chemistry midterm review

Thompson

Discovered electrons via Cathode Ray Experiment.

Cathode Ray Experiment

Showed atoms emit particles in electromagnetic fields.

Plum Pudding Model

Proposed atom contains subatomic particles uniformly.

Rutherford

Discovered nucleus and empty space in atoms.

Gold Foil Experiment

Demonstrated nucleus by deflecting alpha particles.

Dalton

Father of chemistry; proposed atomic theory.

Dalton's Atomic Theory

Atoms are indivisible; different elements are distinct.

Law of Conservation of Matter

Matter cannot be created or destroyed.

Law of Definite Proportions

Compounds have fixed ratios of elements.

Law of Multiple Proportions

Elements can combine in different ratios.

Bohr Model

Electrons exist in specific energy levels.

Electron Cloud Model

Describes probable locations of electrons around nucleus.

Energy Level

Regions around the nucleus where electrons occupy defined energy states only.

Energy Sublevel

A division within an energy level of an atom, characterized by a specific shape of electron orbitals

Orbital

Regions where up to two electrons are likely to reside. Has certain shapes like circular or propeller.

Spin

Direction of electron's rotation on its axis.

Alpha Decay

Emission of a "α" particle from an unstable nucleus

Beta Decay

High-energy radiation from the emission of a "β" particle.

Gamma Radiation

High-energy electromagnetic radiation.

Half-Life

Time taken for half of a radioactive sample to decay.

Aufbau Principle

Electrons fill lowest energy levels first.

Hund's Rule

Unpaired electrons will maximize same spin.

Pauli Exclusion Principle

Orbitals can hold at most 2 electrons , and the must have opposite spin.

Emission Spectra

Light emitted when electrons fall to lower energy levels.

Absorption Spectra

Black lines shown when an electron moves to a higher energy level.

Group

Columns in the periodic table.

Period

Rows in the periodic table.

Metal

Elements left of the staircase, excluding hydrogen.

Nonmetal

Elements right of the periodic table, excluding hydrogen.

Metalloid

Elements located on the staircase of the periodic table.

Alkali Metal

Group 1 elements, highly reactive.

Halogen

Group 17 elements, known for reactivity.

Alkaline Earth Metals

Group 2 elements, reactive metals.

Noble Gases

Group 18 elements, inert and non-reactive.

Transition Metals

Elements in groups 3-12

Inner Transition Metals

Elements between groups 2 and 3, lanthanides and actinides.

Valence Electrons

Electrons in the outermost shell of an atom.

Ionization Energy

Energy required to remove an electron from an atom.

Atomic Radius/Atomic Size

Distance between two atomic nuclei.

Electronegativity

Atom's ability to attract electrons in a bond.

Continuum of Bonding

States no compound is entirely covalent or ionic, shows cutoff values for nonpolar, polar, and ionic bonds.

Nonpolar

0.0 -> 0.4

Polar

0.4 -> 2.0

Ionic

2.0 -> 4.0

Ionic Bond

Transfer of electrons between metals and nonmetals.

Metallic Bond

Attraction between metal nuclei and delocalized electrons.

Covalent Bond

Sharing of electrons between two nonmetals.

Polar Bond

Unequal sharing of electrons between atoms.

VSEPR Theory

Predicts molecular shapes based on electron repulsion.

Molecular Orbital Theory

Explains how sharing occurs, includes sigma and pi bonds.

Ammonium

NH₄+

Acetate

C₂H₃O₂−

Cyanide

CN-

Hydroxide

OH-

Nitrate

NO₃-

Nitrite

NO₂-

Carbonate

CO₃²⁻

Sulfate

SO₄²⁻

Sulfite

SO₃²⁻

Phosphate

PO₄³⁻

Organic Compound

Contains carbon, often with hydrogen and other elements.

Atomic Mass

Weighted average mass of an element's isotopes.

Atomic Mass Formula

(isotope mass x percent abundance/100) + (isotope mass x percent abundance/100)

Isotope Notation

Representation of isotopes with mass number and element.

Ion

Atom with a net electric charge due to loss/gain of electrons.

Cation

Positively charged ion, loses electrons.

Anion

Negatively charged ion, gains electrons.

Radioisotope

Isotope that is unstable and undergoes radioactive decay.

Radioactive Decay

Process by which unstable nuclei lose energy.

Electron Configuration

Distribution of electrons in an atom's orbitals.

s Sublevel Shape

Spherical shape, one orbital.

p Sublevel Shape

Dumbbell shape, three orbitals.

s orbital

1 orbital, can hold 2 electrons.

p orbital

3 orbitals, can hold 6 electrons.

d orbital

5 orbitals, can hold 10 electrons.

f orbital

7 orbitals, can hold 14 electrons.

Lewis Structure

Diagram showing bonds and lone pairs in a molecule.

Molecular Shape

Geometric arrangement of atoms in a molecule.

Hybridization

Mixing of atomic orbitals to form new hybrid orbitals.

Nomenclature

Systematic naming of chemical compounds.

Binary Acid

Acids with monoatomic ions (ex: HCl, Hydroflouric Acid)

Oxyacid

Acids with polyatomic ions (ex: H₂CO₃, Carbonic Acid)

Hydrocarbon

Compound concisting of only hydrogen and carbon

Zn²⁺, Cd²⁺, and Ag⁺

Transition metals that only have one charge

H₂O

Water

NH₃

Ammonia

CH₄

Methane

Organic Prefix for 1

-metha

Organic Prefix for 2

-etha

Organic Prefix for 3

-propa

Organic Prefix for 4

-buta

Linear

180°

Trigonal Planar

120°

Bent (1 lone pair)

<120°

Tetrahedral

109.5°

Trigonal Pyramidal

<109.5°

Bent (2 lone pairs)

<<109.5°

Trigonal Bipyramidal

120°,90°

Seesaw

180°,120°

T-Shaped

90°