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MCAT > chemistry fundamentals > Flashcards

chemistry fundamentals Flashcards

1
Q

What unit of volume is used for laboratory work? Why?

A

Litres, because, the cubic meter, the SI unit for volume is too large for most lab work

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2
Q

1 L = ? cubic meter

A

1000

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3
Q

Define Molarity

A

Moles of solute per litre of solution

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4
Q

1mL = 1 cm3, True or False?

A

True

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5
Q

Litre is used to express the volume of solids? T/F

A

False, litre is used to express the volumes of liquid and gases

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6
Q

Define Angstrom?

A

Angstrom is a unit of length equals to 10^-10m. The atomic radii and bond lengths are typically around 1 to 3 A

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7
Q

Define Density? What are its SI units?

A

It is mass per volume. m/V. The SI Units of Density is kilogram per cubic meter (kg/m^3)

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8
Q

What is density most often expressed in?

A

Grams per cubic centimeter

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9
Q

What is the density of water?

A

1 g/cm3

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10
Q

Define Molecule

A

When two or more atoms form a covalent bond they create a molecule

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11
Q

How do we know the identity of a compound and the number of its atoms?

A

Using the molecular formula

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12
Q

Are molecular formula and empirical formula the same thing?

A

No, to get the empirical formula divide all subscripts by a common factor. Empirical formula is a reduced form of the molecular formula

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13
Q

What is the formula for ammonium?

A

NH 4 +

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14
Q

Identify H3O+

A

Hydronium

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15
Q

Formula for acetate?

A

CH3CO2-

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16
Q

IDENTIFY HCO3-

A

bicarbonate

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17
Q

FORMULA FOR CYANIDE

A

CN-

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18
Q

IDENTIFY OH-

A

HYDROXIDE

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19
Q

FORMULA FOR NITRATE

A

NO 3 -

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20
Q

FORMULA FOR NITRIDE

A

NO2-

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21
Q

IDENTIFY CL04-

A

PERCHLORATE

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22
Q

FORMULA FOR CARBONATE

A

CO3 2-

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23
Q

IDENTIFY SO4 2-

A

SULFATE

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24
Q

IDENTIFY SO4 2-

A

SULFATE

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25
Q

IDENTIFY SO3 2-

A

SULFITE

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26
Q

FORMULA FOR PHOSPHATE

A

PO 4 2-

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27
Q

Define formula weight? What is the unit?

A

Sum of all the atomic weights of all the atom in the molecule. The unit is atomic mass unit (amu). It says weight but it really is mass

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28
Q

Define one amu?

A

is equal to exactly 1/12 the mass of an atom of carbon -12

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29
Q

What is mole?

A

it is a particular number of things. One mole contains 6.02 x 10 ^23 entities (Avogadro’s number)

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30
Q

What is Avogadro’s number denoted by?

A

N A in the subscript or N 0 in the subscript

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31
Q

What is the link between atomic mass unit and grams?

A

Avogadro’s number

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32
Q

What is the formula to calculate number of moles?

A

mass in grams/ molecular weight

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33
Q

What formula is preferred to calculate percent mass composition?

A

Empirical formula

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34
Q

What expresses the concentration of a solution in terms of moles of solute per volume?

A

Molarity

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35
Q

[ ] = concentration? T/F

A

True

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36
Q

What is Mole fraction?

A

Fraction of moles of a given substance relative to the total moles in the solution

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37
Q

When is it more useful to use mole fraction over molarity?

A

Mole fraction is used when more than one solute is present in a solution and is often used in gases

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38
Q

What is the Law of conservation of mass?

A

the amount of mass does not change in a chemical reaction

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39
Q

Define stoichiometry co-efficients?

A

they tell us in what proportions the reactants react and products are formed

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40
Q

The stoichiometric co-efficients are used to give the ratios by mass. T/F

A

False, they give ratios of the number of molecules or moles

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41
Q

What is a limiting reagent?

A

Whatever reactant runs out first to limit the production of a product is called limiting reagent

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42
Q

Is limiting reagent a reactant that is present in the smallest amount? T/F

A

No, Limiting reagent is the reactant that is consumed first and not necessarily the reactant that’s initially present in the smallest amount.

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43
Q

A product can be a limiting reagent? T/F

A

False, only the reactants are limiting reagent.

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44
Q

What are the four states used in chemical formula?

A

Solid, liquid, gas and aqueous

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45
Q

Define catalyst.

A

A substance that increases the rate of reaction without being consumed

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46
Q

What does the oxidation state of an atom indicate?

A

Atom’s oxidation state or number indicates how the atom’s ownership of its valence electrons changes when it forms a compound.

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47
Q

What results in a positive vs negative oxidation state?

A

Giving up the ownership of valence electrons results in positive oxidation and accepting ownership of valence electrons results in a negative oxidation state.

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48
Q

What is the oxidation state of any element in its standard state?

A

0

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49
Q

What is the rule for oxidation state of an atom in a molecule/ion?

A

The sum of the oxidation state of the atoms in a neutral molecule must always be 0. For ions, the sum of the oxidation states of the atoms must always equal to the ion’s charge

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50
Q

What is the oxidation state of Group 1 and Group 2 metals?

A

+1 for Group 1 and +2 for group 2

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51
Q

Oxidative state of Flourine?

A

-1

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52
Q

What are the rules for oxidation state of hydrogen

A

+1 when bonded to something that is more electronegative than carbon, -1 when bonded to something less electronegative than carbon and 0 when it is bonded to carbon

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53
Q

What is the oxidative state of oxygen and atoms that belong to the oxygen family?

A

-2

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54
Q

What oxidative state do halogens have?

A

-1

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55
Q

When does O2 not have a oxidative state of -2?

A

In peroxides, oxygen is in a -1 oxidative state.

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56
Q

What is the pneumonic for the order of electronegativity?

A

FONClBrISCH (Fawn-cull-brish), most to least

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57
Q

In FONClBrISCH, identify the oxidative states of H.

A

Anything before C is +1 for H, anything that is not here is -1 for H.

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58
Q

A metal will never assume a negative oxidation state? T/F

A

True

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59
Q

What is an atom? What are some parts of the atom?

A

The smallest unit of any element is called an atom. All atoms contain a nucleus and outside the nucleus, there are electrons.

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60
Q

What is inside a nucleus?

A

Protons and nuetrons

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61
Q

What are nucleons?

A

Protons and neutrons are collectively known as nucleons

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62
Q

Charge of protons

A

+1

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63
Q

Charge of nucleus

A

0

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64
Q

charge of electrons

A

-1

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65
Q

How are electrons held inside an atom?

A

Through the electrostatic attraction of positively charged nucleus

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66
Q

every neutral atom has equal charge of electrons outside the atom and protons inside the atom? T/F

A

True

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67
Q

What is atomic number (Z)?

A

The number of protons in a nucleus of an atom

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68
Q

Are protons represented by a subscript or a superscript?

A

Subscript

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69
Q

What is the mass of protons and neutrons?

A

1.66 x 10 ^-27

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70
Q

What is the mass of an electron?

A

0.05 % mass of the protons or the neutrons

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71
Q

What contributes to all the mass of an atom

A

Nucleus

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72
Q

How do you find the mass number of an atom? (A)

A

of protons + # of neutrons

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73
Q

What are isotopes?

A

Two atoms of same element differ in their number of neutrons. (Protons stay the same)

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74
Q

What is an atomic weight of an element?

A

weighted average of all the masses of its naturally occurring isotopes

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75
Q

Define ions, anions, cations?

A

When a neutral atom gains or loses electrons, it becomes charged and the resulting atom is called an ion. A negatively charged atom is called anions. A positively charged atom is called cation.

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76
Q

Charge of an atom when it gains an electron

A

-1

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77
Q

Charge of an atom when it loses an electron

A

+1

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78
Q

How are neutrons and protons held together?

