enriched chem midterm terms

Chapter 1 (Modern Chemistry)

Chemistry (1.1): the study of matter and the changes it undergoes

Atom (1.2): ??

Compound (1.2): 2 or more elements that are chemically bound

Molecule (1.2): a combination of 2 or more pure substances that are not chemically bound

Element (1.2): a pure substance with one type of atom

Solid (1.2):

Liquid (1.2):

Gas (1.2):

Change of State (1.2):

Physical & Chemical Change (1.2):

Physical: no new substances formed/new molecules, no bonds broken or formed, still the same substance just a different arrangement

Ex: boiling, melting, freezing, condensation, evaporation

Chemical: new substances are formed, bonds are either broken or formed or both

Ex: color/odor change, explosion, transparent to opaque, precipitate formed, bubbles, heat increased or decreased,

Groups (1.3): columns

Periods (1.3): rows

Metal (1.3):

Nonmetal (1.3):

Semimetals (metalloids) (1.3):

  • conductive but brittle

  • combo of the components of metal and nonmetal
    Distillation: separation based on different compounds’ boiling points

Filtration: particle size based separation

Chapter 2 (Modern Chemistry)

Quantity (2.2)

SI Base Units (2.2)

Measurement (2.2): always approximate one digit!! the last digit tells the uncertainty!!

Scientific Notation (2.3):

Units (2.2):

Metric System (2.2):

Volume (2.2):

Mass (1.2):

Density (2.2)

Conversion Factors (2.2)

Accuracy (2.3)

Precision (2.3)

Significant Figures (2.3): aka sig figs

sig fig rules:

  • estimate one place beyond the smallest division

  • the last digit contains the uncertainty

  • counting zeros: # of digits = # of sig figs

  • zeros are counted when:

    • they are between non zeros or sig figs

    • they are after a non zero and a decimal is present

Rounding (2.3)

Dimensional Analysis (2.2)

Chapter 3 (Modern Chemistry)

Atom (3.2): a fundamental particle; cannot be broken down further

Atomic Mass Unit (3.3): AMU — ½ the mass of a C-12 atom

Atomic Number (3.3): the number of protons

Average Atomic Mass (3.3): the average mass of an atom of a given element — abundance a x mass a + abundance b x mass b

Avogadro’s Number (3.3): 6.02× 10²³

Isotopes (3.3): different atoms of the same element with a different number of neutrons

Mass Number (3.3): the number of protons and neutrons combined

The Mole (3.3): the amount of substance that contains as many particles as there are in 12.0 g of C-12

Mole Conversions (Dimensional Analysis) (3.3)








Chapter 6 (Modern Chemistry)

Chemical Bond (6.1)

Ionic Bonds (6.1): when one electron is transferred from one atom to another resulting in oppositely charged particles with an attraction to each other (between a metal and a nonmetal)

Covalent Bond (6.1): 2 elements of similar strength and electronegativity, nonmetals (some metalloids) sharing electrons

Nonpolar Covalent (6.1): electrons shared equally

Polar Covalent (6.1): electrons shared unequally

Bond Polarity (6.1)

Molecular Compound (6.1)

Chemical Formula (6.1)

Formula Unit (6.3): ratio of atoms

Octet Rule (6.2): atoms must have a complete valence shell (8 electrons)

Lewis Dot Structures (6.2)

Single Bond (6.2): longest, weakest, 2 electrons

Double Bond (6.2): inbetween strength and length, 4 electrons

Triple Bond (6.2): 6 electrons, shortest, strongest

Lattice Energy (6.3): energy released when 1 mole of an ionic compound forms from its ions

VSEPR Theory (6.5)

Molecular Shapes (6.5)

Polarity/Dipoles (6.5)

Intermolecular Forces (6.5)

Dipole-Dipole Interactions (6.5):

  • middle

  • all polar molecules have these

London Dispersion Forces (6.5):

  • only IMF for nonpolar molecules

  • weakest

Hydrogen Bonding (6.5):

  • strongest

  • between an H and a very attractive element (eg. N, F, O)

  • for polar molecules

Chapter 4 (Modern Chemistry)

Bohr (4.1): simple model of the atom based on understanding of the sharp line emission spectra of excited atoms

Orbitals (4.1)

Continuous Spectrum (4.1)

Electromagnetic Radiation (4.1)

Electromagnetic Spectrum (4.1):

  • higher frequency = high energy

  • long wavelength = low energy

  • red (700 μm) → violet (400 μm

  • radio → micro → infrared → visible (red → violet) → x-ray → gamma

  • low energy/frequency/wavelengths → high energy/frequency/wavelengths

Excited State (4.1):

  • emit light of certain wavelengths — this determines its color and depends on the element

Frequency/Energy (4.1)

