CH 101 Final NCSU

Law of conservation of matter

Total mass stays the same in a chemical reaction

Law of definite proportions

Elements of a compound are present in fixed proportions by mass

Law of multiple proportions

When 2 different compounds are formed from the same 2 elements, the masses of one element that combine with a fixed mass of the other are in a ratio of small whole numbers

Compounds

Combinations of atoms of different elements; integral and constant

Elements

Composed of atoms which all have same chemical properties

Gay-Lussac Law of Combining Volumes

Volumes of reacting gasses are in simple, whole number ratios

Avogadro's law

Equal volumes of gas at the same temperature and pressure contain an equal number of molecules; The volume of a fixed amount of gas at constant temperature and pressure is directly proportional to the number of moles of gas: V=k(P,T)n

Molecule

2 or more atoms combine

Diatomic

Molecule of an element with 2 atoms and only when it's a free molecule and never in compounds

Mole

Conversion factor between the microscopic and macroscopic scale

Atomic mass

The mass of one atom

Molecular mass

The sum of the atomic masses of the elements that make up the molecule

Molar mass

The mass of any atom or molecule is equal to the atomic or molecular mass expressed in gram instead of amu

Limiting reactant

The reactant completely converted to products during a reaction

Energy

Capacity to do work (kinetic energy + potential energy)

Electromagnetic forces

When charged particles are brought together, they are either attracted or repelled by each other

Isotopes

Atoms of the same element which have the same number of protons and electrons, but different number of neutrons

Ions

Not an atom because it's not neutral- aka has a charge

Cation

Positive ion

Anion

Negative ion

Electromagnetic radiation

Light

Frequency

v; how many maxima pass a given point in one second

Continuous spectrum

All wavelengths merge into each other continuously

Blackbody radiation

The amount of energy in a wave is quantized

Photoelectric effect

Increasing the frequency of the light increases the kinetic energy: KE=h(v-v(initial))

Photoelectrons

Ejected electrons

Emission

Electrons move from higher level to lower level

Absorption

Electromagnetic moves from lower level to higher level

Uncertainty principle

We cannot know exactly where the electron is, but we can predict the probability that it will be found in some region of space

Electromagnetic configuration of an atom

The listing of the occupied sublevels and the number of electrons in each sublevel

Pauli exclusion principle

No two electrons can have the same set of 4 quantum numbers

Hund's rule

In order for an orbital to be in it's ground state, each orbital in a subshell must be filled one electron at a time all with the same spin direction. Electrons may be paired only after each orbital only has one electron in it. If these aren't followed it's in the excited state

Core electrons

All electrons in full sublevels except the outermost S sublevel; strongly held by the the nucleus, not involved in chemical bonding

Valence electrons

Electrons in partially filled sublevels and in outermost S sublevel; high energy electrons involved in chemical bonding

Ionization energy

Energy needed to remove the highest energy electron

Electronegativity

The measure of the ability of an atom o attract bonding electrons

Paramagnetic atoms

Have unpaired electrons; attracted by an applied magnetic field

Diamagnetic atoms

Have no unpaired electrons and are not attracted by an applied magnetic field

Group

Vertical row, or family; exhibits similar chemical properties

Horizontal row

As one progresses across the elements in a period, properties change gradually from one extreme to another

Photon

Particle of light

Ionic compound

In general, metal + nonmetal; form extended solids with no identifiable molecules

Isoelectronic

Have same electron configuration

Anion

A nonmetal that gains electrons to fill partially filled sublevels

Cation

A metal that loses some or all electrons to empty partially filled sublevels

Oxidation state

When part of a bonded atom is its charge, if it's bond was ionic when it's a monatomic ion it's simply it's charge

Polyatomic ions

Charged groups of covalently bound atoms

Oxoanions

Negative ions that contain oxygen atoms covalently bound to another element

Protonated anions

When ions that have charges of -2 and -3 pick up up protons

Covalent bonds

Nonmetal- nonmetal bonds (generally)

Bond length

Separation between 2 bound atoms at the position of minimum energy

Bond energy

The amount by which the energy of the two atoms is reduced by forming the bond i.e. How much energy is required to break the bond

Orbital energy

The amount by which the energy of the electrons is reduced by binding to the nucleus

Bond polarity

Bonding electrons are shared in a covalent bond, electronegativity differences between the bound atoms can result in unequal sharing, so, bonds in which the electrons are not shared equally are said to be polar bonds

Lewis structure

Used to determine the structure and bonding in a covalent molecule

Octet rule

Lewis structure of molecules are obtained by giving each atom (except hydrogen) an octet (8) of valence electrons as bonds -2 electrons per bond

Shared pairs

The number of bonds in the molecule (1/2(ER-VE))

