כימיה מלא

empirical formula

simplest ratio of a compound

Pauli Exclusion Principle

within an atom no 2 e- can have the same set of quantum numbers; if an orbital has 2 e-, they must have different spins

quantized energy levels

-e- can only exist at specific energy levels
-as they get farther from the nucleus, their potential energy increases

Coulomb's law

-can calculate the energy an atom has based on its distance from the nucleus
-greater the charge of the nucleus, the more energy the e- will have
-the closer the e- to the nucleus, the more energy the e- will have

Bohr Model

-electrons are present in specific energy levels
-when e- gain energy, they move up energy levels, while e- release energy as they move down

ionization energy (aka binding energy)

-amount of energy necessary to remove an e- from an atom; related to effective nuclear charge
-↑ across a period
-↓ down a group
-second i.e. \> first i.e.
-i.e. gradually increases each successive time, until the outer shell is empty, then it increases a lot

kinetic energy

-energy of motion

electron configuration

-use the periodic table to do this; note that the p's start w/ 2, d's w/ 3, and f's with 4
-shorthand method: write noble gas to stand for the configuration up to that element

metals

-left hand side of periodic table
-give up e- in ionic bonds

non-metals

-upper right hand of periodic table
-gain e- in ionic bonds

atomic radius

-↓ across a period
-↑ down a group
-cations < atoms
-anions \> atoms

electronegativity

-atom's ability to pull electrons toward itself when involved in a chemical bond
-↑ across a period
-↓ down a group

ionic bond

-electrostatic attraction between ions (e- are given up, creating ions)
-creates a lattice structure; greater the charges and smaller the ions, greater the lattice energy
-high melting/boiling points

metallic bonds

-metals can bond with themselves, forming a sea of electrons
-metals can bond with other metals, forming alloys; alloys are interstitial if between atoms w/ vastly different radii or substitutional when between atoms w/ similar radii

covalent bonds

-2 atoms share electrons

conductors: ionic compounds

-ionic solids do not conduct electricity
-ionic liquids do conduct electricity

sigma (σ) bond

-first covalent bond

pi (π) bond

-2nd and 3rd bonds in a covalent compound

as the number of covalent bond increases...

-...the bond length decreases
-...the bond energy increases

polar covalent bond

-e- are unequally shared b/c atoms have different electronegativities; greater the difference in electronegativities, more polar the bond
-molecule has dipole moment

Intermolecular forces

-aka IMF's
-only exist in covalently bonded molecules
-includes network covalent bonds, hydrogen bonds, dipole-dipole forces, and london dispersion forces

hydrogen bond

-like a dipole-dipole, but stronger
-hydrogen end of molecule attracted to F/O/N
-high melting/boiling points

dipole-dipole force

-positive end of one polar molecule attracted to negative end of another
-greater the polarity, greater the dipole-dipole
-low melting/boiling points

London Dispersion Forces (LDF's)

-weak attraction due to e- movement that forms a temporary dipole
-larger the molecule, larger the LDF
-low melting/boiling points

How to Draw a Lewis Structure

1. count up number of total valence e- (add e- for anions; subtract for cations); this is how many e- should be in your final answer
2. draw molecule w/ bonds so that each molecule has full octet

electron-deficient

-Boron--\> only needs 6 e- to be stable (BF₃)
-Beryllium--\> only needs 4 e- (BeCl₂)

exceeding octet rule

-element must be in 3rd period or higher
-exceeds by using d orbital
-ex--\> SF₆, PCl₅, XeO₃, ICl₄⁻, ClF₃

resonance form

-occurs when 2+ Lewis structures can be made for a molecule
-can be flipped to resemble each other
-'real' molecule is an average of these structures

formal charge

-(normal valence e- \# for atom) - ((\# of lone pair e-) + 1/2(\# of shared e-))
-sum of the formal charges must

