Chem units 1-7 midterm review

Properties of ionic solids

Soluble in polar solvents

Conduct electricity only when molten or dissolved in a polar solvent

High melting points

Very hard

Low volatility

Brittle

Ions line up in a repetitive pattern that maximize attractive forces

Not mailable or ductle

Properties of molecular solids

Molecular solids do not conduct electricity

Individuals molecules have no net charge

Molecular solids are held together by IMFs

High VP

Low melting and boiling points

Vapor pressure

Rate increases as temp increases

Weaker IMFs make things evaporate quicker

Boiling point

A liquid could when it’s VP is equal to atmospheric pressure

sublimation

Solids can evaporate and they have very low VP because their IMFs are so strong

Covalent network solid

Always composed of one or two nonmetals

Carbon group elements often form covalent covalent network solids as they can from four covalent bonds

The highest melting points and normally very hard, as atoms are covalent bonded with fixed angles

Diamonds

Many carbon atoms bound together with sp³

Each carbon mass a single covalent bond with 4 other carbon atoms

Very hard and high melting point

Graphite

Each carbon forms sp² that bond with 3 other carbon atoms

Form sheet like structures linked by \pi bonds and LDFs

Electrons will flow through \pi bonds

High melting point

Quartz (SiO2)

Every silicon atom is covalent bonded with 4 O atoms

Every O is covalent bonded with 2 Si atoms

Properties of synthetic polymers

Plastics are generally flexible solids or viscous liquids

Heating plastic increases flexibility and allows them to be molded

Metallic solids

Bonding is not covalent

Bonding results from the attractions between nuclei and delocalized valence electrons moving throughout the structure

Bond strength increases as the number of bonding electrons increases

Saturated solution

• When the solvent has dissolved the maximum amount of solute possible at a certain temperature, and some solid particles remain undissolved.

• This is an equilibrium system where solid particles continually dissolve in the solvent and dissolved particles fall out of solution.

Liquid-liquid solutions

Only mix if they both have the same polarity

Miscible solutions never become saturated.

Differences in intermolecular forces can cause the solution's volume to differ from the sum of the volumes before mixing.

Solids-liquid

• Many ionic compounds dissolve in polar solvents. (ion-dipole).

• Polar solids, dissolve in polar solvents.

• Non-polar solids, dissolve in non-polar solvents.

Gas-gas

Gases are infinity soluble with each other

Gas-solid

Gases can occupy the spaces between some metal

Light formula

c=\lambda\upsilon

Planks equations

E=h\upsilon

The photoelectric effect

intense low frequency light dont eject any electrons, even if it shines for some time

When the threshold frequency is reached, electrons are ejected immediately.

Increasing the intensity of the light will increase the rate of ejection However, all ejected electrons share the same velocity.

Increasing the frequency of the light increases the velocity of the electrons.

Electromagnetic Spectrum

is a Continuous Spectrum of Light

Atomic emission spectrum

Every element release specific lights

Ion-dipole solubility

• Some ionic compounds do not dissolve in water

• If cation-anion attractions are stronger than ion-dipole attraction, the compound will not be soluble.
  Ionic compounds do not dissolve in non-polar solvents, as non-polar solvents do not carry permanent dipoles.

Fractional distillation

The separation of volatile liquids in a liquid-liquid solution on the basis of boiling points.

The condensed solution has a higher concentration of the component with the higher vapor pressures

Gas solubility and temperature

The solubility of most gases decreases as temperature increases.

• As the kinetic energy of particles within a solution increases, aqueous particles break free from these weak attractions and re-enter the gas phase.

Henry’s law

The solubility of a gas is directly proportional to the partial pressure of that gas above the solution.

UV/vis spectroscopy

Shining these types of lights cause electrons to be excited to go to a different orbital, releasing light. This allows us to identify the element.

Absorption spectrum

The peaks of this light absorption spectrum represent the wavelengths of light that correspond to the energy of an electron transitions from ground to excited state orbitals.

The tallest peaks represents the wavelength most absorbed by the electrons

Beer-lambert law (beers law) equation

Used to find concentration graph

A = \xi bi

A = Absorbance how much light was absorbed by the solute

\xi = absorbtivity, slope (m)

b= length of solution (X)

C= length of solution pathway

Infrared spectroscopy

Examines the Vibrations of the bonds. All covalent bonds in molecules are vibrating

detects the presence of different types of bonds so as to identify molecules.