Bond order

Number of shared pairs; increasing x-y bond order strengthens and the shortens the x-y bond (''tis only applies to bonds between the same 2 atoms)

Resonance structure

Structures that differ only in the placement of electrons

Formal charge

Electrons in free atom minus electrons in bonded atom

VSEPR

Valence shell electron pair repulsion; model of molecular shapes

Region of negative charge

Can be a single bond, a double bond, a triple bond, or a lone pair; repel each other and love as far away from each other as possible

Linear shape

One central atom, 2 bonding atoms, no lone pairs; 180 degrees

Bent shape

Central atom with 2 bonding atoms, and one or 2 line pairs; ~120 degrees if one lone pair; ~109 degrees if 2 lone pairs

Planar shape

Central atom with 3 bonding atoms, no lone pairs; 120 degrees

Pyramidal shape

Central atom with 3 bonding atoms and one lone pairs; ~109 degrees

Tetrahedral shape

Central atom with 4 bonding atoms and no lone pairs; 109 degress

Trigonal bipyramidal shape

Central atom with 5 electron regions

Octahedral shape

Central atom with 6 electron regions; lone pairs are always situated opposite to each other

Valence bond theory

Each bond results from the overlap of 2 atomic orbitals on adjacent atoms; a single bond is composed of 2 bonding electrons; the bonding electrons are localized in the region between the 2 atoms

Sigma bond

When atomic orbitals overlap head on; when electron density is located along the internucleus axis

Pi bond

When atomic orbitals overlap side-on; formed when electron density lies above and below the internucleus axis, no on it

Hybridization

The mixing of 2 or more atomic orbitals to make hybrid orbitals form; accounts for molecular geometries

Molecular orbital theory

Valence bond theory allowed us to rationalize observed molecular shapes by linking hybridization of atomic orbitals prior to bound formation. The hybridization model, however, leaves unexplained some molecular properties and other aspects of bonding; molecular orbitals are formed by combination of atomic orbitals from different atoms in the molecule; associated with the entire molecule, not with a single atom

In phase overlap

Of atomic orbitals generates a "bonding molecular orbital" (increased density between the nuclei)

Out of phase overlap

Generates an "antibonding" molecular orbital (decreased electron density between the nuclei)

Nodal plane

Region where the electron density is 0

Paramagnetic

Attracted to a magnetic field

HOMO

Highest occupied molecular orbital

LUMO

Lowest unoccupied molecular orbital

Kinetic molecular theory

1. The volume of the particles is negligible compared to the volume of the container. On average, the distances between gas particles are large compared to the size of the particles are large compared to the size of the particles; 2. The particles are in constant random motion; 3. Particles in the gas phase do not interact with one another; 4. The average kinetic energy of the particles in a gas depends only on the temperature

Boyle's law

Relating pressure and volume of gases; P1V1=P2V2

Charle's law

The volume of a fixed amount of gas at constant pressure is proportional to its temperature expressed in kelvins: (V1/T1)=(V2/T2)

Polar molecules

Molecules with permanent asymmetric charge distribution; if there are polar bonds present and dipole vectors don't cancel

Hydrogen bonding

Occurs between molecules with H-F, H-O, and H-N bonds; strong dipole interaction between the H from one molecule and F, O, or N from another; this is because H is small and can penetrate close to the F, O, N while F,O,N are also small, and highly electronegative

The ideal gas law

PV=nRT, R=0.0821 L atm/mol K

Dalton's law

Law of partial pressures; P(mixture)= Pb + Pb= [na/V]RT + [nb/V]RT

Thermal energy of gasses

Average kinetic energy ~ RT; R= 8.3154 J/K•mol

Intermolecular forces

Forces between different molecules; holds different molecules together; weaker

Intramolecular forces

Forces within molecules; holds the molecules together; stronger

Dipole-dipole forces

Present between polar molecules

Dispersion forces

Results from the interaction of induced dipoles; present in all molecules; more important in large molecules where the electron clouds are more easily deformed (more polarizable) so dispersion forces increase with molar mass

Viscosity

Resistance to flow; higher the stronger the intermolecular forces; lower the higher the temperature; strongly depends on the shape of the molecules (ability to tangle); result of cohesive forces

Surface tension

Energy required to increase the surface area of a liquid by a fixed amount; higher the stronger the intermolecular forces; result of cohesive forces

Cohesive forces

Forces between like molecules

Adhesive forces

Forces between unlike molecules

Concave shape

Result of strong, adhesive forces between substance and container

Convex shape

The result of weak adhesive forces between the substance and container

Phase diagram

A map showing which phase is the most stable at different pressure and temperatures

Vaporization

The change of state from liquid to vapor