VSEPR Model

-model used to predict molecular geometry
-double/triple bonds treated same way as single bonds
-lone e- pairs occupy more space than bonding pairs

linear

-2 e- groups
-180° apart

hybridization

-mixing atomic orbitals to from special orbitals for bonding
-individual atom (normally center atom in a molecule) does this
-based on how many e- areas (lone pair groups, bonds)are around the atom; for ea. e- area, one orbital added
-ex--\>O in H₂O has 4 e- areas, so: sp³
C in CO₂ has 2 e- areas, so: sp
P in PCl₅ has 5 e- areas, so: sp³d

trigonal planar

-3 e- groups
-120° apart

bent

-2 e- groups
-normally in the molecular shape b/c two lone pairs were lost from tetrahedral (or one pair from trigonal planar)

tetrahedral

-4 e- groups
-109° apart

trigonal pyramidal

-3 e- groups
-normally in the molecular shape b/c one lone pair from the tetrahedral was lost

molecular shape

-the electronic shape minus the lone pairs

electronic shape

-all e- counted; lone pairs counted like bonds

trigonal bipyramidal

-5 e- groups
-120° (in triangle), 90° (up/down elements)
-when lone pairs are removed for molecular, they are removed from the trigonal planar part b/c bond angle is larger--\> can become see-saw, T-shaped, or linear

octahedral

-6 e- groups
-90°
-when lone pairs removed shapes are: square pyramid, square planar

phase and IMF's

-molecules w/ weak IMF's tend to be gases at room temperature
-molecules w/ strong IMF's tend to be solids @ room temp.

Kinetic Molecular Theory

1. volume of gas particles

Variation of Ideal Gas Equation

-@ constant temp: PV

Dalton's Law

-total pressure of a mixture of gases is the sum of all the partial pressures of gases

Partial Pressure

Deviation of Ideal Gas

-happens at low temp. or high pressure
-this is b/c the volume of gas molecules become relevant, raising the volume and gas molecules can start attracting each other, lowering the pressure

Molarity

moles of solute / liters of solution

mole fraction

moles of substance A / total moles of solution

synthesis reaction

-simple compounds combined to form one, more complex, compound

decomposition reaction

-a single compound is split into 2+ compounds

acid-base reaction

-an acid reacts with a base to form water and a salt
-ex→ HCl + NaOH → H₂O + NaCl

oxidation-reduction (redox) reaction

-oxidation states of some participating molecules
-ex→ Cu²⁺ + 2e- → Cu

precipitation reaction

-2 aqueous solutions mix and an insoluble salt is created

net ionic equation

-spectator ions that stay aqueous on both the reactants and products side are not included

solubility rules

-compounds w/ an alkali metal are soluble
-compounds w/ an NH₄⁺ are soluble
-compounds w/ an NO₃⁻ are soluble

limiting reactant problems

1. pick one product, and see how many moles of that product each reactant would make
2. the reactant making the least amount of the product is the limiting reactant

enthalpy change (∆H)

\-

energy diagrams

-displays the reaction as a graph

catalyst

-lowers the activation energy by displaying an alternate path
-in a rxn mechanism, it would be in reactants of 1st step and products of last step, but not in overall equation

oxidation states--covalent compound

-if they are identical atoms, the e- are split equally (o.s. of 0)
-if they are different atoms, the e- are given to the atom with stronger attraction to e-

an atom is oxidized

-when the atom's o.s. decreases as it gains e-

an atom is reduced

-when the atom's o.s. increases as it loses e-

oxidation states--special cases

-Fluorine→o.s. of -1
-Oxygen→o.s. of -2 in covalent; -1 in peroxides; +2 in OF₂
-Hydrogen→o.s. of +1 in covalent w/ nonmetals

oxidation states--atom/one element

-is always 0
-ex→F₂, Hg

collision theory

molecules must collide with correct orientation and enough energy in order to react

reaction rate increases...