A

by strong nuclear force

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79
Q

Electrical force is stronger than nuclear force? T/F

A

False, nuclear force is stronger than electrical force between charged particles.

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80
Q

What is the limitation of the nuclear force?

A

It only works over extremely short distances

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81
Q

What is a radioactive nuclie?

A

An unstable nuclie

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82
Q

Define radioactive decay?

A

unstable nuclei undergo a transformation to make them more stable, altering the number and ratio of protons and neutrons or lowering their energy,

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83
Q

Difference between a parent nucleus versus the daughter?

A

The nucleus that undergoes the radioactive decay is known as the parent and the resulting more stable nucleus is known as the daughter

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84
Q

3 types of radioactive decay?

A

Alpha, Beta and gamma

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85
Q

Define alpha decay

A

When a large nucleus wants to become more stable by reducing the number of proton and neutrons by emitting an alpha particle.

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86
Q

What is the symbol for an alpha particle and what does it consist of?

A

α (4 superscript and 2 subscript). It consists of 2 protons and 2 neutrons

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87
Q

What atom is the alpha particle similar to?

A

Helium

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88
Q

Alpha decay reduces the parent’s atomic number by —- and mass number by —-

A

2 and 4

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89
Q

What particles cant travel very far and can be stopped by human skin and piece of paper?

A

Alpha Particles

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90
Q

What is beta decay?

A

Beta decay involved conversion of neutrons into protons or vice versa through the action of the weak nuclear force.

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91
Q

3 types of beta decay?

A

β− decay, β+ decay and electron capture

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92
Q

What type of particles can be stopped by aluminium foil and a centimetre of glass and plastic

A

Beta particles

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93
Q

Explain β− decay

A

β− decay is when an unstable nucleus contains too many neutrons, it may convert those into a proton and an electron which is ejected. As a result the atomic number of the daughter nucleus is 1 greater than the parent and the mass stays the same. (coverts into proton and electron, -ve and +ve = neutral,)

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94
Q

Explain β+ decay.

A

When an unstable nucleus contains too few neutrons, it converts a proton into a neutron and a positron is ejected. The atomic number of the daughter nucleus is one less than the parent nucleus but the mass stays the same. (protons gets converted into neutron that is neural, the leftover positive charge is emitted in a form of a positron)

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95
Q

Difference between an electron and positron?

A

Positron is antiparticle of electron. It is identical to an electron but has a positive charge.

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96
Q

What is the most common type of Beta decay?

A

β- decay.

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97
Q

Explain electron capture

A

An unstable nucleus increases its number of neutrons by capturing an electron from the closest shell (n=1) which is then used in the conversion of protons into neutrons ( -ve charge from electrons and +ve from protons = neutral neurones). The atomic number is reduced by 1 and mass stays the same.

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98
Q

Explain gamma decay

A

After the nucleus has undergone alpha or beta decay, it is in an excited state of energy. It can relax to its ground state by emitting energy in the form of one or more photons of electromagnetic radiation. This radiation changes neither the atomic number nor the mass number

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99
Q

What are the characteristics of gamma photons?

A

They have very high frequency and energy . They do not have mass or charge so they can penetrate most matter effectively.

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100
Q

Order in terms of least to highest degree of penetration, beta, alpha and gamma

A

Alpha < beta < gamma

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101
Q

True or False, alpha, beta and gamma decay change the identity of the nucleus

A

False, gamma doesn’t change the identity. Gamma decay is simply the expulsion of energy

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102
Q

Define half-life

A

Time it takes for one-half of some sample of the substance of decay.

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103
Q

True or False. Shorter the half life, the faster the decay?

A

True

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104
Q

What is nuclear binding energy?

A

The energy that was released when individual protons and nucleus were bound together by strong force to form a nucleus. This is the same energy that would require to break a nucleus into nucleons.

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105
Q

What is the relationship between the binding energy and stability?

A

direct

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106
Q

Explain mass defect?

A

When nucleons bind together to form a nucleus, some mass is converted into energy, so the mass of the nucleus when you combine all the nucleons is less than the sum of individual nucleon masses. The difference is called mass defect

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107
Q

Mass defect and nuclear binding energy are different, True or false

A

True, but they are equivalent through Einstein’s equation for mass energy equivalence E = mc 2

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108
Q

What is the characteristic of a mass defect of a stable nucleus

A

It is always positive

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109
Q

How would you convert kg to joules?

A

Multiply by 9 x 10 ^16

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110
Q

Formula for nuclear binding energy?

A

E = mc2

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111
Q

In a nuclear domain, energy is expressed in volts?

A

Electronvolts, 1 ev = 1.6 10 ^ -29

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112
Q

What is c?

A

Speed of light (3.0 x 10^8)

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113
Q

Formula for nuclear binding energy in MeV?

A

m in amu x 931.5 ev, 1 emu = 1.66 x 10 ^ -27

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114
Q

What does emission spectrum display?

A

different wavelengths and energies of light for the gas passing through it

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115
Q

Formula for energy of a photon?

A

E = hf, h is planck’s constant (6.63 x 10^-34)

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116
Q

Frequency in terms of speed of light and wavelength formula?

A

F = c/wavelength

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117
Q

Explain Bohr’s model of atom?

A
  1. electrons in an atom orbited the nucleus like planets did the sun
  2. distance from the nucleus is related to the energy of the electrons, electrons with large energy orbited at greater distances.
  3. electrons in the atom cannot assume any arbitrary energy, but have quantized energy states.
  4. they only orbit at certain allowed distances from the nucleus.
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118
Q

Explain the phenomenon of jumping and dropping of an electron?

A

if an electron absorbs energy that is exactly equal to the different in energy between its current level and the one available at a higher level, it jumps to that higher level.

If an electron emits a photon with an energy exactly equal to the difference between its current level and the one at the lower energy.

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119
Q

Bohr’s model predicted that elements have line spectra instead of a continuous spectra. True/ False

A

True

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120
Q

An electron can absorb any amount of energy?

A

False, electron can only gain or lose specific amount of energy due to their quantized nature of energy levels

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121
Q

Define ground state? excited state?

A

It is where electron is initially, lowest possible energy level n = 1
when electron absorbs a photon and it jumps to a higher energy level

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122
Q

Excited electrons relax to the ground level in large jumps? t/f

A

False, they gradually lose energy and come to the ground state

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123
Q

Formula for energy levels of electrons?

A

E n = (-2.178 x 10 ^ -18) / n 2

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124
Q

What is Bohr’s atom

A

an atom that contains only one electron

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125
Q

Limitation of Bohr’s model of atom

A

cannot predict the line spectra of multiple electrons, cannot explain the electron-electron interactions

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126
Q

What does quantization of atoms describe?

A

Shell, sub-shell, orbital and spin number

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127
Q

higher shell electron has a greater amount of energy? T/F

A

True

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128
Q

Define orbital

A

3-D region around the nucleus where the electron is most likely to be found.

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129
Q

Define subshell

A

A subshell in an atom contains one or more orbitals with the same energy. It also describes the shape and energy of the orbitals. Subshells are s, p, d, f

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130
Q

what are degenerate orbitals

A

one or more orbitals of the same energy

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131
Q

orbitals of same energy have same 3-D orientation in space? T/F

A

False. The number of orientation increases by two in each successive subshells. 1 in s, 3 in p and so on

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132
Q

The number of orientations increases by two in each successive subshell? T/F

A

True, s subshell contains one orientation and p contains three

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133
Q

Describe the shape of s and p orbitals

A

s subshell has one spherical symmetrical orbital and p has three orbitals with different spacial orientations

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134
Q

describe electron spin states

A

two possible spin states (electron’s intrinsic magnetism). orbital with two electrons, one spins up and one spins down

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135
Q

3 rules of electron configuration?

A

Aufbau principle: electrons occupy the lowest energy orbitals available
Hund’s rule: electrons in the same sub shells occupy available orbitals singly before pairing up
Pauli principle: there can be no more than 2 electrons in a given orbital

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136
Q

maximum numbers of electrons in s, p, d, f subshell?

