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Description
Flashcards by Kelsey Hopland, updated more than 1 year ago
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Created by Kelsey Hopland
almost 9 years ago
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| Question | Answer |
| atom | basic building block of matter, representing the smallest unit of a chemical structure |
| subatomic particles | 1. proton 2. neutron 3. electron |
| nucleus | core of the atom |
| element | composed of atoms alike in size, mass, and chemical properties |
| Dalton's Atomic Theory | theorized in the early 1800s by English scientist John Dalton, formulated theory of invisible building blocks of matter which came to be known as atoms |
| compound | composed of atoms of more than one element |
| chemical reaction | involves the separation, combination, or rearrangement of atoms, however NOT the creation or destruction of such |
| (amu) | unified atomic mass unit, which is equivalent to one Dalton (Da) |
| atomic number (Z) | equivalent to number of protons found in an atom of that element |
| proton | subatomic particle contained in the nucleus possessing a single charge in its mass of 1 amu determines atomic number of an element |
| neutron | a subatomic particle contained in the nucleus contributing 1 amu, yet no charge |
| electron | subatomic particle contained in orbital shells contributing a single negative charge with negligible mass |
| ion | a positively or negatively charged atom |
| mass number (A) | equal to total number of nucleons contained within in the atom |
| nucleon | corresponds to subatomic particles contained within the nucleus, namely protons and neutrons |
| molecular weight | weight in grams per one mole of a given element |
| mole | unit used to count particles, represented by Avogadro's number which is equivalent to 6.02 * 10^23 |
| isotopes | species which possess the same number of protons but a different number of neutrons, resulting in different masses |
| standard atomic weight | weighted average of all isotopes of an element found naturally on Earth |
| Ernest Rutherford | demonstrated in 1911 than an atom has a dense, positively charged nucleus that accounts for only a small portion of the volume of an atom |
| Max Planck | developed the first quantum theory in 1900, proposing that energy emitted as electromagnetic radiation from matter comes in discrete bundles called quanta |
| energy value of a quantum | E=hf |
| Planck's constant | designated as h value is equivalent to 6.626 * 10^-34 J-s |
| angular momentum of an electron (equation) | L = nh / 2π |
| energy of an electron (equation) | E = -Ry / n^2 |
| Rhydberg energy | 2.18 * 10^-18 J/e- |
| relationship between energy and orbital radius | smaller the radius, the lower the energy state of the electron |
| ground state | lowest energy state of an electron (n=1) |
| electromagnetic energy of photons (equation) | E = hc / λ |
| line spectrum | quantized energies of light emitted under excited conditions each line on an emission spectrum corresponds to a specific electronic transition |
| atomic emission spectra | each element's unique excitation energy level |
| energy change between differing energy states (equation) | E = hc / λ = -Ry (1/n2i - 1/n2f) |
| absorption spectra | used in the identification of elements present in a gas sample by measuring the wavelengths of emission and relating the conclusions to wavelengths of absorption |
| orbital | representation of the probability of finding an electron within a given region in the space surrounding the nucleus |
| Heisenberg Uncertainty Principle | states that it is impossible to simultaneously determine with perfect accuracy the momentum and position of an electron |
| quantum numbers | describes the position and energy of an electron denoted by: n, l, ml, ms |
| Pauli Exclusion Principle | states that no two electrons in a given atom can possess the same set of four quantum numbers |
| principal quantum number | denoted by "n," represents the electron shell maximum "n" occurs at the ground state, which corresponds to element period (row) maximum number of electrons in shell = 2n^2 |
| Azimuthal quantum number | denoted by "l," describes angular momentum and shape of electron orbital l = n - 1, maximum electrons = 4l + 2 greater l = greater energy 0 = S, 1 = p, 2 = d, 3 = f |
| magnetic quantum number | designated as ml, describes the orientation in space of the electron shell possible values: -l through +l, including 0 |
| spin quantum number | denoted by ms, describes spin (intrinsic angular momentum) only two possibilities: ±1/2 |
| paired electrons | electrons with opposite spins in the same orbital |
| parallel electrons | electrons in different orbitals with the same spin |
| electron configuration | describes patter by which subshells are filled and the number of electrons within each principal level and subshell |
| Aufbau Principle | states that subshells must be filled from lowest to highest energy, each subshell will fill completely before electrons enter the next one |
| Hund's Rule | states that within a given subshell, orbitals are filled such that there are a maximum number of 1/2 filled orbitals with parallel spins |
| paramagnetic | describes a material with unpaired electrons which will be aligned in the presence of a magnetic field slight magnetic attraction |
| diamagnetic | describes a material with no unpaired electrons, which will be slightly repelled by a magnetic field |
| valence electrons | the electrons which occupy the outermost energy shells and or are available for chemical bonding |
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