...as concentration of reactants increases
-as temperature increases
-as surface area increases
-as volume decreases
-when a catalyst is added

heat

-total energy due to molecular motions in a substance (not the same as temperature)

temperature

-measurement of the average kinetic energy of a substance (not the same as heat)

types of energy transfer

-heat (energy goes from a warm object to a cold one)
-work (substance is stirred, raising its energy)

first law of thermodynamics

-energy of the universe is constant
-energy can't be created or destroyed, only converted

second law of thermodynamics

-if a reaction happens spontaneously (on its own) in one direction, it won't happen spontaneously in the reverse direction
-entropy of universe increases during spontaneous reaction

state functions

-enthalpy (H), entropy (S), and free-energy (G) change
-don't depend on the actual process of the reaction
-depend only the initial and final states

standard state conditions

-are true when you see a °, like H°
-include: gases @ 1atm, pure solids/liquids. 1M substances, element @ normal states has energy of formation (∆H°ƒ) of 0

Heat of formation, ∆H°

-change in energy when 1 mole of a compound is formed from pure elements
-exothermic → H is negative
-endothermic → H is positive
-pure element → H is zero

bond energy

-energy required to break a bond
-always endothermic and positive
-∆H°

Hess's Law

-if a reaction happens in multiple steps, you can add the ∆H values of the steps together
-if you flip the equation, flip the sign of the ∆H
-if you multiply/divide the equation, multiply/divide the ∆H

vapor pressure

-pressure of the molecules as they escape from the surface
-water boils when vapor pressure

heat of fusion

-energy that must be put into a solid to melt it

heat of vaporization

-energy needed to turn a liquid into a gas

phase diagram

-in water (when solid is less dense than liquid) line between solid/liquid slopes downward

heat capacity

-measure of how much the temperature of an object is raised when is absorbs heat
-large

specific heat

-amount of heat needed to raise 1g of a substance 1°C

heat added/ calorimetry equation

q

entropy (∆S)

-measure of the randomness
-entropy of solid < liquid < gas
-two moles have more entropy than one

Gibbs free energy (∆G)

-measure of whether a process will proceed w/o outside energy
-∆G positive→ won't happen
-∆G negative → will happen
-∆G

equilibrium

when the rate of the forward reaction is equal to the rate of the reverse reaction

Different K's

-Kc → constant for molar concentrations
-Kp → constant for partial pressures
-Ksp → solubility product
-Ka → acid constant for weak acids
-Kb → base constant for weak bases
-Kw → water ionization

Le Chatelier's Law

-whenever stress is placed on a reaction @ equilibrium, the equilibrium will shift to relieve the stress
-stress can be concentration, temp., pressure, volume

Le Chatelier's Law--Concentration

-if the concentration increases equilibrium will shift away from that substance
-if the concentration decreases equilibrium will shift towards that substance

Le Chatelier's Law--Volume

-if the volume decreases (increases) then equilibrium will shift toward the side with less (more) moles of gas molecules

Le Chatelier's Law--Temperature

-exothermic → heat is a product
-endothermic → heat is a reactant
-treat like concentration problem

Le Chatelier's Law--Pressure

-if the pressure decreases (increases) then equilibrium will shift toward the side with more (less) molecules of gas
-if a inert gas is added, there will be no change

reaction quotient, Q

-determined just like equilibrium constant, K, but using initial conditions

comparing K and Q

-if K \> Q, then more products need to be made
-if Q \> K, then more reactants are needed
-if Q

Ksp

-measure of how much a salt disassociates in a solution
-higher the Ksp, more soluble the salt

Arrhenius Acid-Base Definition

-acid → substance that produces H⁺ ions
-base → substance that produces OH⁻ ions

Bronsted-Lowry Acid-Base Definition

-acid → proton donor
-base → proton acceptor

conjugate base

-used to be part of an acid; now acts as a base because it will accept an H⁺
-ex → Cl⁻ from HCl

conjugate acid

-used to be part of an base; now acts as a acid because it will donate an H⁺
-ex → NH₄⁺ from NH₃

strong acid

-completely dissociates in water
-HCl, HBr, HI, HClO₄, HNO₃, H₂SO₄