A

s has 1 possible orbital orientation, and there can be no more than 2 electrons in any given orbital
2 x 1 = 2 elections
p has 3 possible orbital orientation,
2x 3 = 6 electrons
d has 5
2 x 5 = 10
f has 7
2 x 7 = 14

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137
Q

Define diamagnetic

A

an atom that has its all of the electrons paired is called diamagnetic. these electrons cancel out each other’s magnetic field leaving the atom with no magnetic field. As a result, the atom will be repelled by an externally produced magnetic field

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138
Q

Define paramagnetic

A

when an atom’s electrons are not all spin-paired. these atoms are attracted to external magnetic field

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139
Q

difference between a period and group on the periodic table

A

period is horizontal row and group is the vertical column

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140
Q

difference between a period and group on the periodic table

A

period is horizontal row and group is the vertical column

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141
Q

Additional rules for d and f period while counting the configuration

A

for d block of period, subtract 1 from period number, outermost electrons are in 3d instead of 4d and for f period block, subtract 2 from the period number, instead of 8f it would be 6f

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142
Q

Name some atoms with anomalous electron configurations. (anticipated electron configuration is not one observed)

A

anomalous atoms can achieve a lower state of energy by having a filled or half filled d shell. These are chromium group, copper group atoms, pg -73

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143
Q

define isoelectronic?

A

F - has the same electron config. as Ne gas after gaining an electron. Atoms with the same configurations said to be isoelectronic.
When an atom becomes an anion, move right on the periodic table
When an atom becomes a cation, move left on the periodic table

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144
Q

Define the process of ionization (electrons being removed) from d group

A

To form a cation, atoms will always lose their valence electron first from s before they lose it from d

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145
Q

List special names for a group in periodic table

A

Group I : Alkali metals ( reducing agents)
Group II: Alkaline earth metals ( reducing agents)
Group VII: Halogens (oxidizing agents)
Group VIII: Noble gases
s, p blocks: representative elements
d blocks: transition metals
f blocks: rare earth metals

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146
Q

Define metalloids

A

Elements that posses the qualities of both metals and non metals. Example: B, Si, Ge, As, Sb, Te, Po

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147
Q

Define shielding

A

Each filled shell between the nucleus and the valence electrons shields the valence electrons from the full effects of the positively charged protons in the nucleus. Z eff: effective nuclear charge

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148
Q

Describe the atomic trends in the periodic table (atomic radius, ionization energy, electron affinity, electronegativity, acidity)

A

Atomic radius:
left to right: atomic radius decreases because more protons are applied higher effective nuclear charge on the electrons
top to bottom: atomic radium increases because more shells are being added and shielding increases
cation < neutral atom < anion

ionization energy: energy required to remove an electron. IE 1 for energy to remove the least tightly bound electron. IE2: energy to remove least tightly bound from the cation. IE2 > IE1
left to right and bottom to top - energy increases because the valence electrons are more tightly bound.
noble gases have large IE

Electron affinity: energy associated with add an electron. energy released: negative, energy absorbed: positive
halogens have large negative EA, more stable: release of energy
noble gases and alkaline earth metals have positive EA as adding an electron destabilizes them
more negative as we right and up (there are exceptions like noble gases)

Electronegativity : tendency to attract electrons
left to right: ENV increases
top to bottom: decreases
F>O>N=Cl>Br>I>S>C=H

Acidity: how well the compound donates protons
left to right: increases because more ENV more stable anion
top to bottom: increases, as the size increases

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149
Q

What does formal charge tell us? What is the formula?

A

If the atoms are sharing their electrons in the best way possible. This happens when the formal charge is as low as possible. FC = valence e - 1/2 bonding e - lone paired e (individual). Negative formal charges on the most electronegative elements

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150
Q

Describe resonance structure

A

These structures are often needed when there are double or triple bonds in molecules along with one or more lone pairs of electrons.

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151
Q

Define resonance hybrid

A

an average of the equivalent or non-equivalent lewis structures. The dotted line in the hybrid signifies a bond and a half

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152
Q

Define bond length

A

the distance between two nuclie that are bonded to one another

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153
Q

Define bond dissociation energy

A

energy required to break the bond homolytically (where one electron of the bond being broken goes to each fragment of the molecule and two radicals form)

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154
Q

Define heterolytic bond cleavage

A

both electrons that make up that bond end up on the same atom forming an anion and a cation.

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155
Q

Define bond order?

A

number of bonds between adjacent atoms, single bond = 1, double = 2, triple = 3

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156
Q

relationship between bond length and bond order

A

higher the bond order, shorter and stronger the bond

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157
Q

Carbon carbon bond lengths can be compared to carbon-oxygen bond lengths? T/F

A

False, bevause of the varying atomic radii the bond length and dissociation energy comparisons should be between similar bonds. Carbon - carbon can be compared to carbon-carbon only

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158
Q

describe the relationship between bond length and s and p charaters

A

the greater the s character, the shorter the bond because s-orbitals are closer to nucleus. p character leads to a longer bond that is farther from the nucleus

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159
Q

How is covalent bond formed?

A

Covalent bond is formed between atoms when each contributes one or more of its valence electrons. These electrons are shared by the atoms to help complete the octet.

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160
Q

When is a covalent bond polar and non-polar?

A

Polar bond when the electron density between two nuclei is uneven or when there is a greater difference between the ENV of the bonding atoms.

Non-polar when the electron density is even and there is a little difference between the ENV of two atoms that are bonded to each other.

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161
Q

what is a coordinate covalent bond?

A

one atom donating both the shared electrons in a bond. example NH3 bonding with BF3 in which N donates both of its lone pair to make a bond

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162
Q

what is Lewis base and Lewis acid?

A

a molecule that donates a pair of electrons is called Lewis base aka ligand aka nucleophile (nucleus loving). base is a diner

a molecule that accepts a pair of electrons is called Lewis acid aka electrophile (electron loving). acid is a acceptor

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163
Q

When a covalent bond breaks, the electrons from the ligand split between two molecules. T/F

A

False, electrons that come from a ligand leaves with a ligand.

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164
Q

How do you identify Lewis base?

A

Lewis base is a molecule with a pair of non-bonding electrons

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165
Q

What is an ionic bond

A

when two molecules are held together by the electrostatic attraction between a cation and anion.

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166
Q

What is one rule of VSEPR THEORY

A

SINCE ELECTRONS REPEL ONE ANOTHER, ELECTRON PAIRS MOVE FAR AWAY FROM EACH OTHER WHETHER THEY ARE BONDING OR NON-BONDING

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167
Q

what are some trends of ionic bonds

A

metal and non mental should have a big difference between their ENV to make an ionic bond
strength of an ionic bond is related to the charges on the ion. bigger charge, greater strength
the strength decreases as the ionic radii increase

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168
Q

CHARACTERISTICS OF A LINEAR SHAPE

A

ANGLE IS 180

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169
Q

WHAT ARE ELECTRON GROUPS

A

ANY TYPE OF BOND AND LONE PAIRS

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170
Q

WHAT DOES TOTAL NUMBER OF ELECTRON GROUP ON THE CENTRAL ATOM DETERMINES

A

THE BOND ANGLES AND ORBITAL GEOMETRY

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171
Q

IS ORBITAL GEOMETRY SAME AS MOLECULAR GEOMETRY?

A

NO, MOLECULAR GEOMETRY AKA SHAPE CHANGES WITH LONE PAIR INTO BENT AND IS NOT SAME AS THE ORBITAL GEOMETRY

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172
Q

What is the value of s character and p character in so, sp2 and sp3 hybridization?

A

sp - 50% s and 50% p
sp2 - 33% - s and 67 % p
sp3 - 25 % s and 75% p

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173
Q

what are the bond angles of three hybridization groups?

