chemistry exam review!
What is a solution!
-A solute dissolved in a solvent
-A physical change not a chemical one
-Usually reversible
Solvent
Solute dissolves in it!
Stays in its own phase.
Solute
Dissolves in solvent
Changes phase
Universal solvent?
Water!!!!
-Bent
-Polar
-Small
-Surrounds solute and forms attractions
-This is called hydration
Predict solubiltiy between molecular compounds
Like dissolves like!
-Polar will dissolve in polar, non-polar will dissolve in non-polar
-Size is also a factor, smaller molecules dissolve easier
Ionic compounds do what in water?
Dissasociate - be freeeeee
Lattice enthalpy
Strength between ions in an ionic compound - affected by the charges betweent the ions and the distance between the ions
Lattice enthalpy and solubility
More lattice enthalphy, less soluble!
Electrolytes
Ionic compounds after having dissasociated in water
Solutions with electrolytes conduct electricity
(Gatorade is one)
Factors that affect solubility
-Solids have higher solubility at higher temperatures
-Liquids stay the same
-Gasses are less soluble at higer temperatures
Pressure affects the solublity of gas, but nothing else
(solubility of gas is porportional to pressure exerted on it, more pressure more solubility)
Factors that will alter speed of solubility
-temperature (heat)
-agitation (stir or shake)
-paticle size (cut up)
States of a solution
Unsaturated- could use more solute
Saturated - At max amount
Supersaturated, limit has been exceeded, preciptate has formed
Solubility curve
Temperature in C on the bottom
Concentration on the side (g/100ml)
Dissasociation equations
-start with balanced equation
-break it down!
-keep only the stuff that forms the precipitate
Qualitative analysis of a solution methods
-colour!
-Flame test, flame colour
-Sequential analysis , what does it react with and how
Calculating concentrations of solutions (solid)
-Calculate the mass of solute you require
-Mass out the solid in beaker or on a weighing paper
-Dissolve the solid in approximately half of the total volume of the solution in a beaker
-Transfer the contents of the beaker to a volumetric flask
-Rinse beaker and stirring rod, transfer water to volumetric flask
-Add water until the miniscus reaches the etched line
-Stopper and invert to mix
Calculating of concentrations of solutions (liquid)
-Calculate volume of liquid required to do solution
-Measure the volume using a pipette
-Empty pipette into beaker, add approximately half of the solutions volume and stir to mix
-Rinse beaker and stirring rod, transfer liquid to volumetric pipette
-Add water until miniscus reaches etched line
-Stopper and invert to mix
Standard solution
-A solution you made
-You know the accurate concentration
Standardization
Determining the conentration of a solution by reacting it with a known amount of standard solution
stock solution
A concentrated solution you make or buy that will then be used to make more diluted solutions
Acids
-Got H^+ ions in them that create the acidic properties
Bases
Generally have OH^- in them
Can also be organic molecules like NH3
Arrhenius
Defined acids as things that left H+ in the water and bases as things that left OH-
Arrhenius revisted
An acid reacts with water to increase the
hydronium ion concentration
A base dissolves in water to increase the
hydroxide concentration
The Bronstead-Lowery concept
-Exchange of H+ protons - an acid gives on and a base recives one
-A compound can only be a bronstead-lower base or acid in a specific reaction, it is not a consistent property of the compound
Conjugate pairs
An acid and a base which differ only by the presence or absence of a proton are called a conjugate acid-base pair.
Amphiproteric
- A substance that can act as either a conjugate acid or base
Acid and base reactions
-Acids and bases neutralize each other,
forming water and an ionic compound
-Acids may be synthesized by adding a
non-metal oxide to water
-Acids react with metals to produce hydrogen
gas and a new ionic compound
-Acids react with carbonate compounds to
produce an ionic compound, carbon dioxide
and water
-Acids react with hydrogen carbonate
compounds to produce an ionic compound,
carbon dioxide and water
-Bases may be synthesized by the reaction of a
metal oxide and water.
-If you place a metal oxide in an acid, it will
react with water first, then neutralization will
proceed, based on solubility!
