nf- final energy level
ni- initial energy level
Defintions:
Orbital: A region in space around the nucleus that the probability of finding an electron
is maximum. It represents the 3D motion of an electron and doesn’t specify a specific
path since the electron can be anywhere in the orbital region. This is part of the
uncertainty principle. They can’t accommodate more than 2 electrons.
Electron Energy Diagrams: A diagram which displays the characteristics of an atom’s
electrons and can help explain the behaviour of an atom.
Polarity: A measure of a molecule’s electrostatic difference within the molecule. To
determine polarity of a molecule, you have to find the polarity of it’s individual bonds.
Dipole: The separation of charge between 2 atoms within a molecule.
Diamagnetic: All electrons are paired; opposed/repelled by magnetic field.
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Bohr’s Theory
• Bohr’s understanding of an atom was formed by the evidence that there’s a
periodicity of the physical and chemical properties of the elements, that there are
2 elements in the first period and 8 elements in the second period on the periodic
table, as well as the evidence that emission and absorption line spectra, but not
continuous spectra, exist for gaseous elements.
• Bohr believed that this evidence represented that electrons travel in the atom in
circular orbits with quantized energy, there’s a maximum number of electrons
allowed in each orbit, and electrons “jump” to a higher level when a photon is
absorbed and a photon is emitted when the electron “drops” to a lower level.
• Bohr’s first postulate was that electrons don’t radiate energy as they orbit the
nucleus. Each orbit corresponds to a state of constant energy, referred to as a
stationary state. Bohr’s second postulate was that electrons can change their
energy only by undergoing a transition from one stationary state to another.
• The main problem with his theory is that it worked well for only the spectrum for
hydrogen atoms or ions with only one electron. His theory can only explain the
results of experiments on atoms with single electrons.
• He thought that electrons orbited the nucleus in circular paths whereas in the
modernized concept, atomic electron structure is more alike 3D standing waves.
Electrons have probability distribution around the nucleus that form different
shaped orbitals.
Pauli Exclusion Principle
• No 2 electrons in an atom can have the same 4 quantum numbers.
• No 2 electrons in the same atomic orbital can have the same spin.
• Only 2 electrons with opposite spins can be in any one orbital.
Aufbau Principle
• Each electron is added to the lowest energy orbital available in an atom or ion.
They build up, filling orbitals from lowest to highest.
Hund’s Rule
• Each orbital will fill with one electron prior to the others filling with a second of the
opposite spin.
Electron Configuration Diagrams (Elements and Ions)
• Draw the atom first
• When losing electrons, they usually come from the s orbitals. The concentration
around the nucleus because of all of the overlapping s orbitals creates more
repulsion. There’s so much electron density near the nucleus because of the
overlapping s orbitals so there’s extreme repulsion, making electrons likely to
leave the highest s orbital to escape the extreme repulsion. They found that for
metals, electrons are usually lost from the highest s orbital. There’s more s orbital
density than d or p, so they usually come from the s orbital, even if it’s a lower
charge.
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Element: Sulphur (S), 16 electrons Ion: Nitride (N-3), 10 electrons
s- 2 p-6 d-10 f-14
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VSEPR Theory
Predicting and explaining the shape of molecules.
1. Draw the Lewis structure including bonding and lone pairs of electrons.
2. All double and triple bonds are treated like single bonds.
3. Add the pairs of lone electrons and then bonding electrons.
4. Use the VSEPR table to find out the shape of the molecule (lone pairs repel the
bonding pairs forcing the atoms closer together).
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Polarity
Bond Polarity:
• To determine a bond’s polarity, you have to find the change in EN between the 2
bonding atoms.
Nonpolar Bond Polar Bond Ionic Bond
0→ 0.5→ 1.7→
Nonpolar Bonds
• Formed when there’s an equal sharing of electrons, so the electrons would have
a high chance of being found around both atoms’ nuclei.
Polar Bonds
• Formed when there’s an unequal sharing of electrons, so the electrons would
have a higher chance of being found around one atom’s nucleus than the other’s.
• When a polar bond is formed a dipole is produced within the molecule.
Ionic Bonds
• Formed when electrons are removed from one atom and move to another, so the
electrons have a high probability of being found around only one of the atoms’
nucleus.
Molecular Polarity:
• Based on 2 main factors, bond polarity and shape of the molecule. If a molecule
has no has no polar bonds, it would be a nonpolar molecule. If it does have polar
bonds, it’s polarity is based on the shape of the molecule and where the polar
bonds are present. If they’re in opposing directions to each other they cancel out.
If they’re in a similar direction they add up.
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Intermolecular Forces
• Forces between molecules. They’re much weaker than any intramolecular force
(ionic or covalent bonding).
• There are 2 types, hydrogen bonding and van der waals forces.
Hydrogen Bonding
• Exists when a molecule (not a ketone or aldehyde, because no oxygen and
hydrogen attached) has a highly electronegative atom such as O or N bonded to
a hydrogen atom. The high polarity between the hydrogen and O or N creates a
relatively strong bond between molecules.