A

sp - 180
sp2 - 120
sp3 - 109.5

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174
Q

review khan academy notes on hybridization and steric number

A
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175
Q

What results in a polar molecule

A

When there are unsymmetrical polar bonds present in a molecule so they don’t get cancelled out. The atoms connecting the the central atom should have unequal ENV

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176
Q

List two forces that hold liquids and solids

A

Intermolecular molecular forces such as dipole-dipole and London dispersion forces

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177
Q

define intermolecular forces

A

weak interactions that take place between neutral molecules

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178
Q

define ion dipole forces

A

polar molecules are attracted to ions and create ion-dipole forces

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179
Q

define dipole dipole interactions

A

attraction between positive end of the polar molecule and the negative end of the polar molecule

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180
Q

define dipole induced dipole force

A

a permanent dipole in a molecule may induce a dipole in a neighbouring non polar molecule producing dipole induced dipole interactions

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181
Q

Define LDF

A

an instantaneous dipole in a non polar molecule induces a dipole in neighbouring non polar molecule. This results in LDF which are very weak and a default force in all atoms or molecules

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182
Q

How is molecular weight related to polarity

A

as the molecular weight increases, the polarity also increases

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183
Q

what is the relationship between dispersion forces and atomic size

A

direct, as the atomic size increases the dispersion forces also increase

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184
Q

How would stronger intermolecular forces effect meeting points, boiling points, viscosities and vapour pressure

A

high melting and boiling points, high viscosity, and lower vapour pressure

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185
Q

what are van Der Waals forces collectively and individually referring to

A

collectively - dipole-dipole, hydrogen bonding and LDF
individually - LDF

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186
Q

What are the 2 conditions of a hydrogen bonding

A
  1. a molecule must have a covalent bond between H and either NOF
    2, the another molecule must have a lone pair of electrons on NOF atom
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187
Q

The boiling points of NOF bonded with H containing compounds is similar to other elements found in the similar groups?
T/F

A

False, the boiling point of the hydrogen containing compounds with NOF is much higher than the atoms found in the same group.

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188
Q

what is vapour pressure

A

pressure exerted by the gaseous phase of the liquid that evaporated from the exposed surface of the liquid

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189
Q

relationship between vapour pressure and IMF and boiling point

A

weaker the forces, higher the vapour pressure and more easily it evaporates
indirectly related to boiling point. the higher the boiling point the Lower the vapour pressure, it doesn’t boil easily.

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190
Q

what substances are said to be volatile

A

substances with high vapour pressure

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191
Q

relationship between vapour pressure and temperature

A

Direct

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192
Q

Define ionic solids

A

solid held together by electrostatic attraction between cation and anion in a lattice structure.
Strong and solid at room temp
intermolecular forces are similar to intramolecular forces

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193
Q

In an ionic solid, what determines the strength of the bond: charge of the ion or the size of the ion?

A

Charge of the ion

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194
Q

Define a network solid

A

atoms are connected in a lattice of covalent bonds (all interaction between the atoms are covalent bonds)
only has intramolecular forces
hard solid, very strong at RT

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195
Q

define metallic solids and conduction electrons

A

covalently bound lattice of uncle or inner shell electrons with at least one freely roaming valence electron called conduction electron.
excellent conductors of heat and electricity
malleable and ductile
almost all are solid at RT

196
Q

Molecular solids

A

particles are molecules that are Held together by intermolecular interactions - D-D, hydrogen bonding and LDF.
weaker than ionic, network or metallic bonds
liquid gasses at RT, solids if their strength of intermolecular force increases

197
Q

What is thermodynamics?

A

Transformation of the energy from one form to the other

198
Q

describe the zeroth law of thermodynamics

A

if two systems are in thermal equilibrium with a third system, then the two initial systems are in thermal equilibrium with one another.
when systems are in thermal equilibrium with one another, their temperatures must be the same
when bodies of different temperature are Brought together, the heat will flow from body with the higher temperature to into the body with lower temperature.

199
Q

Describe the first law of thermodynamics

A

the total energy of the universe is constant. energy may be transformed from one form to another but it cannot be created or destroyed
isolated system has constant energy - first law
work can be put into a system to increase its overall energy.

200
Q

Describe three distinct designations to describe energy flow

A

system - the thing we are studying
surroundings - everything else
thermodynamic universe - system +surroundings

201
Q

when the energy flows into the system from the surroundings, the energy of the system — and the energy of the surroundings —-.

A
201
Q

when the energy flows into the system from the surroundings, the energy of the system — and the energy of the surroundings —-.

A
202
Q

when the energy flows into the system from the surroundings, the energy of the system — and the energy of the surroundings —-.

A

increases, decreases

203
Q

when the energy flows out of a system into the surroundings, the energy of the system — and the energy of the surroundings —-.

A

decreases, increases

204
Q

Define enthalpy

A

measure of heat energy that is released or absorbed when bonds are broken or created during a reaction under constant pressure

205
Q

2 principles of enthalpy

A

when bonds are formed, energy is released change in H < 0
when bonds are created, energy is absorbed change in H >0

206
Q

how do you find enthalpy of the reaction?

A

Enthalpy of products - enthalpy of reactants

207
Q

Is enthalpy also known as heat of reaction (delta H)? t/f

A

True

208
Q

Define exothermic and endothermic reactions

A

exothermic: products have a stronger bond than reactants, more energy is released in making the product bonds than was put in to break the reactant bonds. energy from the overall system is released.

Endothermic: products have a weaker bond than reactants, more energy is put in to break the reactant bonds than it is released.

209
Q

Are products and reactants higher or lower in energy in an endothermic reaction? exothermic reaction?

A

Endothermic: products are higher in energy state than reactants, H is positive

exothermic: products are low in energy than reactants, H is negative

diagram on pg - 124

210
Q

Which of the following does not contribute to the change in enthalpy?
phase change, formation of stronger intermolecular forces, breaking covalent bonds or a catalyst

A

A catalyst lowers the activation energy but doesn’t effect the enthalpy in any way

211
Q

What is standard heat of formation?

A

It is the amount of energy required to make one mole of a compound from its constituents elements in their natural or standard state. It is donated by ΔHfo. The ΔHfo of the elements in their standard form is 0

212
Q

ΔHfo for C (S) (graphite) is zero. True or false

A

True

213
Q

What is ΔHfo for diatomic elements? (O2, H2, CL2)

A

Zero

214
Q

Why is it that ΔHfo for O2 is zero but for O ΔHfo = 349 Kj/mol?

A

Because it takes energy to break the double bonds of O2

215
Q

What does it mean if ΔHfo is positive? negative?

A

Positive : input of heat is required to make that compound from its constituents
negative: making the compound from its constituents gives off energy

216
Q

How do you find ΔHo reaction?

A

subtract heat of reaction from heat of products

217
Q

Why is ΔHfo for Br2 (g is not zero?

A

Because under standard conditions, br2 is liquid and not gas

218
Q

What is hess’s law of heat summation

A

if the reaction occurs in several steps, then the sum of energy absorbed or given off in all the steps will be the same as that of overall reaction

219
Q

ΔH is dependent on the pathway of the reaction? True or False

A

False, ΔH is independent of the pathway of the reaction due tp Hess’s law pg 127

220
Q

Enthalpy is energy stored in the chemical bonds of the compounds? T/F

A

True

221
Q

Describe the phenomenon to create and break a bond (what should be endothermic and exothermic)

A

breaking a bond should be endothermic
making a bond should be exothermic

if it is the opposite then the bonds would be unstable and immediately dissociate

222
Q

What is bond dissociate energy

A

energy required to break a bond homolytically. every bond has characteristic enthalpies to denote this1

223
Q

if a question provides bond enthalpies, how would you find ΔH reaction?

A

ΔH rxn = sum of BDE bonds broken - sum of BDE bonds formed

224
Q

if stronger bonds are being made than those broken the energy would be positive? T/f

A

False, it would be negative

225
Q

if stronger bonds are being broken are being broken then the reaction would be exothermic? t/f

A

false, it would be endothermic

226
Q

reaction enthalpy by using heat of formation and bond dissociation energy is not the same. Why?

