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MCAT Gen. Chem(David) > General Chemistry > Flashcards

General Chemistry Flashcards

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1
Q

Periodic table families - show

A
  1. Alkali metals
  2. Alkali earth metals
  3. Halogens
  4. Noble gases

d-block Transition elements

4f Lanthanide

5f Actinide

LINK

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2
Q

Ionisation energy - define and graph

A

Energy required to REMOVE the most loosely held electron of 1 mol of gaseous atoms

LINK

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3
Q

Periodic table characteristics (2 of 2)

halogen family

noble gases

A

halogen family - most reactive nonmetals,

never found uncombined in nature,

halogens combined with metals form salts



noble gases - normally unreactive,

also called inert gases,

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4
Q

Periodic table

  1. Name the elements 19-36
A

table

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5
Q

Reaction types

Name and explain in symbols (4)

A

Combination A + B -> C

Decomposition C -> A + B

Single displacement A + BC -> B + AC

Double displacement AB + CD -> AD + CB

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6
Q

Atomic orbitals

Name the orbitals

how many electrons they hold

where they appear on the periodic table

A

s p d f

s 2

p 6

d 10

f 14

s groups 1 and 2

p groups 3 - 8

d transition elements

f lanthanide and actinide series

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7
Q

Draw a table and classify salts as acidic, neutral or basic

Strong parent acid weak parent acid

Strong parent base

Weak parent base

A

??? means if Ka > Kb then acidic salt

Kb > Ka then basic salt

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8
Q

Redox

  1. What is the difference between a galvanic cell and an electrolytic cell?
A

Galvanic: turns chemical energy into electrical energy

Electrolytic Cell: turns electrical energy into chemical energy

The most common form of Electrolytic cell is the rechargeable battery (cell phones, mp3’s, etc) or electroplating.

While the battery is being used in the device it is a galvanic cell function (using the redox energy to produce electricity).

While the battery is charging it is an electrolytic cell function (using outside electricity to reverse the completed redox reaction).

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9
Q

Le Chatellier

  1. What effect does adding an inert gas to a gas reaction have?
A

padlet

  1. None - doesn’t change the partial pressures of the reactants and products so is not a disturbance
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10
Q

Effusion - define

effusion rates for mix of gas A and B - formula

Diffusion - define

Diffusion rates for mix of gas A and B - formula

A

Effusion is the process in which a gas escapes through a small hole

Diffusion - is the process by which molecules intermingle as a result of their kinetic energy of random motion.

Rates are the same for effusion and diffusion - see graphic and beware of inverse

A

A

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11
Q

Acid

  1. Explain the terms conjugate acid and conjugate base
A
  1. Acids always react with bases so in an equation both sides have acids and bases

The LHS reactants are called acid and base

The RHS products are called the conjugate acid and conjugate base

An acid base reaction involves the donation of a proton (H+) from an acid to a base.

The species which loses H+ is the acid
The species which gains H+ is the base
The conjugate base is what becomes of the acid after it loses H+
The conjugate acid is what becomes of the base after it gains H+

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12
Q

Chemical yield

Formula

A

% Yield = Actual yield / Theoretical yield

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13
Q

Redox

  1. Is a rechargable bettery a galvanic cell or an electrolytic cell?
A

While the battery is being used in the device it is a galvanic cell function (using the redox energy to produce electricity).

While the battery is charging it is an electrolytic cell function (using outside electricity to reverse the completed redox reaction).

Galvanic: turns chemical energy into electrical energy

Electrolytic Cell: turns electrical energy into chemical energy

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14
Q

Acid

  1. Name seven strong bases for MCAT
A

hydride ion H-

amide ion NH2-

sodium oxide Na2O

calcium oxide CaO

sodium hydroxide NaOH

potassium hydroxide KOH

calcium hydroxide Ca(OH)2

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15
Q

Solutions - units of concentration

  1. moles of solute / volume of solution =
  2. moles of solute / kg of solvent =
  3. moles of solute / (total moles of all solutes & solvent) =
  4. 100 x mass of solute / total mass of solution =
  5. 106 x mass of solute / total mass of solution =
A
  1. Molarity M = moles of solute / volume of solution
  2. Molality m = moles of solute / kg of solvent
  3. Mole fraction χ = moles of solute / (total moles of all solutes & solvent)
  4. Mass percentage mass % = 100 x mass of solute / total mass of solution
  5. ppm = 106 x mass of solute / total mass of solution
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16
Q

Redox

Rule 4: The oxidation number of an alkali metal (IA family) in a compound = ..?..

the oxidation number of an alkaline earth metal (IIA family) in a compound = ..?..

A

Rule 1: The oxidation number of an element in its free (uncombined) state is zero — for example, Al(s) or Zn(s). O2

Rule 2: The oxidation number of a monatomic (one-atom) ion = the charge on the ion, for example: Na+ = +1

Rule 3: The sum of all oxidation numbers in a neutral compound =zero.

The sum of all oxidation numbers in a polyatomic ion = the charge on the ion

This rule often allows chemists to calculate the oxidation number of an atom that may have multiple oxidation states, if the other atoms in the ion have known oxidation numbers.

Rule 4: The oxidation number of an alkali metal (IA family) in a compound is +1;

the oxidation number of an alkaline earth metal (IIA family) in a compound is +2.

Rule 5: The oxidation number of oxygen in a compound is usually –2. peroxides (for example, hydrogen peroxide), then the oxygen has an oxidation number of –1.

If the oxygen is bonded to fluorine, the number is +1.

Rule 6: The oxidation state of hydrogen in a compound is usually +1.

If the hydrogen is part of a binary metal hydride (compound of hydrogen and some metal), then the oxidation state of hydrogen is –1.

Rule 7: The oxidation number of fluorine is always –1. Chlorine, bromine, and iodine usually have an oxidation number of –1, unless they’re in combination with an oxygen or fluorine.

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17
Q

Draw a Phase diagram for water

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18
Q

Redox

Rule 7: The oxidation number of fluorine is always ..?..

Chlorine, bromine, and iodine usually have an oxidation number of ..?..

unless they’re in combination with an oxygen or fluorine.

A

Rule 1: The oxidation number of an element in its free (uncombined) state is zero — for example, Al(s) or Zn(s). O2

Rule 2: The oxidation number of a monatomic (one-atom) ion = the charge on the ion, for example: Na+ = +1

Rule 3: The sum of all oxidation numbers in a neutral compound =zero.

The sum of all oxidation numbers in a polyatomic ion = the charge on the ion

This rule often allows chemists to calculate the oxidation number of an atom that may have multiple oxidation states, if the other atoms in the ion have known oxidation numbers.

Rule 4: The oxidation number of an alkali metal (IA family) in a compound is +1;

the oxidation number of an alkaline earth metal (IIA family) in a compound is +2.

Rule 5: The oxidation number of oxygen in a compound is usually –2. peroxides (for example, hydrogen peroxide), then the oxygen has an oxidation number of –1.

If the oxygen is bonded to fluorine, the number is +1.

Rule 6: The oxidation state of hydrogen in a compound is usually +1.

If the hydrogen is part of a binary metal hydride (compound of hydrogen and some metal), then the oxidation state of hydrogen is –1.

