Atomic Theory - Dalton & Thomson

• Unit 3 Test Review covers multiple choice, thinking, application, and communication.

Dalton's Atomic Theory (1805)

• This theory created the modern understanding of atoms.
• Matter consists of indestructible, indivisible atoms.
• Atoms of the same element are identical but differ from atoms of other elements.
• Each atom possesses a specific combining capacity.
• Some atoms can have more than one combining capacity.
• Atoms are neither created nor destroyed.
• Law of definite composition: elements combine in a characteristic mass ratio.
• Law of multiple proportions states there may be more than 1 mass ratio.
• Law of conservation of mass states total mass remains constant.

Thomson Atomic Theory (1897)

• Atoms can gain or lose electrons to form ions in a solution.
• Atoms and ions gain or lose a specific number of electrons.
• Electricity is composed of negatively charged particles.
• Matter is composed of atoms with electrons embedded in a positive material.
• The kind of element is characterized by the number of electrons in the atom.

Rutherford Atomic Theory (1911)

• Positive charge and most of an atom's mass are confined to a small central region called the nucleus.
• Few positive alpha particles deflect at large angles when fired at gold foil.
• The positive charge in the atom is concentrated in a small volume.
• Materials are stable and do not break down because the nucleus has strong forces
• A strong nuclear force holds the positive charges within the nucleus.
• Most alpha particles pass straight through gold foil, indicating most of the atom is empty space.

Bohr Model

• Electrons orbit the nucleus in fixed energy levels.
• Electrons can absorb or emit energy when transitioning between orbits.
• There is a nucleus (positive charge) and electrons orbit it.

Electromagnetic Radiation

• Visible light is a small fraction of the electromagnetic radiation spectrum.
• This includes x-rays, ultraviolet radiation, microwaves, and radio waves.

Photons

• A photon is a packet, or quantum, of electromagnetic energy.
• A photon's energy depends on its wavelength and frequency.
• Electromagnetic radiation with long wavelengths and low frequency is composed of lower-energy photons.
• Short-wavelength radiation with high frequencies is composed of higher-energy photons.

Quantum Theory

• Energy and matter behave in discrete and quantized units.
• Particles, such as electrons, have both wave and particle properties.
• Their behavior is only predictable probabilistically.
• Electromagnetic energy is not indefinitely sub-divisible.
• Energy exists as packets or quanta called photons.
• Photon intensity only changes the number of electrons released, not the energy of the electrons.

Planck's Quantum Hypothesis

• Solids glow when heated.
• Energies of oscillating atoms in the heated solid are multiples of a small quantity of energy.
• Light emitted is quantized, occurring in bursts, packages, or bundles, not continuously.
• As temperature increases, the proportion of larger quantum increases.

The Photoelectric Effect

• When light strikes a material, typically metal, electrons are released.
• Energy from the light is absorbed by electrons when light shines on a metal surface.
• These escaping electrons are called photoelectrons.

Emission Spectra

• Emission Spectra produce bright lines of specific colors.
• Emission Spectra occur when excited electrons fall to lower energy levels and release energy as light.

Absorption Spectra

• Absorption Spectra show dark lines or bands.
• Absorption Spectra occur when light has been absorbed by atoms or molecules.
• Absorption Spectra occur when electrons absorb specific energies and jump to higher energy levels.

Orbits

• Orbits are a 2D path.
• Orbits are at a fixed distance from the nucleus.
• Orbits have a circular or elliptical path.

Orbitals

• Orbitals are a 3D region in space.
• Orbitals are a valuable distance from the nucleus.
• Orbitals have no defined path, but have a varied shape of region.

Quantum Numbers

• Quantum numbers describe the quantum mechanical properties of orbitals.
• Quantum numbers are like addresses for locating electron positions.

Principal Quantum Number (n)

• It describes the size and energy of an atomic orbital.
• Relates to how far the electron is from the nucleus.
• The bigger the number indicates increases the distance of the electron from the nucleus.
• A higher n value signifies a larger energy level and greater probability of finding electrons further from the nucleus.
• When an electron moves from a higher shell to a lower shell, the energy difference is released as a photon.

