Gen Chemistry Ch. 1: Atomic Structure Flashcards
Proton
Positive charge, mass around 1 amu
Neutron
No charge, mass around 1 amu
Electron
Negative change, negligible mass
Nucleus of an atom
Contains the protons and neutrons, electrons move around the nucleus
Atomic number (Z)
Number of protons in a given element, unique identifier because elements defined by number of protons
Mass number (A)
Sum of an element’s protons and neutrons, can be different within an element because of isotopes
Atomic mass
Essentially equal to the mass number, the sum of an element’s protons and neutrons
Isotopes
Atoms of a given element (same atomic number) that have different mass numbers, differ in number of neutrons
Isotopes of hydrogen
Protium, deuterium, and tritium
Atomic weight
Weighted average of the naturally occurring isotopes of an element, listed by periodic table
Rutherford model
The atom has a dense, positively charged nucleus that makes up only a small fraction of the volume of the atom
Bohr model of the atom
A dense, positively charged nucleus is surrounded by electrons revolving around the nucleus in orbits with distinct energy levels
Planck model
The energy difference between energy levels is called a quantum
Quanta
Discrete bundles of energy that are emitted as electromagnetic radiation from matter
Quantization
There is not an infinite range of energy levels available to the electron; electrons can exist only at certain energy levels
Atomic absorption spectrum
Unique to each element, for an electron to jump from a lower energy level to a higher one, it must absorb an amount of energy precisely equal to the energy different between the two levels
Atomic emissions spectrum
Unique to each element, when electrons return from an excited state to the ground state, they emit an amount of energy that is exactly equal to the energy difference between the two levels, sometimes the EM energy emitted corresponds to a frequency in the visible light range
Quantum mechanical model
Electrons do not travel in defined orbits but rather are localized in orbitals
Orbital
A region of space around the nucleus defined by the probability of finding an electron in that region of space
Heisenberg uncertainty principle
It is impossible to know both an electron’s position and momentum exactly at the same time
Quantum numbers
Completely describe any electron in an atom
Four quantum numbers
n, l, ms, ml
Principal quantum number
n - described the average energy of a shell
Azimuthal quantum number
l - aka angular momentum quantum number, describes the subshells within a given principal energy level (s, p, d, or f), values range from 0 to n-1, s = 0, p=1, d = 2, f =3
Magnetic quantum number
ml - specifies the particular orbital within a subshell where an electron is likely to be found at a given moment in time, values from -l to l
Spin quantum number
ms - indicates the spin orientation of an electron in an orbital (either +1/2 or -1/2)
Electron configuration
The pattern by which subshells are filled, as well as the number of electrons within each principal energy level and sub shell, uses spectroscopic notation
Spectroscopic notation
Combining the n and l values as a number and letter, respectively to designate the location of electrons
n+1 rule
Determines how electrons fill the principal energy levels and subshells according to increasing energy, the lower the sum of n+l, the lower the energy, if n+l is equal than the sub shell with the lowest n value will fill first
Hunds rule
Describes how electrons fill orbitals, states that subshells with multiple orbitals (p, d, and f) fill electrons so that every orbital in a subshell gets one electron before any of them gets a second
Paramagnetic materials
Have unpaired electrons that align with magnetic fields, attracting the material to a magnet, causes electrons to have parallel spins
Diamagnetic materials
Have all paired electrons, which cannot easily be realigned, they are repelled by magnets
Valence electrons
Electrons in the outermost shell available for interaction (bonding) with other atoms, least electrostatic pull from nucleus
Fundamental unit of charge value
e = 1.6 e -19 C
Attraction between nucleus and electrons
Masses so small that electrostatic force of attraction is far greater than the gravitational force of attraction
Cation
Positively charged atom
Anion
Negatively charged atom
Mole
Number of “things” equal to avogadros number
Avogadro’s number
Na = 6.02 E 23
Energy of a quantum
Described by Planck relation, E = h*f
Plancks constant value
6.626 E -34 J*s
Bohr angular momentum equation
L = n*h/2/pi
Bohr energy of an electron equation
E = -Rh/(n^2)
Rydberg unit of energy value
Rh = 2.18 E -18 J/electron
Ground state
An atom in the state of lowest energy
Excited State
When at least one electron has moved to a sub shell of higher than normal energy
Electromagnetic energy of a photon equation
E = h*c/lambda
Relation between wavelength and frequency equation
C = 3E8 = f*lambda
Line spectrum
Spectrum where each line on the emission spectrum corresponds to a specific electron transition
Lyman series
Group of hydrogen emission lines corresponding to transitions from energy levels n>=2 to n=1, highest energy
Balmer series
Group of hydrogen emission lines corresponding to transitions from energy levels n>=3 to n=2, includes 4 wavelengths in the visible region
Paschen series
Group of hydrogen emission lines corresponding to transitions from energy levels n>=4 to n=3, highest wavelength
Energy associated with energy change equation
E = hc/lambda = -Rh[1/ni^2 - 1/nf^2]
Pauli exclusion principle
No two electrons in a given atom can possess the same set of four quantum numbers
Maximum number of electrons within a shell equation
num = 2*n^2
Maximum number of electrons within a sub shell equation
num = 4*l + 2
Paired electrons
Electrons in the same orbital that has opposite spins
Parallel spins
Electrons in different orbitals that have the same spine
Aufbau principle
aka building-up principle; electrons fill from lower to energy subshells and each upswell will fill completely before electrons begin to enter the next one
Electron configuration of cations
Start with neutral atom, remove electrons from the subshells with the highest n first; if multiple are tied with highest n value, remove from the highest l among them
Electron configuration exceptions
Chromium and copper groups because half filled and fully filled orbitals are more stable so they will take 1 electron from the s orbitals to create either a half filled or fully filled d orbital