Rule 7: The oxidation number of fluorine is always –1. Chlorine, bromine, and iodine usually have an oxidation number of –1, unless they’re in combination with an oxygen or fluorine.

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19
Q

VSEPR

  1. What molecular geometries are predicted with five electron dense areas?

(include all possible lone pairs options)

A
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20
Q

Redox

Oxidation is what ..?.. would have done if it was present

A

Oxidation is what oxygen would have done if it was present

so in H2O hydrogen is oxidised

in HF hydrogen is oxidised

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21
Q

Redox

    1. What does a positive standard reduction potential indicate?
A

The standard reduction potential is the tendency for a chemical species to be reduced, and is measured in volts at standard conditions. The more positive the potential is the more likely it will be reduced.

It is written in the form of a reduction half reaction. eg Cu2+ + 2e- -> Cu

Standard reduction or oxidation potentials can be determined experimentally using a SHE (standard hydrogen electrode).

Universally, hydrogen has been recognized as having reduction and oxidation potentials of zero. Therefore, when the standard reduction and oxidation potential of chemical species are measured, it is actually the difference in the potential from hydrogen. By using a galvanic cell in which one side is a SHE, and the other side is half cell of the unknown chemical species, the potential difference from hydrogen can be determined using a voltmeter.

Standard reduction potentials are used to determine the standard cell potential. The standard reduction cell potential and the standard oxidation cell potential can be combined to determine the overall Cell Potentials of a galvanic cell

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22
Q

VSEPR

  1. What molecular geometries are predicted with six electron dense areas?

(include all possible lone pairs options)

A
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23
Q

Redox

  1. What is the relationship between cell potential (E) and the equilibrium constant (K)?
A

nF E∘cell= RT lnK

We can use the relationship between ΔG∘ and the equilibrium constant K, to obtain a relationship between E∘cell and K.

Recall that for a general reaction of the type aA+bB→cC+dD, the standard free-energy change and the equilibrium constant are related by the following equation:

ΔG°=−RTlnK

Given the relationship between the standard free-energy change and the standard cell potential ΔG∘=−nFE∘cell , we can write

−nFE∘cell=−RTlnK

Rearranging this equation gives above

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24
Q

Redox

Rule 2: The oxidation number of a monatomic (one-atom) ion = ..?..

for example: ..?..

The sum of all oxidation numbers in a polyatomic ion = ..?..

A

Rule 1: The oxidation number of an element in its free (uncombined) state is zero — for example, Al(s) or Zn(s). O2

Rule 2: The oxidation number of a monatomic (one-atom) ion = the charge on the ion, for example: Na+ = +1

Rule 3: The sum of all oxidation numbers in a neutral compound =zero.

The sum of all oxidation numbers in a polyatomic ion = the charge on the ion

This rule often allows chemists to calculate the oxidation number of an atom that may have multiple oxidation states, if the other atoms in the ion have known oxidation numbers.

Rule 4: The oxidation number of an alkali metal (IA family) in a compound is +1;

the oxidation number of an alkaline earth metal (IIA family) in a compound is +2.

Rule 5: The oxidation number of oxygen in a compound is usually –2. peroxides (for example, hydrogen peroxide), then the oxygen has an oxidation number of –1.

If the oxygen is bonded to fluorine, the number is +1.

Rule 6: The oxidation state of hydrogen in a compound is usually +1.

If the hydrogen is part of a binary metal hydride (compound of hydrogen and some metal), then the oxidation state of hydrogen is –1.

Rule 7: The oxidation number of fluorine is always –1. Chlorine, bromine, and iodine usually have an oxidation number of –1, unless they’re in combination with an oxygen or fluorine.

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25
Q

Acid

  1. What is an amphoteric substance?
  2. Give an example
A
  1. Can act as either an acid or a base depending upon its environment
  2. Water
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26
Q

Redox

What is the Effect of Concentration (Q) on Cell Potential (Ecell and E0cell)

The Nernst Equation

A

The Effect of Concentration on Cell Potential:

The Nernst Equation - arguably the most important relationship in electrochemistry

nF Ecell= nF E∘cell −RT lnQ

Recall that the actual free-energy change for a reaction under nonstandard conditions, ΔG, is given as follows:

ΔG=ΔG°+RTlnQ

We also know that ΔG = −nFEcell and ΔG° = −nFE°cell.

Substituting and dividing both sides of this equation by −nF we obtain above

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27
Q

Ionic naming

Anions

A

1- …ide

With OXYGEN

Cl O- ….hypo chlor ite

Cl O2- ….chlor ite

Cl O3- …chlor ate

Cl O4- …per chlor ate

With HYDROGEN

H CO3-hydrogen carbonate

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28
Q

Redox

  1. How is standard reduction potential determined?
  2. What is it used for?
A

The standard reduction potential is the tendency for a chemical species to be reduced, and is measured in volts at standard conditions. The more positive the potential is the more likely it will be reduced.

It is written in the form of a reduction half reaction. eg Cu2+ + 2e- -> Cu

Standard reduction or oxidation potentials can be determined experimentally using a SHE (standard hydrogen electrode).

Universally, hydrogen has been recognized as having reduction and oxidation potentials of zero. Therefore, when the standard reduction and oxidation potential of chemical species are measured, it is actually the difference in the potential from hydrogen. By using a galvanic cell in which one side is a SHE, and the other side is half cell of the unknown chemical species, the potential difference from hydrogen can be determined using a voltmeter.

Standard reduction potentials are used to determine the standard cell potential. The standard reduction cell potential and the standard oxidation cell potential can be combined to determine the overall Cell Potentials of a galvanic cell

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29
Q

Acid

  1. How do polyprotic acids differ from monoprotic acids in titration?
A
  1. They will have more than one equivalence point and half equivalence point
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30
Q

Model thermodynamic transformations

There are 4 processes:-

name them

formula (From the first law of thermodynamics)

describe terms in the formula

A

The first law of thermodynamics ΔU=Q+Wwork done <u><strong>on </strong></u>the system

  1. Isothermal - Constant temperature and thus constant internal energy, Heat flow is compensated by work

ΔU = 0 Q = -W

  1. Adiabatic - no heat flow, internal energy change = work performed

Q = 0 ΔU = -W

  1. Isovolumetric - no work occurs, internal energy change = heat flow

W = 0 ΔU = Q

  1. Isobaric - Constant pressure

W = P ΔV ΔU = Q - P ΔV

The first law of thermodynamics

It is typical for chemistry texts to write the first law as ΔU=Q+W. It is the same law, of course - the thermodynamic expression of the conservation of energy principle. It is just that W is defined as the work done <strong>on</strong> the system instead of work done <strong>by </strong>the system.

n the context of physics, ΔU=Q-W the common scenario is one of adding heat to a volume of gas and using the expansion of that gas to do work, as in the pushing down of a piston in an internal combustion engine. In the context of chemical reactions and process, it may be more common to deal with situations where work is done on the system rather than by it.