A

bond energies are reported as an average of many example of that type of bonds whereas heat of formation are determined for each individual chemical compound.

227
Q

What is the second law of thermodynamics

A

disorder of the universe increases in a spontaneous process. All processes tend to run in the direction of maximum disorder.

228
Q

Define entropy

A

entropy measures disorder or randomness. direct relation. represented by S

229
Q

Change in entropy formula?

A

S products - S reactants

230
Q

When is 🔼 S positive or negative

A

S is positive when the randomness increases and order decreases
S is negative when the randomness decreases or the order increases

231
Q

As the molecule becomes more organized, the entropy decreases? T/F

A

True
H2CO3 ⇌ CO2 + H2O
the entropy of forward reaction is positive/increasing
for reverse reaction, entropy is decreasing

232
Q

State any general predictions for entropy

A

gasses have more entropy than solids and liquids
liquids have more entropy than solids
two moles have more entropy than one mole
particles in a solution have more entropy than in an undissolved solution
the value of △ S for forward reaction is same as the reverse but opposite

233
Q

review examples from pg-130 to 132

A
234
Q

Describe third law of thermodynamics

A

absolute zero is a state of zero-entropy. at absolute zero, thermal energy is absent and only at least energetic thermodynamic state is available. There is no randomness to the system S=0

235
Q

What is Gibbs free energy

A

it is the energy that is available to do useful work from a chemical reaction

236
Q

how do you determine spontaneity of the reaction

A

using entropy, enthalpy, and Gibbs free energy

237
Q

Predict spontaneity if △H is positive and △S is negative

A

non spontaneous. products have more energy and less disorder

238
Q

Predict spontaneity if △H is negative and △S is positive

A

Spontaneous, products have less energy and more disordered

239
Q

If the signs are same for △H and S, then G is calculated. True/False

A

True

240
Q

Formula for △G

A

△G = △H - T△S

241
Q

Predict spontaneity
△G < 0
△G = 0
△G > 0

A

negative = spontaneous in forward direction
0 = equilibrium
positive = non spontaneous in forward direction

If △G for forward direction is positive, then the reverse reaction is negative therefore the reverse reaction is spontaneous.

242
Q

relationship between △G and temperature

A

direct

243
Q

At low temperature, the entropy has the same effect on free energy

A

False, at low temperatures entropy doesn’t have that much influence

244
Q

in general universe tends towards increasing disorder and stable bonds, what is the sign for H and S

A

Increasing disorder - positive △S
Stable bonds - negative △H

245
Q

What is the use of reaction energy diagrams

A

the graph plots the free energy of the total reactions versus the conversion energy of reactants into the products

246
Q

define activation energy

A

the energy that reactants require to overcome the activation barrier

247
Q

higher activation energy = slow speed of reaction towards equilibrium. T/F

A

True and vice versa

248
Q

A large Ea means a large K eq. T/f

A

True

249
Q

Does a spontaneous reaction mean the reaction would taking place rapidly

A

No, spontaneous reaction means that under standard conditions the formation of a compound is favoured. Rea pg -134 last para

250
Q

A reverse reaction has the same magnitude for all thermodynamic values but the opposite sign. name these values

A

△H. △G, △S

251
Q

if you are given the graph of a forward reaction, how would you derive the graph of the reverse reaction

A

it is the mirror image of the forward reaction

252
Q

Name one value that would different for forward and reverse reactions

A

Ea would different. forward rxn - Ea = G of reactants - activated complex
but in reverse rxn - Ea = G of products - activated complex

253
Q

define temperature

A

measure of internal kinetic energy that molecules have

254
Q

Relationship between kinetic energy and entropy

A

Direct, more kinetic energy more entropy

255
Q

define phase change

A

breaking or forming of intermolecular interactions

256
Q

diagram for phase transitions

A

pg 145
heat absorbed, kinetic energy increases, entropy increases during fusion (melting of solid), boiling and sublimation

Heat is released, entropy decreased, kinetic energy decreased during condensation, crystallization (liquid to solid) and deposition (gas to solid)

257
Q

define heat of transition

A

the amount of energy required to complete a phase transition

258
Q

define heat of fusion

A

heat that must be absorbed to change a solid into a liquid

259
Q

define heat of vaporization

A

energy absorbed when a liquid Changes into gas

260
Q

formula for heat required for a phase transition

A

q = n x △H phase change

261
Q

heat of transition can be in joules and cal/g. t/f

A

True

262
Q

define calorie

A

amount of heat required to raise the temperature of 1g of water by 1 degree C

263
Q

1 cal =? J

A

4.2 J

264
Q

two things that a change of phase depends upon

A

type of substance and amount of substance

265
Q

when a substance absorbs or releases heat, what two things can happen

A

either its temperature changes or it will undergo a phase change but not both at the same time

because the energy absorbed is used to break the bonds between the molecules during the phase change and not to increase the kinetic energy of the sample, so the temperature doesn’t change `

266
Q

relation between heat capacity constant and specific heat

A

C =mc
q = CT
q = mct

267
Q

define specific heat

A

it is an intrinsic property of the substance that tells how resistant it is to changing its temperature

268
Q

relationship between temperature and specific heat capacity

A

inverse t = q/mc.

269
Q

substance with high specific heat will undergo high change in temperature than a substance with lower specific heat. T/F

A

False it the opposite

270
Q

The numerical value for specific heat is different in kelvin and Celsius. T/f

A

False

271
Q

draw phase transition diagram

A

pg 149

272
Q

what are the horizontal and vertical axes of the heating curve

A

horizontal - heat aded
vertical - temperature

273
Q

What does a long flat line represent on the heating curve

A

high value of heat of transition

274
Q

substance’s heat of vaporization is smaller than heat of fusion?T/F

A

False, it is the opposite

275
Q

What is the slope of temperature change lines on the heating curve

A

1/C, reciprocal of substance heat capacity at the phase

276
Q

pressure influences the phase of a substance. T/F

A

True

277
Q

what does a phase diagram show

A

how phases are determined by temperature and pressure pg-151

278
Q

in the phase diagram, when are two phases in equilibrium

A

at the boundary lines

279
Q

what phase is favoured at low temp and high pressures

A

solid

280
Q

what phase is favoured at high temperature and low pressure

A

gas

281
Q

What is a triple point in phase diagrams

A

where all three phases exist simultaneously in equilibrium and therefore, all phase changes are happening simultaneously

282
Q

what is a critical point on phase diagrams

A

it marks the end of liquid gas boundary beyond which a substance is called supercritical fluid (gas and liquid phases are not distinct) and no amount of increased pressure can force this substance into liquid phase

283
Q

What is different about water phase diagram

A

water is denser in liquid phase than solid phase therefore the solid liquid boundary line has a slight negative slope.
For water, increase in pressure favours the liquid phase and not the solid phase

284
Q

What are the most important properties of gas

A

pressure, volume and temperature

285
Q

What are 4 laws of kinetic-molecular theory

A
  1. the modules of gas are so small compared to the average spacing between them that the molecules themselves take up no volume
  2. The molecules of gas are in constant motion, moving in straight lines at constant speeds and in random directions between collisions. The collisions of the molecules with the walls of the container define the pressure and all these collisions are elastic
  3. since each molecule moves at a constant speed and all the collisions are elastic. the molecule of gas experiences no intermolecular forces.
  4. The molecules of a gas span a distribution of speed and the average kinetic energy of the molecules is directly proportional to the absolute temperature
286
Q

If a gas satisfies all the conditions of kinetic molecular theory, what is it called

A

ideal gas

287
Q

Most real gases behave like ideal gas. t/f

A

True

288
Q

What is SI unit for volume in chemistry

A

cm3 and L

289
Q

1 cm3 = ?L

A

1000L

290
Q

What unit is used to measure gas temperature

A

Kelvin

291
Q

T in K = T in C +—–

A

273.15

292
Q

At 0 Kelvin, what is the entropy of gas?