Strong acids
-Ionize completely in water
-Halogen containing, accept flourine
Weak acids
All other acids
Strong bases
Group 1 hydroxides (ex NaOH)
Group 2 hydroxides (ex Ca(OH)2 )
Weak bases
Organic molecues that contain nitrogen
Lewis acids and bases
Lewis acid is an electron pair acceptor
Lewis base is an electron pair donor
Forms a coordinate bond
pH scale
0- very acid
7- neutral
14 - basic
pH formulas
pH=−log[H3O+]
pOH=−log[OH-]
Kw=[H3O+][OH−]=1.0×10−14
pKw=pH+pOH=14.
Titrations
Indicators tell us when the solution has neutralized - acids and bases have canceled out
Titrant
Solution of known concentration used in titration
Analyte
Substance being analysed in titration
Equivalence point
enough titrant added to react exactly with the analyte
End point
The indicator changes color so you can
tell the equivalence point has been reached.
Titration calculation procedure
If you want to know the molarity of an acid
-Place a known volume of the acid in the flask
-Add a pH indicator
-Add base of a known concentration until the end
point has been reached
-Record the volume of base added
-Perform stoichiometric calculations to determine
the concentration of H3O+ in the acid
Ionic bond
A bond between a metal and non-metal
Covalent bond
A bond between two nonmetals
Coordinate bond:
A bond between two nonmetals wherein both electrons shared are from one atom
Polarity
Electrons shared between two atoms in a covalent bond, negative charge of electron can be equally or not equally distributed. Something a bond is considered more polar when an atom can take the electron away completely from the other, such as in an ionic bond. Whether a covalent bond is polar depends on the electronegativity difference of the two atoms.
Electronegativity:
The characteristic of the atoms that determines the distribution of the negative charge from the shared electron.
Determining the polarity of a bond
Use electronegativity values to determine the polarity of a bond
0-0.4 nonpolar covalent
0.4-1.7 polar covalent
>1.7 likely ionic
Only ever take the electronegativity value of the same atoms once.
Dipole:
When atoms in a covalent bond share an electron unequally, this happens when the atoms have an electronegativity difference.
Steps to make Lewis Structures
-Find total valence electrons
-Draw a skeleton structure using single bonds
-Assign remaining valence electrons
-Make multiple bonds if all the octets aren’t filled
-Lines indicate bonds, dots indicate electrons, square brackets if it has a charge, dotted line indicates resonance hybrid.
Resonance:
Resonance: When you have a bond that could go multiple places equally and are unsure where to put it.
Resonance hybrid:
When you use the dotted line.
Resonance structures:
You make lewis diagrams of all the possible structures and connect them with double sided arrows between them.
Bond order:
a lewis structure with one line has a bond order of 1, 2 lines 2, 3 lines 3, and so on. If there is resonance, the bond order could be 1.33.
Use formal charge to evaluate if a Lewis structure is valid
When there are multiple ways to structure a Lewis structure, we need to decide which one is most valid. We assign a formal charge to each atom, and the structure with the lowest formal charges is most correct.
Formal Charge:
Formal charge = (group number) - (number of covalent bonds )- (number of electrons in lone pairs (number of electrons visible)
We do this for every atom in the molecule, then add up the results. The sum of the formal charges must equal the total charge on the molecule.
If a lewis structure that doesn’t equal the total charge exists, the structure with the most negative or zero formal charges will be most right.
VSEPR:
3d adaptation of the lewis structure, used for determining and visualizing polarity.
PREDICT VSEPR:
Count number of connections to the central atoms, count number of lone pairs. (Ignore double bonds, count them as one single bond)
Lone pairs are sort of triangular blobs, connections are lines.