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31
Q
  1. Name all the six strong acids

(all other acids are weak)

A

HCl

HBr

HI

HNO3

HCLO4

H2SO4

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32
Q

Gibbs free energy implications for reactions

G

G = 0

G > 0

A

G

G = 0 equilibrium

G > 0 Not spontaneous (needs a bit of energy put in to work

The Gibbs free energy is defined as: G(p,T) = U + pV - TS

which is the same as: G(p,T) = H - TS

where: U is the internal energy (SI unit: joule) p is pressure (SI unit: pascal) V is volume (SI unit: m3) T is the temperature (SI unit: kelvin) S is the entropy (SI unit: joule per kelvin) H is the enthalpy (SI unit: joule)

Gibbs energy is a thermodynamic potential that measures the “usefulness” or process-initiating work obtainable from a thermodynamic system at a constant temperature and pressure (isothermal, isobaric).

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33
Q

Redox

  1. What is the relationship between the potential of an electrochemical cell (E) and ΔG (the change in free energy)?
A

ΔG0 = -nF E0cell

The maximum amount of work that can be produced by an electrochemical cell (wmax) is equal to the product of the cell potential (E∘cell) and the total charge transferred during the reaction (nF = moles of electrons x Coulombs per mole of electrons)

The change in free energy (ΔG) is also a measure of the maximum amount of work that can be performed during a chemical process (ΔG = wmax).

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34
Q

Reaction rate

define reaction rate

formula & explain all terms in the reaction A + B -> C

A

reaction rate = - Δ [reactants] / Δtime []=concentration

reaction rate = + Δ [products] / Δtime

k is the rate constant for the forward reaction

m is the order of reactant A

n is the order of reactant B

m + n is the overall order of the reaction

see other cards for graphs of the different orders - amount vs time

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35
Q

Acid

  1. Define polyprotic acids
  2. Define diprotic acids
A
  1. Can donate more than one proton
  2. Can donate just two protons
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36
Q

Redox

  1. What does an electrolytic cell require to initiate a redox reaction?
A

An electrolytic cellis a cell which requires an outside electrical source to initiate the redox reaction.

The process of how electric energy drives the nonspontaneous reaction is called electrolysis.

Whereas the galvanic cell used a redox reaction to make electrons flow, the electrolytic cell uses electron movement (in the source of electricity) to cause the redox reaction.

In an electrolytic cell, electrons are forced to flow in the opposite direction

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37
Q

Chemical bonding

name the bond types

describe them

list strongest to weakest

A

STRONG BONDS

  1. Metallic bonds
  2. Ionic bonds
  3. Covalent bonds

WEAK BONDS

  1. London dispersion forces - induced dipole-induced dipole attraction

The London dispersion force is a temporary attractive force that results when the electrons in two adjacent atoms occupy positions that make the atoms form temporary dipoles. London forces are the attractive forces that cause nonpolar substances to condense to liquids and to freeze into solids when the temperature is lowered sufficiently.

  1. Dipole-dipole attractions - electrostatic interactions between permanent dipoles in molecules
  2. Hydrogen bonds (specially strong form of dipole-dipole)
  3. Ion-dipole bonds

Bond type Dissociation energy (kcal/mol)[9]
Ionic Lattice Energy 250–4000 [10]
Covalent Bond Energy 30–260
Hydrogen Bonds 1–12 (about 5 in water)
Dipole–Dipole 0.5–2
London Dispersion Forces

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38
Q

Periodic table

  1. Name the first four Halogens
A

table

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39
Q

Gas laws

Boyles law - formula

Charles law - formula

The ideal gas law - formula

A

P1V1 = P2V2 constant temperature

V1/T1 = V2/T2 constant pressure

PV = nRT n is mol; R is universal gas constant

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40
Q

Sigma and pi bonding

explain what they are and when they occur

their relative strengths

A

sigma bond - direct s or p orbital overlap (in line)

pi bond - after a sigma bond; p orbital overlap; weaker than sigma

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41
Q

Acid

  1. What are the factors in molecular structure that affect acid strength (3)?
  2. And how?
A
  1. Strength of bond holding hydrogen to the atom - weaker means more likely to lose so more acidic

increasing acidity with electronegativity across the table

increasing acidity with atom size down the table

  1. Polarity of the bond - increasing polarity increases the strength of the acid
  2. The key factor in determining acidity is the stability of the conjugate base.

Any factor which makes the conjugate base more stable will increase the acidity of the acid.

Usually, it means stabilizing negative charge since the conjugate base will always be one unit of charge more “negative” than the acid.

First, by bringing the charge closer to the positively charged nucleus

Second, by spreading charge out over a larger volume

A conjugate base is what you obtain when you remove a proton (H+) from a compound. For instance, HO(-) is the conjugate base of water. O(2-) is the conjugate base of HO(-). Conversely, conjugate acids are what you obtain when you add a proton to a compound. The conjugate acid of water is H3O(+).

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42
Q

Enthalpy of Solution

Explain and give direction of change

A

Enthalpy of Solution

The heat generated or absorbed when a certain amount of solute is dissolved in a certain amount of solvent.

Dissolution by most gases is exothermic. That is, when a gas dissolves in a liquid solvent, energy is released as heat, warming both the system (i.e. the solution) and the surroundings.

Often the dissolved state is a lower energy state so the excess energy is given off as heat.

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43
Q

Acid base definitions

  1. Define an Arrhenius acid (oldest)
  2. Define a Bronsted Lowry acid
  3. Define a Lewis acid
A
  1. Aqueous solution only -

Arrhenius acid = anything that produces hydrogen ions in aqueous solution

base = hydroxide ions

  1. Bronsted Lewis acid = anything that donates a proton

base = accept

  1. Lewis acid = anything that accepts a pair of electrons

base = donate

(most general definition and need to be labelled Lewis acids specifically as they are not generally called acids)

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44
Q

Redox

Assuming the triangle between change in free energy (ΔG0)

cell potential (E0cell) and

the equilibrium constant (K) and the

  1. Show the values for reaction:

spontaneous in forward direction

spontaneous in reverse direction

no net reaction = equilibrium

A

Unfortunately, these criteria apply only to systems in which all reactants and products are present in their standard states, a situation that is seldom encountered in the real world

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45
Q

Acid

  1. water autoionization constant Kw =
A

The following equation describes the reaction of water with itself (called autoprotolysis):

H2O + H2O H3O+ + OH¯

The equilibrium constant for this reaction is written as follows:

Kc = ( [H3O+] [OH¯] ) / ( [H2O] [H2O] )

However, in pure liquid water, [H2O] is a constant value. To demonstrate this, consider 1000 mL of water with a density of 1.00 g/mL. This 1.00 liter (1000 mL) would weigh 1000 grams. This mass divided by the molecular weight of water (18.0152 g/mol) gives 55.5 moles. The “molarity” of this water would then be 55.5 mol / 1.00 liter or 55.5 M.

The solutions studied in introductory chemistry are so dilute that the “concentration” of water is unaffected. So 55.5 molar can be considered to be a constant if the solution is dilute enough.

Cross-multiplying the above equation gives:

Kc [H2O] [H2O] = [H3O+] [OH¯]

Since the term Kc [H2O] [H2O] is a constant, let it be symbolized by Kw, giving:

Kw = [H3O+] [OH¯] same as [H+] [OH¯]

This constant, Kw, is called the water autoprotolysis constant or water autoionization constant

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46
Q

What is the general method for determining molecular structure from the molecular formula?

