Unit 2 Flashcards
Planck’s quantum hypothesis
Energy is emitted in small bursts or brackets called a quantum of energy
Photon
A quantum of light energy
The photo electric effect (Maxwell)
Mid 19th C
Light is an electromagnetic wave
The photoelectric effect (Hertz)
1887
The energy of light is dependant on the frequency, not the intensity or brightness of the light
The photoelectric effect (Einstein)
Electrons can absorb foxed amounts of light energy, these energy pockets are called photons
Bohr model
Electrons can be excited to a higher energy level by absorbing a photon and return to a ground state by emitting a photon
“uneven stair case”
The bigger the difference, the more energy the photon has
2n^2
Described the concept of quantized energy levels
Bohr faileurs
His model only worked for the hydrogen atom
Wave or particle
Electrons have properties of both waves AND particles
This idea helped Schrodinger create his wave equation
Orbital
Region around the nucleus where there is high probability of finding an electron. Can only hold 2 electrons with opposite spins
Heisenbergs uncertainty principal
It is impossible to simultaneously know the exact position and speed of an electron within an orbit
Quantum numbers
The energy of each subshell within a shell are not equal
Pauli exclusion principal
No two electrons can have the same 4 quantum values
Aufbau principle
Each electron is added to the lowest energy orbital available
Hund’s rule
One electron occupies each of several orbitals at the same energy level before a second electron can occupy the same orbital
Principle quantum number (n)
Represents the main shell of electrons (1….)
Secondary Quantum number
Divides the shell unto subtle that have slightly different energy levels and tells the shape of the subtle (0- N-1)
Magnetic quantum number (ml)
Describes the orientation of several subshells (-e - e)
Spin quantum number
In each orbital, there is one electron spinning at +1/2 and one spinning at -1/2
Origin of the quantum theory
Energy is not continuous
Stability- Hund’s rule
The lowest energy configuration for an atom is the one with all the orbitals of an energy level filled with one electron before adding the second. Electrons will choose to be by themselves if possible
Diamagnetic
Atoms with no unpaired electrons
Paramagnetism
The spin of the unpaired electrons creates a magnetic dipole moment, they act like tiny magnets
An external magnetic field will cause the spins to align
Isoelectronic ions
Have the same number of electrons
Anomalies in electron configurations
Cr: [Ar] 4s1 3d5
Cu: [Ar] 4s1 3d10
Lewis theory of bonding
Atoms are the most stable when they have the same electron configuration (isoelectric) as the noble gasses
Ionic compounds
Ionic formation is generally endothermic
Ionic compounds form when the lattice energy is more exothermic
The overall potential energy much be lowered when forming a compound
Square brackets
Show that there was a movement of electrons
Bond length
The distance between two nuclei in a covalent bond (average is 75pm)
Bond energy
The energy required to separate two nuclei in a covalent bond (average is 435kJ/mol)
Coordinate covalent bond
The shared e- come from one atom
Super octet
Occurs in the d orbital
C,N,O and F can not form super octets
Less than an octet
Beryllium is most stable with 2 bonds
Boron is most stable with 3 bonds
Formal charges
Lewis structures with the smallest formal charges are the most stable
VSEPR
Valance shell electron pair repulsion theory. Uses the idea that valance shell electron pairs will repel each other
AX2
Linear
AX2E
Bent
AX3
Trigonal planar *
Bond angle= 120
AX2E
Bent
Bond angle= less than 120
AX4
Tetrahedral
Bond angle= 109.5
AX3E
Pyramidal
Bond angle of NH3 is 107
AX2E2
Bent
Bond angle of water is 109.5
AXE3
Linear
AX4E
See saw
AX3E2
T- shaped
AX2E3
Linear
AX6
Octahedral
Bond angle= 90
AX5E
Square pyramidal
In this case, it does not matter which bond you take off
Ax4E2
Square planar
Valance bond theory
Covalent bonds form when two orbitals overlap to produce a new combined orbital containing two electrons of opposite spins
Hybridization theory
Mixing together two of more orbitals on the same atom to form new orbitals that have equal energy
Sigma bonds
A bond that is formed when two orbitals overlap and the electron density is concentrated along the axis between the nuclei (single bonds)
Pi bond
A bond that is formed when orbitals are parallel to each other and they overlap to create two regions of electron density on opposite sides of the axis connecting the two nuclei (double and triple bonds)
What does the formation of Pi bonds do
Prevents the rotation of the molecules along the sigma bonds
Rotation of double or triple bonds require breaking Pi bonds therefore the energy is higher
Hybridization- Lone pairs
Lone pairs of electrons can be accommodated in hybrid orbitals as well
They have stronger repulsion with bonding pairs so they cause a slight decrease to the bond angles
Bond angle is calculated from sigma to sigma bond
sp shape
linear
sp2
trigonal planar
sp3
tetrahedral
sp3d
tetrahedral
sp3d2
octahedral
Delocalized molecular orbitals
For compounds is might be easier to share electrons by overlapping multiple p orbitals to form a large Pi bond and a delocalized Pi bond
Resonance
Two possible lewis structures but neither is correct. Real molecule is a mixture of the tw
Thompson
1897
Positive matter with negative pockets (plumb pudding)
Rutherford
1911
Gold foil experiment- discovered in the nucleus
Positive charge concentrated in a nucleus with empty, negative space surrounding
Chadwick
1932
Discovered in the nucleus
Ionic crystals
Hard, brittle, can conduct electricity when dissolved in water
Metallic crystals
Metals, shiny, silver, flexible solids that can conduct electrical and thermal energy
Electron sea provides electrostatic glue holding the atom enters together
Molecular crystals
Non metals
Most are soft crystals with low melting points that can’t conduct electricity
Intermolecular forces hold them together
Covalent network crystals
Metalloids/ carbon
Hard brittle solids with very high melting points that do not dissolve and can’t conduct electricity
Network of covalent bonds