• The strongest type of intermolecular forces.
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Van der Waals Forces
• London Forces: The weakest intermolecular force. The very weak attraction
between the positive particles (protons) that make up a molecule and the
negative particles (electrons) that make up a molecule. The more electrons
present, the stronger the London forces.
• Debye Forces: Second weakest intermolecular force. More important when
dealing with intermolecular forces between 2 different types of molecules
(solubility). Attraction of molecules with polarity to the positive and negative
particles of non-polar molecules.
• Keesom (dipole-dipole): Third weakest or second strongest intermolecular
force. If 2 polar molecules are present (same or different molecule) they would
have a stronger force of attraction.
Lewis Structures
1) Find electrons has (have) and needs (need).
2) If the compound is charged, add or subtract the number of electrons needed to
give it the charge (if negative then add, if positive then subtract).
3) Make a symmetrical arrangement of atoms.
4) Put the electrons needed on and then add the rest of the electrons (have).
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5) Find the formal charge of each atom by adding the total lone pairs plus half of the
bonding electrons around an atom and subtracting it from the valence electrons
that the atom started with.
6) Shift electrons so that the lowest magnitude of formal charges is produced and
the negative formal charge is on the more electronegative atom (whatever’s
closer to F has a higher electronegativity). Make sure there’s no formal charge on
as many atoms as possible. If it has to be charged, positive goes on the lower En
atom and negative goes on the higher EN atom.
7) If a charged compound, redraw with stick bonds, showing lone pairs, and
brackets with charge.
Ex, PO4-3
Valence Bond Theory
• Pauling’s theory that when 2 half-filled orbitals overlap, they bond. The electrons
would have opposing spins.
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The p orbitals of N have 1 electron each so the hydrogens, that also only have on
electron, bond covalently to share electrons.
• Hybrid orbitals are identical spontaneously formed orbitals that only form when
bonding happens. Each orbital has an equal energy level and are usually
represented by a petal shape.
Atomic Orbitals vs. Hybrid Orbitals
• Atomic Orbitals: s, p d, and f.
• Hybrid Orbitals: Orbitals of different shape and energy that form when
molecules form (they’re from different atomic orbitals).
Ex, CH4
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Multiple Bonds
• Valence bond theory is used to explain multiple bonds.
• There are 2 ways orbitals can overlap and that partial hybridization can happen
which leaves remaining orbitals as they are.
• The first way is end-to-end overlapping. This creates a sigma bond (an example
of this was NH3). They’re found in single, double, and triple bonds.
• The second way is side-by-side overlapping which creates a pi bond. They’re
found in both double and triple bonds.
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Atomic Spectra Lab Information
• Small transition (1 shell back)- lower energy of wavelength (red light).
• Big transition (2+ shells back)- higher energy of wavelength (violet light).
• When energy is added to a particle (i.e. gases) its electrons get “excited” and
jump to a higher energy level (further shell from the nucleus).
o The light we see is after this happens when the electrons transition back
down to a lower energy level (shell closer to nucleus). Light is released.
The closer to the nucleus, the higher the wavelength (violet-higher, blue-
lower).
• Only specific energy levels an electron can have, can’t have halfway.
• When electrons transition from an outer shell to the first are emitting UV light.
𝐻2 =
1
𝑠
= 𝑠−1
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Crystal Structures
• Ionic Crystals
o Solid structures made of ions/ionic compounds.
o Formed from metal and non-metal ions.
o At room temperature, they’re fairly hard but brittle.
o High melting points due to strong ionic bonds between ions.
o If dissolved in water, will conduct electricity but not as undissolved solids.
o If heated to point of melting, will conduct electricity.
• Metallic Crystals
o Solid structures made of metal atoms.
o At room temperature, range from soft to hard and are malleable.
o Differing ranges of melting points.
o Believed to form from indirect electron attraction to the nuclei of
surrounding atoms.
o Loosely held ‘cloud’ of electrons that surrounds the nuclei are believed to
be the reason why they’re shiny, flexible and able to conduct electricity.
• Molecular Crystals
o Formed from solid, covalently bonded molecules.
o Relatively soft and brittle.
o Relatively low melting points due to weak intermolecular bonds holding
together (van der Waals and hydrogen bonding).
o They don’t have freely moving electrons and aren’t ions so they don’t
conduct electricity.
• Covalent Network Crystals
o Formed from metalloids (B, Si, Ge, As, Sb, Te, Po) and crystals of carbon.
o Extremely hard because of atoms forming covalent bonds with
surrounding atoms.
o Brittle because they aren’t malleable, although difficult to break.
o Due to strong bonds, have very high melting points than the other 3
crystals.
o Don’t have free moving electrons and don’t conduct electricity.
o Most common include carbon, graphite and diamond, but other gemstones
would also be covalent networks too.
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