Gen Chem 1

intensive

property

A property such as density that is independent of the amount of the given substance

extensive property
A property that depends on the amount of given substance, such as mass
isotope
atoms of the same element with the same number of protons but different numbers of neutrons and consequently different masses
mass number
(A) the mass of the element
atomic number (Z)

The atomic number of the element (Z) that determines the number of protons and electrons in an atom

Natural abundance (percent abundance)

the relative percentage of a particular isotope in a naturally occuring sample with respect to other isotopes of the same element
anode
The electrode is an electrochemical cell where oxidation occurs electrons flow away from the anode
cathode
The electrode is an electrochemical cell where reduction occurs, electrons flow toward the cathode
alkali metals
Highly reactive metals in Group 1A of the periodic table
alkaline earth metals
Fairly reactive metals in group 2A of the periodic table
Law of Conservation of Mass
A law stating that mass is neither created nor destroyed in a chemical reaction

Law of Definite Proportion

A law stating that all samples of a given compound have the same proportions as their constituent elements

Law of Multiple Proportions

A law stating that when two elements (A and B) form two different compounds, the masses of element B that combines with one gram of element A can be expressed as a ratio of small whole numbers

oxoanions

A polyatomic atom containing a nonmetal covalently bonded to one or more oxygen atoms
formula unit
The smallest electrically neutrol collection of ions in an ionic compound
empirical formula
A chemical formula that shows the simplest whole number ratio of atoms in a compound
molecular formula

The chemical formula that shows the actual number of atoms of each element in a molecule of a compound

oxidation number

A positive or negative whole number that represents the "charge" an atom in a compound would have if all shared electrons were assigned to the atom with a greater attraction for those electrons

limiting reagent

The reactant that has the smallest stoichiometric  amount in a reactant mixture and consequently limits the amount of product in a chemical reaction
reducing agent
A substance that causes the reduction of another substance ; a reducing agent loses electrons and is oxidized
reactant in excess
The reactant where there is some left over at the end of the reaction
partial pressure
The pressure due to any individual component in a gas mixture
ideal gas
The proportionality constant of the ideal gas law (r = .0821)
ideal gas law
The law that combines the relationship of Boyle’s, Charles, and Avogadro’s laws into one comprehensive equation of state with the proportionality constant of R in the form PV = nRT.
kinetic molecular theory
A model of an ideal gas as a collection of point particles in constant motion undergoing completely elastic collisions

electromagnetic radiation

A form of energy embodied in oscillating electric and magnetic field
emission spectrum
The range of wavelengths emitted by a particular element used to identify the element
photon
The smallest possible packet of electromagnetic radiation with an energy equal to hv
photoelectric effect
the observation that many metals emit electrons when light falls upon them
de Broglie wavelength
The observation that the wavelength of a particle is inversely proportional to it’s momentum

Heinsburg Uncertainty Principle

The principle stating that due to the wave-particle duality, it is fundamentally impossible to precisely determine both the position and veleocity of a particle at any given time
principal quantum number (n)
An integer that expresses the overall size and energy of an orbital. The higher the quantum number n, the greater the average distance between the electron and the nucleus and the higher in energy
angular momentum quantum number (l)
An integer that determines the shape of an orbital
magnetic quantum number (m)
An integer that specefies the orientation of an orbital
principle shells (levels)
The group of orbitals with the same value of n
sublevels (subshells)
Those orbitals in the same principle level with the same value of n and l
The purpose of line spectra
Spectra are used to identify elements
Bohr’s explanation of spectral lines:
Light has energy equal to the difference in energy of the electron in two orbitals
Electrons possess….
both wave and particle properties
Spin quantum number (ml)
The fourth quantum number, which denotes the electrons spin order as 1/2 (up arrow) or -1/2 (down arrow)
Pauli Exclusion Principle
The principle that no two electrons in an atom can have the same four quantum numbers
Hund’s Rule
The principle stating that when electrons fill degenerate orbitals, they first fill them singly with parallel spins
Aufbau Principle
The principle that indicates the pattern of orbital filling in an atom
electron configuration
A notation that shows the particular orbitals that are occupied by electrons in an atom

Effective Nuclear Charge (Zeff)

The actual nuclear charge experienced by an electron, defined as the charge of the nucleus plus the charge of the shielding electrons

metallic radius

The way of expressing radius for metals
covalent radius (bonding atomic radius)
Defined in nonmetals as one-half the distance between two atoms bonded together, and in metals as one-half the distance between two adjacent atoms in a crystal of the metal
core electrons
those electrons in a complete principle energy level and those in the complete d and f sublevels

ionization energy

The energy required to remove an electron from an atom or ion in its gaseous state
electron affinity
The energy charge associated with the gaining of an electron by an atom in its gaseous state
expanded octet

In a Lewis Structure where the central atom has more than 8 valence electrons

formal charge
The charge that an atom in a Lewis Structure would have if all the bonding electrons were shared equally between the bonded atoms
resonance hybrid
The actual structure of a molecule that is intermediate between two or more resonance structures
coordinate covalent bond
The bond formed when a ligand donates electrons to an empty orbital of a metal in a complex ion
polar covalent bond
A covalent bond between two atoms with significantly different electronegativity values, resulting in an uneven distribution of electron density
nonpolar covalent bond
covalent bond in which electronegativity values are smilar
dipole
A measure of the separation of positive and negative charges in a molecule

valence bond theory

An advanced model of chemical bonding in which electrons result in quantum-mechanical orbitals localized on individual atoms that are a hydridized blend of standard atomic orbitals; chemical bonds result from an overlap of thse orbitals

