Midterm 1 Flashcards
Ionization and Ionization formula
Change in energy from ground state (n=1) to n = infinity in the gas state.
E(ionization) = E(n=infinity) - E(n=1)
Molecule
atoms of different elements connected together by a covalent bond. Ex) methane (CH4)
4 Postulates of Atomic Theory
- All matter consists of atoms
- Atoms of elements cannot be destroyed/created
- Atoms of an element are all the same
- Compounds tend to form molecules in a chemical reaction with specific ratios
3 Mass Laws
- Mass is conserved (cannot be created/destroyed)
- A specific compound is composed of the same elements in the same fractions by mass.
- Specific elements react to form compounds that combine with fixed masses that can be expressed as a ratio of small whole numbers.
Cathode Ray Experiment
FILL IN WITH SPECIFIC NOTES
H-atom mass ratio
mass of an electron is about 1/1600 the mass of an H-atom
Millikan oil-drop experiment
Determined the charge of an electron by having a drop of oil between a positive and a negative magnet, shooting it with x-rays and seeing how many electrons come off
Charge of an electron
-1.602 x 10^-19 C (coulombs)
Mass of an electron
9.109 x 10^-28 g
What is the early atomic nucleus model and who made it?
“plum-pudding” model by Rutherford
What is AMU?
Atomic mass units aka the number of protons and neutrons in an element
Classical Mechanics
Waves and particles behave differently, describes large objects
F=ma p=mv
Quantum Mechanics
Small particles behave as both waves and particles, uncertainty about momentum and position, described based on probability
Spectroscopy
study of the interaction light with matter
Frequency
Number of cycles that the wave makes over a unit of time
Photons speed
2.998 x 10^8 ms^-1
Formula to calculate photon’s energy
E(photon) = hv
h = Planck’s constant
v = frequency
Formula to calculate the energy level of an atom
E(n) = -R(H) = (Z^2 / n^2)
E(n) = energy level
-R(H) = Rundberg Constant
Z = number of protons
n = energy level
Ground state
Energy level of a proton where n = 1
Ionization and Formula
change in energy from a ground state to n = infinity in the gas phase
H -> H+ + e-
E(ionization) = E(n=infinity) - E(n = 1)
Ionization and Formula
change in energy from a ground state to n = infinity in the gas phase
H -> H+ + e-
E(ionization) = E(n=infinity) - E(n = 1)
Emission for ionization
n(initial) > n(final)
energy of electrons is going down
Excitation
n(initial) < n(final)
Energy of the electron i going up
Orbitals
Allowed energies and regions of space electrons occupy
Paramagnetic properties
Unpaired electrons and attracted to magnetic fields
Diamagnetic Properties
All electrons are paired and ion is weakly attracted by a magnetic field
Mendeleev’s periodic law
similar properties repeat periodically when elements are arranged in order increasing Z
Atomic Radius
Size of an atom defined as the 1/2 distance between the bonds
Atomic Radii Trends
Across a period, atomic radius decreases.
Down a group, atomic radius increases
Isoelectronic
Species with the same number of electrons
Ionizatin enengy/cation formation trends
Across a period, ionization energy increases
Down a group, ionization energy decreases
Electron Affinity
the energy change associated with the addition of an electron to a neutral atom in the gas phase. ANION FORMATION
Anhydrides
Oxides without water in acid-base reactions
Basic Oxide
metaloxygen + water»_space; Metal + OH-
Acidic oxide
nonmetaloxygen + water»_space; oxyacid
Why do bonds form?
Electrostatic attractions between cations and anions in ionic bonds or between electrons and the nucleus of another atom in covalent bonds
Quantum mechanical reasons of good orbital overlap and a balance of interactions
Properties of Metals
Few valence electrons
Low ionization energy
Less negative electron affinity
Tend to lose electrons
Properties of Nonmetals
Many valence electrons
High ionization energy
More negative electron affinity
Tend to gain elections to become anions
Can form covalent bonds by sharing electrons
Heteroatomic Bond
Polar covalent bond where electrons are not shared equally, leaning towards being ionic
Non-polar bond
Covalent bond
Ionic Bonds
metals and nonmetals with electrons transferred completely from one atom to another
Extended Structures
What ionic bonds between cation and anions form, held together by attractive electrostatic forces in a 3D array
Lattice Energy
Energy release when the gaseous ions combine to form ionic solid, proportional to the electrostatic energy.
Negative lattice energy is favourable
Physical Properties of Ionic Solids
Hard and rigid from strong attractive electrostatic forices
Brittle from external sources forces moving the charges too close together
Do NOT conduct electricity when solid because charges don’t move. Only conduct when molten or dissolved in solution
Molecule
Structure of covalent bonded atoms
Metallic Bonds physical Characteristics
malleable, ductile, moderately high melting point and boiling point, and conduct electricity
Greatest Force in Covalent Bonds
attraction of nucleus of one atom to electron of the other, and the e-e and nucleus-nucleus repulsion
Optimal Distance
Balance between attractive and repulsive forces in covalent bonds, creating the ideal bond length
Bond Energy
The amount of energy it takes to go from one mol of gaseous molecules to their respective individual atoms aka overcome the attraction of a covalent bond
Bond Breaking
requires energy to break bonds, endothermic
Bond Making
Releases energy, negative bond energy, exothermic
Electronegativity
The relative ability of an atom, covalently bonded within a molecule, to attracted shared electrons to itself.
Partial negative charge
Electronegativity atom taking a greater share of the bonding electrons requires partial negative charge
Partial poitive charge
A less electronegative atom taking the lesser share of electrons in a covalent bond
Hund’s Rule
When orbitals of equal energy are available, maximize unpaired spins
Aufbau Process
Process of building electron configurations