Linear:
-2 bonded atoms
-0 lone pairs
-180
Trigonal planar:
-3 bonded atoms
-0 lone pairs
-120
Bent:
-2 bonded atoms
-1 lone pair
-<120
Tetrahedral:
-4 bonded atoms
-0 lone pairs
-109.5
Trigonal pyramidal:
-3 bonded atoms
-1 lone pair
-107.3
Bent (2.0):
-2 bonded atoms
-2 lone pairs
-104.5
Trigonal bipyramidal:
-5 bonded atoms
-0 lone pairs
-120, 90
Butterfly/Seesaw:
-4 bonded atoms
-1 lone pair
-<120, 90
T-shaped (not actually t-shaped):
-3 bonded atoms
-2 lone pairs
-90
Linear (2.0):
-2 bonded atoms
-3 lone pairs
-180
Octahedral:
-6 bonded atoms
-0 lone pairs
-90
Square pyramidal:
-5 bonded atoms
-1 lone pairs
-90
Square planar:
-4 bonded atoms
-2 lone pairs
-90
Predict the hybridization of central atoms of simple molecules, including describing the shape of sp, sp2 and sp3 orbitals
Linear - sp
Trigonal planar & bent - sp2
Tetrahedral, Trigonal pyramidal & bent - sp3
Trigonal bipyramidal, seesaw/butterfly, t-shaped, linear - sp3d
Octahedral, square pyramidal, square planar - sp3d2
Describe sigma and pi bonds, and explain which type may exist for a given molecule
Sigma - simple overlap of orbital resulting in the sharing of 2 electrons is a sigma bond
Pi - area of electron density, double bonds that result from the overlap of one unhybridized orbital per atom, double bonds are sp2, triple bonds are sp
The 3 types of intermolecular forces
-Dipole-dipole bonds, negative end attracts positive (very strong)
-Hydrogen bonds (like dipole-dipole bonds, just super polar so even stronger)
-London dispersion forces (weak, short burst of polarity due to instantaneous random formation of dipoles)
IMF affect on melting/boiling point
Stronger IMF, higher boiling point, higher melting point
why Zn is not considered a transition metal
In main block, incomplete d-subshell, ability to create stable actions with other atoms that have incomplete d subshells, creates coloured compounds, variable oxidation states, magnetic character, dense and hard, high boiling and melting points, catalytic activity
What is a ligand?
ligand is an ion or molecule that bonds to a central metal atom to form a coordinate complex.
Electromagnetic spectrum
Hot - Higher frequency/Shorter waves to long waves low frequency colder
Gamma--X-Rays--UV Rays--Visible light--Infrared Rays--Radar--FM--TV--Shortwave--AM
Bohr’s idea of electron shells, why we know they exist
-Bohr - the MVP connected the ideas of spectra to quanta, the idea that electrons only absorb energy at discrete levels, can give them energy to make them go up, then they fall and emit photons of the difference between energy levels.
The concept of quantum jumps
Give energy electron goes up, and then down, emits energy, lyman and Balmer series, 1 further from 2 than 2 is from 3, to infinity
How emission and absorption spectra work.
Elements get heated and give off light which can be separated into components on emission spectrum, each color pattern is distinct to the element, kind of like a fingerprint. - Called SPECTROSCOPY by Bunsen and Kirchkoff
Energy, frequency and wavelength formula
c = λν
c = 3.00 x 108 m/s
λ = wavelength (m)
ν= frequency (s-1)
Small λ = bigν
quanta
energy packets theorized by planck
Energy and frequency of light formula
E = hν
E = energy (joules, J)
h = Planck’s constant = 6.626 x 10-34 J s
v = frequency (s-1)
Bohr
Bohr connected the line spectra to the idea of quantized energy – that electrons only absorb energy of certain wavelengths
Bohr thought that electrons could only possess certain amounts of energy
From this came the planetary model of the atom with electrons existing only at discrete, energy levels
Electrons did not lose energy while staying in those levels
Balmer series
6-2 (visible light spectrum)
Lymen series
6-1
Convergence limit
You need to know the data for the convergence limit to calculate ionization energy and vise versa
S, P, D and F subshells
1 lobe, 2 lobes, 4 lobes, 8 lobes, funky ballon shaped items
Double slit experiment
Shows electrons can be both waves and particles
Heisenburg Uncertainty principle
We can never know both the location and speed of a particle at the same time
Aufbau principle
Electrons fill lower levels of subshells first and the go to higher ones
Pauli exclusion principle
within an orbital the two electrons are spining opposite ways
Hunds rule
You add the up arrows first, three of them for the p orbital, before going to the down ones
Subshell anomalies
unexpected configurations in the d block (transition metals)
Cr and Cu are evil (4s1 instead of filling it completely)
Mendeleev (period table contribution?)