Secondary Quantum Number (l)

• Secondary Quantum Number describes the shape of an atomic orbital.
• l = 0 is an s orbital, which is sphere shaped and sharp.
• l = 1 is a p orbital. which is dumb bell shaped and principal.
• l = 2 is a d orbital, which is pear shaped lobes and diffuse.
• l = 3 is an f orbital, which is fundamental.
• The range can be from 0 to n-1.
• The shape of the orbital changes depending on the value of l.
• Beyond l = 3, the letters start at g for l = 4 and continues alphabetically.
• Energy level differences between subshells are much less than energy level differences between shells.

Magnetic Quantum Number (ml)

• Ml describes an orbital's orientation in 3D space.
• If there are multiple shapes, this number indicates which way they are pointing.
• The number of values for ml is the number of independent orientations of orbits possible.
• If orbits are oriented in different planes, the energies of the orbits are different when the atom is near a strong orbit.
• If l = 1, ml can be -1, 0, or +1, so there are 3 orbits with the same energy and shape, but differing in their orientation

Spin Quantum Number (s)

• The spin can be either "up" or "down".
• This number describes the direction of electron spin.
• S = +½ (clockwise)
• S = -½ (counterclockwise)
• If two electrons are in the same orbital, they have opposite spins.

Quantum Numbers Overall

• n = energy level
• l = orbital type
• ml = specific orbital
• ms = spin

Energy Level Diagrams

• Energy Level Diagrams indicate which orbital energy levels are occupied by electrons for a particular atom or ion.
• Lines represent electron orbitals within subshells.
• Higher n# = greater energy of electrons
• For n#, sublevels increase in energy in the order s

Aufbau Principle - Electron Configuration

• An energy sublevel must be filled before moving onto the next higher sublevel.

Pauli Exclusion Principle

• No two electrons in an atom can have the same four quantum numbers.
• This means that in one orbital, the two electrons must have opposite spins.
• Electrons (arrows) are placed into orbitals by filling the lowest energy orbitals first.

Hund's Rule

• When several orbitals within the same energy level (p, d, or f) are available, one electron is placed into each before adding a second electron to any orbital.

Chemical Bonds

• Chemical Bonds are the electrical attraction that holds atoms or ions together within a molecular element or a compound.
• Atoms combine to form molecular elements or compounds.
• Bonded atoms have lower energy than single, uncombined atoms.
• The bonding potential of an atom is dictated by the number of valence electrons it has.

Ionic Bonds

• Ionic Bonds involves electrostatic attraction between negative and positive ions.
• Ionic Bonds occur in compounds formed via the transfer of valence electrons from one atom to another.
• Ionic Bonds form between a metal and a non-metal.

Covalent Bonds

• Covalent Bonds are the attractive force between two atoms.
• Covalent Bonds occur when valence electrons are shared between atomic nuclei.
• Covalent Bonds form between two non-metals.

Lewis Theory of Bonding (1916)

• Atoms and ions are stable if they have a full valence shell of electrons.
• Electrons are most stable when paired.
• Atoms form chemical bonds to achieve a full valence shell of electrons.
• A full valence shell can be achieved by exchanging electrons between metal and non-metal atoms.
• A full valence shell may be achieved by atoms sharing of electrons between non-metal atoms
• Sharing of electrons results in a covalent bond

Lewis Structures

• Lewis Structures are a diagram representing the arrangement of covalent electrons and bonds in a molecule or polyatomic ion.
• Duet rule applies to hydrogen and period 2 metals in bonding where a complete outer shell of valence electrons is observed.

Steps for Drawing Lewis Structures

1. Identify the central atom, usually the element with highest bonding capacity
2. Add up the number of valence electrons available in each atom of the elements.
3. Place one pair of electrons between each adjacent pair of atoms which forms a single covalent bond.
4. Place pairs of the remaining valence electrons as lone pairs on the surrounding atoms following the duet rule for H atoms and the octet rule for all other atoms.
5. Put any remaining electrons onto the central atom.
6. Move lone pairs from the surrounding atoms into a bonding position between those atoms and the central atom until all octets are complete.
7. Draw Lewis structure replacing bonds with dashes.
8. For polyatomic ions add large square bracket and charge

Valence Bond Theory

• Atoms combine to form covalent bonds with overlapping orbitals of their own electrons.
• Orbitals with unpaired electrons move toward each other causing their orbitals to overlap.
• Electrons are shared by both orbitals, becoming concentrated in the overlap region to form a covalent bond.
• This overlap reduces the potential energy of both atoms, forming the bond.
• A covalent bond will form if attractive forces between inter-acting atoms are greater than repulsive forces.
• A single electron cloud forms that is attracted to both nuclei.