A

0

293
Q

What is the SI unit for pressure

A

Pascal, but it very small for measurements so other units are used

294
Q

What are some other units that are used for pressure

A

1 atm = 760 tor = 760 mm hg = 101.3 KPa

295
Q

Define Standard temperature and pressure

A

STP means temperature is 0 degree Celsius (273.15 K) and a pressure of 1 atm

296
Q

What would be the reading of a barometer filled with a liquid of lower density than Hg if at that moment another nearby Hg barometer reads 752 mm Hg?

A

Look the the picture of a barometer, greater than 752mmhg

297
Q

How do pressure, temperature and volume of an ideal gas relate to one another

A

Pv=nrt, as most gas behave as ideal gas this law can be applied to them

298
Q

What is the value of R in pv =nrt

A

0.0821

299
Q

Describe Charles’s Law

A

Value is proportional to temperature. If the pressure is to remain constant, then a gas will expand when heated and contract when cooled. If the temperature of the gas increases, the molecules will move faster and hit the wall of the container wth more force

300
Q

What is Boyle’s law

A

Pressure is indirectly proportional to volume. if the volume decreases, the molecules have less space to move around =, they will collide with the walls of the container more often and the pressure increases

301
Q

Presssure is proportional to temperature. T/F

A

True. if the temperature increases, the molecules will move faster and collide with the walls more often

302
Q

In Charles law, name the variable that is constant

A

n and P

303
Q

In Boyle’s law name the variable that is constant

A

n and T

304
Q

What is combined gas law

A

combined Boyle’s and Charles’s law where n is constant

305
Q

define Avogadro’s law

A

If two equal volume containers hold gas at the same temperature and pressure, then they contain the same number of particles regardless of the identity of the gas.

306
Q

How would you put Avogadro’s law into an equation

A

V/n = k

307
Q

What volume does an ideal gas occupy at STP where T = 0 degree Celsius and P = 1atm

A

22.4 L , see the calculation at pg 165

308
Q

review how to find the limiting reagent

A
309
Q

What two assumptions from the kinetic molecular theory deviate from the idea gas behaviour.

A
  1. the particle of gas experience no intermolecular forces. P real < P ideal under fixed volume
  2. the volume of the gas is negligible compared to the volume of the container. V real < V ideal under fixed pressure
310
Q

When does a gas behave ideally, in terms of pressure, temperature and volume, intermolecular forces and weights

A

high temperature, low pressure, high volume, low intermolecular forces and smallest weights

311
Q

What is Van der Waals equation

A

equation accounts for the differences in observed real gas behaviours while maintaining the same gas law
pg 168 for equation. (P + an2/v2) (V- nb) = nrt where A is the intermolecular forces and B is the volume

312
Q

Van der Waal constants, a and b, are smaller for gases with large intermolecular forces and molecular weights? T/F

A

False

313
Q

Describe Dalton’s law of partial pressure

A

total pressure is the sum of the partial pressure of all the constituents gases.

314
Q

How do you find the mole fraction?

A

gas molecule of individual gas / total gas molecules

315
Q

how do you find the partial pressure of the gas given the total pressure

A

mole fraction x Total pressure

316
Q

what is effusion?

A

the escape of a gas molecule through a very tiny hole into an evacuated region is called effusion

317
Q

What is rms while finding masses from average kinetic energy

A

root mean square, the square root of the mean of the square of speed

318
Q

What is graham’s law of effusion?

A

rate of effusion of Gas A / rate of effusion of Gas B = ⎷molar mass of Gas B / Molar mass of gas A

319
Q

molecule of gas at the same temperature would have the same average speed but different average kinetic energy? T/F

A

False, same KE but different average speed

320
Q

What is a reaction intermediate

A

a substance that is produced in one elementary step and consumed in the other

321
Q

Intermediates are part of the overall balanced reaction?

A

False

322
Q

What is the rate determining step of an reaction?

A

the slowest step in the process determines the rate

323
Q

If the last step of the reaction is slow, the intermediates would pile up?

A

False, the intermediaries will go back and forth between reactants and products

324
Q

What is the formula for rate in general

A

rate = △[A]/△t

325
Q

Define rate of the reaction

A

how fast the reactant are being consumed or how fast the products are being made

326
Q

Name three factors that effect the reaction rate

A

how frequently the reactant molecules collide
the orientation of the colliding molecules
their energy

327
Q

What is activation energy

A

the minimum energy required by the reactants during the molecular collision for the reaction to proceed to products

328
Q

what is activated complex

A

when reactants have the necessary activation energy they reach a high energy but short lived transition state called activated complex

329
Q

draw a reaction coordinate diagram

A
330
Q

What is the energy maximum in the reaction coordinate diagram

A

Activated complexes

331
Q

An intermediate has more energy than reactants and products?

A

True

332
Q

What are the three variables that increase the rate of reaction

A

temperature, lower activation energy and high concentration of reactants

333
Q

G has no effect on the rate fo reaction because thermodynamic factors and kinetic factors do not effect each other?

A

true

334
Q

How does a catalyst speed up a reaction

A

by speeding up the rate determined step or lowering the activation energy of the rate determining step (energy of the highest transition state). Ea of forward and reverse reaction is lowered

335
Q

Catalyst changes with the reaction?

A

it undergoes a temporary change but always returns back to its initial state so it remains unchanged

336
Q

catalysts are included in the overall reaction?

A

False

337
Q

Catalyst only changes the activation energy and not the equilibrium or thermodynamic quantities of the reaction?

A

True

338
Q

How would you find the rate law of a reaction?

A

rate = k [A] ^x [B]^y, k = the rate constant

339
Q

You can figure out the orders of reactants by looking at the balanced equation

A

False

340
Q

Formula to find the rate constant

A

k = rate / [A][B]^2

341
Q

What is Arrhenius equation

A

k = Ae ^ - (Ea/RT) or k = ln - (Ea/RT)

342
Q

k = ln - (Ea/RT), how would adding a catalyst effect the Arrhenius equation

A

adding the catalyst, lowering the Ea, increase the T and Increase k

343
Q

rate increases by a factor of 2 to4 for every 10 degree Celsius increase in temp?

A

True

344
Q

rate = k [a][b], how do the units of k change with order

A

rate has a unit of M/s, first order reaction for both, M/s /M/M = 1/M.s

345
Q

What is equilibrium

A

When the forward reaction and reverse reaction come into a balance

346
Q

At equilibrium forward and reverse reaction occur at different rates?

A

False

347
Q

The overall conc of reactants and products remain the same but at molecular level they are different?

A

True

348
Q

Equilibria occurs for open systems?

A

False, Closed systems

349
Q

equation for equilibrium constant

A

k eq = [C]^c [D]^d / [A]^a [B]^b

350
Q

What is mass action ration

A

Where square brackets represent the molar conc. at equilibrium [C]^c [D]^d / [A]^a [B]^b

351
Q

What is not included in the equilibrium equation

A

solids, pure liquids and solvents

352
Q

The value of K eq is constant at a given temperature

A

True

353
Q

What does keq>1, keg<1 keq=1 mean?

