Gen Chem 1 Flashcards
Protons
In the nucleus of atom mass=1 amu, charge=+1
Neutrons
neturons= A-Z
In the nucleus of atom mass=1 amu, charge=0
Electrons
electrons = #protons (if atom is neutral)
Circulate around the nucleus
mass=0 amu, charge=-1
strong nuclear force
Hold together the subatomic particles (nucleons) within the nucleus
-One of the strongest of nature’s 4 basic forces
Nature’s 4 basic forces
Strong nuclear force
Weak nuclear force
Gravity
Electromagnetic force
Electromagnetic force
Attraction between + and - or repulsion between like charges.
- Acts within the nucleus as repulsion between positively charged protons.
- The strong nuclear force overcomes this repulsion
Atomic Number (Z)
protons -> fingerprint of an atom, never changes for that element
Mass Number (A)
Protons + #Neutrons
- The actual weight of any atom is approximately equal to its mass number in amu or Daltons (1 amu = 1 Da = 1.66x10^-27 kg)
- The weight of 1 mole of any element is equal to its mass number in grams
Isotope
same # protons, different #neutrons (isotopes of any element differ only in the #neutrons; i.e. C-12 and C-13 -> C-12 has 6 neutrons while C-13 has 7.
Ion
Atom with a net-zero charge because it has gained or lost e- from its neutral state.
Cation (+): has lost e-
Anion (-): has gained e-
(Nuclear) binding energy
The energy required to break the nucleus down into its constituent parts.
-This amount of energy required was released when the nucleus formed from individual nucleons
Mass Defect
The mass of an intact nucleus is less than the sum of the masses of the nucleons that make it up. This difference is the mass defect.
-This amount of mass was converted to energy and released when the nucleus formed (nuclear binding energy) and these quantities are related by E=mc^2
Excited state vs. ground state of an e-
- Electrons are usually at their “ground state” energy level but can absorb energy and be promoted to an “excited state” of higher energy.
- When an e- falls from the excited state back to the ground sate, it releases energy (same amount of energy as was required to excite it).
Absorption spectrum
- Shine light through a substance and it will absorb certain specific frequencies (or colors) of that light.
- It is visualized as dark bands on a rainbow spectrum. The dark bands show which frequencies of light were absorbed.
Emission spectrum
- Pass light through a substance and certain frequencies of light are emitted. It is visualized as bright bands (the colors of the bands represents which frequencies of light were emitted) on a dark surface.
- Light is emitted from the substance when an atom transitions from a high energy state (due to an excited electron) to a low energy/ground state.
- The difference in energy between these two states is exactly equivalent to the energy of the photon released.
- Many different electron transitions possible for each atom, there are many different energies of photons (colors) released.
- “fingerprint” of an atom
Energy of Photon equation
E=hv = h(c/wavelength)
h=6.63x10^-34 Js
Bohr atom
- Simplified model of an atom in which electrons orbit the nucleus in a circular path.
- The distance of the e- from the nucleus is related to the energy of the e- (further away = higher energy e-; because negative e- want to be close to the positively charged nucleus).
- electrons can only possess discrete energy quantities (quantized energy states) and can “jump” between these discrete energy levels.
- transitions between energy levels are accompanied by an input of energy (if e- goes to higher energy level) or a release of energy (if e- goes to lower energy level)
Calculate energy of an e- at the nth energy level
E(n) = - (1/n^2) (13.6 eV)
-The value will be negative because it’s relative to a free electron (0 eV because n= infinity), and bound e- are more stable than free e-.
Larger - # = more stable
E(1)
-13.6 eV
Calculate energy absorbed/emitted when an e- transitions between energy levels
delta E= E(nf) - E(ni)
Reality of Bohr model
- only works for an atom with on electron, because the presence of more e- results in repulsion and the model no longer depicts reality.
- the idea that e- occupy certain “shells” or “levels” is still used. But the idea that those e- have fixed circular orbit was rejected.
Heisenberg’s Uncertainty Principle
The exact position or speed of an e- cannot be calculated, only the probability of it being in a certain region can be calculated.
-Modern quantum mechanical model
Electron orbitals and configurations
We can estimate the general “location” of an electron by noting its principle quantum number (the energy shell) and the orbital type (the energy subshell)
Principle quantum number (n)
Tells you the energy shell of the electron and ranges from 1 onwards (whole numbers only)
- Higher n values = further from the nucleus = higher in energy
- The value of n also tells you how many subshells there are in that energy level (n=1 has 1 subshells, n=2 has 2 subshells: s and p; n=3 has subshells: s, p, d).
Subshell
Describes the shape of the region of space in which an electron may exist. The possible subshells are s,p,d, and f. Each subshell type has a set number of individual orbitals.
s<p></p>
Electron spin
Each e- has a spin number which describes whether its spin is up or down. In each orbital, one e- has an up spin and one e- has a down spin.
Aufbau’s Principle
Electrons occupy the lowest energy level orbital first (1s), then fill in order of increasing energy.
Order of electron filling
Can be found by tracing across the rows (periods) of the PT.
-“n” value is the same as the period number for every block except d-block (row#-1) and f-block (row#-2)
Pauli Exclusion principle
There can be no more than 2e- per orbital, one with spin up and one with spin down
Hund’s Rule
Electrons in the same subshell (s, p, d, or f) occupy available orbitals singly first before pairing up
Valence shell
Highest energy (least stable) shell of that element
- Except for transition metals, #e is counted by column (skipping d-block)
- Noble gases are inert (unreactive) and very stable bc they have a full valence shell
Losing e-
Almost all elements will lose e- in the reverse order as they are filled.
Exception: Transition metals will lose e- from the s-subshell before the d-subshell
Diamagnetic
All e- are spin paired (each orbital has 2e-).
-Have no internal magnetic field but are repelled by the presence of an external magnetic field.
Paramagnetic
Not all e- are spin paired.
-Attracted to an external magnetic field.
Excited vs. Ground state
An atom may achieve an excited state by promoting an e- to a higher orbital than it usually occupies in the ground state.
The Photoelectric Effect
If a light shines on certain metals, e- can be ejected from the metal (now called photoelectrons).
-The incident light has a certain amount of energy (Ephoton), which is used to overcome the work function of the metal and to give e- some kinetic energy (KEe-) once it is ejected.
As long as Ephoton > work function of metal, then the KE of the photoelectrons will increase as the frequency of incident light increases.
Increasing the amplitude of the incident light (more photons) results in increased outgoing current (more e- ejected) but no change in KEe-.
