Chemistry Fundamentals and Periodic Trends Flashcards
Metric Units to measure:
Length Mass Time Amount Temperature Electrical Current
Meter (m) Kilogram (kg) Second (s) Mole (mol) Kelvin (K) Ampere (A)
Order of Magnitude
Factor of 10
Density
Mass/Volume (m/v)
Molecular Formulas
Give identities and number of atoms in a molecule
Empirical Formulas
Smallest whole numbers that give the same ratio of atoms in a molecule
Formula Weight
Sum of the atomic weights of all atoms in an IONIC compound
Molecular Weight
Term for molecules
Atomic Mass Unit (AMU)
Unit for atomic weight. One AMU is 1/12th mass of an atom of Carbon-12
Mole
A particular number of things, collection of 6.022x10^23 of something (Avogadro’s Number)
moles = mass (g)/molecular weight
Molar Mass
Same as molecular weight but in grams/mol
Percentage Composition by Mass
Use empirical formula to find molecule’s percent mass composition
Molarity (M)
Measure of concentration that calculates the moles of solute in liters of a solution
Molarity = # moles/#liters
Mole Fraction
Fraction of moles of a given substance
Mole fraction = Xs = # moles of substance S/total # moles in a solution
Law of Conservation of Mass (Matter)
Amount of matter, and therefore mass, does not change in a chemical reaction.
Modify stoichiometric coefficients to balance Eq and have equal mass on both sides of arrow
Limiting Reagent
Limits the extent of a reaction (which reactant you are going to run out of first given an eq)
Notation of States of Matter
s
l
g
aq
Phases of substances of reaction
solid
liquid
gas
aqueous
Oxidation States
Atom’s ownership of its valence electrons change when it forms a compound
Rules for assigning Oxidation States (7)
- Oxidation of any element in standard state is 0
- Sum of oxidation states of atoms must equal molecule/ion’s overall charge
- Group 1 metals have +1 state, group 2 metals have +2
- F has -1 oxidation state
- H has a +1 state when with an element more electronegative than carbon, -1 with less electronegative, 0 with carbon
- Oxygen has -2 oxidation state (exception of peroxides with -1)
- Rest of halogens have -1 state, oxygen family has -2 oxidation state
Order of Electronegativity
F>O>N>Cl>Br>I>S>C>H
Atoms
Smallest unit of any element contain a central nucleus with nucleons (protons +1/neutrons 0) and surrounding electron cloud -1.
Number Neutrons
= Mass Number - Atomic Number (Number Protons)
Isotopes
Varying number of neutrons between two atoms
Charged Ions
Anion (negatively charged)
Cation (positively charged)
Atomic Weight
Weighted average of masses of naturally occurring isotopes
Radioactive decay - 3 types
Radioactive unstable nuclear undergo transformation to make them more stable
Alpha/beta/gamma
Alpha Decay
Large nucleus wants to be more stable by reducing number of protons and neutrons - emits alpha particle (2 protons and 2 neutrons)
4
a
2
Beta Decay
beta- decay = unstable nucleus that contains too many neutrons, may convert neutron into a proton and electron which is ejected
0
B
-1
Positron Emission
B+ decay = unstable nucleus containing too few neutrons, converts proton into a neutron and positron which is ejected
0
B
+1
Electron Capture
Capture electron from closest electron shell and use it in the conversion of a proton to a neutron
0
e-
-1
Gamma Decay
Nucleus in excited energy state (after any sort of alpha or beta decay) can relax to its ground state by emitting energy in the form of one or more photons of electromagnetic radiation (called gamma photons)
No change in atomic number/mass number, the gamma ray is emitted.
Half-life
Time it takes for one-half of some sample of the substance to decay. Shorter the half-life, the faster the decay. Decay is exponential over time
Decay Constant
Inversely proportional to half life
Nuclear Binding Energy
Energy released when individual nucleons (Protons/neutrons) bound together to form the nucleus. Also equal to energy required to break up the intact nucleus. Greater binding energy equals more stable
Mass Defect
When nucleons bind together, some mass converted to energy, so the mass defect is the difference between total mass of separate nucleons and the mass of hte nucleus
Einstein’s Equations for Mass- Energy Equivalence
E = mc^2
Emission Spectrum
Provides identity for a particular element through separation into component wavelengths
Gives energetic ‘fingerprint’
Energy
Ephoton = hf = h(c/wavelength)
h = Planck's constant f = frequency c = speed of light
Planck’s Constant
h = 6.63 x 10^-34
Bohr Model
Electron orbit nucleus in a circular path. Distance between n and n+1 decreases with increasing n
Photons
Absorbed to make an electron jump from ground to excited state (positive energy change)
Emitted to create a negative energy change (lower energy level, electron emits photon)
Quantum Model
Accounts for electron-electron interactions that exist in many-electron atoms
Energy Shell
(n) of an electron is analogous to circular orbits of Bohr model. higher the shell the higher the energy
Energy Subshell
Changes from circular orbits to 3D region around nucleus wherein electron likely to be found.
spdf (shape and energy described)
Orbital Orientation
Each subshell contains one or more orbitals of same energy (degenerate orbitals)
Electron Spin
Every electron has two possible spin states (up and down), every orbital can only acount for 2 electrons one up and one down.
Afbau Principle
Electrons occupy lowest energy orbitals available
Hund’s Rule
Electrons in same subshell occupy available orbitals singly before pairing up
Pauli Exclusion Principle
There can be no more than two electrons in any given orbital
Diamagnetic
All electrons spin paired
Paramagnetic
Not all electrons spin paired
Electron Configurations
Follow the periodic table (1s, 2s, 2p, 3s, 3p, 4s, 3d, 4p, etc/)
Anomalous Configurations
Some atoms achieve lower energy state by having fully filled or half filled d-shell (5 or 10)
Isoelectronic
Same electron configurations (not necessarily same number of electrons)
Shielding Effect
Also known as nuclear shielding effect - each filled shell between nucleus and valence electrons shields electrons from positive charges in the nucleus
Decrease of effective nuclear charge
Atomic /Ionic Radius
As you go across a period or up a group, atomic radius increases
Cations smaller than neutral atom, anions larger than neutral
Ionization Energy
Removal of an electron that requires energy input
Across period/up group, increases
Second ionization energy always greater than the first
Electron Affinity
Addition of an electron (releases energy)
Across a period and up a group, increases
Electronegativity
Atoms ability to pull electron density towards itself
Increase across period up a group
Acidity
Willingness to donate a proton
High electronegativity and large radius make favorable
With electronegativity trend, each of the following
Atomic Radius
Ionization Energy
Electron Affinity
Acidity
Decreases
Increases
Becomes more negative
Increases across, but decreases up