Arranged elements in order of chemical properties and then by mass, pedicted the devlopment of undisocvered elements and also data for them
-1872
Atomic radius
big at the bottom left, tiny at the top right
Ionization energy trend
Top right (halogens) is easiest to ionize, bottom left is hardest
Electron affinity (opposite of ionization energy, add one)
same as ionization energy weirdly
Electronegativity
same as ionization energy, accept final line is very low
reactivity trend
Same as atomic radii, francium is ready to roll
Ionic crystals
metal and non-metal
crytsal lattice
hard and brittle
conducts electricty
ionic character
soluble in water
Molecular crytsals
non metal compound
weaker intermolecular forces
soft
non conductor
low melting point
insoluble
Giant covalent network crystals
hard
brittle
high mleting point
insoluble
Metallic crystals
Low ionization energy
shiny
flexible
conductive
Fomula for paticles
Number of particles = moles (avagdros number)
Formula for molar mass
mass = number of moles (molar mass)
Percent composition
(% mass of element)= mass of element/ mass of compound (100)
IMPORTANT: Assume you have 1 mol of the compound!
Find the molar mass (M) of the compound
Find the molar mass (M) of each element within the compound
Use the equation above to find the % composition of each element
Empirical formula
Simplest formula
Molecular formula
Actual formula
Determining empirical formulas
IMPORTANT: Assume the total mass of the sample is 100 g. This will mean that the % of each element = its mass (m) in grams!
Determine the # of moles (n) of each element - use n = m / M!
Determine the simplest ratio between elements by dividing each mole amount by the smallest mole amount.
Determining molecular formulas
NOTE:
You will need the empirical formula AND the molar mass of the molecular formula (MFM)!
Determine the molar mass of the empirical formula (MEF) - use the PT!
Determine the mass multiple by dividing the MMF by the MEF.
Multiple each subscript in your empirical formula by the mass multiple to determine the molecular formula.
Percent yield
% Yield = actual yield/theoretical yield (100)
Percent error
% Error = (actual yield - theoretical yield )/theoretical yield x 100 %
Atom economy
% atom economy = molar mass of yield/ molar mass of all reactants
Concentration formula
c1v1 = c2v2
Gas laws (all of them!!!!!!!)
1 atm = 760 mm Hg
1 atm = 101.325 kPa
P1V1 = P2V2 (Boyles law)
V1/T1 = V2/T2 (charles law)
P1V1/T1=P2V2/T2
P1/T1=P2/T2 (Gay-Lussacs law)
Gas assumptions
-The particles themselves have no volume
-There are no IMF between the particles
-The gas will not condense under pressure
-Relationships of pressure, temperature and
volume are perfectly linear
-Gas volume at absolute zero is zero
Ideal gas law
PV = nRT
R = proportionality constant = 8.31 dm3 kPa Κ−1 mol−1
P = pressure in kPa
V = volume in dm3
n = moles
T = temperature in Kelvin
STP
P = 100 kPa
T = 273 K
The molar volume of an ideal gas is 22.7 dm3
mol-1 at STP
SATP
P = 100. kPa
T = 298 Κ
The molar volume of an ideal gas is 24.8 dm3
mol-1 at SATP
Organic starters
Ethy, methy, prop, but, pent, hex, hept, oct, non
Organic chemistry
Alkane, ane, plain
Alkene, ene, double
Alkyne, yne, triple
alcohol, anol, oh
Aldehyde, anal, O (only on one end)
Ketone, one, O double bond
Ether, ane, O in the middle
Ester, ate, O double bond plus O in the middle
Carboxylic acid, oic acid, O and OH
Amines, amine, NH3
Aromatic, benzene, benzene