(Hint - counting electrons to give four numbers)

A

Available electrons = (valence electrons)

Needed electrons = (to fill shells)

Shared electrons = (needed - available)

Bonds = (1/2 shared electrons)

Draw the bonds and fill in lone pairs as required

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47
Q

Redox

  1. What is the standard reduction potential?
  2. What are its units?
A

The standard reduction potential is the tendency for a chemical species to be reduced, and is measured in volts at standard conditions. The more positive the potential is the more likely it will be reduced.

It is written in the form of a reduction half reaction. eg Cu2+ + 2e- -> Cu

Standard reduction or oxidation potentials can be determined experimentally using a SHE (standard hydrogen electrode).

Universally, hydrogen has been recognized as having reduction and oxidation potentials of zero. Therefore, when the standard reduction and oxidation potential of chemical species are measured, it is actually the difference in the potential from hydrogen. By using a galvanic cell in which one side is a SHE, and the other side is half cell of the unknown chemical species, the potential difference from hydrogen can be determined using a voltmeter.

Standard reduction potentials are used to determine the standard cell potential. The standard reduction cell potential and the standard oxidation cell potential can be combined to determine the overall Cell Potentials of a galvanic cell

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48
Q

VSEPR

  1. What does VSEPR stand for?
  2. What is it used to predict?
A
  1. Valence shell electron pair repulsion theory
  2. The geometries or 3-D shapes of molecules
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49
Q

Dalton’s law of partial pressures

formula =

A

Ptotal = P1 + P2 ….

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50
Q

Redox

Rule 3: The sum of all oxidation numbers in a neutral compound = ..?..

The sum of all oxidation numbers in a polyatomic ion = ..?..

A

Rule 1: The oxidation number of an element in its free (uncombined) state is zero — for example, Al(s) or Zn(s). O2

Rule 2: The oxidation number of a monatomic (one-atom) ion = the charge on the ion, for example: Na+ = +1

Rule 3: The sum of all oxidation numbers in a neutral compound =zero.

The sum of all oxidation numbers in a polyatomic ion = the charge on the ion

This rule often allows chemists to calculate the oxidation number of an atom that may have multiple oxidation states, if the other atoms in the ion have known oxidation numbers.

Rule 4: The oxidation number of an alkali metal (IA family) in a compound is +1;

the oxidation number of an alkaline earth metal (IIA family) in a compound is +2.

Rule 5: The oxidation number of oxygen in a compound is usually –2. peroxides (for example, hydrogen peroxide), then the oxygen has an oxidation number of –1.

If the oxygen is bonded to fluorine, the number is +1.

Rule 6: The oxidation state of hydrogen in a compound is usually +1.

If the hydrogen is part of a binary metal hydride (compound of hydrogen and some metal), then the oxidation state of hydrogen is –1.

Rule 7: The oxidation number of fluorine is always –1. Chlorine, bromine, and iodine usually have an oxidation number of –1, unless they’re in combination with an oxygen or fluorine.

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51
Q

First law of thermodynamics

The first law of thermodynamics is a version of the law of ..??.. , adapted for ..??.. systems.

The first law of thermodynamics states two things - describe in words

Give the formula and describe what the terms mean

A

The first law of thermodynamics is a version of the law of conservation of energy, adapted for thermodynamic systems.

The law of conservation of energy states that the

1. total energy of an isolated system is constant;

  1. energy can be transformed from one form to another, but cannot be created or destroyed.

The first law is often formulated by stating that the change in the internal energy of a closed system is equal to the amount of heat supplied to the system, minus the amount of work done by the system on its surroundings.

deltaU = Q - W

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52
Q

Kinetic theory

formula for kinetic energy relating to temperature

formula for internal energy relating to temperature

what is gas molecule speed proportional to?

What does this relationship directly predict?

A

Kinetic energy = (3/2) kT k is Boltzmann’s constant

webpage LINK

Internal energy U = (3/2) nRT n is mol; R is universal gas constant

gas molecule speed proportional to (mass)-1/2

Important - speed-mass relationship at constant temperature directly predicts comparative efusion and diffusion rates

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53
Q

Redox

  1. Draw the triangle between cell potential (E0cell) and the equilibrium constant (K)

and the change in free energy (ΔG0)

[Include the three equations]

A

ΔG0 = -nF E0cell

nF E∘cell= RT lnK

ΔG0 = RT lnK

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54
Q
  1. If I know the Ka how do I find the Kb?
A
  1. Kw = 1.00x10-14 = Ka Kb
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55
Q
  1. What does putting a prefix p on a quantity indicate in chemistry?
A
  • log
    e. g. pH = -log [H+]

pKa = -log Ka

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56
Q

Acids

Naming convention

A

…ide –> hydro…ic H2S hydro sulf ic

with OXYGEN

H2 SO3 Sulfur ous acid SO32- Sulf ite ion

H2 SO4 Sulfur ic acid SO42- Sulf ate ion

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57
Q

Acid

  1. What is the equivalence point in a titration?
  2. What is the half equivalence point
A
  1. The equivalence point, or stoichiometric point, of a chemical reaction is the point at which chemically equivalent quantities of acid and base have been mixed. It can be found by means of an indicator, most often phenolphthalein.

The equivalence point on the graph is where all of the starting solution (usually an acid) has been neutralized by the titrant (usually a base). It can be calculated precisely by finding the second derivative of the titration curve and computing the points of inflection (where the graph changes concavity);

  1. When a reaction in titration reaches to the half way, at this point the value of equivalence point is termed as half equivalence point.

For example, the titration of 2M mono protic acid reach to the half-equivalence point when 1 M of the acid has been made into its corresponding base, and 1 M is still left by Stoichiometric. <br></br><br></br>The half-equivalence point is can be used to determine the acid dissociation and pKa of the acid used in titration. In acid-base titration the ratio between the acid and corresponding base is exactly 1:1 at the half-equivalence point.

Video webpage image LINK

58
Q

Collision model of reactions - explain including activation energy

A

The Effect of Molecular Orientation on the Reaction of NO and O3.

Most collisions of NO and O3 molecules occur with an incorrect orientation for a reaction to occur.

Only those collisions in which the N atom of NO collides with one of the terminal O atoms of O3 are likely to produce NO2 and O2, even if the molecules collide with

E > Ea activation energy

59
Q

Thermodynamics

zeroth law

A

The zeroth law of thermodynamics states that

if two thermodynamic systems are each in thermal equilibrium with a third, then they are in thermal equilibrium with each other.

The physical meaning of the law was expressed by Maxwell in the words: “All heat is of the same kind”

Two systems are said to be in the relation of thermal equilibrium if they are linked by a wall permeable only to heat, and do not change over time

60
Q

Reactions

  1. Arrhenius equation k =

from his direct observations of the plots of rate constants vs. temperatures

61
Q

Thermodynamics

Draw the four thermodynamic processes on a PV diagram - cyclic process

identify work

A

webpage LINK

note the adiabatic (no heat energy in or out) or isothermal (no change in temperature) can go in either direction

62
Q

What is the trend of Ka for polyprotic acids?