Ground State (4.1): lowest possible energy

Speed of Light: 3.00 × 10^8 m/sec

Planck (4.1):

Line-emission Spectrum (4.1): when atoms absorb energy, that energy is often released as light energy — when that energy is passed through a prism, a pattern is seen that is unique to that type of atom (this is the emission spectrum)

  • falls = emits ⬇

  • jumps = absorbs ⬆

  • to 1st = UV light

  • to 2nd = visible

  • to 3rd = IR

Photon (4.1): electromagnetic energy with the properties of waves and particles

  • energy of a photon = hc/λ

Quantum (4.1): a discrete amount of energy released when an electron rests

Wavelength (4.1): λ

Photoelectric Effect (4.1)

Heisenberg Uncertainty Principle (4.2): light starts to fan out when it reaches a certain point

Orbital (4.2)

Quantum Theory (4.2)

Schrödinger (4.2)

Aufbau Principle (4.3): electrons will occupy the lowest potential energy orbital possible

Electron Configuration (4.3)

Hund’s Rule (4.3): aka the roommate rule — all orbitals of equal energy must have one electron before any have 2 electrons

Noble Gases (4.3):

  • full valence shell (8)

  • rare gases

  • stable

Noble-gas configuration (4.3)

Pauli Exclusion Principle (4.3): electrons in the same orbital must have opposite spins (beds facing opposite ways)

Quantized: electrons could only have very specific amounts of energy and traveled in orbits that were at fixed distances from the nucleus

Chapter 5 (Modern Chemistry)

Mendeleyev (5.1): made the periodic table — predicted there would be new elements in the spots he had left blank, predicted densities based on properties of other atoms around blank ones

Moseley (5.1): established the concept of atomic numbers — essentially discovering protons, ordering of the wavelengths of the x-ray emissions of the elements coincided with the ordering of elements based on atomic numbers

Actinide (5.1)

Lanthanide (5.1)

Periodic Law (5.1): states that the properties of elements are periodic functions of their atomic number

Periodic Table (5.1)

Alkali Metals (5.2):

  • lowest ionization energy in period

  • most reactive metal

Alkaline-earth Metals (5.2):

  • 2 valence electrons '

  • less reactive than alkali’s

  • form bases when reacting with water

  • harder, denser, higher melting points than alkali

Halogens (5.2):

  • salt formers (with alkali’s)

  • brinclhoff’s: gases found diatomically in nature: Br, I, N, Cl, H, O, F

Main-group Elements (5.2)

Transition Elements (5.2): aka transition metals: good conductors, harder, denser, higher melting point than s and p block elements, less reactive than s block elements

metals lose electrons to form positive: cations

nonmetals gain electrons to form negative: anions

Shielding Effect: the repulsion of outer electrons by inner electrons

  • F = (+ charge) (- charge) / r2

Effective Nuclear Charge (Z*): net positive charge that is attracting a particular electron

  • protons - inner electrons

Atomic Radius (5.3): size of the atom

  • get smaller as you go from left to right

  • the more electrons you add, the more protons are attracted

  • the bigger they are, the weaker they hold their electrons

Ionic Radius (5.3)

Ionization Energy (5.3): energy that an atom needs to absorb in order to lose an electron

Successive Ionization Energies: each successive removal of electron requires more energy than the last with it having a sizeable energy cost for the inner shell electrons

Valence Electrons (5.3): electrons in the highest energy levels s and p subshells (lose or gain to complete shell — except for noble gases)

Electronegativity (5.3): the strength an atom has when attracting electrons in a bond

  • increases as you go up and to the right

Reactivity (metals vs. nonmetals):

  • metals have a low Ionization energy and lose electrons to become smaller cations, reactive, conductive, not brittle

  • nonmetals have high electron affinities and gain electrons to become larger anions, brittle not conductive

Metallic Character



Chapter 7 (Modern Chemistry)

Naming Ionic Compounds (7.1)

Polyatomic Ions (7.1): more than one atom, with a charge, bonded covalently

Formula-writing for Ionic compounds (7.1)

Chapter 16 (Modern Chemistry)

Temperature (16.1):

Heat (16.1)

Calorimetry (16.1)

q=mc∆T problems (16.1):

  • Q = heat transferred (joules)

  • M = mass (grams)

  • C = specific heat (j/g°c)

  • ΔT = change in temperature (final - initial) (°c)

Endothermic

Exothermic

Specific heat (16.1): the energy required to raise the temperature of 1 gram of a substance by 1°c

symbol

name

conversion

k

kilo

1 kilo = 103 base

h

hecto

1 hecto = 102 base

da

deca

1 deca = 10 base

d

deci

1 base = 10 deci

c

centi

1 base = 102 centi

m

milli

1 base = 10³ milli