The VSEPR Theory

• VSEPR Theory is a method to determine molecular geometry.
• Geometry is based on electron pairs positioning themselves as far apart as possible.

Electron-Pair Repulsion

• Bonded and lone pair electrons position themselves as far apart as possible in a molecule to minimize the repulsive forces between them.

Hybridization

• Hybridization explains how atoms form bonds in molecules.
• Atoms mix atomic orbitals.
• Normally, atoms have certain types of orbitals (like s or p); they blend their orbitals to make bonding connections.
• This blending is hybridization and helps explain molecule shapes and bond angles.

Hybrid Orbital

• Hybrid Orbital is an orbital that forms from the combination of at least two different orbitals.

Molecular Polarity

• Molecular polarity is determined by adding together all polar bonds in a molecule.
• Molecular polarity is important because many physical properties are affected by it, such as melting point, boiling point, and solubility.
• Overall polarity is determined by the molecule's dipole moment.

Polar Covalent Bond

• Describes bonds with a En from 0.1 to 1.69.
• Has unequal sharing of electrons.

Non-Polar Covalent Bond

• Bond in which the bonding electros are shared equally between atoms.
• If the bonds are non-polar, then the molecule is non-polar

Rules for Polarity

1. Symmetrical molecules will be non-polar regardless of polar bonds, dipoles cancel out.
2. Non-symmetrical molecules will be polar if bonds are polar, dipoles do not cancel out.
3. If a covalent bond is non-polar, the molecule stays non-polar.

Sigma Bond

• A bond that is formed when the lobes of 2 orbitals directly overlap.
• Every single bond in a molecule is a sigma bond.
• Present in single, double, and triple bonds.
• Visualize Sigma Bonds like a handshake between two people facing each other.

Pi Bond

• A bond that is formed when the sides of the lobes of 2 orbitals overlap.
• Only present in double and triple bonds:
  • Double bond = 1 sigma + 1 pi
  • Triple bond = 1 sigma + 2 pi
• Think of it like atoms giving a side hug by sharing more electrons above and below the handshake.

Intramolecular Forces

• A chemical bond WITHIN a molecule.

Intermolecular Forces

• An attraction BETWEEN molecules.
• Strength determines physical properties of molecular compounds:
  • Physical state of a compound at a specific temperature and pressure.
  • Melting point.
  • Boiling point.
  • Surface tension.
  • Hardness and texture.
  • Solubility in various solvents.

Types of Intermolecular Forces

• The are three types of Intermolecular Forces:

1. Dipole-Dipole Forces
2. London Dispersion Forces (weakest)
3. Hydrogen Bonds (strongest)

Dipole-Dipole Forces

• Forms between the slightly positive end of one polar molecule and the slightly negative end of an adjacent polar molecule.
• Occur between all polar molecules.
• Strength depends on the polarity of a molecule, greater the difference, stronger the force.
• Weaken rapidly as distance between dipoles increases.

London Dispersion Forces

• Weak attractive force between all entities including non-polar molecules and unbonded atoms.
• They are caused by temporary imbalance of electrons within entities.
• The larger the molecule, the more electrons and protons attract each other, which increases the strength of the forces and the melting point.
• Large molecules have a greater tendency to stay together. Therefore, as you go down a group in the periodic table, london dispersion forces increase.

Hydrogen Bonds

• Strong dipole-dipole force between a H atom attached to a highly electronegative atom (N, O, or F).
• Reasons for H-bonding include large electronegativity differences and the small size of the H atom, which means the positive pole is highly concentrated.
• Hydrogen Bonds make the molecules strongly attract the negative pole of a nearby molecule