A

Keq >1 - reaction favours the product
k <1 rxn favours the reactants
k = 1 reaction has equal amounts of reactants and products

354
Q

What is the reaction quotient (Q)

A

if the reaction is not at equilibrium then the same equation for Keq is known as Q

355
Q

How does comparing Q and Keq tell us how the reaction would proceed

A

K = Q, equilibrium
Q < K -forward
Q > K - reverse

Q → K ← Q
Q< k, rxn forward towards k
Q>K rxn reverse, towards k

356
Q

What is le chatelier principle

A

system at equilibrium will try to neutralize any imposed change or stress to reestablish equilibrium

357
Q

How does change in volume of the reaction container effect the equilibrium of the gas

A

If the volume is reduced, the pressure increases and the reaction favours the side with small amount of total moles
if the volume is increased, the pressure decreases and the net reaction favours the side with greater total moles
only applicable to gas equilibrium

358
Q

Describe changing the temperature with respect to le chatelier principle

A

Temp can be seen as a reactant or product in the Rxn based on whether it is exo or end
removing heat favours exothermic rnx
adding heat favours endothermic tnx

359
Q

Changing temperature, pressure and concentration will change the Keq of the reaction? T/F

A

Not all, only temp will change the Keq

360
Q

According to le chatelier principle how does adding an inert gas effect the equilibrium

A

does not effect the equilibrium

361
Q

What is the case when adding an inert gas would change the equilibrium according to the le chatelier principle

A

inject inert gas into constant pressure container one with movable piston, some extra gas particles will push against the piston. they will increase the volume and decrease the partial pressure and would cause a shift in equilibrium

362
Q

Would adding a catalyst to the reaction effect equilibrium

A

No

363
Q

What is a solution

A

when one substance dissolves into another

364
Q

what is dissolution

A

process of dissolving

365
Q

Define solvation

A

when solvent molecules surround the solute molecules

366
Q

What is hydration

A

When water is used as a solvent and the molecules surround the solutes

367
Q

Solutions can be made out of any three phases of matter?

A

True

368
Q

Define aqueous solution

A

when water is the solvent

369
Q

How can you tell what solutes are soluble to what solvents

A

like dissolves like (intermolecular forces)

370
Q

Ionic substances dissociate into ions?

A

True

371
Q

What are electrolytes

A

Free ions in a solution that conducts electricity

372
Q

Do all ions dissociate fully?

A

No

373
Q

What are strong and weak electrolytes

A

solutes that dissociate completely are strong and those who remain ion-paired are weak

374
Q

What are non-electrolytes

A

Covalent bonds that don’t dissociate at all

375
Q

what is van’t Hoff factor/ ionization factor

A

how many ions one unit of substance will produce

376
Q

what is the ionization factor for C6H1206, NACL, HNO3, CACL2

A

1 Because non-electrolyte
2 na and cl
2, h no3
3, ca and 2 cl-

377
Q

At molar solubility precipitation and dissolution occur at the same time

A

True

378
Q

What is molarity

A

n / volume

379
Q

What is solubility

A

Amount of solute that will saturate a solvent

380
Q

What are the phase solubility rules

A

solid solubility liquid in water with temp
gas solubility decreases in liquid with increasing temp
gas solubility increases in liquids with increasing pressure

381
Q

What are salt solubility rules

A

All group 1 are soluble
all nitrates, percolates and acetates are soluble
lead, silver and mercury are insoluble unless they are nitrate, percolate or acetate

382
Q

Define solubility product constant

A

the extend to which a solid will dissolve in water

383
Q

What is molar solubility

A

the number of moles of a salt that will saturate a solution

384
Q

What is the ion product of the solubility reaction

A

reaction quotient when solution is not at equilibrium

385
Q

Qsp < K
Qsp >K
Qap = K

A

more salt can be dissolved
precipitation
solution is saturated

386
Q

Adding a ligand to complex ions would increase/decrease the solubility of the ions

A

Increase the solubility and forms a salt

387
Q

One biological application of complex ion formation (review ICE tables khan academy)

A

Metal chelation therapy where metal chelation agent called (one word )– ehtylene diamine tetra acetic (EDTA) is used for lead poisoning

388
Q

What is the equation that related Q/K to G

A

△G = △G degree + RT ln Q
0 = △G degree + RT ln K, G = 0 AT EQUILIBRIUM

389
Q

WHAT IS THE DIFFERENCE BETWEEN G DEGREE AND G

A

G DEGREE IS THE RELATIVE PROPORTION OF REACTANTS AND PRODUCTS AT EQUILIBRIUM
G IS THE SPONTANEITY OF THE REACTION

390
Q

WHEN G DEGREE < 0, K > 1, WHAT IS FAVOURED
G DEGREE = 0, K = 1,
G DEGREE >0, K < 1,

A

PRODUCTS
EQUILIBRIUM
REACTANTS

391
Q

Define Bronsted Acid and Base

A

Bronsted acid are proton donors and bases are proton acceptors

392
Q

Define Lewis base and acid

A

Lewis base are electron base donor and acid are electron pair acceptor

393
Q

No matter what definition of acid or base you use, the acid and base remain the same? Lewis acid would be the same as the bronsted acid

A

True

394
Q

what are conjugate acids and bases

A

when Bronsted acid donates a proton, the remaining structure is called conjugate base of the acid.
when the base accepts a pair the new species is called conjugate acid of the base

394
Q

what are conjugate acids and bases

A

when Bronsted acid donates a proton, the remaining structure is called conjugate base of the acid.
when the base accepts a pair the new species is called conjugate acid of the base

395
Q

how do you differentiate between strong and weak acid

A

strong acid dissociates in water completely

396
Q

Describe the acid ionization constant value with respect to 1 to describe a strong or weak acid
Ka>1
ka<1

A

Ka > 1, strong acid
Ka <1 weak acid

397
Q

Name 6 strong acids

A

HI, HCL, HBr, HClO4, H2so4, HNO3
So I Brought No Clean Clothes.

398
Q

A strong acid would not have a stable conjugate base?

A

False, more stronger the acid, more stable the base

399
Q

A strong acid would not have a stable conjugate base?

A

False, more stronger the acid, more stable the base

400
Q

Why is HF not a strong acid

A

the conjugate base, F- has a smaller radius making it less stable, the weaker acids have less stable conjugated bases

401
Q

How can you tell if something is a strong acid/base

A

if K b has a greater value the base is strong

402
Q

Name 4 categories of strong bases

A

group 1 oxides, group1 hydroxides, group 2 with 4 s orbital and metal amides

403
Q

name two categories of weak bases

A

one that contain ammonia and amines

404
Q

Strong acid do not have a reverse reaction

A

True

405
Q

The conjugate base of a strong acid has no basic properties in water

A

True

406
Q

4 statements describing the strength of conjugate acid and bases

A

Conjugate base of a strong acid has no basic properties in water
conjugate acid of a strong base has no acidic properties in water
conjugate base of a weak acid is a weak base
conjugate acid of a weak base is a weak acid

407
Q

Weaker the acid/base, stronger the conjugate base/acid

A

True

408
Q

The weaker the acid/base, the less reverse reaction is favoured

A

False, more favoured q

409
Q

Weaker the conjugate acid, weaker the base?

A

False, stronger the base

410
Q

What is an amphoteric substance

A

the substance that can act as an acid or base

411
Q

What is polyprotic acid

A

substance that has more than one proton to donate

412
Q

The conjugate base of a weak polyprotic acid is always amphoteric

A

True, polyprotic acid such as H2co3 has a base of HCO3- where this conjugate base can donate or accept another proton

413
Q

Every time a polyprotic species donates a proton, it becomes stronger

A

False

414
Q

Water is amphoteric

A

True

415
Q

What is auto ionization of water

A

water reacts with itself acting as base and acid

416
Q

What is the equilibrium expression for auto ionization of water

A

k w = [h3o+] [OH-]
kw = 1.0 x 10^-14
[H3O+] = 1.0 x 10 ^ -7
[OH-] = 1.0 x 10^-7

417
Q

All equilibrium constants increase as the temperature increases

A

True

418
Q

What is a pH scale

A

It measures the conc of H3O ions in a solution

419
Q

Write the expression for pH

A

pH = -log [H3O+]
H3O+ = 10^-pH

420
Q

When is a solution acidic, neutral and basic based on value of Ph

A

pH > 7 basic
pH = 7 neutral
pH< 7 acidic

421
Q

What is pOH?

A

Measure the basicity of a solution

422
Q

formula for POH

A

pOH = -log [OH-]
10 ^(-pOH) = [OH-]

423
Q

What is the relationship bw [oh-] [h3o+]

A

Inverse

424
Q

PH + POH = ?

