Midterm 1 Flashcards
What is electromagnetic radiation?
It exhibits wavelike properties of electric and magnetic fields, and all electromagnetic radiation travels at the same velocity.
Characteristics of waves
Wavelength, frequency, and amplitude
Amplitude
Intensity, height of the wave maximum from centre
Wavelength
Number of waves passing a given point per unit of time
Blackbody radiation
Hot objects emit light and the wavelength depends on the temperature of the substance and based on quanta.
Quanta
Energy can only be absorbed/released in certain amounts
Not explained by classic theory
Blackbody radiation, photoelectric effect, line spectra
Quantum
The smallest amount of energy that can be emitted/absorbed as electromagnetic radiation
Photoelectric Effect
The emission of electrons from a metal when light shines on the metal that increases with brightness and frequency of incoming light
Line Spectra
The emission of light only a specific wavelength. Proves the existence of quantized energy
Line Spectrum
Line spectra emitted from excited gaseous elements, when separated by a prism splits light into component wavelengths
Continuous Spectrum
When radiation from a light source is separated into its different wavelength components but it is not true for excited gaseous elements.
Ground State
Lowest energy state
Quantized Energy
Electrons can only posses certain energy values; values between are not permitted. Electrons move to lower energy as light is emitted in the form of wavelength and move to higher energy states as light is absorbed.
Emission Lines
Caused from excited electrons dropping into lower energy orbitals, colours depend on the energy change during emission.
Uncertainty Principle
Cannot determine exact position, direction of motion or speed of particles simultaneously because the measuring process interferes with what is being measured.
Principle Quantum Number
n, describes main energy level and specifies electron shell
Angular Quantum Number
l, describes shape based on subshell (s, p, d, f)
Magnetic Quantum Number
ml, designates a specific orbitals and specifies orientation, based on order of electron being placed in orbitals. Each electron can only have one magnetic quantum number. -l…0…+l
Spin Magnetic Quantum Number
ms, describes the spin of the electron and based on the order of electrons placed in orbitals. = +/- 1/2
Node
Region where the probability of finding an electron is zero. The number of nodes is given by n-1
S-Orbitals
Spherical shaped and starts from 1s. Only have radial nodes and can hold up to 2 electrons. (1 set of 2) n-l-1 radial node(s), n phases
P-Orbitals
Dumbbell shaped and starts from 2p. Can hold up to 6 electrons. (3 sets of 2) n=l angular nodes and n-l-1 radial nodes
D-Orbitals
Four-leaf clover shaped, starts from 3d. Can hold up to ten electrons. (5 sets of 2) n=l angular nodes and n-l-1 radial nodes
F-Orbitals
Starts from 4f. Can hold up to 14 electrons. (7 sets of 2) n=l angular nodes and n-l-1 radial nodes
Pauli’s Exclusion Principle
No 2 electrons can have the same set of 4 quantum numbers and electrons in the same orbitals must have opposite spins.
Degenerate Orbitals
Orbitals with the same energy. The lowest energy is attained when the number of electrons with the same spin are maximized.
Electron Configuration Exceptions
A more stable configuration of electrons is when a set of p and d orbitals are either filled or half-filled, so Cr, Cu, Mo, Pd, Ag, Gd, Th, Pa, U, Np, Cm, Bk
Atomic Size
Increases down and to the left of the periodic table. Nuclear charge increases as atoms get smaller.
Effective Nuclear Charge (Zeff)
Charge experienced by an electron on a many-electron atom. The electron is attracted to the nucleus, but repelled by electrons that shield or screen it from the full nuclear charge. Higher Zeff means more charge and stronger bonds. Increases to the right of the periodic table and slightly down. Calculated by Zeff = Z-S, where Z is the atomic number and S is called the screening/shielding constant, represents the portion of the nuclear charge that is screened from the valence electron by other electrons in the atom. Also the number of core electrons based on the electron(s) being calculated.
Electron Screening
Core electrons are more effective at screening outer electrons from the full charge of the nucleus. Electrons in the same shell do not screen each other effectively. The reason Zeff becomes slightly large down a group is because screening is not perfect.
Cations
Positively charged ions. Smaller than their parent atom. Electrons are removed from the largest orbital and Zeff increases.
Anions
Negatively charged ions. Larger than their parent atom. Electrons are added to the largest orbital and the total electron-electron repulsions have increased.
Ionization Energy
Minimum energy required to remove an electron from the ground state of the isolated gaseous atom or ion. I(1) is the energy required to remove an electron from a gaseous atom, while I(2) is the energy required to remove an electron from a gaseous cation. The larger the ionization energy, the more difficult it is to remove the electron. It will always be a positive value as energy is required to remove electrons (endothermic). Increases up and to the right of the periodic table. Small atoms have high ionization energy because it is easier to remove an electron from the most spatially extended orbital. Ionization energy greatly increases when a core electron is removed.
Ionization Energy Exceptions
When you remove the first and fourth p electron it is easier than expected due to electron-electron repulsion.
Valance Electrons
Electrons in s and p orbitals; non core electrons. Electrons involved in bonding.
Ionic Bonds
Electrostatic attraction between oppositely charged cations and anions. A metal reacts with a non-metal by losing electrons to the non-metal, until both achieve the electron configuration of the noble gas (full octet) with the closet atomic number.