A

Ka for the second proton to be donated is much less than for the first and so on

63
Q

Acid

  1. How does the strength of an acid affect the strength of its conjugate base?
  2. How does the strength of a base affect the strength of its conjugate acid?
A
  1. The stronger the acid the weaker its conjugate base
  2. The stronger the base the weaker its conjugate acid
64
Q

chemical reaction rate constant

define the rate constant k =

how does k vary with [reactant] ?

draw graph for large k, medium k and small k

A

k = - (slope of reactant concentration with time graph) / [reactant]

k does not vary within a reaction

large k - very steep (big slope)

small k - shallow (small slope)

Video webpage LINK

65
Q

Chatelier’s principle

define

A

Any change in status quo prompts an opposing reaction in the responding system

When a system at equilibrium is subjected to change in concentration, temperature, volume, or pressure, then the system readjusts itself to (partially) counteract the effect of the applied change and a new equilibrium is established.

or whenever a system in equilibrium is disturbed the system will adjust itself in such a way that the effect of the change will be nullified. (in short)

This principle has a variety of names, depending upon the discipline using it (see homeostasis, a term commonly used in biology). It is common to take Le Châtelier’s principle to be a more general observation,[1] roughly stated:

In chemistry, the principle is used to manipulate the outcomes of reversible reactions, often to increase the yield of reactions.

66
Q

Acid

  1. Name seven strong acids for MCAT
A

hydroidoic acid HI

hydrobromic acid HBr

hydrochloric acid HCl

Nitric acid HNO3

Chloric acid HClO3

perchloric acid HClO4

Sulfuric acid H2SO4

67
Q

Enthalpy

Hess’s law - state

A

The total enthalpy change during the complete course of a reaction is the same whether the reaction is made one one step or several steps

68
Q
  1. What is the difference between Binary acids and oxyacids?
A

Binary acids are two elements e.g HCl

Oxyacids contain oxygen e.g HClO4

69
Q

Living cell reactions

  1. What do many reactions in living cell involve?
  2. What does the rate of such reactions depend upon?
A
  1. Transfer of a proton
  2. pH [H+]
70
Q

Redox

Rule 6: The oxidation state of hydrogen in a compound is usually ..?..

If the hydrogen is part of a binary metal hydride (compound of hydrogen and some metal), then the oxidation state of hydrogen is ..?..

A

Rule 1: The oxidation number of an element in its free (uncombined) state is zero — for example, Al(s) or Zn(s). O2

Rule 2: The oxidation number of a monatomic (one-atom) ion = the charge on the ion, for example: Na+ = +1

Rule 3: The sum of all oxidation numbers in a neutral compound =zero.

The sum of all oxidation numbers in a polyatomic ion = the charge on the ion

This rule often allows chemists to calculate the oxidation number of an atom that may have multiple oxidation states, if the other atoms in the ion have known oxidation numbers.

Rule 4: The oxidation number of an alkali metal (IA family) in a compound is +1;

the oxidation number of an alkaline earth metal (IIA family) in a compound is +2.

Rule 5: The oxidation number of oxygen in a compound is usually –2. peroxides (for example, hydrogen peroxide), then the oxygen has an oxidation number of –1.

If the oxygen is bonded to fluorine, the number is +1.

Rule 6: The oxidation state of hydrogen in a compound is usually +1.

If the hydrogen is part of a binary metal hydride (compound of hydrogen and some metal), then the oxidation state of hydrogen is –1.

Rule 7: The oxidation number of fluorine is always –1. Chlorine, bromine, and iodine usually have an oxidation number of –1, unless they’re in combination with an oxygen or fluorine.

71
Q

chemical reaction order

how is the order determined

A

Experimentally

vary the quantity of one reactant (at a time) and see how the reaction rate varies

if the rate doubles - order = 1

if the rate x4 - order = 2

if the rate x9 - order = 3

etc

72
Q
  1. What is the difference between strong acids and weak acids in terms of the reaction arrow used?
  2. And what does that mean?
A

Strong acids use a single arrow to reflect that they dissociate completely in solution

Weak acids use a double arrow to reflect that they partially dissociate so it is an equilibrium reaction

73
Q

Periodic table characteristics (1 of 2)

alkali family

alkaline earth

transition metals

rare earth elements

A

alkali family - group 1, hydrogen is not a member.

good conductors of heat and electricity, most reactive metals, never found in nature uncombined

alkaline earth - group 2, also very reactive

transition metals - group 3-12, largest group on periodic table, good conductors of heat and electricity, compounds with these elements are usually brightly colored,

rare earth elements - Lanthanoids- soft malleable metals, high luster and conductivity

Actinoids- radioactive

74
Q

Molecular geometry

  1. What is an electron domain?
  2. How is it used?
A

single bond = 1 electron domain

Lone pair = 1 electron domain

Double bond = 1 electron domain

Triple bond = 1 electron domain

  1. Electrons in a single domain cannot be separated spatially => bond angles
75
Q

Colligative

  1. osmotic pressure (pressure to stop osmosis) pi =
A

Pi = i M R T

M is molarity

R is the gas constant and we will be using the same value as in the gas laws unit: 0.08206 L atm/mol K.

i is the van’t Hoff factor equal to the number of particles the solute
dissociates into. For nonelectrolytes, i = 1. For electrolytes, i = number of ions per
formula.

T is the thermodynamic (absolute) temperature

PV=nRT gives

76
Q

Activation energy

define

draw graph for exothermic reaction

draw graph for endothermic reaction

draw graph for reaction with catalyst

A

Activation energy - the minimum energy which must be available to a chemical system with potential reactants to result in a chemical reaction.

Catalyst lowers the activation energy (lower peak)

77
Q

Buffer solution

  1. What are the two ways of making a buffer solution?
A

Weak acid and its conjugate base

Weak base and its conjugate acid

Key is that there is a substantial amount of acid and base available at equilibrium to soak up any added acid or base

78
Q
  1. What does ICE mean?

2, With what sort of reactions can you use ICE tables?

A

Initial concentration

Change concentration

Equilibrium concentration

  1. Equilibrium reactions
79
Q

Reaction order

  1. zeroth order reaction rate =
  2. First order reaction rate =
  3. Second order reaction rate =
A
80
Q

Acid

  1. What is the endpoint in a titration?
A
  1. The endpoint (related to, but not the same as the equivalence point) refers to the point at which the indicator changes color in a colorimetric titration.
81
Q

chemical equilibrium define

reaction quotient Q define

reaction direction - equilibrium constant K vs Q

A

In a chemical reaction, chemical equilibrium is the state in which both reactants and products are present in concentrations which have no further tendency to change with time.

Usually, this state results when the forward reaction proceeds at the same rate as the reverse reaction.

The equilibrium constant is a ratio of the concentration of the products to the concentration of the reactants. If the K value is less than one the reaction will move to the left and if the K value is greater than one the reaction will move to the right

82
Q

Le Chatellier

  1. When volume is increased in a gas reaction -

which side of the reaction is favoured?

A

padlet

  1. The side with more moles of gas
83
Q

Bond energy

Draw graph

A
84
Q

Ionic naming

Cations

A

The preferred method is to use the metal name followed in parentheses by the ionic charge written as a Roman numeral: Iron(III).

But an older naming method, which is still in use, is to use -ous and -ic endings. The ion with the lower oxidation state (lower numerical charge, ignoring the + or -) is given an -ous ending, and the ion with the higher oxidation state (higher numerical charge) is given an -ic ending.