Vibrations and infrared

covalent bonds have a vibrational frequency that corresponds to the frequency of light in the infrared light spectrum.

When this exact IR frequency is absorbed by the molecule, the atoms vibrate more rapidly

Vibrate in all directions

Infrared use

IR spectroscopy is used to identify:

bong types

Each atom has a specific viberation that can be used to identify it

Microwave spectroscopy

This absorbed light changes the rotations bonded atoms. This Tells us the Locations molecule.

Solids states of motions

Solids can only vibrate

Liquids and gases in motions

They can rotate vibrate, and transition (move from one place to another)

Boron specialness

Doesn’t need a complete octet

Expanded octets

atoms between period 3-7 can expand an and bond with more than 8 electrons because of Ds

Formal charge

Valence electrons - electrons assigned = formal charge

VSEPR

Valence shell electron pair repulsion theory

Charge clouds repel each other

Terminal atoms move as far away from each other as possible

Results in distinctive shapes

Linear

2 charge clouds, 2 bonds, 0 lone pairs

180°

Triangle planar

3 charge clouds, 3 bonds, 0 lone pairs

120°

Bent

3 charge clouds, 2 bonds, 1 lone pair

120°

Tetrahedral

4 charge clouds, 4 bonds, 0 lone pairs

109.5°

Trigonal pyramidal

4 charge clouds, 3 bonds, 1 lone pairs

109.5°

Super bent

4 charge clouds, 2 bonds, 2 lone pairs

109.5°

Trigonal bipyramidal

5 charge clouds, 5 bonds 0 lone pairs

120°

Seesaw

5 charge clouds, 4 bonds, 1 lone pair

T-shaped

5 charge clouds, 3 bonds, 2 lone pairs

90°

Super linear

5 charge clouds, 2 bonds, 3 lone pairs

180°

Octahedral

6 charge clouds, 6 bonds, 0 lone pairs

90°

Square pyramidal

6 charge clouds, 5 bonds, 1 lone pair

90°

Square planar

6 charge clouds, 4 bonds, 2 lone pairs

90°

Mass percent formula

X = ((# of atoms)(element atomic mass))/formula weight of compound

Mass % in mixture

(Mass of compound)/total mass of substance

Empirical formula steps to solve

Assume 100 gram sample and convert grams to moles

Divide all moles by the smallest mole count of the elements

Coulombs Law equation

F = k(q1q2)/d²

F=force of attraction

K= constant

q = magnitude of charge with a particle - electrons and protons

d = distance

Coulomb’s law explanation

The force of attraction decreases as the distance between the outermost electron and the protons increases

Valence electron location

Valence electrons are located on the outermost shell of an atom

Shielding effect

Electrons farthest away from the nucleus is partially shielded by the inner core electrons due to repulsion

This reduces the electrostatic attraction between the nucleus and the outer electrons

First ionization energy

The minimum amount of energy that is required to remove an outermost, least tightly held, electron from an atom in gas state

First ionization energy periodic trends

As shells count decreases the more energy is require

As valance electrons increase the more energy is required

Electron configuration d’s and f’s knowledge

The d’s are always one shell down, and the f’s are always two shells down

ex: (U) 1s² 2s² 2p⁶ 3s² 3p⁶ 4s² 3d¹⁰ 4p⁶ 5s²4d¹⁰ 5p⁶ 6s² 4f¹⁴ 5d¹⁰ 6p⁶ 7s² 5f³ 6d¹

Effective nuclear charge

Z(eff) = Z - sigma

Z(eff) - the charge experienced by an electron

Z - the actual nuclear charge (atomic number of element)

Sigma - shielding constant (0<sigma<Z)

Repulsive forces caused by shielding effect reduce the effective nuclear charged by outer electrons

Atomic radius periodic trend

Increases as shells increase

Decreases as proton count increases

Ionic radius cations

Cations are smaller than neutral atoms

Ionic radius anions

Anions are larger than neutral atoms

Electron affinity

The energy change that occurs when an electron is added to a gaseous atom to form a negative ion

it is a measure of how much an element wants to accept another electron