Bond Order

1/2 (bonding electrons – antibonding electrons) in MO theory
Ion-dipole forces
An intermediate force between an ion and the oppositely charged end of a polar molecule
dipole-dipole forces
An intermediate force exhibited by polar molecules that results from the uneven charge distribution
ion-induced dipole forces
An intermediate force between an ion and the oppositely charged end of;a polar molecule

dispersion forces (London forces)

Intermediate force that is present in all molecules
hydrogen bonding
A strong dipole-dipole attractive force between a hydrogen bonded to O,N,or F and one of these electronegative atoms on a neighboring molecule
vaporization
The phase transition from liquid to gas
condensation
The phase transition from gas to liquid

Melting

The phase change from solid to liquid
freezing
the phase change from liquid to solid
sublimation
The phase transition from solid to gas
deposition
The phase change from gas to solid
vapor pressure
The partial pressure of a vapor in dynamic equilibirum with a liquid
critical pressure
The pressure required to bring about a transition; to a liquid at the critical temperature
critical temperature
The temperature above which a liquid cannot exist, regardless of pressure
critical point
The temperature and pressure above which a supercritical fluid exists
Triple Point
The unique set of coordinates at which all three phases of a substance are equally stable and in equilibrium
phase diagram
The map of the phase of a substance as a function of pressure and temperature
The prefixes:

Mega (M) = 10^6

Kilo (k) = 10^3

Deci (d) = 10^-1

Centi (c) = 10^-2

Milli (m) = 10^-3

micro (u) = 10^-6

nano (n) = 10^-9

pico (p) = 10^-12

Enery of a photon

E = hv

h = plank’s constant (6.626 x 10^-34)

 

Restrictions on the quantum numbers
Shapes of the orbitals
Avogadro’s Number
6.022 x 10^23 particles
Strong Acids

HCl

HBr

HI

HNO3

H2SO4

Strong Bases

Metal salts of….

OH

O2-

Molarity (M)

moles of solute

liters of solution

The Ideal Gas Law
PV = nRT

Initial and Fianl Gas Problems

P1V1 = P2V2
Sig Fig Rules

Determining Oxidation State

Properties of a Gas
Paramagnetic vs. Diamagnetic

Paramagnetic = there are unpaired orbitals

;

Diamagnetic = all the orbitals are paired up

Trends in Atomic Radius

Atomic radius decreases across a period.

;

Atomic radius increases down a group

Trends in Ionic Radius

The radius of a cation is much smaller than that of the corresponding atom.

;

;

The radius of an anion is much larger than that of the corresponding atom.

Trends in Ionization Energy

Ionization energy increases acorss a period.

;

Ionization energy decreases down a group.

Trends in Electronegativity

Electron affinity increases across a period.

;

Electron affinity decreases down a group.

Using electronegativity to pedict whether a bond is ionic, polar covalent, or nonpolar covalent.

Ionic = elements with very different electronegativities (more than 2)

Nonpolar covalent bonds = elements with very similar electronegativities

Polar covalent bond = those with intermediate electronegativity (less than 2)

;

Florine is the most electronegative element

Formal Charge

The formal charge of an atom in a Lewis Structure is the charge the atom would have if all bonding electrons were shared equally between bonding atoms.

;

Formal charges should add up to be zero.

Bond Length
Single bonds are longer than double bonds and double bonds are longer than triple bonds.
Bond strength
Triple bonds are stronger than double bonds and dobule bonds are stronger than single bonds.
Atoms with one bond:

H

F

Cl

Br

I

Atoms with two bonds:

O

S

Se

Te

Be

Atoms with three bonds:

N

P

As

B

Al

Atoms with four bonds:

C

Si

Ge

Oxidation Number:

– In its compounds, the oxidation state of flourine is -1

– In its compounds, the oxidation state of hydrogen is +1 (except when hydrogen is bonded to metals, its oxidation state is -1)

-In its compounds, the oxidation state of oxygen is +2 (when oxygen atoms are bonded to each other, its oxidation state is -1)

– Group 7A = -1

– Group 6A = -2

– Group 5A = -3

Sigma vs. Pi bonds

Single bond = 1 sigma bond

Double bond = 1 sigma + 1 pi

Triple bond = 1 sigma + 2 pi

Properties of A Liquid
Intermolecular Forces

Dispersion Forces: The weakest of the intermolecular forces, are present in all molecules and atoms and increase with increasing with increasing molar mass. These forces are always weak in small molecules but can be significant in molecules with high molar masses.

Dipole-Dipole Forces: present in polar forces

Hydrogen Bonds: The strongest of the intermolecular forces that can occur in pure substances (second only to ion-dipole forces), are present in molecules containing hydrogen bonded directly to flourine, oxygen, or nitrogen

Ion-dipole Forces: present in mixtures of ionic compounds and polar compounds. These are very strong and are especially important in aqueous solutions of ionic compounds.

Trends in Vaporization

– The rate of vaporization increases with increasing temperature

The rate of vaporization increases with increasing surface area

– The rate of vaporization increases with decreasing strength of intermolecular forces

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