Element Cation Preferred Name Other Name
copper Cu+ copper(I) cuprous
Cu2+ copper(II) cupric
iron Fe2+ iron(II) ferrous
Fe3+ iron(III) ferric
lead Pb2+ lead(II) plumbous
Pb4+ lead(IV) plumbic
mercury Hg22+ mercury(I) mercurous
Hg2+ mercury(II) mercuric
tin Sn2+ tin(II) stannous
Sn4+ tin(IV) stannic

1+ …ous

2+ …ic

Non-metal …ium eg ammonium

85
Q

Hydride

  1. What is a hydride?
  2. How does their acidity change in the periodic table?
A
  1. A binary molecule containing hydrogen
  2. Acidity increasing left to right; and down
86
Q

Le Chatellier

  1. When the temperature is increased in an exothermic reaction - which side of the reaction is favoured?
A

padlet

  1. Reactants
87
Q

Redox

Reduction is a reduction of ..?..

(ie going more ..?..)

A

Reduction is a reduction of charge

(ie going more negative)

88
Q

Colligative properties

  1. Define
  2. What are the colligative properties of solutions?
A
  1. Colligative properties are the physical changes that result from adding solute to a solvent.

Colligative Properties depend on how many solute particles are present as well as the solvent amount, but they do NOT depend on the type of solute particles.

2a. Vapour pressure Vapor pressure or equilibrium vapor pressure is the pressure of a vapor in thermodynamic equilibrium with its condensed phases in a closed container. All liquids and solids have a tendency to evaporate into a gaseous form, and all gases have a tendency to condense back to their liquid or solid form.
2b. Boiling point
2c. Freezing point
2d. Osmotic pressure Osmosis is the diffusion of a fluid through a semipermeable membrane. When a semipermeable membrane (animal bladders, skins of fruits and vegetables) separates a solution from a solvent, then only solvent molecules are able to pass through the membrane. The osmotic pressure of a solution is the pressure difference needed to stop the flow of solvent across a semipermeable membrane. The osmotic pressure of a solution is proportional to the molar concentration of the solute particles in solution.

89
Q

VSEPR

  1. What molecular geometries are predicted with four electron dense areas?

(include all possible lone pairs options)

A
90
Q

Redox

Draw a galvanic cell Zn Cu

A
91
Q

VSEPR

  1. What molecular geometries are predicted with three electron dense areas?

(include all possible lone pairs options)

A
92
Q

Electronegativity - define and graph

A

Tendency to attract a bonding pair of electrons

Max top right with F 4.0

LINK

93
Q

standard molar volume =

A

1 mol gas at STP = 22.4L

94
Q

Redox

  1. Draw an electrolytic cell

Cd Cu

95
Q

Redox

  1. Write the half reactions for

Fe + 2HCl -> FeCl2 + H2

  1. Identify what has been oxidised and by what
  2. Identify what has been reduced and by what
  3. Identify the reducing agent
  4. Identify the oxidising agent
A
  1. Write the half reactions for

Fe0 - 2e- -> Fe+2(write it his way!)

2H+ + 2e- -> H20

  1. Fe oxidised by H OIL
  2. H reduced by Fe RIG
  3. Fe is the reducing agent
  4. H is the oxidising agent
96
Q

Atomic radius by element - graph

A

Added protons pull in all electrons a bit more.
New shell causes big jump in size.

LINK

97
Q
  1. What is the periodic trend for binary acids

ie increasing acidity

A
98
Q

What elements exist as diatomic molecules (7)?

hint - rock band

A

HON and the halogens

Hydrogen, Oxygen, Nitrogen

F Cl Br I

99
Q

Redox

  1. Oxidation is ..?.. high energy electrons
  2. Reduction is ..?.. high energy electrons
A
  1. Oxidation is losing high energy electrons

Oil Rig

  1. Reduction is gaining high energy electrons
100
Q

Quantum numbers

Draw table

A
101
Q

Reaction order

  1. Draw straight line graph for zeroth order reaction
  2. Draw straight line graph for firstorder reaction
  3. Draw straight line graph for second order reaction
102
Q

Solutions - units of concentration

  1. Molarity M =
  2. Molality m =
  3. Mole fraction χ =
  4. Mass percentage mass % =
  5. ppm =
A
  1. Molarity M = moles of solute / volume of solution
  2. Molality m = moles of solute / kg of solvent
  3. Mole fraction χ = moles of solute / (total moles of all solutes & solvent)
  4. Mass percentage mass % = 100 x mass of solute / total mass of solution
  5. ppm = 106 x mass of solute / total mass of solution
103
Q

Gibbs free energy ΔG = ΔH - TΔS

Draw a 2x2 grid showing combinations of ΔH and ΔS >0 and

explain the implications for spontaneous reactions (use ice/water as example)

A

ΔH (exothermic) ΔH > 0 (endothermic)

ΔS > 0 Spontaneous ΔG at high temperature

ΔS at low temperature NOT Spontaneous ΔG>0

temperature o > for water freezing / ice melting

ΔG = ΔH - TΔS

Video at 2min 50 for explanation if required:

104
Q

Calorimeters

  1. Describe the two main types of calorimeter
  2. and what parameter they each keep constant
A
  1. Coffee cup and bomb
  2. Coffee cup - constant pressure

Bomb - constant volume

105
Q

Solubility

  1. Solubility - define
  2. Solubility product Ksp - define
A

Solubility is a solute’s tendency to dissolve in a solvent

the maximum moles of the solute that can dissolve in the solution

Solubility product Ksp is the equilibrium constant for solvation

106
Q

Specific heat

  1. define specific heat (and units) c
  2. define molar specific heat (and units) C
  3. Define heat capacity Q =
A
  1. Specific heat is in units of J/K/kg , and is the amount of heat needed (in Joules) to raise the temperature of 1 mole of something, by 1 Kelvin (assuming no phase changes).

2. Molar specific heat is in units of J/K/mol , and

is the amount of heat needed (in Joules) to raise the temperature of 1 mole of something, by 1 Kelvin (assuming no phase changes).

107
Q

Reaction rate

for the reaction 4A + 3B -> 2C + 6D the reaction rate for C is 1 mol/Ls

what is the reaction rate for A B and D

A

the reaction rate for A is 2 mol/Ls

the reaction rate for B is 1.5 mol/Ls

the reaction rate for D is 3 mol/Ls

video LINK

108
Q

Collision theory

  1. Why does raising the concentration increase reaction rate?
  2. Why does raising the temperature increase reaction rate?
A
  1. There are more ways that molecules can collide so they are more likely to do so
  2. Increasing the temperature changes the distribution of velocites.

More will move to a higher velocity and cross the activation energy required for the reaction.

Reactants need a certain kinetic energy when they collide to break bonds and make the reaction possible. Below this KE they bounce.