A

14

425
Q

IF YOU ARE TOLD THAT [H+] = 6.2 X 10 ^-5, HOW WOULD PREDICT THE PH

A

PH WILL BE IN BETWEEN (5-1 = 4) AND 5

425
Q

IF YOU ARE TOLD THAT [H+] = 6.2 X 10 ^-5, HOW WOULD PREDICT THE PH

A

PH WILL BE IN BETWEEN (5-1 = 4) AND 5

426
Q

THE ACID WITH A HIGH PKA VALUE IS A STRONG ACID

A

FALSE, ACIDS AND BASES WITH SMALL PKA VALUE ARE STRONGER

427
Q

Ka x Kb = ?
where Ka is acid and K b is a conjugate base

A

1 x 10^-14

428
Q

pKa+pkb = ?

A

14

429
Q

`What is ⎷10?

A

3.2

430
Q

Weak acid at equilibrium has a greater conc of undissociated ion than hydrogen ions at the product side.

A

True

431
Q

When is x dded or subtracted negligible in ICE tables

A

When Ka < 10^-4

432
Q

what is a neutralization process

A

when acid and base react together to make water and salt

432
Q

what is a neutralization process

A

when acid and base react together to make water and salt

433
Q

Taking an antacid to relieve excess stomach acid is an example of what type of reaction

A

neutralization

434
Q

Are neutralization process ends or exothermic

A

exothermic

435
Q

What is the formula for neutralization

A

a x [A] x Va = b x [B] x Vb
a is number of hydrogen and b is the number of hydrogen a base can accept

436
Q

what is a salt

A

ionic compound consisting of cation and anion

437
Q

What are the two predictions for the cation and anion in salt to predict whether solution would be acidic or basic

A

cation that doesn’t react with water (group1 and 2 cations) or is a stronger acid than water (nh4 +, Be2+, Cu 2+. Fe 3+, al3+, zn2+, cr3+)
anion that doesn’t react with water (conjugate base of a strong acid) or is a stronger base than water (conjugate base of a weak acid)

438
Q

what is a hydrolysis reaction

A

reaction of the substance with water

439
Q

if water has a strong conjugate base and a strong acid, how would you predict the pH?

A

compare the ka and Kb values

440
Q

What is a buffer solution

A

solution that resists changing pH when small amount of acid or base is added. contains a weak acid and its conjugate base or weak base and its conjugate acid

441
Q

How would you know if the buffer solution is effective

A

if the strong acid or bases added cause less change in pH

442
Q

What is the Henderson-hasselblach equation for buffer solutions for acid and base

A

pH = Pka + log ( [conjugate base]/ [weak acid])
poH = Pkb + log ( [conjugate acid]/ [weak base])

443
Q

An ideal buffer is solution is where weak acid conc = conjugate base conc. so pka = pH (Desired pH)?

A

True

444
Q

What is an indicator

A

a weak acid that undergoes colour change when it’s converted to its conjugate base

445
Q

Write an expression with indicator’s equilibrium constant and concentration of acidic reactants to predict indicator’s colour change

A

HA + H2O —–> H3O+ + A-
Ka =[H3o+] [A-] / [HA]
[H3o+] / Ka = [HA]/ [A-]

H3O+ > Ka then we will see colour of HA
if H3O = Ka, then we willl se a mix of colour of HA and A-
if H3O +< Ka, then we will see colour of A-

446
Q

What is a general pH range that an indicator changes colours

A

2 units pKa -1, pKa, Pka +1
read page 249 last para

446
Q

What is a general pH range that an indicator changes colours

A

2 units pKa -1, pKa, Pka +1
read page 249 last para

447
Q

What is an acid base titration

A

experimental technique used to determine the identity of an unknown weak acid or weak base by determining their Pka or Pkb

448
Q

what is a titrant during titration

A

a strong acid or base with known concentration and identity

449
Q

explain the procedure of titration

A

titrant is added to a solution of unknown base or acid in small amounts and Ph of the solution is recorded.

450
Q

You can titrate strong base with base or an acid with acid

A

False

451
Q

What is an equivalence point

A

during titration when the acid completely neutralizes the base or vice versa

452
Q

state the equivalence point for different titrations
weak acid and strong base
weak base and strong acid
strong acid and strong base

A
  1. ph > 7
  2. ph <7
  3. ph = 7
453
Q

at the beginning of the titration the solutions acts like a buffer

A

True

454
Q

What is a buffering domain?

A

the section of titration curve where pH changes very gradually

455
Q

The titration curve for polyprotic acids will have one equivalence point

A

False

456
Q

What is oxidation number

A

how many electrons a molecule is donating or accepting

457
Q

What are redox reactions

A

the reaction in which oxidation numbers of any reactants change are called redox reactions

458
Q

in a redox reaction, the total number of electrons for a reaction changes

A

False

459
Q

define oxidation and reduction

A

oxidation is when an atom loses an electron, the oxidation number increases
reduction is when an atom gains electrons and the oxidation number decreases

460
Q

nemonic to remember redox reactions

A

LEO says GER

461
Q

What is a reducing agent and oxidation agent

A

An oxidized atom is a reducing agent and a reduced atom is an oxidizing agent

462
Q

What is galvanic cell or voltaic cell

A

Generates electric current using the work done by redox reactions

463
Q

What are the components of the galvanic cell

A

anode, cathode, ammeter (measures electric current), salt bridge, electrolyte solution, 2 half cells and wire

464
Q

In the galvanic cell, the metal at anode gets reduced and cathode gets oxidized?

A

False, anode is oxidized and cathode is reduced
an ox, red cat

465
Q

Write a cell diagram equation where anode - zinc, cathode is copper, inSO 4 solution

A

zn (S)⎪zn 2+(aq) ⎪⎪cu2+ (aq)⎪cu(S)

466
Q

how do you convert a reduction potential to an oxidation potential

A

flip it, reverse the sign

467
Q

write the equation for free energy change of a redox reaction

A

G = -nFE, where F = 96500 C

468
Q

When is a redox reaction spontaneous

A

When G is negative and Voltage is positive

469
Q

While balancing a redox half reaction, you have to multiply the reductio potential as well?

A

False

470
Q

What determines a good or poor oxidizing agent or reducing agent

A

if reduction potential is largely negative, then the reactant is a weak oxidizing agent and products are a strong reduction agent
if the reduction potential is positive, the reactants are strong oxidizing agent and products are weak reduction agent

471
Q

When do we use the Nernst equation for redox

A

under non standard conditions where temperature and conc of reactants deviate

472
Q

What is a concentration cell

A

galvanic cell that has identical metal electrodes but have different ion concentrations of the solutions. The electrons go from low conc of the ionic solution to a high conc cell. Low conc is anode and high conc is cathode. the reactions stops when the conc of the solutions are equal

473
Q

a concentration cell has standard conditions?

A

False, because the solution is the Half cells in not 1 M

474
Q

What are electrolytic cells

A

these cells use an external voltage source to create an electric current that forces non -spontaneous redox reactions to occur

475
Q

The electrons flow from cathode to anode?

A

False, anode to cathode

476
Q

What are electroplating

A

placing a thin layer of metal on top of another material

477
Q

What are 4 differences between galvanic and electrolytic cells

A
  1. electrolytic is non-spontaneous
  2. the E is negative
  3. Cathode is negative in electrolytic cells
  4. anode is positive in electrolytic cells
478
Q

Name 5 similarities voltaic and electrolytic cell

A

reduction at cathode
oxidation at anode
cations migrate to cathode
anions migrate to anodes
e flows from anode to cathode

479
Q

What are two common rechargeable batteries through redox

A

lead-acid battery and nickel cadmium battery

480
Q

What is Faraday’s law of electrolysis

A

the amount of chemical change is proportional to the amount of electricity that flows through the cell

MCAT (2 decks)

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