Electron Affinity
Energy change when a gaseous atom gains an electron to form a gaseous ion. Energy is released when this happens, so electron affinity energy is negative for stable ions (exothermic). Increases to the right of the periodic table and slightly up, but not by much because the attraction of the added electrons to the nucleus is much less - but so is electron-electron repulsion. Attraction of electrons to the nucleus goes up with increasing charge on the nucleus.
Covalent Bonds
Mutual attraction of atoms for a “shared” pair of electrons. Involves attraction between electrons and nuclei, repulsions between nuclei and repulsion between electrons.
Metallic Bonds
Electrons delocalized throughout a lattice of close-packed metal atoms.
Octet Rule
An octet consists of 8 electrons and full s and p subshells. Atoms tend to gain, lose or share electrons until they are surrounded by eight electrons - four electron pairs. Two electrons form one bond.
Equilibrium Bond Length
Minimum distance at which attractive forces between opposite charges balance repulsive forces between like charges.
Non-Polar
Electrons are shared equally. Less than 0.5 EN.
Polar
Electrons are shared unequally. Greater than 0.5 (0.4) EN to 2. Ionic is polar, but ranges from 2 to 4 EN.
Electronegativity
The ability of an atom in a molecule to attract electrons to itself. Explains how electrons react and related to ionization energy and electron affinity. The greater the difference in electronegativity between two atoms, the more polar the bond is. If the difference is large enough, a polar bond forms. Electronegativity increases up and to the right of the periodic table.
Formal Charges
The difference between the number of valence electrons in an atom and the number of electrons assigned to an atom.
Exceptions to the Octet Rule
Odd numbers electrons, less than an octet, or more than eight valance electrons (expanded octet).
Odd Number of Electrons
All atoms except one have an octet.
Fewer than Eight Electrons
Look out for B and Al. If filling an octet of the central atom results in a negative charge on the central atom and a positive charge of the more electronegative outer atom, don’t fill the octet of the central atom.
More than eight electrons
Atoms from the third period of the periodic table and beyond can accommodate more than an octet.
Resonance structures
Structures that occur when it is possible to draw two or more valid Lewis structures that have the same number of electron pairs for a molecule or ion.
Bond enthalpy
The strength of a bond is measured by determining how much energy is required to break the bond. Positive because bond breaking is an endothermic process. Multiple bonds are stronger than single bonds. Calculated by the sum of bonds broken subtracted by the bonds formed.
Lattice energy
Energy required to completely separate a mole of a solid ionic compound into its gaseous ions. Increases as the charges between two ions increases and distance decreases. Positive values to break a bond.
Electron domain
Also called an electron cloud, a region occupied by electrons. Each nonbonding pair (lone pair), single bond or multiple bond produces an electron domain about the central atom.
VSEPR
Predicts that the best arrangement of electron domains is one that minimizes the repulsions among them.
Nonbonding pairs
Two paired valence electrons that don’t participate in a chemical bond, physically larger than bonding pairs. Lone pairs should be placed as far apart as possible and always in the least crowded equatorial positions.
Repulsions
Electrons in nonbonding pairs and multiple bonds repel more than electrons in single bonds. (Increase or decreases angles)
Larger molecules
More complicated but are usually formed of tetrahedral, trigonal planar, bent and linear molecular geometry.
Molecular polarity
The uneven distribution of molecular charge.
Dipole moments
Point from higher electronegativity to lower, vectors, polar molecules have dipole moments. Orientation of the individual bond dipole moments determines whether a molecule has an overall dipole moment.
2 electron domains, 0 lone pairs
Linear electron domain geometry and linear molecular geometry. 180 degrees.
3 electron domains, 0 lone pairs
Trigonal planar electron domain geometry and trigonal planar molecular geometry. 120 degrees.
3 electron domains, 1 lone pair
Trigonal planar electron domain geometry and bent molecular geometry. <120 degrees.
4 electron domains, 0 lone pairs
Tetrahedral electron domain geometry and tetrahedral molecular geometry. 109.5 degrees.
4 electron domains, 1 lone pair
Tetrahedral electron domain geometry and trigonal pyramidal molecular geometry. <109.5 degrees.
4 electron domains, 2 lone pairs
Tetrahedral electron domain geometry and bent molecular geometry. <109.5 degrees.
5 electron domains, 0 lone pairs
Trigonal bipyramidal electron domain geometry and trigonal bipyramidal molecular geometry. 90, 120 degrees.
5 electron domains, 1 lone pair
Trigonal bipyramidal electron domain geometry and see-saw molecular geometry. <90, <120 degrees
5 electron domains, 2 lone pairs
Trigonal bipyramidal electron domain geometry and T-shaped molecular geometry. <90, <120 degrees.
5 electron domains, 3 lone pairs
Trigonal bipyramidal electron domain geometry and linear molecular geometry. <90, <120 degrees.
6 electron domains, 0 lone pairs
Octahedral electron domain geometry and octahedral molecular geometry. 90 degrees.
6 electron domains, 1 lone pair
Octahedral electron domain geometry and square pyramidal molecular geometry. <90 degrees, 90 degrees
6 electron domains, 2 lone pairs
Octahedral electron domain geometry and square planar molecular geometry. <90 degrees, 90 degrees.
Dipole
If two charges equal in magnitude and opposite in sign are separated by distance, then a dipole is established. Vector quantity.
Polar Molecules
Binary compounds are polar if their centers of negative and positive charge do not coincide.
Dipole Moments
The orientation of the individual bond dipole moments determines whether a molecule has an overall dipole moment.
All 5 parent AB(n) molecular geometries are…
non-polar