109
Q

chemical reaction order

Define

A

First order - rate proportional to [reactant]

Second order - rate proportional to [reactant]2

Third order - rate proportional to [reactant]<strong>3</strong>

doesn’t have to be integer - eg order 1.5 is valid <strong>1.5</strong>

110
Q

Gibb’s free energy - thermodynamics

Formula

describe

A

The Gibbs free energy is defined as: G(p,T) = U + pV - TS

which is the same as: G(p,T) = H - TS

where: U is the internal energy (SI unit: joule) p is pressure (SI unit: pascal) V is volume (SI unit: m3) T is the temperature (SI unit: kelvin) S is the entropy (SI unit: joule per kelvin) H is the enthalpy (SI unit: joule)

Gibbs energy is a thermodynamic potential that measures the “usefulness” or process-initiating work obtainable from a thermodynamic system at a constant temperature and pressure (isothermal, isobaric).

Just as in mechanics, where potential energy is defined as capacity to do work, similarly different potentials have different meanings. The Gibbs free energy (SI units kJ/mol) is the maximum amount of non-expansion work that can be extracted from a thermodynamically closed system (one that can exchange heat and work with its surroundings, but not matter); this maximum can be attained only in a completely reversible process. When a system changes from a well-defined initial state to a well-defined final state, the Gibbs free energy change ΔP equals the work exchanged by the system with its surroundings, minus the work of the pressure forces, during a reversible transformation of the system from the initial state to the final state.

Gibbs energy is also the chemical potential that is minimized when a system reaches equilibrium at constant pressure and temperature. Its derivative with respect to the reaction coordinate of the system vanishes at the equilibrium point. As such, it is a convenient criterion for the spontaneity of processes with constant pressure and temperature.

111
Q

Acid

  1. Base ionisation constant Kb =
A

Basic Information

1) Weak bases are less than 100% ionized in solution.<br></br>2) Ammonia (formula = NH3) is the most common weak base example used by instructors.

The following equation describes the reaction between ammonia and water:

NH3 + H2O NH4+ + OH¯ Note that it is an equilibrium condition.

The equilibrium constant for this reaction is written as follows:

Kc = ( [NH4+] [OH¯] ) / ( [NH3] [H2O] )

However, in pure liquid water, [H2O] is a constant value. To demonstrate this, consider 1000 mL of water with a density of 1.00 g/mL. This 1.00 liter (1000 mL) would weigh 1000 grams. This mass divided by the molecular weight of water (18.0152 g/mol) gives 55.5 moles. The “molarity” of this water would then be 55.5 mol / 1.00 liter or 55.5 M.

The solutions studied in introductory chemistry are so dilute that the “concentration” of water is unaffected. So 55.5 molar can be considered to be a constant if the solution is dilute enough.

Moving [H2O] to the other side gives: Kc [H2O] = ( [NH4+] [OH¯] ) / [NH3]

Since the term Kc [H2O] is a constant, let it be symbolized by Kb, giving:

Kb = ( [NH4+] [OH¯] ) / [NH3]

This constant, Kb, is called the base ionization constant.

112
Q

Antibonding

Explain why it exists

How does it explain why He2 is not a stable molecule

A

The number of molecular orbitals has to equal the number of atomic orbitals. Bringing together two atomic orbitals will always result in two new molecular orbitals, one bonding (“constructive overlap”)

one antibonding (“destructive overlap”)

Antibonding orbitals have higher energy than the separated atoms orbitals which in turn have higher energy level than the bonding orbitals.

He has a full s shell of two electrons.

So when two He atoms bond there are two electrons in the bonding orbital and two in the non bonding orbital.

This is a higher energy state than the two separate atoms so they would prefer to be apart.

113
Q

Binary molecules

Naming convention

A

Lowest left in periodic table FIRST

eg N2 O4

di nitrogen tetr ox ide

114
Q

Name the following metals

Sn

Ag

Au

Ru

Hg

Pb

Ba

Pt

Pd

U

A

Sn Tin

Ag Silver

Au Gold

Ru Ruthenium

Hg Mercury

Pb Lead

Ba Barium

Pt Platinum

Pd Palladium

U Uranium

115
Q

Chemical kinetics

Define the rate determining step

A

In chemical kinetics, the overall rate of a reaction is often approximately determined by the slowest step, known as the rate determining step

116
Q
  1. What is the periodic trend for oxyacids

ie increasing acidity

A Change of halogen (and why)

B Number of oxygen atoms (and why)

A

A increases up (opposite to binary acids)

The increasing electronegativity pulls the electron more towards the halogen making the hydrogen looser

B Adding oxygens pulls the electrons more towards the additional oxygens making the hydrogen looser

117
Q

Redox

Rule 1: The oxidation number of an element in its free (uncombined) state is ….?..

for example, …?..

A

Rule 1: The oxidation number of an element in its free (uncombined) state is zero — for example, Al(s) or Zn(s). O2

Rule 2: The oxidation number of a monatomic (one-atom) ion = the charge on the ion, for example: Na+ = +1

Rule 3: The sum of all oxidation numbers in a neutral compound =zero.

The sum of all oxidation numbers in a polyatomic ion = the charge on the ion

This rule often allows chemists to calculate the oxidation number of an atom that may have multiple oxidation states, if the other atoms in the ion have known oxidation numbers.

Rule 4: The oxidation number of an alkali metal (IA family) in a compound is +1;

the oxidation number of an alkaline earth metal (IIA family) in a compound is +2.

Rule 5: The oxidation number of oxygen in a compound is usually –2. peroxides (for example, hydrogen peroxide), then the oxygen has an oxidation number of –1.

If the oxygen is bonded to fluorine, the number is +1.

Rule 6: The oxidation state of hydrogen in a compound is usually +1.

If the hydrogen is part of a binary metal hydride (compound of hydrogen and some metal), then the oxidation state of hydrogen is –1.

Rule 7: The oxidation number of fluorine is always –1. Chlorine, bromine, and iodine usually have an oxidation number of –1, unless they’re in combination with an oxygen or fluorine.

118
Q

Periodic table

  1. Name the first six Noble gases
A

table

119
Q
  1. What is the percentage ionisation of an acid A?
A

100 [H+] /[HA]

120
Q
  1. What is the generic form for the conjugate acid of a weak base B in water?

Hydrolosis of a cation

A

BH+ + H2O <-> H3O+ + B

121
Q

Electron affinity - define and graph

A

Electron affinity is defined as the change in energy (in kJ/mole) of a neutral atom (in the gaseous phase) when an electron is added to the atom to form a negative ion. In other words, the neutral atom’s likelihood of gaining an electron.

X + e− → X− + energy

Half filled subshells (eg N P) are more difficult to add to as the new electron has to go close to an existing one.

LINK

122
Q

Enthalpy

Define + formula

How is it used in chemistry?

A

Enthalpy is defined as a thermodynamic potential, designated by the letter “H”, that consists of the internal energy of the system (U) plus the product of pressure (p) and volume (V) of the system:

H = U + pV

Since U, p and V are all functions of the state of the thermodynamic system, enthalpy is a state function.

The unit of measurement for enthalpy in the International System of Units (SI) is the joule

In reactions ΔH = Hproducts - Hreactants

Exothermic = heat given off Hproducts reactants

Endothermic = heat absorbed Hproducts > Hreactants

123
Q

Chemical reaction order

What is the overall reaction order for this reaction

A

order of overall reaction = m + n

124
Q

Redox

  1. What does the mnemonic RedCat AnOx stand for?
A

Galvanic cell

Reduction at the Cathode

Oxidation at the Anode

125
Q

Redox

Rule 5: The oxidation number of oxygen in a compound is usually ..?..

peroxides (for example, hydrogen peroxide), then the oxygen has an oxidation number of ..?..

If the oxygen is bonded to fluorine, the number is ..?..

A

Rule 1: The oxidation number of an element in its free (uncombined) state is zero — for example, Al(s) or Zn(s). O2

Rule 2: The oxidation number of a monatomic (one-atom) ion = the charge on the ion, for example: Na+ = +1

Rule 3: The sum of all oxidation numbers in a neutral compound =zero.

The sum of all oxidation numbers in a polyatomic ion = the charge on the ion

This rule often allows chemists to calculate the oxidation number of an atom that may have multiple oxidation states, if the other atoms in the ion have known oxidation numbers.

Rule 4: The oxidation number of an alkali metal (IA family) in a compound is +1;

the oxidation number of an alkaline earth metal (IIA family) in a compound is +2.

Rule 5: The oxidation number of oxygen in a compound is usually –2. peroxides (for example, hydrogen peroxide), then the oxygen has an oxidation number of –1.

If the oxygen is bonded to fluorine, the number is +1.

Rule 6: The oxidation state of hydrogen in a compound is usually +1.

If the hydrogen is part of a binary metal hydride (compound of hydrogen and some metal), then the oxidation state of hydrogen is –1.

Rule 7: The oxidation number of fluorine is always –1. Chlorine, bromine, and iodine usually have an oxidation number of –1, unless they’re in combination with an oxygen or fluorine.

126
Q
  1. The thermodynamic equilibrium constant is defined as: Keq=
127
Q

Redox

Cathode is the _??_ve terminal and the ..??.. of electrons

Anode is the ??ve terminal and the ..??.. of electrons

A

Cathode is the +ve terminal and the sink of electrons

Anode is the -ve terminal and the source of electrons

padlet

128
Q

Acid

  1. Acid dissociation constant Ka =
A

Basic Information

1) Weak acids are less than 100% ionized in solution.<br></br>2) Acetic acid (formula = HC2H3O2) is the most common weak acid example used by instructors.<br></br>3) Another way to write acetic acid’s formula is CH3COOH.<br></br>4) A common abbreviation for acetic acid is HAc, where Ac¯ refers to the acetate polyatomic ion.

The following equation describes the reaction between acetic acid and water:

HAc + H2O H3O+ + Ac¯ Note that it is an equilibrium condition.

The equilibrium constant for this reaction is written as follows:

Kc = ( [H3O+] [Ac¯] ) / ( [HAc] [H2O] )

However, in pure liquid water, [H2O] is a constant value. To demonstrate this, consider 1000 mL of water with a density of 1.00 g/mL. This 1.00 liter (1000 mL) would weigh 1000 grams. This mass divided by the molecular weight of water (18.0152 g/mol) gives 55.5 moles. The “molarity” of this water would then be 55.5 mol / 1.00 liter or 55.5 M.

The solutions studied in introductory chemistry are so dilute that the “concentration” of water is unaffected. So 55.5 molar can be considered to be a constant if the solution is dilute enough.

Moving [H2O] to the other side gives: Kc [H2O] = ( [H3O+] [Ac¯] ) / [HAc]

Since the term Kc [H2O] is a constant, let it be symbolized by Ka, giving:

Ka = ( [H3O+] [Ac¯] ) / [HAc]

This constant, Ka, is called the acid ionization constant.

129
Q

Redox

Explain the purpose and operation of the salt bridge in a galvanic cell Zn Cu

A

The salt bridge is viscous (a gel) so would stay in the bridge if only subject to gravity.

The Cl- gets attracted into the +ve Zn side and keeps the solution neutral

Likewise K+ to the Cu side

As Zn jumps into solution the solution would become more and more positively charged. Eventually the current would stop flowing to the Cu and go back into the Zn solution.

130
Q

First 20 elements of periodic table

with electronegativity of relevant Organic chemistry elements

A

LINK

131
Q

Colligative

  1. How does the freezing point of a solution compare to that of the pure solvent?
  2. ΔTf =
A
  1. The freezing point of a solution is always lower than that of the pure solvent.
  2. The freezing point depression is given as
    ΔTf = iKfm
    where Kf is the molal freezing point depression constant,

m is the molality of the solution, and

i is the van’t Hoff factor equal to the number of particles the solute
dissociates into. For nonelectrolytes, i = 1. For electrolytes, i = number of ions per
formula.

132
Q

Effective nuclear charge

What is it?

A

In an atom with one electron, that electron experiences the full charge of the positive nucleus. In this case, the effective nuclear charge can be calculated from Coulomb’s law.

However, in an atom with many electrons the outer electrons are simultaneously attracted to the positive nucleus and repelled by the negatively charged electrons. The effective nuclear charge on such an electron is given by the following equation:

Zeff = Z - S

where

Z is the number of protons in the nucleus (atomic number), and
S is the average number of electrons between the nucleus and the electron in question (the number of nonvalence electrons).

133
Q

What is the rule of thumb for identifyng strong bases and weak bases?

A

Strong bases contain the hydroxide anion while

most weak bases will be an ammine compound (contain an N as a central atom).

134
Q

Entropy

Define (2) S =

What happens to entropy with dispersal?

give examples

A

Entropy - a measure of the amount of energy that is unavailable for work

Entropy - a measure of disorder

Entropy increases with dispersal eg

  1. phase change (melting ice),
  2. stoichiometric products (if I have more moles after the reaction A + 2B -> 3C + 2D),
  3. increasing gas volume,
  4. increasing gas temperature(spreads out)
135
Q

Redox

How does oxidation state relate to ionic bonds and partial charged covalent bonds?

A

The oxidation state rounds up the partial charge and pretends that all the bonds are ionic.

So you can see where the electrons have gone

136
Q

Acid

  1. In a series of oxyacids what indicates a stronger acid?
  2. and why?
A
  1. More oxygens
    2a. the oxygens draw electrons to one side increasing polarity
    2b. the oxygens in the conjugate base can share the electrons spreading it over a larger volume thus stabilising the conjugate base
137
Q

pH

  1. pH =
  2. Name some substances in the pH range (table image)
A
  1. pH = -log[H+]
138
Q

Solubility

Explain the relationship between pressure and solubility using a soda can

  1. when popping the tab
  2. if heated
A
  1. Solubility of a gas is proportional to the vapour partial pressure

so when the can is popped the pressure goes down, so does the solubility of CO2 gas and the thing fizzes

  1. As temperature increase the solubility of gases decreases - so the CO2 is released from solution and the can explodes
139
Q

VSEPR

  1. What molecular geometries are predicted with two electron dense areas?

(include all possible lone pairs options)

140
Q

Acid-base

  1. Derive the Henderson-Hasselbalch approximation

starting with a buffer solution

HA H+ + A¯

pH = pKa + log ([A−] / [HA])

A

The Henderson-Hasselbalch approximation allows us one method to approximate the pH of a buffer solution

pH = pKa + log ([A−] / [HA])

HA is the acid, A- is the conjugate base

padlet

derivation

MCAT Gen. Chem(David) (2 decks)

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